Ultimate Guide: IB Chemistry (HL)

Unit 1: Stoichiometric Relationships

Unit 1.1: Particulate nature of matter

Elements

• all substances are made up of one or more elements

• over 100 known elements

• the smallest part of an element is called an atom

Compounds

• contains more than one element combined chemically in a fixed ratio

• eg. water (two hydrogen atoms and one oxygen atom)

• compounds have varying chemical and physical properties from their respective component elements

Mixtures

• components may be elements or compounds

• are not chemically bonded together

  • retain individual properties

• all in the same stage = homogeneous

• eg. air

States of Matter

Solid state

• fixed shape and volume

• particles held together by intermolecular forces

• particles will vibrate at a fixed point but do not have translational velocity

• turns into gaseous state by sublimation

• turns into liquid state by melting

Liquid state

• fixed volume

• takes up the shape of the container

• held together by intermolecular forces

• turns into solid state by freezing

• turns into gaseous state by boiling/vaporization/evaporation

Gaseous state

• takes the space to completely fill container

• widely spaced particles

• intermolecular forces between particles are negligible

• pressure is due to gaseous particles colliding with the walls of the container

• turns into liquid state by condensation

• turns into solid state by deposition

Unit 1.2: The mole concept and chemical formulas

Avogadro’s Constant

• a single atom of an element has a small mass

• this mass is known as Avogadro’s constant (N_A)

• the mass of one mole of any substance is known as the molar mass (M)

• relative atomic mass (A): weighted mean of all the naturally occurring isotopes of the element

  • find the average

  • for molecules: relative molecular mass

  • for ionic compounds: relative formula mass

Formulas of Compounds

• empirical formula: in simplest whole number the ratio of atoms of each element in a substance

  • obtained by knowing the mass of each element

• molecular formula: shows the actual number of atoms in a substance

• structural formula: shows the arrangement of atoms and bonds within a molecule

Unit 1.3: Chemical reactions and equations

Properties of Chemical Reactions

• new substances are formed

• bonds are broken and formed resulting in an energy change

• fixed relationship between the number of particles and reactants

Chemical Equations

• reactants are written on the left hand side

• products are written on the right hand side

• number of moles of each element must be the same on both sides in a balanced chemical equation'

• single arrow shows reaction goes to completion

• state symbols

  • s = solid

  • l = liquid

  • g = gas

  • aq = aqueous solution

• coefficient refers to the number in front of the reactants and products in the equation

  • shows information on the molar ratio

Ionic Equations

• ionic compounds completely dissociate in solution

• spectator ions do not need to be written in the reaction

Unit 1.4: Mass and gaseous volume relationships

Measurement of molar quantities

• solids: measure in mass (weighed)

• liquids: weighed or volume recorded

  • density = mass/volume (units is g cm^-3)

• gases: mass or volume

• solutions: units is litre, dm³, cm³)

  • solute (dissolved substance) in is a known volume of solution (solute plus solvent)

Fixed mass of gas

• P = 1/ V

• P = T

• V = T

• Ideal gas equation: PV = nRT

  • R = 8.31 (J K-1 mol-1)
    
    

Unit 1.5: Molar volume of a gas

• ideal gas equation depends on the amount of gas but not on the nature of the gas

• one mole of any gas will occupy the same volume at the same temperature and pressure

• at 273K and 100kPa, the volume is 22.7 dm3
  Calculations from equations

  1. write down the correct formulas and all the reactants and products

  2. balance the equation

  3. work out the limiting reagent. the limiting reagent is the maximum yield to any of the products that can be determined

  4. work out the amount (in mol) of the substance

  5. convert to required unit

  6. express solution in the correct number of significant digits and units

Unit 2: The Periodic Table 

What is the Periodic Table?

• The periodic table is an arrangement of all the elements known to man by their increasing atomic number and recurring chemical properties. 

• They are assorted in a tabular arrangement wherein a row is a period and a column is a group. Elements are arranged from left to right and top to bottom in the order of their increasing atomic numbers. 

• Thus,

  • Elements in the same group will have the same valence electron configuration and hence, similar chemical properties.

  • Whereas, elements in the same period will have an increasing order of valence electrons. 

    • Therefore, as the energy level of the atom increases, the number of energy sub-levels per energy level increases.

• The first 94 elements of the periodic table are naturally occurring, while the rest from 95 to 118 have only been synthesized in laboratories or nuclear reactors. 

  • The modern periodic table, the one we use now, is a new and improved version of certain models put forth by scientists in the 19th and 20th century. 

• Dimitri Mendeleev put forward his periodic table based on the findings of some scientists before him like John Newlands and Antoine-Laurent de Lavoisier. 

  • However, Mendeleev is given sole credit for his development of the periodic table.

Acceptance of Mendeleev Periodic Table:

• Dimitri Mendeleev, widely referred as the father of the periodic table put forth the first iteration of the periodic table similar to the one we use now. 

  • Mendeleev’s periodic law is different from the modern periodic law in one main aspect.

• Mendeleev modeled his periodic table on the basis of increasing atomic mass, whereas, the modern periodic law is based on the increasing order of atomic numbers.

• Even though Mendeleev’s periodic table was based on atomic weight, he was able to predict the discovery and properties of certain elements. 

• During his time only around half of the elements known to us now were known, and most of the information known about the elements were inaccurate. 

  • Mendeleev’s Periodic Table was published in the German Journal of chemistry in 1869.

Genesis of Periodic Classification:

Dobereiner’s Trend:

• German chemist Johann Wolfgang Dobereiner attempted to classify elements with similar properties into groups of three elements each. 

  • These groups were called ‘triads’. 

• Dobereiner suggested that in these triads, the atomic mass of the element in the middle would be more or less equal to the mean of the atomic masses of the other two elements in the triad.

  • An example of such a triad would be one containing lithium, sodium, and potassium. 

  • The atomic mass of lithium is 6.94 and that of potassium is 39.10. 

  • The element in the middle of this triad, sodium, has an atomic mass of 22.99 which is more or less equal to the mean of the atomic masses of lithium and potassium (which is 23.02).

• The Limitations of Dobereiner’s Triads are:

  • All the elements known at that time couldn’t be classified into triads.

  • Only four triads were mentioned – (Li,Na,K ),  (Ca,Sr,Ba) ,  (Cl,Br,I) , (S,Se,Te).

Newland’s Octaves:

• English scientist John Newlands arranged the 56 known elements in increasing order of atomic mass in the year 1866. 

• He observed a trend wherein every eighth element exhibited properties similar to the first. 

• This similarity in the properties of every eighth element can be illustrated as follows.

• Newland’s Law of Octaves states that when the elements are arranged in increasing order of atomic mass, the periodicity in properties of two elements which have an interval of seven elements in between them would be similar.

• Limitations of Newland’s octaves are:

  • It was only up to calcium that the classification of elements was done via Newland’s Octaves.

  • The discovery of noble gases added to the limitations of this method since they couldn’t be included in this arrangement without disturbing it completely.

Mendeleev’s Periodic Table:

• Russian chemist Dmitri Ivanovich Mendeleev put forth his periodic table in 1869. 

  • He observed that the properties of elements, both physical and chemical, were periodically related to the atomic mass of the elements.

• The Periodic Law (also referred to as Mendeleev’s Law), states that the chemical properties of elements are a periodic function of their atomic weights.

• The advantages of Mendeleev’s Periodic table are:

  • The inclusion of these newly discovered elements did not disturb the periodic table. Examples include germanium, gallium, and scandium.

  • It was used to correct the wrong atomic weights in use at that time.

  • A variance from the atomic weight order was provided by Mendeleev’s table.

• The limitations of Mendeleev’s Periodic table are:

  • Hydrogen’s position was in the group of alkali metals but hydrogen also exhibited halogen like qualities.

  • Isotopes were positioned differently since this type of classification of elements was done by considering the atomic weight of the element. 

    • Therefore – protium, deuterium, and tritium would occupy varying positions in Mendeleev’s table.

  • An anomalous positioning of a few elements showed that the atomic masses did not increase regularly from one element to the next. 

    • An example of this would be the placement of cobalt (atomic mass of 58.9) before nickel (atomic mass of 58.7).

• These methods were the foundation on which the modern periodic table was built. 

  • However, the greatest contributor to the modern periodic table was Dmitri Mendeleev. 

• Mendeleev is also known as the Father of the Modern Periodic Table. 

  • The modern periodic law is also called Mendeleev’s Law to honour him.

Modern Periodic Table:

• In the year 1913, English physicist Henry Moseley studied the wavelength of the characteristic x-rays. 

• By using different metals as anti cathode and showed that the square root of the frequency of the line is related to the atomic number. 

• On the basis of the above observations Moseley gave the modern periodic law which states that :

  • “Physical and chemical properties of the elements are the periodic function of their atomic numbers”.

  • The atomic mass of an element is due to the mass of protons and neutrons present in the nucleus of its atom. 

  • Since the nucleus is located inside an atom, it is not very much linked with the properties of the element, particularly the chemical properties.  

    • These are related to the number of electrons and also the distributions of the electrons in the different energy shells. 

  • The elements with different electronic arrangements of atoms possess different chemical properties. 

  • As the number of electrons in an atom is given by the atomic number and not by the mass number, therefore atomic number should form the basis of the classification of the elements in the periodic table and not atomic mass as predicted by Mendeleev. 

• Repetitions of the similar properties of the elements placed in a group and separated by certain definite gap of atomic number are known as Periodicity.

Classification of elements in modern periodic table

• The modern periodic table consists of 18 vertical columns, called the groups(1-18) and 7 Horizontal rows, called periods.

  • The first period contains two elements, Hydrogen and Helium.

  • The second period contains eight elements, from Lithium to Neon.

  • The third period contains eight elements, from Sodium to Argon.

  • The fourth period contains eighteen elements, from Potassium to Krypton.

  • The fifth period contains eighteen elements, from Rubidium to Xenon.

  • The sixth period contains thirty-two elements.

  • The seventh period is incomplete.

• On the basis of electronic configuration, elements are classified into four Blocks known as s, p, d and f- blocks.

  • 1st and 2nd group elements are called s-block elements. The general electronic configuration is ns1-2.

  • 13th to 18th group elements are called p-block elements. The general electronic configuration is ns2 np1-6.

  • 3rd to 12th group elements are called d-block elements. The general electronic configuration is (n-1)d1-10 ns1-2.

  • Lanthanides and actinides elements are called f-block elements. The general electronic configuration is (n-2)f1-14 (n-1)d0-1 ns2.

• Periodic properties and their trends

  • The periodic properties may be defined as: 

• The properties of the elements are directly or indirectly related to the electronic configuration of their atoms and show gradation (increases or decreases) in moving down a group or a longer period.

• The common physical properties of the elements are melting points, boiling points, density, enthalpy of fusion and vaporization etc. 

  • But we shall focus our attention mainly on the properties which are based on electronic configuration  these are:

    • Atomic and ionic radii 

    • Ionization enthalpy

    • Electrons gain enthalpy

    • Electronegativity

Atomic and Ionic Radii:

Atomic Radii:

• The atomic radius may be defined as the distance from the centre of the nucleus to the outermost shell containing electrons. Depending upon the nature of bonding in the atoms these are:

  • Covalent radii: One-half of the distance between the centres of the nuclei of two adjacent similar atoms joined to each other by a single covalent bond is known as covalent radii. 

    • Eg Cl-Cl bond distance=198 pm covalent radius of Cl= 99 pm.

  • Van Der Waals Radii: Half of the internuclear distance between two similar adjacent atoms belonging to the two neighbouring molecules of the same substance in the solid state is known as Van Der Waals Radii.

  • Metallic radii: Half the distance between the centre of the  nuclei of two adjacent atoms  in the metallic crystal is known as metallic radii 

• As we move from left to right in a period, the atomic radius decreases due to an increase in effective nuclear charge (Zeff). 

  • Along the group, as we move from top to bottom, atomic radius increases  due to increase in principal quantum number which causes an increase in the number of shells and increases in shielding effect.

Ionic Radii:

• The ions formed by the loss of one or more electrons from the neutral atom are known as cation (positive ion) when the electrons added to the neutral atom form an anion (negative ion). 

  • The effective distance from the centre of the nucleus of the ion upto which it exerts its influence on the electron cloud is known as the ionic radii.

• The ionic radii change in the same trend as atomic radii. 

  • It decreases along the period from left to right and increases down the group from top to bottom. size of cation and anion of any natural atom as: cation< neutral atom < anion 

• Ionization enthalpy:

  • The amount of energy required when an electron is removed from the outermost orbit of an isolated gaseous atom is known as Ionisation Enthalpy (IE).

• Generally left to right in period IE increases whereas on moving down the group it decreases but half-filled orbital and fully filled orbitals are highly stable and thus have high IE.

Electrons gain enthalpy:

• The electron gain enthalpy is defined as the change in enthalpy which takes place when a gaseous atom gains an extra electron to form a monovalent anion in the gaseous state.

  • Electron gain enthalpy increases across the periods while it decreases down the group.

    • Chlorine has the highest electron affinity than fluorine.

Electronegativity:

• Electronegativity is the tendency of an atom to attract the shared pair of electrons towards itself in a covalent bond. 

  • Fluorine is the most electronegative element while Cesium is the least. 

• In the periods left to right electronegativity increases.

  • In the groups while moving down the groups electronegativity decreases.

Unit 3: Periodicity

Unit 3.1 The Periodic Table & Periodic Trends

3.1.1 The Periodic Table

Structure of the Periodic Table

• The Periodic Table = A list of all known elements arranged in order of increasing atomic number (from 1 to 118)

• Elements are in rows and columns

  • Atoms with the same # of shells are placed in the same row

• Period = A row of elements

  • n = Period # = The outer energy level that is occupied by electrons

• Group = Elements aligned vertically in columns

  • These elements share a similar outer-shell electron configuration

  • Valence Electrons = electrons in the outer shell

The Blocks of the Periodic Table

All elements are categorized into 4 main blocks

• s-block = elements with only s electrons in the outer shell

• p-block = elements with at least 1 p-electron in the outer shell

• d-block = elements with at least 1 d-electron and at least 1 s-electron but NO f or p electrons in the outer shell

• f-block = elements with at least 1 f-electron and at least 1 s-electron but NO d or p electrons in the outer shell

Deducing Electron Configurations

• You can use the position of an element in the periodic table to help you figure out its electron configuration

• ex. Germanium = [Ar] 3d¹⁰4s²4p²

  • Ge is in Group 4 → Tells you that there are 4 Valence Electrons

  • Ge is in Period 4 → Tells you that the 4 VE are in the 4th shell

  • The 2nd position in the p-block → Tells you that 2 electrons are in the p-subshell

3.1.2 Periodic Trends: Physical - Atomic & Ionic Radius

Atomic & Ionic Radius

Atomic Radius

• Atomic Radius= the size of an atom = the distance between the nucleus of an atom and the outermost electron shell

  • Decreases across each period

  • Increases down each group

Electron Shell Theory

• As you move across a period, the amount of protons increases, thus increasing the pull of the nuclei on the electrons, so the atoms become smaller

• As you move down a group, more electrons are being added, so the radii increase

  • However, the outer electrons don’t feel the pull of the positive nuclei charge as much because the inner electrons repel them

Ionic Radius

• Ionic Radius = the size of an ion

  • Trend down a group is the same as atomic radius: as the number of shells increases, the radius increases

  • However, the trend across a period depends on whether an ion is positive or negative

• In general:

  • As the negative charge increases, the ionic radii increase

  • As the positive charge increases, the ionic radii decrease

Electron Shell Theory

• Negative ions are formed when an atom accepts extra electrons which experience repulsion between the other valence electrons

  • So that causes a larger ionic radius

• Positive ions are formed when an atom loses electrons

  • Since there are fewer electrons while the nuclear charge remains the same, each remaining electron feels the pull of the nucleus more strongly than before

  • Thus, the ionic radius decreases

3.1.3 Periodic Trends: Physical - Ionization Energy

First Ionization Energy

• Ionization Energy (IE) = The amount of energy required to remove 1 mole of electrons from 1 mole of atoms (in the gaseous state) to form 1 mole of gaseous ions

  • Measured in kilojoules per mole (kJ mol^-1)

• First Ionization Energy = the amount of energy to remove the first mole of electrons

Ionization Energies: Trends

• Increases across a period

• Decreases down a group

Trends influenced by 4 factors:

1. Size of the Nuclear Charge

   • As the atomic number (# of protons) increases

   • → the nuclear charge increases

   • → greater attractive forces between positive nucleus and outer electrons

   • → more energy is required to overcome these strong attractive forces when trying to remove an outer electron

2. Distance of Outer Electrons from the Nucleus

   • Electrons in shells further from the nucleus feel its pull less

   • → So the further out electrons are, the lower ionization energy they have

3. Shielding Effect of Inner Electrons

   • = When the electrons in the inner shells repel those in the outer shells, so they don’t feel the attractive forces of the nucleus as much

   • The greater the shielding, the lower the ionization energy of the outer electrons

4. Spin-Pair Repulsion

   • Paired electrons in the same atomic orbital in a subshell repel each other more than electrons in different atomic orbitals, so it’s easier to remove an electron

   • Tip: This is the reason why the first ionization energy is always the lowest

Ionization Energy (IE) Across a Period

• IE increases across a period because:

  • The nuclear charge across a period increases

  • The distance between the nucleus and outer electrons is the same

  • The shielding of the outer electrons by the inner is the same

Exceptions

• IE decreases suddenly between the last element in one period and the first element in the next period because:

  • The distance between the nucleus and the outer electrons increases

  • The inner electrons’ shielding increases

  • These 2 factors outweigh the increased nuclear charge

• IE slightly decreases between beryllium and boron since the 5th electron in boron is in the 2p subshell (which is further away from the nucleus than the 2s subshell of beryllium)

  • beryllium = 900 kJ mol^-1

  • boron = 801 kJ mol^-1

• IE slightly decreases between nitrogen and oxygen because there is spin-pair repulsion in the 2p subshell of oxygen

  • nitrogen = 1402 kJ mol^-1

  • oxygen = 1314 kJ mol^-1

Ionization Energy (IE) Down a Group

• As you go down a group, IE decreases because

  • The distance between the nucleus and the outer electrons increases

  • Electron shielding increases

  • The effective nuclear charge decreases as shielding increases

Successive Ionization Energies of an Element

• Successive ionization energies (the amount of energy when you keep taking off electrons) increases because removing the next electron always takes more energy than the previous

  • In other words, it is harder to remove an electron from an already positive ion than a neutral atom

• For several reasons:

  • As you remove more electrons, the shielding decreases, so the attractive forces increase along with an increased proton to electron ratio (so each electron feels the pull of the nucleus more and more)

• Note: increases in successive ionization energy is NOT constant, but dependent on each atom’s electron configuration

3.1.4 Periodic Trends: Physical - Electron Affinity

• Anions = negative ions formed when electrons gain electrons

• Electron Affinity (EA) = The amount of energy released when 1 mole of electrons is gained by 1 mole of atoms of an element in the gaseous state to form 1 mole of gaseous ions

  • Think of it as the opposite of Ionization Energy

  • Measured in kilojoules per mole (kJ mol^-1)

• The periodicity of Electron Affinity has a similar pattern to Ionization Energies (but inverted)

• The strongest “pull” on electrons → The greater amount of energy released when negative ions are formed

• EA decreases down a group because:

  • As the atoms take on more electrons, the attraction for more decreases because the shielding increases and the effective nuclear charge decreases

3.1.5 Periodic Trends: Physical - Electronegativity

Electronegativity: Definition

• Electronegativity = The ability of an atom to attract an electron pair in a covalent bond

  • Caused by the positive nucleus being able to pull the negative electrons closer

  • The Pauling Scale = A way to measure the value of electronegativity for each atom (on an arbitrary scale from 0.0 to 4.0)

Electronegativity: Factors

Nuclear Charge

• The more positively charged protons in the nucleus, the more nuclear attraction between them and the negatively charged electrons

  • Thus, an increase in nuclear charge means an increase in electronegativity

Atomic Radius

• Atomic Radius = The distance between the nucleus and valence electrons

• As the atomic radius increases, the valence electrons are farther away, so they have a decreased attraction to the nucleus and less electronegativity

Electronegativity: Trends

Down a Group

• Electronegativity decreases down a group

• Although more protons are added and you would expect the electronegativity to increase with the increased nuclear attraction, this isn’t the case due to electron shielding:

  • As more electrons are being added, more shells are being filled, so more shells means more distance between the nucleus and valence electrons

  • In this case, the increased distance (that causes a decreased electronegativity) matters more than the increased nuclear charge (that would have caused an increase in electronegativity)

Across a Period

• Electronegativity increases across a period

• As you go across a period, the number of protons increases, while the shielding remains the same across a period, since no new shells are being added

  • Therefore, the increased nuclear charge matters more in this case

• So the increased nuclear charge causes

  • Decrease in atomic radius (pulls electrons closer to nucleus)

  • Increase in attraction between nucleus and valence electrons

  • Increase in Electronegativity

3.1.6 Periodic Trends: Chemical

Metallic & Non-metallic Properties

• What causes these different properties? Trends in:

  • Atomic Radius

  • Ionic Radius

  • Electron Affinity (EA)

  • Ionization Energy (IE)

  • Electronegativity (EN)

• Low EN and low IE of metals allow their valence electrons to be able to move away from the nucleus or “delocalize”

• High EN and high EA of non-metals allows for their electrons to be shared and form covalent bonds

• Similarities in EN for the elements in the diagonal band separating metals from non-metals explains their behavior as" “metalloids”

Unit 3.2 Oxides, Group 1 & Group 17

3.2.1 Periodic Trends: Oxides Across a Period

Oxides Across a Period

• Oxide = a binary compound that contains oxygen and another element

  • ex. Carbon dioxide CO₂

• Since oxides are acid-base they show their chemical trend: they change from basic to amphoteric to acidic as you move across a period

• Amphoteric = having the ability to react chemically as either an acid or a base

• ex. Aluminum oxide (so it can react with an acid like HCl or with a base like NaOH)

Period 3 Oxides
The reason for the different acidic or basic natures is because of their structure, bonding, and electronegativity

• Electrons will be transferred to oxygen when forming oxides and providing an ionic bond because of the highest difference in electronegativity between oxygen and Na/Mg/Al

• On the other hand, Si, P, and S (elements with the least difference in EN) will share the electrons with oxygen to form covalently bonded oxides

3.2.2 Periodic Trends: Group 1 - The Alkali Metals

• Group 1 Metals = “Alkali Metals” = metals that form alkaline solutions with high pH when they react with water

• All elements in group 1 end with the electron configuration: ns¹

Physical Properties

• Soft and easy to cut (more so as you move down the group)

• Shiny

• Conduct heat/electricity

• Low melting points (decreases more going down the group since atomic radius increases so metallic bonding gets weaker)

• Low Density

Chemical Properties

• Reactive with oxygen and water in the air

  • Often kept under oil to prevent reactions

• Highly reactive with water

  • And produces an alkaline metal hydroxide solution and hydrogen gas

3.2.3 Periodic Trends: Group 17 - The Halogens

The Halogens

• Group 17 non-metals

• Poisonous

• Diatomic = forming molecules of 2 atoms

• Have 7 valence electrons

• Form halide ions by gaining one more electron to complete the octet rule

Physical Properties Trends (Down the Group)

• Density, Melting/Boiling Points of the halogens increase going down the group

• Reactivity decreases down the group

  • Why?

    • Halogens electron configurations all end in ns²np⁵

    • When the 7 valence electrons react, they need one more to have a full shell

    • Electron Affinity decreases

    • Atomic Radius increases

      • → # of Shells increases

      • → Shielding increases

      • → Distance to nucleus increases

      • → Weaker electrostatic forces to attract the 8th electron

      • Therefore, as you go down the group, the reactivity decreases (because it is harder to attract the 8th electron)

Reaction of Halogens with Halid Ions in Displacement Reactions

• Halogen Displacement = When a more reactive halogen displaces a less reactive halogen from an aqueous solution of its halide

Halogen Displacement Reactions

Chlorine and Bromine

• Chlorine + colorless potassium bromide = bromine (Orange Solution)

• Since chlorine is above bromine in group 17, it is more reactive

  • So chlorine will displace bromine

2KBr (aq)       +    Cl₂ (aq)  →     2KCl (aq)     +  Br₂(aq)

potassium bromide + chlorine → potassium chloride + bromine

Bromine and Iodine

• Since bromine is higher than iodine in group 17, it is more reactive

• Thus, bromine will displace iodine

  Br₂ (l) +      2NaI (aq)    →    2NaBr (aq)     +    I₂ (aq)

bromine +  sodium Iodide → sodium bromide + iodine

UNIT 4: Chemical Bonding and Structure

Unit 4.1 Ionic & Covalent Bonding

4.1.1 Forming Ions

• Metals are on the left of the Periodic Table

  • Lose electrons from their valence shell

  • Form positively charged cations

• Non-metals are on the right of the Periodic Table

  • Gain electrons

  • Form negatively charged anions

• Ionic Bonds = bonds formed from the transfer of electrons from a metallic element to a non-metallic element

  • Usually results in both the metal and non-metal having a full outer shell

  • Thus, these atoms have electronic configurations that are the same as a noble gas (elements with full outer shells)

Example:

Na (Sodium)

• It is a metal in group 1, so it has 1 valence electron

• After it loses this electron, it becomes Na⁺ (Sodium Cation)

Cation = A positively charged ion

• A sodium ion has the same electronic configuration as neon: [2, 8]

CL (Chlorine)

• It is a non-metal in group 17, so it has 7 valence electrons

• After it gains an electron from a different atom, it becomes Cl^- (Chloride Anion)

Anion = A negatively charged ion

• A chloride ion has the same electronic configuration as argon: [2, 8, 8]

Electrostatic Attractions = Attractions formed between the oppositely charged ions (cations and anions) to form ionic compounds

• Very strong, so it takes a lot of energy to break

• This is why ionic compounds have high melting points

4.1.2 Ionic Compounds

Ionic Lattices

Ionic Lattice = an evenly distributed crystalline structure formed by ions

• Ions arranged in a regular repeating pattern due to electrostatic forces of attraction between cations and anions

• This causes positive charges to cancel out negative charges, causing the final, overall lattice to be electrically neutral

Properties of Ionic Compounds

• Different types of structure/bonding → different physical properties (ie. melting/boiling points, electrical conductivity, and solubility)

Ionic Bonding & Giant Ionic Lattice Structures

• Ionic compounds are strong

  • Their strong electrostatic forces keep ions held together strongly

• Ionic compounds are brittle, so ionic crystals can split apart

• Ionic compounds have high melting and boiling points

  • This is because of their strong electrostatic forces between the ions that keep them held strongly together

  • As charge density of the ions increase, the melting and boiling points increase

    • This is due to the greater electrostatic attraction of charges

    • ex. Mg²⁺O^2- has a higher melting point than Na⁺Cl^-

• Ionic compounds are soluble in water because they can form ion-dipole bonds

• Ionic compounds only conduct electricity when molten or in solution

  • In those two cases, the ions can freely move around and conduct electricity

  • However, as a solid, the ions are fixed, so they are unable to move around

4.1.3 Formulae & Names of Ionic Compounds

Review:

• Ionic compounds are formed from a metal and a nonmetal bonded together

• Ionic compounds are electrically neutral (positive charges = negative charges)

Charges on Positive Ions

• Positive Ions:

  • All metals

  • Some non-metals

    • ex. NH₄⁺ (Ammonia) and H⁺ (Hydrogen)

Charges of ions depend on their position in the Periodic Table:

• Metals in Groups 1, 2, and 13 = charges of 1+, 2+, and 3+

• Charge on ions of the transition metals can vary which is why Roman numerals are often used to indicate their charge

  • ex. Copper (II) Oxide = copper ion has a charge of 2+

  • ex. Copper (I) Nitrate = copper ion has a charge of 1+

Non-metal Ions

• Non-metals in Groups 15-17 have a negative charge and have the suffix “-ide”

  • ex. Nitride, Chloride, Bromide, Iodide

• Elements in Group 17 = gain 1 electron → have a 1- charge

  • ex. Br^-

• Elements in Group 16 = gain 2 electrons → have a 2- charge

  • ex. O^2-

• Elements in Group 15 = gain 3 electrons → have a 3- charge

  • ex. N^3-

• Additionally, there are more polyatomic or compound negative ions (negative ions made up of more than one type of atom)

7 Common Polyatomic Ions

4.1.4 Covalent Bonds

Covalent Bonds

• Covalent Bonding occurs between 2 non-metals

• Involves the electrostatic attraction between the nuclei of 2 atoms and the electrons of their outer shells

• No electrons transferred, but only shared

• 2 atomic orbitals overlap and a molecular orbital is formed (see left side of image, “bonding”)

• Covalent bonding occurs due to the fact that electrons are more stable when attracted to 2 nuclei rather than only 1

  • Why are they more stable?

    • Because sharing electrons allows each of the atoms to achieve an electron configuration similar to a noble gas (octet rule)

• Usually, each atom provides one of the electrons in the bond

• A covalent bond is represented by a short straight line between 2 atoms

  • ex. H-H

• Note: Think about covalent bonds as electrons being in a constant state of motion, or “charge clouds” rather than an electron pair in a fixed position

Predicting Covalent Bonding

• Use differences in electronegativity to predict whether a bond is either covalent or ionic

Electronegativity & Covalent Bonds

• Electron density in diatomic molecules are shared equally

  • ex. H₂, O₂, Cl₂

Coordinate Bonds

• In simple covalent bonds, the 2 atoms involved share electrons

• However, some molecules have a lone pair of electrons that can be donated to form a bond with an electron-deficient atom (an atom that has an unfilled outer orbital)

  • So both electrons are from the same atom

• This type of bonding is called dative covalent bonding or coordinate bond

• An example of a dative bond is in an ammonium ion

  • The hydrogen ion, H⁺ is electron-deficient and has space for two electrons in its shell

  • The nitrogen atom in ammonia has a lone pair of electrons which it can donate to the hydrogen ion to form a dative covalent bond

Multiple Bonds

• Non-metals are able to share more than one pair of electrons to form different types of covalent bonds

• Sharing electrons in the covalent bond allows each of the 2 atoms to achieve an electron configuration similar to a noble gas

  • This makes each atom more stable

• It is not possible to form a quadruple bond as the repulsion from having 8 electrons in the same region between the two nuclei is too great

Bond Length & Strength

Bond Energy

Bond Energy = The energy required to break one mole of particular covalent bond in the gaseous states (in units of kJ mol^-1)

• The larger the bond energy, the stronger the covalent bond is

Bond Length

Bond length = internuclear distance of two covalently bonded atoms

• It is the distance from the nucleus of one atom to another atom which forms the covalent bond

• The greater the forces of attraction between electrons and nuclei, the more the atoms are pulled closer to each other

• This decreases the bondlength of a molecule and increases the strength of the covalent bond

• Triple bonds are the shortest and strongest covalent bonds due to the large electron density between the nuclei of the two atoms

• This increase the forces of attraction between the electrons and nuclei of the atoms

• As a result of this, the atoms are pulled closer together causing a shorter bond length

• The increased forces of attraction also means that the covalent bond is stronger

• Triple bonds are the shortest covalent bonds so they are the strongest

4.1.5 Bond Polarity

• When two atoms in a covalent bond have the same electronegativity the covalent bond is nonpolar (In other words, the difference in electronegativity = 0)

• When two atoms in a covalent bond have different electronegativitiesthe covalent bond is polar and the electrons will be drawn towards the more electronegativeatom

• As a result of this:

  • The negative charge center and positive charge center do not coincide with each other

  • This means that the electron distribution is asymmetric

  • The less electronegative atom gets a partial charge of δ+ (delta positive)

  • The more electronegative atom gets a partial charge of δ- (delta negative)

• The greater the difference in electronegativity the more polar the bond becomes

Dipole moment

• The dipole moment is a measure of how polar a bond is

• The direction of the dipole moment is shown by the following sign in which the arrow points to the partially negatively charged end of the dipole:

4.1.6 Lewis Structures

• Lewis structures are simplified electron shell diagrams and show pairs of electrons around atoms.

• A pair of electrons can be represented by dots
  The Octet Rule = The tendency of atoms to gain a valence shell with a total of 8 electrons

Steps for drawing Lewis Structures

1. Count the total number of valence

2. Draw the skeletal structure to show how many atoms are linked to each other.

3. Use a pair of crosses or dot/cross to put an electron pair in each bond between the atoms.

4. Add more electron pairs to complete the octets around the atoms ( except H which has 2 electrons)

5. If there are not enough electrons to complete the octets, form double/triple bonds.

6. Check the total number of electrons in the finished structure is equal to the total number of valence electrons

Incomplete Octets

• For elements below atomic number 20 the octet rule states that the atoms try to achieve 8 electrons in their valence shells, so they have the same electron configuration as a noble gas

• However, there are some elements that are exceptions to the octet rule, such a H, Li, Be, B and Al

  • H can achieve a stable arrangement by gaining an electron to become 1s², the same structure as the noble gas helium

  • Li does the same, but losing an electron and going from 1s²2s¹ to 1s² to become a Li⁺ ion

  • Be from group 2, has two valence electrons and forms stable compounds with just four electrons in the valence shell

  • B and Al in group 13 have 3 valence electrons and can form stable compounds with only 6 valence electrons

Unit 4.2 Resonance, Shapes, and Giant Structures

4.2.1 Resonance Structures

• Some atoms or elements have structures that don’t seem to fit with what you would expect their typical Lewis Structure to be

• This can be explained by the delocalization of electrons

• Delocalized electrons = electrons in a molecule, ion, or solid metal that are not permanently associated with one atom or covalent bond

Example:

Nitrate (V) Ion

• A molecule with 1 double bond and 2 single bonds

• It has 3 possible Lewis Structures where the double bond moves around and is with each of the three oxygens

• Since there are different possibilities, these Lewis Structures are also called Resonance Structures

• So you would expect each of these Resonance Structures to have explicitly 1 double bond and 2 single bonds, right?

• That’s not the case, however, because studies of the electron density and bond length show that all 3 bonds are equal in length

• In fact, the electron density is spread evenly between the three oxygen atoms

  • The actual bond length for all of them is somewhere between a single and a double bond

  • The actual structure is something between all the resonance structures and is called a resonance hybrid

Resonance Structures of the Nitrate (V) Ion

Steps to Determine a Lewis Structure:

1. Count the # of Valence Electrons

2. Consider the Charge

   • Add more electrons for negative charges

   • Subtract electrons for positive charges

N + 3O + 1

5 + (3 × 6) +1

= 24 electrons

3. Draw a Skeleton

   • Put single bonds between atoms first

   • Generally, the least electronegative atom goes in the center

   • On the ends, put Hydrogens (because Hydrogen can only have 2 valence electrons and can only bond once) and Halogens

4. Subtract Skeletal Electrons from Valence Electrons

   • Use the remaining electrons to create Lone Pairs

   • OR additional bonds (double/triple) as needed

   • Remember the overall goal is to satisfy the Octet Rule

3 structures are possible for Nitrate (V) Ion, with 1 double and 2 single bonds

The negative charge is distributed throughout the ion and is depicted with the negative sign outside of the resonance brackets

Electron pairs rapidly oscillate between different positions, never really staying a single or a double bond for long

• Criteria for forming resonance hybrid structures: molecules must have a double bond that is capable of migrating from one part of a molecule to another

• In other words, when there are adjacent atoms with equal electronegativity and lone pairs of electrons that can move to another position in order for the double bonds to be in other positions

  • Ex. Carbonate Ion, Benzene, Ozone, and the Carboxylate Anion

4.2.2 Shapes of Molecules

• VSEPR (Valence Shell Electron Pair Repulsion) Theory = A theory that predicts molecular shape and the angles between bonds based on the concepts:

  • All electron pairs and all lone pairs arrange themselves as far apart in space as possible

  • Lone Pairs repel more strongly than bonding pairs

  • Multiple bonds (double/triple) behave as single bonds

• Domains = The regions of negative cloud charge

• Steric Number (SN) = # of atoms + # of Lone Pairs around the central atom (same concept as Domains)

• Can also be denoted as:

  • A = Central Atom

  • B = Bonded Pair

  • E = Lone Pair

• ex. SN = 2 → AB₂

• ex. SN = 3 → AB₃

• ex. SN = 3 (but one of the domains is a lone pair) → AB₂E₁

Steric Number = 2

• If SN = 2, then the angle between bonds is 180°

• SN = 2 → AB₂

• Molecular Geometry = “Linear”

  • ex. BeCl₂, CO₂, HC≡CH

Steric Number = 3

• If SN = 3, then the angle between the bonds is 120°

• SN =3 → AB₃

• Molecular Shape = “Trigonal Planar”

  • ex. BF₃ and CH₂CH₂ and CH₂O

AB₂E₁

• If one of the electron domains is a lone pair, then the bond angle is slightly less than 120° since lone pairs repulse more, pushing against the other two bonding pairs closer together

  • ex. SO₂

• Molecular Geometry = “Bent Linear”

Steric Number = 4

• If SN = 4, then the angle between bonds is 109.5°

  • E.g. CH₄, NH₄⁺

• SN = 4 → AB₄

• Molecular Geometry = “Tetrahedral”

AB₃E₁

• If one of the electron domains is a lone pair, the bond angle is slightly less than 109.5° due to increased lone pair repulsion

  • ex. NH₃

• Molecular Geometry = “Trigonal Pyramidal”

AB₂E₂

• If 2 electron domains are lone pairs, bond angle also less than 109.5°

  • ex. H₂O

• Molecular Geometry = “Bent”

4.2.3 Predicting Shapes & Bond Angles

1. Draw Lewis Structure

   • Determine the number of bonding (B) and Lone Pairs (E) around the central atom (A)

2. Apply VSEPR Rules

   • Deduce shape and bond angle

4.2.4 Molecular Polarity

Bond Polarity ≠ Molecular Polarity

• Previously, you learned that bond polarity was determined by the difference in electronegative felt between two bonded atoms

• However, now you can determine if a molecule is polar or not

  • Consider:

    1. The polarity of each bond in the molecule

    2. How the bonds are arranged in the molecule

• Note: Some molecules have polar bonds, yet are overall not molecularly polar since the polar bonds in the molecule are arranged in a way that the individual bond dipole moments cancel each other out

  • ex. CH₃Cl vs. CCl₄

• CH₃Cl

  • Has 4 polar covalent bonds that don’t cancel each other out

  • This means the molecule is polar overall

  • The overall dipole moment is pointing towards the electronegative chlorine atom

• CCl₄

  • Also has 4 polar covalent bonds, BUT the individual bond dipole moments cancel each other out

  • So CCl₄ is a nonpolar molecule

4.2.5 Giant Covalent Structures

Giant Covalent Structures

Covalent Lattices

• Covalent Bonds = Bonds between nonmetals in which electrons are shared between atoms

• Giant Covalent Substances = Sometimes, a substance can't bond like a regular molecule. Instead, the bonds between atoms continue forever, forming a big lattice. There are no separate molecules in this situation, and all the nearby atoms are connected by covalent bonds.

  • ex. C

• Allotrope = Different atomic or molecular arrangements of the same element in the same physical state

• Graphite, diamond, buckminsterfullerene and graphene are allotropes of carbon

Giant Covalent Structures Examples

Diamond

• Diamond is a giant lattice of carbon atoms

• Each carbon is covalently bonded to 4 others in a tetrahedral geometry with a bond angle of 109.5°

• This results in a giant lattice with strong bonds in all directions and causes diamond to be the hardest known substance
  
  

Graphite

• Each carbon atom is bonded to 3 others in a layered structure of hexagons with a bond angle of 120°

• The spare electron is delocalized and moves around in the space between the layers

• All atoms in the same layer are held together by strong covalent bonds while the different layers are held together by weak intermolecular forces

Buckminsterfullerene

• Contains 60 carbon atoms

  • Each atom is bonded to 3 others by single covalent bonds

• The fourth electron is delocalized so the electrons can migrate throughout the structure

  • This allows for the structure to be a semi-conductor

• Has the same shape as a soccer ball, so it is nicknamed the football molecule

Graphene

• Some substances infinitely covalent bond only in two dimensons, forming only layers

  • ex. Graphene

• Graphene is a single layer of carbon atoms bonded in a repeating hexagonal pattern

  • It is so thin, 1 million times thinner than paper, that Graphene is actually considered 2D

Properties of Giant Covalent Structures

• As always, different structures and bonding types have different effects on the physical properties of substances (ie. melting/boiling points, electrical conductivity, and solubility)

Covalent Bonding & Giant Covalent Lattice Structures

• Giant Covalent Lattices:

  • Very High melting and boiling points

  • Large # of covalent bonds

  • A lot of energy is needed to break the lattice

  • Can be hard or soft

    • Hard (difficult to break their 3D network of strong covalent bonds)

      • Diamond

      • Silicon (IV) Oxide

    • Soft (forces between carbon layers are weak)

      • Graphite

    • (Graphene is strong, flexible, and transparent, making it a very useful material)

  • Insoluble in water (Most)

  • Do NOT conduct electricity (Most)

    • The some that do have delocalized electrons:

      • Graphite

      • Graphene

      • Buskminsterfullerene (semi-conductor)

Unit 4.3 Intermolecular Forces & Metallic Bonding

4.3.1 Types of Intermolecular Forces

Intermolecular Forces (IMF)

• Forces of attraction between molecules that hold them together

• 3 main types:

  • Dispersion forces

  • Dipole-dipole

  • Hydrogen bonding

Dispersion Forces

• Electrons in constant motion in an atom

• Electron cloud dispersion can be asymmetrical at any given point in time, creating an instantaneous dipole

• Attraction between partial negative and positive instantaneous dipoles form a dispersion force

• Strength of dispersion forces increases with difference in electronegativity and electron cloud movement

Key Facts about Dispersion Forces

• Present in all atoms and molecules

• Weakest IMF

• Strength depends on the number of electrons

• More electrons = stronger temporary/instantaneous dipole

• Exception: dispersion forces can be stronger than dipole-dipole if the atom is large enough

• Molecules composed of only C and H can only have dispersion forces

Dipole-dipole Attractions

• Attractive force between molecules with permanent dipoles

• Stronger than dispersion forces

• Only for small molecules with the same number of electrons

Hydrogen Bonding

• Strongest IMF

• Special type of dipole-dipole attraction

• Conditions for hydrogen bonding to occur:

  • A species with a very electronegative atom (O, N, or F) that has a lone pair of electrons

  • A hydrogen attached to the O, N, or F

• Hydrogen becomes partially positively charged and can form a bond with the lone pair on another molecule

Hydrogen Bond Donors and Acceptors

• Every hydrogen bond has two components

• A molecule can be both the donor and acceptor, able to hydrogen bond with itself

• Hydrogen bond acceptor only requires an available lone pair, not a hydrogen atom

Hydrogen Bonding in Water and Ammonia

• Water can form a maximum of two hydrogen bonds per molecule

• Ammonia can form a maximum of one hydrogen bond per molecule

• Number of hydrogen bonds possible is restricted by the number of

4.3.2 Deducing Intermolecular Forces

• The structure and chemical formular of the molecules will indicate the types of intermolecular forces present

  • Structure and Symmetry → Is molecule polar or not? (See Section 4.2.4)

  • Chemical Formula → How electronegative are the elements in the molecule?

    • Helps to tell you polar bonds

    • Also tells you if hydrogen bonds are possible when there is N, O, or F

4.3.3 Properties of Covalent Compounds

• Types of intermolecular forces indicate physical properties of molecular covalent compounds (melting/boiling point, solubility, and conductivity)

Melting and Boiling Point

• Changing the state means overcoming the intermolecular forces

• The stronger the forces, the more energy is needed to break the attraction between molecules

• Substances with low melting and boiling points = “volatile”

• As the intermolecular forces increase in strength:

  • The size of the molecule increases

  • The polarity of the molecule increases

Solubility

• “Like dissolves like” = non-polar substances dissolve in non-polar solvents while polar substances dissolve in polar solvents

• However, as the size of a covalent molecule increases in size at a certain point, their solubility can decrease

  • This is because the polar part of the molecule remains the smaller part of the overall structure (In other words, the ratio of polar to non-polar decreases)

  • Ex. alcohols (ethanol is soluble yet hexanol isn’t)

• Giant Covalent substances don’t dissolve in any solvents

  • This is because the energy required to overcome their strong covalent bonds from the lattice structure is too great

Conductivity

• Usually, covalent substances can’t conduct electricity in solid or liquid states since they don’t have any free-moving charged particles

• Only in some cases, polar covalent molecules can ionize and conduct electricity

• Other exceptions are Giant Covalent Structures and they can conduct electricity because they have delocalized electrons (the free-moving charged particles required for conductivity) (See Section 4.2.5)

4.3.4 Metallic Bonding

Metallic Bonding

• Metal atoms tend to pack together in lattice structures

  • This causes their outer electrons to be able to move freely throughout the entire structure = “delocalizes electrons”

• Once their valence electrons are delocalized, the metals gain a positive charge, which repel each other, keeping the entire structure neatly arranged in a lattice

• Metallic Bonding involves strong electrostatic forces of attraction between the metal centers and delocalized electrons

Properties of Metals

• Metals are malleable

  • This is because when a force is applied, the metal layers can slide over each other (the attractive forces between the metal ions and the delocalized electrons act in all directions)

  • So, when the layers slide, the metallic bonds can re-form in a new shape and the lattice is not broken

• Metallic compounds are strong and hard

• Due to the strong attraction between the metal cations suspended in a sea of delocalized electrons

• This also causes metals to have a high melting and boiling point

Conductivity

• Unlike non-metals, metals are able to conduct electricity when in the solid or liquid state

  • Because in both states, they have mobile electrons that can move around and conduct electricity (remember: electric current = flow of electrons)

Strength of Metallic Bonds

Not all metallic bonds have equal strength; there are several factors that affect it:

1. Charge on the Metal Ion

The greater the charge on the metal ion,

→ the greater number of electrons in the sea of delocalized electrons

→ the greater the charge difference between ions and electrons

→ the greater the electrostatic attraction

→ the stronger the metallic bond

2. Radius of the Metal Ion

Metal ions with a smaller ionic radii exert a greater attraction on the sea of delocalized electrons

→ requires more energy to break

→ stronger metallic bond

4.3.5 Trends in Melting Points of Metals

• An increase in the strength of electrostatic attraction is caused by:

  • Increasing the # of delocalized electrons in each metal atom

  • Increasing the positive charges on the metal centers in the lattice structure

  • Decreasing the size of the metal ions

Melting Points of Metals Across a Period

• Ex. Compare the electron configuration of sodium, magnesium, and aluminum and observe the # of valence electrons (do they increase or decrease?)

Na = 1s²2s²2p⁶3s¹

Mg = 1s²2s²2p⁶3s²

Al  = 1s²2s²2p⁶3s²3p¹

• Since aluminum ions are smaller in radius than magnesium or sodium ions

  • So considering that aluminum has the most electrons AND has the smallest radius, it has a stronger metallic bonding → higher melting point

• So as you go across a period, the metallic bonding is stronger and the melting points increase

Melting Points of Metals Down a Group

• As you go down a periodic group, the size of the metal cations increase, thus decreasing the attraction between the negative valence electrons and the positive metallic lattice, so the melting point decreases

4.3.6 Alloys & their Properties

• Alloys = mixtures of metals (the metals are mixed together physically but are not chemically combined)

  • Alloys can also be a mixture of metals and non-metals (ex. with carbon)

• The different metal ion mix is spread evenly throughout the lattice (not clumped together) and are bound together by their delocalized electrons

• Alloys are able to form due to the fact that metallic bonds are non-directional by nature

So why are Alloys made?

• They have distinct and desirable properties since the cations are structured differently in the lattice

Alloy Properties

• Greater strength,

• Harder

  • Since the mixture of atoms in an alloy are different sizes, this distorts the regular arrangement of cations

  • So the layers in a lattice structure have a more difficult time sliding over each other, causing the alloy to be harder than a pure metal

• Higher resistance to corrosion/extreme temperatures

Unit 5: Energetic/Thermochemistry

Measuring Energy Changes

Energetics is the study of heat changes in chemical reactions. Heat is a form of energy, and in chemistry, we measure how much heat is absorbed or released during reactions. For example, when you burn wood, heat is released into the surroundings.

Temperature and Energy

• Temperature: Measures how hot or cold something is. It’s related to the average kinetic energy (energy of motion) of particles in a substance. For example, boiling water has higher kinetic energy than ice.

• Heat: Measures the total energy content of a substance. A beaker of boiling water has more heat energy than a few drops of boiling water, even though both are at the same temperature.

• First Law of Thermodynamics: Energy cannot be created or destroyed, only transformed. For instance, when you burn fuel, chemical energy is transformed into heat and light.

• Second Law of Thermodynamics: For a reaction to happen on its own (spontaneously), there must be an overall increase in entropy (disorder). For example, ice melting into water increases entropy because water molecules are more disordered.

• Energy: The ability to do work. Energy cannot be created or destroyed, only transformed (First Law of Thermodynamics).

Enthalpy (ΔH)

Enthalpy is the heat energy stored in a substance, mainly in its chemical bonds. We can’t measure the total enthalpy of a substance, but we can measure the change in enthalpy (ΔH) during a reaction.

• Units: kJ/mol.

• System vs. Surroundings:

  • System: The reaction itself (e.g., a test tube with reactants).

  • Surroundings: Everything outside the system (e.g., the air around the test tube).

Standard Conditions

• Pressure: 100 kPa.

• Concentration: 1 mol/dm³ for solutions.

• Temperature: Usually 298 K (25°C).

• Standard State: Each substance is in its most stable form (solid, liquid, or gas).

Types of Standard Enthalpy Changes

• Standard Enthalpy Change of Reaction (ΔHᵣ): Enthalpy change when reactants react to form products under standard conditions.

• Standard Enthalpy Change of Formation (ΔHᶠ): Enthalpy change when 1 mole of a compound is formed from its elements under standard conditions.

• Standard Enthalpy Change of Combustion (ΔH_c): Enthalpy change when 1 mole of a substance is burned in excess oxygen under standard conditions.

• Standard Enthalpy Change of Neutralization (ΔH_neut): Enthalpy change when 1 mole of water is formed by reacting an acid and alkali under standard conditions.

Exothermic vs. Endothermic Reactions

• Exothermic Reactions: Release heat to the surroundings. ΔH is negative. Example: Combustion of propane:

  • C3H8+5O2→3CO2+4H2O(ΔH=−2220 kJ/mol)

  • The products are more stable than the reactants.

• Endothermic Reactions: Absorb heat from the surroundings. ΔH is positive. Example: Photosynthesis:

  • 6CO2+6H2O→C6H12O6+6O2(ΔH=+2800 kJ/mol)

  • The products are less stable than the reactants.

Measuring Enthalpy Changes

To measure enthalpy changes experimentally, we use the formula:

Q=mcΔT

• Q: Heat energy (in kJ).

• m: Mass of water (in kg).

• c: Specific heat capacity of water (4.18 kJ/kg·K).

• ΔT: Change in temperature (in K).

Then, the enthalpy change (ΔH) is calculated as:

ΔH= Q/n

​

• n: Number of moles of the limiting reactant.

Calorimetry

Measuring Enthalpy Changes

• Calorimeter: A device used to measure heat changes in reactions. Can be made from a polystyrene cup, vacuum flask, or metal can.

• Specific Heat Capacity (c): The energy needed to raise the temperature of 1 g of a substance by 1 K. For water, c = 4.18 J g⁻¹ K⁻¹.

• Equation: q=m×c×ΔT

Worked Example

• Problem: 0.01 mol of propan-1-ol was burned, heating 250 g of water from 298 K to 310 K. Calculate the enthalpy of combustion.

• Solution:

  • q=250×4.18×12=12,540 J

  • ΔH=q/n=12,54/00.01=1,254,000 J/mol=−1,254 kJ/mol

Calorimetry Experiments

Enthalpy Changes in Solution

• Principle: Reactants are mixed in solution, and the temperature change is measured.

• Assumptions:

  • Specific heat capacity of the solution = 4.18 J g⁻¹ K⁻¹.

  • Density of the solution = 1 g/cm³.

  • Heat losses are negligible.

Temperature Correction Graphs

• Purpose: To account for heat loss during slow reactions.

• Method:

  1. Record temperature before adding reactants.

  2. Add reactants and continue recording temperature.

  3. Extrapolate the cooling section of the graph to find the maximum temperature change.

Enthalpy of Combustion Experiments

• Principle: Heat from combustion is used to heat water.

• Sources of Error:

  • Heat loss to surroundings.

  • Incomplete combustion.

Worked Example

• Problem: 1.023 g of propan-1-ol was burned, heating 200 g of water by 30°C. Calculate the enthalpy of combustion.

• Solution:

  • q=200×4.18×30 = 25,080 J

• ΔH=q/n=25,080/0.01702=1,473,560 J/mol=−1,474 kJ/mol

Hess’s Law

Hess’s Law states that the total enthalpy change for a reaction is the same, no matter the route taken. For example, if you want to find the enthalpy change for:

A→C

You can use intermediate steps:

A→B(ΔH1)B→C(ΔH2)

Then:

ΔHtotal=ΔH1+ΔH2

​

Bond Enthalpies

• Bond Breaking: Requires energy (endothermic). Example: Breaking H-H bonds in hydrogen gas.

• Bond Making: Releases energy (exothermic). Example: Forming H-O bonds in water.

The enthalpy change for a reaction can be calculated using bond enthalpies:

ΔH=Bond Enthalpies of Reactants−Bond Enthalpies of Products

Energy Cycles

Born-Haber Cycle

The Born-Haber Cycle is used to calculate the lattice enthalpy of ionic compounds. It breaks down the formation of an ionic compound into steps:

1. Atomization: Convert solid metal to gaseous atoms.

2. Ionization: Remove electrons from metal atoms to form cations.

3. Electron Affinity: Add electrons to non-metal atoms to form anions.

4. Lattice Formation: Combine gaseous ions to form the solid lattice.

For example, to calculate the lattice enthalpy of LiF:

ΔHlat=ΔHatom+ΔHIE+ΔHEA−ΔHf

​

Entropy and Spontaneity

Entropy (S)

Entropy measures the disorder of a system. The more ways energy can be distributed, the higher the entropy. For example:

• Solids have low entropy (ordered).

• Gases have high entropy (disordered).

The change in entropy (ΔS) is calculated as:

ΔS=Sproducts−Sreactants

​

Gibbs Free Energy (ΔG)

Gibbs free energy combines enthalpy and entropy to predict whether a reaction is spontaneous:

ΔG=ΔH−TΔS

ΔG=ΔH−TΔS

• If ΔG is negative, the reaction is spontaneous.

• If ΔG is positive, the reaction is not spontaneous.

Example: At high temperatures, even endothermic reactions (ΔH > 0) can be spontaneous if ΔS is large

Unit 6:Chemical Kinetics

What is Chemical Kinetics?

• Chemical kinetics, also called reaction kinetics, helps us understand the rates of reactions and how it is influenced by certain conditions.

  • It further helps to gather and analyse information about the mechanism of the reaction and define the characteristics of a chemical reaction.

Rate of Formations and Disappearances:

• In any chemical reaction, as the reaction proceeds, the amount of reactants decreases, whereas the amount of products increases. 

  • One has to understand that the rate of the overall reaction depends on the rate at which reactants are consumed or the rate at which the products are formed.

• If a graph is plotted between the concentration of reactants and products and time, the rate of formation of products and the rate of disappearance of reactants can be easily calculated from the slope of curves for products and reactants. 

  • The overall rate of the reaction may or may not be equal to the rate of formations and disappearances.

    • Product concentration is zero at time t = 0

    • at time t = 0, both reactants and products are present.

  • From the graph, it is understood that the slope of the reactants curve is negative and that for product curve is positive, indicating the concentration of reactants and products decreases and increases, respectively. 

    • We will take a simple reaction as an example to illustrate how the rate of overall reactions, rate of disappearances of reactants and rate of formation of products are related.

    • Let us take the reaction of the formation of water.

      • 2H2 + O2 → 2H2O

    • From the balanced equation, it is under that for one mole of O2 consumed, 2 moles of H2 will be consumed, and 2 moles of H2O will be formed. Say, the reaction proceeds for 10 mins, taking 1 mole of H2 and O2 each in the reaction mixture.

      • 2H2 + O2 → 2H2O

        • t = 0 1 1 0

        • t = 10 mins 1 – 0.5 1 – 0.25 0.5

      • Say after 10 minutes, 0.5 moles of H2 is consumed, and according to stoichiometry, 0.25 moles of O2 is consumed, and 0.5 moles of H2O is formed. Now, let us calculate the rates for H2, O2 and H2O for the first 10 minutes.

    • Rate of disappearance of  H2:
      

    • Rate of disappearance of O2:
      

    • Rate of formation of H2O:
      

  • From the above calculations, we can see that rate at which H2 is consumed is twice the rate at which O2 is consumed. 

    • So, the stoichiometry of the reaction relates rates of formation and disappearances of different reactants and products as follows:
      

Average and Instantaneous Rate:

• The rate of reaction can be classified into average and instantaneous rates depending on the amount of time period. If the time period taken is finite, then it’s called the average rate and is represented as,

• 𝑟_avg=Δ[𝐶]Δ𝑡

  • Δl → change in concentration

  • Δt → change in time

  • r_avg → average rate

• The average rate doesn’t give exact information in most cases about the completion of the reaction.

  • For example, let us consider the hydrolysis of esters to form acid and alcohol.

    • 𝑅𝐶𝑂𝑂𝑅′→𝐻2𝑂𝑅𝐶𝑂𝑂𝐻+𝑅′𝑂𝐻

  • Say, at time t = 0, and there was 1 M solution of ester which becomes 0.5 M in 30 mins.

    • So, it is logical for us to assume that the reaction will be 100% completed in 1 hour.

    • But in reality, the reaction takes more than 3 hours to reach completion.

    • So, to get a broader insight into the time taken for completion and other purposes, “Instantaneous rate is used, which is represented as,

      • 𝑟_inst=limΔ𝑡→0Δ[𝐶]Δ𝑡=𝑑[𝐶]𝑑𝑡

      • From above, it’s understood that the time period taken is almost zero, from t = 0 to t = 0.0000 …1 second.

        • This eventually comes out to be the differential of change in concentration with respect to time.

• For all practical purposes, the instantaneous rate is used, which can be calculated from the concentration in the time graph by finding a tangent at a point.
  
  

  • The unit of rate is Moll-1 s-1 because it is concentration/time, and concentration is expressed in terms of Molarity/mol l-1). 

    • It can also be Nm-2/s if the active mass is used in terms of partial pressures. 

    • Depending on other units of time, it can also be mol l-1 min-1 or mol l-1 hour-1 etc.

Factors Affecting the Reaction Rate:

• The rate of a reaction can be altered if any of the following parameters are changed.

Concentration of Reactants:

• According to collision theory, reactant molecules collide with each other to form products.

  • If the concentration of reactants is increased, the number of colliding particles will increase, thereby increasing the rate of reaction.

Nature of the Reactants:

• The reaction rate also depends on the types of substances that are reacting. If we consider acid/base reactions, salt formation and ion exchange, they are mostly fast reactions. 

  • During the formation of a covalent bond between the molecules that results in the formation of larger molecules, the reaction that takes place is usually slower. 

  • Furthermore, the nature and strength of bonds in reactant molecules significantly affect the rate of their transformation into products.

Physical State of Reactants:

• The physical state of a reactant, whether it is solid, liquid or gas, can greatly affect the rate of change. 

  • To discuss it further, if reactants are in the same phase, let’s say they are in an aqueous solution, here the thermal motion will bring them together. 

  • If they are in different phases, then the reaction will be limited to the interface between the reactants.  

    • The reaction mainly occurs only at their area of contact, in the case of a liquid and a gas, at the surface of the liquid.

Surface Area of Reactants:

• If we take two solids, the particles that are at the surface will take part in the reaction. 

  • Likewise, if we want to crush a solid into smaller parts, more particles will be present at the surface. 

    • What it means is that the frequency of collisions between these and reactant particles will most likely increase. 

  • As a result, the reaction will occur more rapidly.

• When two or more reactants are in the same phase of fluid, their particles collide more often than when either or both are in the solid phase or when they are in a heterogeneous mixture. In a heterogeneous medium, the collision between the particles occurs at an interface between phases. 

  • Compared to the homogeneous case, the number of collisions between reactants per unit time is significantly reduced, and so is the reaction rate.

Temperature:

• If the temperature is increased, the number of collisions between reactant molecules per second (frequency of collision) increases, thereby increasing the rate of the reaction. 

  • But depending on whether the reaction is endothermic or exothermic, an increase in temperature increases the rate of forward or backward reactions, respectively.

• In a system where more than one reaction is possible, the same reactants can produce different products under different temperature conditions.

  • At 100 ⁰C in the presence of dilute sulphuric acid, diethyl ether is formed from ethanol.

    • 2CH₃CH₂OH → CH₃CH₂OCH₂CH₃+H₂O

  • At 180 ⁰C in the presence of dilute sulphuric acid, ethylene is the major product.

    • CH₃CH₂OH → C₂H₄+H₂O

Effect of Solvent:

• The nature of the solvent also depends on the reaction rate of the solute particles.

• Example:

  • When sodium acetates react with methyl iodide, it gives methyl acetate and sodium iodide.

    • CH₃CO₂Na_(sol)+CH₃I_(liq)→CH₃CO₂CH_3(sol)+NaI_(sol)

    • The above reaction occurs faster in organic solvents such as DMF (dimethylformamide) than in CH₃OH (methanol) because methanol is able to form a hydrogen bond with CH₃CO₂ ^– but DMF is not possible.

Catalyst:

• Catalysts alter the rate of the reaction by changing the reaction mechanism.

  • There are two types of catalysts, namely, promoters and poisons, which increase and decrease the rate of reactions, respectively.

Positive Catalysts:

• Catalysts that increase the rate of a chemical reaction are positive catalysts. 

  • It increases the rate of reaction by lowering the activation energy barriers such that a large number of reaction molecules are converted into products, and thereby the percentage of yield of products increases.

• Positive catalyst example: 

  • In the preparation of NH3 by Haber’s process, iron oxide acts as a positive catalyst and increases the yield of ammonia in spite of less reaction of nitrogen.

Negative Catalysts:

• Catalysts that decrease the rate of reaction are negative catalysts. 

  • It decreases the rate of reaction by increasing the activation energy barrier, which decreases the number of reactant molecules to transform into products, and hence the rate of reaction decreases.

• Negative catalyst example: 

  • The decomposition of hydrogen peroxide into water and oxygen is retarded by using acetanilide, and this acts as a negative catalyst to decrease the rate of decomposition of hydrogen peroxide.

Unit 7: Equilibrium

Equilibrium Reactions in Chemistry

Many chemical reactions are reversible, meaning they can move in both directions: converting reactants into products (forward reaction) and converting products back into reactants (reverse reaction).

Reversible Reaction and Equilibrium:

• In a reversible reaction, a state of dynamic equilibrium can be reached where:

  • The rates of the forward and reverse reactions are equal.

  • Reactants and products are continuously interconverted, but their overall concentrations remain unchanged over time.

• This state of balance is reached in a closed system where no additional reactants or products are introduced or removed.

Example: Water in a Sealed Container:

• When water in a sealed container is heated, some molecules in the liquid phase gain enough energy to escape into the gas phase (evaporation).

• Since the container is sealed, some gas-phase molecules lose energy through collisions and return to the liquid phase (condensation).

Process to Equilibrium:

• Initially: More liquid molecules evaporate than condense, as there is a high amount of liquid compared to vapor.

• As Vapor Builds Up: The rate of condensation increases as more gas molecules are available to return to the liquid state.

• At Equilibrium: The rates of evaporation and condensation are equal, resulting in no net change in the amounts of liquid and vapor. Molecules continue to transition between phases, but the total amounts of each phase remain constant.

Dynamic Equilibrium:

• At equilibrium:

  • The forward and reverse reactions occur continuously at equal rates.

  • The concentrations of reactants and products stay constant, even though individual molecules keep interconverting.

  • Observable, or macroscopic, properties (such as color or density) remain unchanged, making the system appear static.

Indications of Equilibrium:

• Single Arrow (→): Used for non-reversible reactions that proceed only in one direction.

• Double Arrow (⇌): Used for reversible reactions where a dynamic equilibrium can be established.

Conditions for Chemical Equilibrium:

• Closed System: To maintain equilibrium, reactants and products must be contained, allowing forward and reverse reactions to continue without interference. Examples include:

  • A gaseous reaction in a sealed container

  • A liquid, solid, or aqueous reaction in a closed beaker

  • Phase equilibrium (such as liquid-gas) in a sealed container

• Constant Observable Properties: When a system is undisturbed at equilibrium, macroscopic properties (e.g., color, pressure, concentration) appear constant because opposing processes occur at equal rates.

• Equal Forward and Reverse Rates: A reversible reaction at equilibrium has forward and reverse reaction rates that are balanced, ensuring that the concentration of reactants and products remains stable.

Temperature Control (Isothermal Systems):

• Equilibrium systems are often maintained at a constant temperature, or isothermal conditions, by adding or removing heat as necessary to prevent changes in reaction rates due to temperature fluctuations.

Chemical Systems

Many reactions in chemistry reach a state called equilibrium, where both the forward and reverse reactions continue to occur.

• The rates of these reactions depend on several factors, including:

  • Physical conditions: Temperature and pressure

  • Chemical conditions: Concentrations of reactants and products

  • Presence of a catalyst: Catalysts can speed up both the forward and reverse reactions but do not change the position of equilibrium.

Example: NO₂ and N₂O₄ Equilibrium

• Nitrogen dioxide (NO₂) and dinitrogen tetroxide (N₂O₄) are gases that exist in a reversible equilibrium:

  • N₂O₄ (g)⇋2NO₂ (g)

• N₂O₄ is a colorless gas, while NO₂ is a brown gas.

How Equilibrium Is Reached

• Initially, the concentration of N₂O₄ is high, so its decomposition into NO₂ is rapid.

• As NO₂ forms, some molecules collide and recombine into N₂O₄.

• Over time, as N₂O₄ decreases and NO₂ increases, the rate of decomposition of N₂O₄ slows down, while the rate of NO₂ recombining into N₂O₄ speeds up.

• Eventually, the forward reaction (N₂O₄ decomposing to NO₂) and the reverse reaction (NO₂ recombining into N₂O₄) occur at equal rates.

Reaching Dynamic Equilibrium

• At dynamic equilibrium:

  • The rates of the forward and reverse reactions are equal.

  • There is no further change in the concentrations of N₂O₄ and NO₂.

  • Macroscopic properties, such as color and density, remain constant.

Characteristics of Dynamic Equilibrium

• Both reactions continue to occur, but the overall concentrations of reactants and products do not change.

• The equilibrium can be approached from either direction:

  • Starting with all N₂O₄, or all NO₂, the system will reach the same equilibrium concentrations.

• If any conditions (temperature, pressure, concentration) change, the equilibrium position will shift.

Equilibrium Law

The law of chemical equilibrium states that at a given temperature, the ratio of the concentrations of products to reactants is constant when the system has reached equilibrium. This ratio is known as the equilibrium constant (denoted as Kₓ), and it varies depending on the temperature of the system.

General Form of the Equilibrium Constant Expression

For a reversible chemical reaction, such as:

The equilibrium constant expression, K_C​, is written as:

Where:

• [A],[B],[C],[D] represent the concentrations of the reactants (A, B) and products (C, D), respectively.

• The concentrations are raised to the power of their corresponding stoichiometric coefficients (a, b, c, and d) from the balanced equation.

Key Points About the Equilibrium Constant Expression

• Numerator: The concentrations of products raised to the power of their coefficients.

• Denominator: The concentrations of reactants raised to the power of their coefficients.

• If there are multiple reactants or products, the terms are multiplied together.

Example: H₂(g) + I₂(g) ⇋ 2HI(g)

The equilibrium expression for this reaction is:

Inverse Reaction

For the reverse reaction, the equilibrium constant expression is inverted:

• Original reaction: aA + bB ⇋ cC + dD

• Reverse reaction: cC + dD ⇋ aA + bB

Types of Equilibria

• Homogeneous Equilibrium: All reactants and products are in the same phase (usually gas phase). For example, reactions occurring entirely in the gas phase.

• Heterogeneous Equilibrium: Reactants and products exist in more than one phase (solid, liquid, gas). For example, a reaction between a solid and a gas.

Important Considerations

• Pure Solids and Liquids: In equilibrium expressions, pure solids and pure liquids are not included because their concentrations do not change and are considered constant.

• Effect of Temperature: The equilibrium constant K_c​ changes with temperature. A higher K_c (greater than 1) suggests that products are favored at equilibrium, whereas a smaller K_c​ (less than 1) indicates that reactants are favored.

Vapor Pressure and Phase Equilibrium

• Vapor Pressure is the pressure exerted by the gas phase of a substance when it is in equilibrium with its liquid phase.

  • In phase equilibrium (e.g., liquid ↔ gas), the system reaches a point where the rate of evaporation equals the rate of condensation.

  • The vapor pressure is a characteristic of the liquid and depends on temperature.

  • The size and shape of the container do not affect the equilibrium vapor pressure.

Effects of Conditions on Equilibrium Constant

When a system is at equilibrium, the position of equilibrium remains stable as long as temperature and pressure stay constant. However, certain experimental changes can affect the equilibrium position or shift the balance of reactants and products. While the equilibrium position may shift in response to changes, the equilibrium constant Kc​ itself only changes with temperature.

1. Effect of Concentration Changes

• Adding or removing reactants or products will shift the equilibrium position to counter the change. This is in line with Le Chatelier’s Principle, which states that a system at equilibrium will adjust to counteract an external change.

  • Adding reactants shifts equilibrium towards the products (forward reaction) to consume the added substance.

  • Removing reactants shifts equilibrium towards the reactants (reverse reaction).

• Equilibrium constant Kc​ remains unchanged because it is only influenced by temperature.

2. Effect of Pressure Changes (for Gaseous Reactions)

• In reactions involving gaseous reactants and products, changes in pressure can impact the equilibrium position.

• Increasing pressure will shift the equilibrium toward the side with fewer gas molecules, which helps reduce overall pressure in the system.

  • For example, in a reaction where 4 moles of gaseous reactants produce 3 moles of gaseous products, an increase in pressure will favor the forward reaction to reduce the number of moles.

• Decreasing pressure will favor the side with more gas molecules.

• Kc remains constant because pressure does not directly alter the equilibrium constant; it only shifts the position.

3. Effect of Temperature Changes

Temperature is unique in that it directly affects both the equilibrium position and the value of Kc​:

• Exothermic Reactions:

  • Heat is a product in exothermic reactions. Adding heat (increasing temperature) causes the equilibrium to shift left, favoring the reactants, to reduce the added heat.

  • Decreasing temperature shifts equilibrium right, favoring the products.

  • Kc changes based on the temperature change:

    • For exothermic reactions, increasing temperature decreases Kc​ as the system favors reactants.

    • For endothermic reactions, increasing temperature increases Kc as the system favors products.

• Endothermic Reactions:

  • Heat is a reactant in endothermic reactions. Adding heat shifts equilibrium to the right to absorb the added energy.

  • Kc increases with temperature for endothermic reactions because the products are favored.

4. Effect of a Catalyst

• A catalyst lowers the activation energy of both the forward and reverse reactions, speeding up the time it takes to reach equilibrium but not changing the equilibrium position or Kc​.

• It does not change the equilibrium constant or shift the balance of products and reactants.

Summary of Temperature Effects on Equilibrium Direction

Note: In any reaction where the enthalpy change ΔH=0, temperature changes have no effect on Kc since no heat is absorbed or released.

Le Chatelier’s Principle

Le Châtelier’s Principle explains how a system at equilibrium responds to changes in conditions (concentration, temperature, or pressure) to maintain or restore balance. When an external change disturbs a system in equilibrium, the system will adjust itself to offset the change and re-establish equilibrium. This principle helps predict the direction in which equilibrium shifts when the system is altered.

Key Influencing Factors on Equilibrium

1. Concentration Changes

   • When the concentration of a reactant increases or a product decreases, the system is no longer at equilibrium. To offset this, the forward reaction is favored to convert some of the added reactants into products, moving the system back toward equilibrium.

   • When the concentration of a product increases or a reactant decreases, the reverse reaction is favored, shifting equilibrium toward the reactants to counteract the imbalance.

   • Kc (equilibrium constant) remains unchanged by concentration changes, as only temperature affects its value.

2. Temperature Changes

   • The direction of the equilibrium shift due to temperature depends on whether the reaction is endothermic (absorbs heat) or exothermic (releases heat).

3. Endothermic Reactions (ΔH>0\Delta H > 0ΔH>0):

   • Increasing temperature adds heat, which acts like a reactant in endothermic reactions, so the forward reaction is favored, and equilibrium shifts toward the products.

   • Decreasing temperature removes heat, favoring the reverse reaction, shifting equilibrium toward the reactants.

4. Exothermic Reactions (ΔH<0):

   • Increasing temperature adds heat, acting as an excess product, so the system responds by favoring the reverse reaction and shifting equilibrium toward the reactants.

   • Decreasing temperature removes heat, favoring the forward reaction and shifting equilibrium toward the products.

   • Effect on Kc​: Unlike other factors, temperature changes the equilibrium constant (Kc) because heat is either absorbed or released, impacting the system’s balance.

5. Pressure Changes (for Reactions with Gaseous Species)

   • For reactions involving gases, pressure changes affect equilibrium if there is a difference in the number of gas molecules (moles) between the reactants and products.

   • If products have fewer gas molecules:

     • Increasing pressure favors the side with fewer gas molecules to reduce pressure, shifting equilibrium toward the products.

     • Decreasing pressure favors the side with more gas molecules, shifting equilibrium toward the reactants.

   • If reactants have fewer gas molecules:

     • Increasing pressure shifts equilibrium toward the reactants, favoring the reverse reaction.

     • Decreasing pressure shifts equilibrium toward the products, favoring the forward reaction.

   • Effect on Kc​: Pressure changes do not alter Kc​, as it is only dependent on temperature.

6. Effect of a Catalyst

   • A catalyst speeds up both the forward and reverse reactions equally by lowering their activation energy, allowing the system to reach equilibrium more quickly.

   • Equilibrium Position: The catalyst does not shift the equilibrium position since it affects both reactions equally.

   • Equilibrium Constant Kc​: Adding a catalyst does not change Kc​ because it does not affect the relative concentrations of reactants and products at equilibrium.

Le Châtelier’s Principle helps predict how a system in equilibrium will respond to various changes, guiding the direction of reaction shifts to maintain or re-establish equilibrium.

Reaction Quotient

The concept of the reaction quotient (denoted by Q) and how it compares to the equilibrium constant (Kc) is essential for understanding the progress of a chemical reaction as it moves toward equilibrium. Let me explain it step by step to ensure clarity.

Reaction Quotient (Q)

The reaction quotient Q is a mathematical expression similar to the equilibrium constant Kc​, but the key difference is that Q is calculated using the concentrations of reactants and products at any point in time, whether or not the system is at equilibrium.

where [C],[D],[A], and [B] are the molar concentrations of the chemical species C, D, A, and B at a given moment in time. This expression is exactly the same as the equilibrium expression for Kc​, but instead of equilibrium concentrations, Q uses the concentrations at any moment.

Comparing Q and Kc​

1. When Q=Kc​:

   • The reaction is at equilibrium.

   • At equilibrium, the forward and reverse reactions occur at the same rate, so the concentrations of reactants and products do not change. The system is balanced, and no net change in the concentrations will occur unless the conditions (e.g., temperature or pressure) change.

2. When Q>Kc​:

   • This means that, at this point in time, the concentration of products is higher than it would be at equilibrium.

   • Since the reaction tends to favor the direction that reduces the concentration of products and increases the concentration of reactants to reach equilibrium, the reverse reaction will be favored.

   • The system will shift toward the left (toward the reactants) to decrease the concentration of products and increase the concentration of reactants, moving toward equilibrium.

3. When Q<Kc​:

   • This indicates that the concentration of reactants is higher than it would be at equilibrium.

   • To reach equilibrium, the system will favor the forward reaction (the production of products).

   • The reaction will shift to the right (toward the products) to decrease the concentration of reactants and increase the concentration of products, moving toward equilibrium.

How Q Helps Predict the Reaction's Direction

The value of Q gives you a snapshot of the system's current state relative to equilibrium. By comparing Q to Kc​, you can predict:

• If Q<Kc​, the reaction will move forward (toward the products) to reach equilibrium. The system needs to produce more products and consume more reactants.

• If Q>Kc​, the reaction will move backward (toward the reactants) to reach equilibrium. The system needs to decrease the amount of products and generate more reactants.

• If Q=Kc​, the system is already at equilibrium, and no net change will occur in the concentrations of reactants and products.

Significance of the Reaction Quotient

• Helps predict the shift: By knowing the value of Q compared to Kc​, you can determine whether the reaction needs to proceed forward (to make more products) or backward (to make more reactants) to achieve equilibrium.

• Indicates the direction of change: The reaction will always move in the direction that brings the system closer to equilibrium. If there are too many products (when Q>Kc), the system will shift toward reactants. If there are too many reactants (when Q<Kc​), the system will shift toward products.

Example

Let's take an example to illustrate how Q works.

Consider the following reaction:

The equilibrium expression for this reaction is:

Now, suppose at a given time the concentrations are as follows:

The reaction quotient at this point would be:

If the equilibrium constant Kc​ = 1.0, we see that:

Since Q<Kc​, the concentration of reactants (N₂ and H₂) is greater than at equilibrium, and the system will shift forward to produce more NH₃​, moving toward equilibrium.

Unit 8: Acids and Bases

Equilibrium Reactions in Chemistry

Many chemical reactions are reversible, meaning they can move in both directions: converting reactants into products (forward reaction) and converting products back into reactants (reverse reaction).

Reversible Reaction and Equilibrium:

• In a reversible reaction, a state of dynamic equilibrium can be reached where:

  • The rates of the forward and reverse reactions are equal.

  • Reactants and products are continuously interconverted, but their overall concentrations remain unchanged over time.

• This state of balance is reached in a closed system where no additional reactants or products are introduced or removed.

Example: Water in a Sealed Container:

• When water in a sealed container is heated, some molecules in the liquid phase gain enough energy to escape into the gas phase (evaporation).

• Since the container is sealed, some gas-phase molecules lose energy through collisions and return to the liquid phase (condensation).

Process to Equilibrium:

• Initially: More liquid molecules evaporate than condense, as there is a high amount of liquid compared to vapor.

• As Vapor Builds Up: The rate of condensation increases as more gas molecules are available to return to the liquid state.

• At Equilibrium: The rates of evaporation and condensation are equal, resulting in no net change in the amounts of liquid and vapor. Molecules continue to transition between phases, but the total amounts of each phase remain constant.

Dynamic Equilibrium:

• At equilibrium:

  • The forward and reverse reactions occur continuously at equal rates.

  • The concentrations of reactants and products stay constant, even though individual molecules keep interconverting.

  • Observable, or macroscopic, properties (such as color or density) remain unchanged, making the system appear static.

Indications of Equilibrium:

• Single Arrow (→): Used for non-reversible reactions that proceed only in one direction.

• Double Arrow (⇌): Used for reversible reactions where a dynamic equilibrium can be established.

Conditions for Chemical Equilibrium:

• Closed System: To maintain equilibrium, reactants and products must be contained, allowing forward and reverse reactions to continue without interference. Examples include:

  • A gaseous reaction in a sealed container

  • A liquid, solid, or aqueous reaction in a closed beaker

  • Phase equilibrium (such as liquid-gas) in a sealed container

• Constant Observable Properties: When a system is undisturbed at equilibrium, macroscopic properties (e.g., color, pressure, concentration) appear constant because opposing processes occur at equal rates.

• Equal Forward and Reverse Rates: A reversible reaction at equilibrium has forward and reverse reaction rates that are balanced, ensuring that the concentration of reactants and products remains stable.

Temperature Control (Isothermal Systems):

• Equilibrium systems are often maintained at a constant temperature, or isothermal conditions, by adding or removing heat as necessary to prevent changes in reaction rates due to temperature fluctuations.

Chemical Systems

Many reactions in chemistry reach a state called equilibrium, where both the forward and reverse reactions continue to occur.

• The rates of these reactions depend on several factors, including:

  • Physical conditions: Temperature and pressure

  • Chemical conditions: Concentrations of reactants and products

  • Presence of a catalyst: Catalysts can speed up both the forward and reverse reactions but do not change the position of equilibrium.

Example: NO₂ and N₂O₄ Equilibrium

• Nitrogen dioxide (NO₂) and dinitrogen tetroxide (N₂O₄) are gases that exist in a reversible equilibrium:

  • N₂O₄ (g)⇋2NO₂ (g)

• N₂O₄ is a colorless gas, while NO₂ is a brown gas.

How Equilibrium Is Reached

• Initially, the concentration of N₂O₄ is high, so its decomposition into NO₂ is rapid.

• As NO₂ forms, some molecules collide and recombine into N₂O₄.

• Over time, as N₂O₄ decreases and NO₂ increases, the rate of decomposition of N₂O₄ slows down, while the rate of NO₂ recombining into N₂O₄ speeds up.

• Eventually, the forward reaction (N₂O₄ decomposing to NO₂) and the reverse reaction (NO₂ recombining into N₂O₄) occur at equal rates.

Reaching Dynamic Equilibrium

• At dynamic equilibrium:

  • The rates of the forward and reverse reactions are equal.

  • There is no further change in the concentrations of N₂O₄ and NO₂.

  • Macroscopic properties, such as color and density, remain constant.

Characteristics of Dynamic Equilibrium

• Both reactions continue to occur, but the overall concentrations of reactants and products do not change.

• The equilibrium can be approached from either direction:

  • Starting with all N₂O₄, or all NO₂, the system will reach the same equilibrium concentrations.

• If any conditions (temperature, pressure, concentration) change, the equilibrium position will shift.

Equilibrium Law

The law of chemical equilibrium states that at a given temperature, the ratio of the concentrations of products to reactants is constant when the system has reached equilibrium. This ratio is known as the equilibrium constant (denoted as Kₓ), and it varies depending on the temperature of the system.

General Form of the Equilibrium Constant Expression

For a reversible chemical reaction, such as:

The equilibrium constant expression, K_C​, is written as:

Where:

• [A],[B],[C],[D] represent the concentrations of the reactants (A, B) and products (C, D), respectively.

• The concentrations are raised to the power of their corresponding stoichiometric coefficients (a, b, c, and d) from the balanced equation.

Key Points About the Equilibrium Constant Expression

• Numerator: The concentrations of products raised to the power of their coefficients.

• Denominator: The concentrations of reactants raised to the power of their coefficients.

• If there are multiple reactants or products, the terms are multiplied together.

Example: H₂(g) + I₂(g) ⇋ 2HI(g)

The equilibrium expression for this reaction is:

Inverse Reaction

For the reverse reaction, the equilibrium constant expression is inverted:

• Original reaction: aA + bB ⇋ cC + dD

• Reverse reaction: cC + dD ⇋ aA + bB

Reactions with Changed Stoichiometry

1. Doubling the Coefficients:

   • If the coefficients of the reaction are doubled, the new equilibrium constant expression will have the concentrations raised to twice the original powers. For example:

2. 2aA+2bB⇋2cC+2dD
   The equilibrium constant, K_Cx​
   ​

3. Halving the Coefficients:

   • If the coefficients are halved, the equilibrium constant expression will involve the square root of the original equilibrium constant K_c​:

​​

1. Summing Reactions:

   • If two or more reactions are combined, the equilibrium constant for the overall reaction is the product of the equilibrium constants of the individual reactions:

Types of Equilibria

• Homogeneous Equilibrium: All reactants and products are in the same phase (usually gas phase). For example, reactions occurring entirely in the gas phase.

• Heterogeneous Equilibrium: Reactants and products exist in more than one phase (solid, liquid, gas). For example, a reaction between a solid and a gas.

Important Considerations

• Pure Solids and Liquids: In equilibrium expressions, pure solids and pure liquids are not included because their concentrations do not change and are considered constant.

• Effect of Temperature: The equilibrium constant K_c​ changes with temperature. A higher K_c (greater than 1) suggests that products are favored at equilibrium, whereas a smaller K_c​ (less than 1) indicates that reactants are favored.

Vapor Pressure and Phase Equilibrium

• Vapor Pressure is the pressure exerted by the gas phase of a substance when it is in equilibrium with its liquid phase.

  • In phase equilibrium (e.g., liquid ↔ gas), the system reaches a point where the rate of evaporation equals the rate of condensation.

  • The vapor pressure is a characteristic of the liquid and depends on temperature.

  • The size and shape of the container do not affect the equilibrium vapor pressure.

Effects of Conditions on Equilibrium Constant

When a system is at equilibrium, the position of equilibrium remains stable as long as temperature and pressure stay constant. However, certain experimental changes can affect the equilibrium position or shift the balance of reactants and products. While the equilibrium position may shift in response to changes, the equilibrium constant Kc​ itself only changes with temperature.

1. Effect of Concentration Changes

• Adding or removing reactants or products will shift the equilibrium position to counter the change. This is in line with Le Chatelier’s Principle, which states that a system at equilibrium will adjust to counteract an external change.

  • Adding reactants shifts equilibrium towards the products (forward reaction) to consume the added substance.

  • Removing reactants shifts equilibrium towards the reactants (reverse reaction).

• Equilibrium constant Kc​ remains unchanged because it is only influenced by temperature.

2. Effect of Pressure Changes (for Gaseous Reactions)

• In reactions involving gaseous reactants and products, changes in pressure can impact the equilibrium position.

• Increasing pressure will shift the equilibrium toward the side with fewer gas molecules, which helps reduce overall pressure in the system.

  • For example, in a reaction where 4 moles of gaseous reactants produce 3 moles of gaseous products, an increase in pressure will favor the forward reaction to reduce the number of moles.

• Decreasing pressure will favor the side with more gas molecules.

• Kc remains constant because pressure does not directly alter the equilibrium constant; it only shifts the position.

3. Effect of Temperature Changes

Temperature is unique in that it directly affects both the equilibrium position and the value of Kc​:

• Exothermic Reactions:

  • Heat is a product in exothermic reactions. Adding heat (increasing temperature) causes the equilibrium to shift left, favoring the reactants, to reduce the added heat.

  • Decreasing temperature shifts equilibrium right, favoring the products.

  • Kc changes based on the temperature change:

    • For exothermic reactions, increasing temperature decreases Kc​ as the system favors reactants.

    • For endothermic reactions, increasing temperature increases Kc as the system favors products.

• Endothermic Reactions:

  • Heat is a reactant in endothermic reactions. Adding heat shifts equilibrium to the right to absorb the added energy.

  • Kc increases with temperature for endothermic reactions because the products are favored.

4. Effect of a Catalyst

• A catalyst lowers the activation energy of both the forward and reverse reactions, speeding up the time it takes to reach equilibrium but not changing the equilibrium position or Kc​.

• It does not change the equilibrium constant or shift the balance of products and reactants.

Note: In any reaction where the enthalpy change ΔH=0, temperature changes have no effect on Kc since no heat is absorbed or released.

Le Chatelier’s Principle

Le Châtelier’s Principle explains how a system at equilibrium responds to changes in conditions (concentration, temperature, or pressure) to maintain or restore balance. When an external change disturbs a system in equilibrium, the system will adjust itself to offset the change and re-establish equilibrium. This principle helps predict the direction in which equilibrium shifts when the system is altered.

Key Influencing Factors on Equilibrium

1. Concentration Changes

   • When the concentration of a reactant increases or a product decreases, the system is no longer at equilibrium. To offset this, the forward reaction is favored to convert some of the added reactants into products, moving the system back toward equilibrium.

   • When the concentration of a product increases or a reactant decreases, the reverse reaction is favored, shifting equilibrium toward the reactants to counteract the imbalance.

   • Kc (equilibrium constant) remains unchanged by concentration changes, as only temperature affects its value.

2. Temperature Changes

   • The direction of the equilibrium shift due to temperature depends on whether the reaction is endothermic (absorbs heat) or exothermic (releases heat).

3. Endothermic Reactions (ΔH>0\Delta H > 0ΔH>0):

   • Increasing temperature adds heat, which acts like a reactant in endothermic reactions, so the forward reaction is favored, and equilibrium shifts toward the products.

   • Decreasing temperature removes heat, favoring the reverse reaction, shifting equilibrium toward the reactants.

4. Exothermic Reactions (ΔH<0):

   • Increasing temperature adds heat, acting as an excess product, so the system responds by favoring the reverse reaction and shifting equilibrium toward the reactants.

   • Decreasing temperature removes heat, favoring the forward reaction and shifting equilibrium toward the products.

   • Effect on Kc​: Unlike other factors, temperature changes the equilibrium constant (Kc) because heat is either absorbed or released, impacting the system’s balance.

5. Pressure Changes (for Reactions with Gaseous Species)

   • For reactions involving gases, pressure changes affect equilibrium if there is a difference in the number of gas molecules (moles) between the reactants and products.

   • If products have fewer gas molecules:

     • Increasing pressure favors the side with fewer gas molecules to reduce pressure, shifting equilibrium toward the products.

     • Decreasing pressure favors the side with more gas molecules, shifting equilibrium toward the reactants.

   • If reactants have fewer gas molecules:

     • Increasing pressure shifts equilibrium toward the reactants, favoring the reverse reaction.

     • Decreasing pressure shifts equilibrium toward the products, favoring the forward reaction.

   • Effect on Kc​: Pressure changes do not alter Kc​, as it is only dependent on temperature.

6. Effect of a Catalyst

   • A catalyst speeds up both the forward and reverse reactions equally by lowering their activation energy, allowing the system to reach equilibrium more quickly.

   • Equilibrium Position: The catalyst does not shift the equilibrium position since it affects both reactions equally.

   • Equilibrium Constant Kc​: Adding a catalyst does not change Kc​ because it does not affect the relative concentrations of reactants and products at equilibrium.

Le Châtelier’s Principle helps predict how a system in equilibrium will respond to various changes, guiding the direction of reaction shifts to maintain or re-establish equilibrium.

Reaction Quotient

The concept of the reaction quotient (denoted by Q) and how it compares to the equilibrium constant (Kc) is essential for understanding the progress of a chemical reaction as it moves toward equilibrium. Let me explain it step by step to ensure clarity.

Reaction Quotient (Q)

The reaction quotient Q is a mathematical expression similar to the equilibrium constant Kc​, but the key difference is that Q is calculated using the concentrations of reactants and products at any point in time, whether or not the system is at equilibrium.

For a generic reaction of the form:

The reaction quotient Q is given by:

where [C],[D],[A], and [B] are the molar concentrations of the chemical species C, D, A, and B at a given moment in time. This expression is exactly the same as the equilibrium expression for Kc​, but instead of equilibrium concentrations, Q uses the concentrations at any moment.

Comparing Q and Kc​

1. When Q=Kc​:

   • The reaction is at equilibrium.

   • At equilibrium, the forward and reverse reactions occur at the same rate, so the concentrations of reactants and products do not change. The system is balanced, and no net change in the concentrations will occur unless the conditions (e.g., temperature or pressure) change.

2. When Q>Kc​:

   • This means that, at this point in time, the concentration of products is higher than it would be at equilibrium.

   • Since the reaction tends to favor the direction that reduces the concentration of products and increases the concentration of reactants to reach equilibrium, the reverse reaction will be favored.

   • The system will shift toward the left (toward the reactants) to decrease the concentration of products and increase the concentration of reactants, moving toward equilibrium.

3. When Q<Kc​:

   • This indicates that the concentration of reactants is higher than it would be at equilibrium.

   • To reach equilibrium, the system will favor the forward reaction (the production of products).

   • The reaction will shift to the right (toward the products) to decrease the concentration of reactants and increase the concentration of products, moving toward equilibrium.

How Q Helps Predict the Reaction's Direction

The value of Q gives you a snapshot of the system's current state relative to equilibrium. By comparing Q to Kc​, you can predict:

• If Q<Kc​, the reaction will move forward (toward the products) to reach equilibrium. The system needs to produce more products and consume more reactants.

• If Q>Kc​, the reaction will move backward (toward the reactants) to reach equilibrium. The system needs to decrease the amount of products and generate more reactants.

• If Q=Kc​, the system is already at equilibrium, and no net change will occur in the concentrations of reactants and products.

Significance of the Reaction Quotient

• Helps predict the shift: By knowing the value of Q compared to Kc​, you can determine whether the reaction needs to proceed forward (to make more products) or backward (to make more reactants) to achieve equilibrium.

• Indicates the direction of change: The reaction will always move in the direction that brings the system closer to equilibrium. If there are too many products (when Q>Kc), the system will shift toward reactants. If there are too many reactants (when Q<Kc​), the system will shift toward products.

Example

Let's take an example to illustrate how Q works.

Consider the following reaction:

The equilibrium expression for this reaction is:

Now, suppose at a given time the concentrations are as follows:

The reaction quotient at this point would be:

If the equilibrium constant Kc​ = 1.0, we see that:

Since Q<Kc​, the concentration of reactants (N₂ and H₂) is greater than at equilibrium, and the system will shift forward to produce more NH₃​, moving toward equilibrium.

Unit 9: Redox Processes

Oxidation and Reduction

What is Electron Transfer Theory?

• all reactions are a combination of 2 parts (half-reactions)

• both equations are balanced by mass and by charge (the number electrons lost by one atom is gained by the other)

• Reduction: # of electron is increased

• Oxidation: # of electron is lost

• Reducing Agent: gives or lose electrons

• Oxidizing Agent: gains or accepts electrons

• Oxidation number: the charge that an atom in a compound would have if the electron pair in the bond belonged solely to the more electronegative atom

  • Oxidation numbers always refer to single atoms

Why does LEO say GER?

• loss of electrons = oxidation

• gain of electrons = reduction

Oxidation Number Rules

1. The oxidation number for any atom in an element is zero.

2. The oxidation number of a monatomic ion is equal to the charge on the ion.

3. The oxidation number of each hydrogen atom in most of its compounds is +1, except hydrides (which are -1).

4. The oxidation number of each oxygen atom in most of its compounds is - 2.

5. Peroxides are an exception (they are -1). In OF2 oxygen is + 2.

6. In compounds, the elements of group 1, group 2, and aluminum have positive oxidation numbers of +1, +2, and +3, respectively.

The sum of the oxidation numbers of all the atoms must equal the apparent charge of that particle

• An increase in the oxidation number indicates that an atom has lost electrons and therefore oxidized

• A decrease in the oxidation number indicates that an atom has gained electrons and therefore reduced

Balancing Redox Reactions in Acidic Conditions: The Half-Reaction Method

Writing half-reactions:

1. Assign oxidation numbers

2. Separate into the two half-equations

For each equation:

1. Balance all atoms other than oxygen and hydrogen

2. Balance oxygens by adding H20(1) (if needed)

3. Balance hydrogens by adding H+(aq) ions (if needed)

4. Balance charges by adding electrons to the more positive side

5. Double check each half-equation is balanced in terms of atoms and charge

Combining half-reactions:

1. Make sure the number of electrons in the two-half equations are equal (If not, multiply each half reaction equation by simple whole numbers to balance the electrons gained/lost)

2. Add the two half-reaction equations, canceling out anything that is the same on both sides of the reaction

Predicting Redox Reactions

Electron Tug-of-War

• A redox reaction can be viewed as a competition for electrons between substances

• Example - Zinc and Copper(Il) sulfate

Development of the Reactivity Series

• Some substances are better at oxidizing than others

• Can be placed in order of their oxidizing ability

• Determined by performing various single displacement reactions and examining whether they occur spontaneously or not

• More reactive metals are stronger reducing agents than reactive metals; more reactive non-metals are stronger oxidizing agents than less reactive non-metals

Predicting Spontaneity of Redox Reactions

• A spontaneous reaction occurs only if the oxidizing agent (A) is below the reducing agent (RA) in a table of relative strengths of oxidizing and reducing agents

Electrochemical Cells

Voltaic and Electrolytic Cells

Standard Cell Potentials

• E° cell is the maximum electric potential difference (voltage) of the cell operating under standard conditions (SATP: 25°C, IM)

• It represents the energy difference (per unit charge) between the cathode and the anode

• All reduction potentials are measured with reference to the standard hydrogen electrode

To Calculate

• E° cell = E° (cathode) - E° (anode)

• Where E° represents the "standard electrode potential"- the ability of the half-cell to attract electrons (under standard conditions), thus undergoing a reduction

• If it is a positive value (according to chart in data booklet), then it is more likely to undergo reduction compared to a negative value

How can you use Standard Cell Potential to determine the spontaneity of a chemical reaction?

Uses:

• Electrolysis to produce pure elements from compounds

• Electroplating: plate or coat some object with a metal (ex. gold or silver-plated jewelry)

Standard Hydrogen Electrode

• Can act as both an ANODE as well as cathode in an electrochemical cell

• Used to measure all the other electrodes

Unit 10: Organic Chemistry

Unit 1: Organic Nomenclature & Structure

Homologous Series

• Organic chemistry is the study of carbon-containing compounds

  • Eg: butane

    • molecular formula: C₄H₁₀

    • empirical formula: C₂H₅

    • condensed formula: CH₃CH₂CH₂CH₃

• Homologous Series: Same functional group

  • The boiling point increases as the carbon chain goes up

  • Refer to the Homologous chart

• For skeletal form: each point has a carbon atom

Naming Hydrocarbons

• Alkanes: name (longest), then alkyl groups

  • Arrange names of substituent groups in alphabetical order (ignoring prefixes)
    
    

• Alkenes: alk-x-ene

  • C=C (Double Bond)

    • Double bond starts at the lowest number

• Alkynes: alk-x-yne

  • CC (Triple Bond)

    • Triple bond starts at lowest number

• Halogen Alkane: x-halo alkane
  
  

Naming Oxygen-Containing Organic Compounds

• Alcohol: Alkan-x-ol (Methanol & ethanol no need for #)

• Ether: x-alkoxy alkane (treat alkoxy as substituent which is substituent with the lowest #)

  • give -OH the lowest #
    
    

• Aldehyde: alkanal, CHO is labelled #1

• Ketone: Alkan-x-one, C=O group give the lowest #

  • no need for # for propanone or butane
    
    

• Carboxylic acid: alkanoic acid, COOH is labelled #1

• Ester: carboxylic acid + alcohol

  • Alkyl alkanoate (alkyl for alcohol, alkanoate end for carboxylic acid side)

  • Numbering in the acid starts from C=O & alcohol starts from O-C group
    
    

Structual Isomers

• Primary, secondary & tertiary compounds

  • Depends on what the C attached to OH is attached to

• Structural Isomers: same molecular formula but different structural formula (atoms joined together differently)

• Branched-chain isomers (less surface area) have lower boiling points than straight chain isomers (high surface area, high london-dispersion forces)

Stereoisomers

• Same structural formula but atoms are arranged differently in space

  • Cis means same side

  • Trans means opposite side

• E/Z Priority rules (higher priority to atom attached to C=C with higher atomic #)

• Optical Isomers: When 4 different atoms/groups attached to a single carbon atom (Chiral)

  • Mirrors (enantiomers)
    
    

• Racemic mixture: Equimolar mixture of 2 enantiomers

Unit 2: Organic Reactions & Mechanisms

Reactions of Hydrocarbons

• Complete oxidation of a hydrocarbon

  • hydrocarbon + oxygen = carbon dioxide + water

    • C₆H₁₂ + 9O₂ (g) = 6CO₂ (g) + 6H₂O (l)

  • When oxygen supply is limited, a hydrocarbon will undergo incomplete combustion (only form carbon if O2 extremely limited)

    • C₃H_{8 (g)} + 2O_{2 (g)} = 3C_(g) +4H₂O_(l)

Free radical substitution mechanism

• Substitution of alkanes
  
  

Nucleophilic substitution reaction mechanism

• Nucleophilic substitution of halogenoalkanes

  • Nucleophiles: electron pair donors, attracted to electron deficient carbon atoms

  • Mechanisms

    • 1° halogenoalkane (SN2 Mechanism)
      rate = k [Rx][OH-]

      • inversion of configuration if a nucleophile attacks a chiral centre

    • 3° halogenoalkane (SN1 Mechanism)
      rate = k [RA]

      • not stereospecific - racemix mixture formed

    • 2° halogenoalkanes undergo a mixture of Sn1 and Sn2 mechanisms

    • Factors that affect the rate of nucleophilic substitution reaction

      • Identity of nucleophile (only Sn2 reaction affected)

        • Anions more reactive than neutral species

      • Identity of halogen (Sn1 & Sn2)

      • 1°, 2°, 3° halogenoalkane

        • Sn2 ratio is 1°>2°>3°

        • Sn1 rate is 3°>2°>1°

      • Choice of Solvent

Electrophilic addition mechanism

• Electrophile: electron-deficient species

  • Attracted to regions of relatively high electron density
    
    

  • Halogenation

    • reaction with Br₂

      • alkanes will have no reaction because it requires heat as a catalyst. therefore it will remain red

      • alkenes will go colourless because it does not require a catalyst

Markovnikov’s Rule

• if more than one product is possible, the more electronegative atom will end up on the carbon atom of the double cond that has fewer hydrogen
  
  

• the carbocation formed is one that has its positive charge on the most substituted carbon

• Addition polymerization of alkenes

Reactions of Oxygen-containing Compounds

Mild oxidation

• (Controlled oxidation of an alcohol to create other functional groups)
  
  

Reduction reactions

• (reverse oxidation reactions)

  • Reducing agents include lithium aluminum hydride (LiAlH₄ (This is stronger) and NaBH₄
    
    

• Reduction of Nitrobenzene
  
  

• Formation of ester (condensation reaction)
  
  

Organic Synthesis

Unit 3: Spectroscopy

IR Spectroscopy

• Index of hydrogen deficiency (IHO)

  • Double bond = 1

  • Triple bond = 2

  • Ring = 1

  • Aromatic ring = 4

  • IHO = ½ (2C + 2 - h - x + n)

• IR is absorbed by certain bonds causing them to stretch or bend

• Bond will only interact with IR radiation if it is polar

  • Match wavenumbers with bonds

• Fingerprint region (1500-650 cm^-1) difficult to interpret, lots of C-C & C-H bond vibrations

Mass Spectrometry

• Measures relative masses of atoms or ions

• Measures mass-to-change ratio of ions

• Ionization causes molecule to break up into different fragments

• Greatest mass peak is parent ion (Molecular mass)
  
  

‘H NMR Spectrometry

• Nuclei in different chemical environments produce different signals in the spectrum

• Signals are measured against the standard signal produced by TMS (8=0 ppm)

• Why TMS?

  • 12 protons, all in the same environment

  • Strong signal even when present in small amounts

  • Chemical shift value very low

• The area under a peak is proportional to the number of proton atoms in that environment

  • Integration trace distance/height of each step is ratio between # of protons in each environment
    
    

  • If there are n H’s on an adjacent atom, the signal for a particular proton will be split into n+1 peaks
    
    

  • Intensities of peaks are given by Pascal’s Triangle

Unit 11: Measurement and Data Processing

11.1 Uncertainties and Errors in Measurements and Results

Qualitative & Quantitative Analysis

Analysis of data can be classified into two types:

1. Qualitative Analysis: Comes from observations and non-numerical methods

2. Quantitative Analysis: Comes from measurements and is always associated with uncertainties determined by either the apparatus or human limitations such as reaction times and sight.

Uncertainties

In science, numerical data is divided into two types:

1. Data with exact numbers (no uncertainty)

2. Data with inexact numbers (Degree of uncertainty involved)

Precision vs Accuracy

Precision: Closeness of agreement between interdependent test results

Accuracy: Closeness of agreement between the result of measurement and the true value

Experimental Error

• Every measurement has a degree of uncertainty called experimental error.

• There are two types of experimental error

  • Systematic error: A flaw in experimental design or methodology. This affects the accuracy of the results.

  • Random error:

  • -

Absolute and Relative Uncertainty

• Absolute Uncertainty is the margin of uncertainty associated with the result of a measurement.

  • Its symbol is given by ΔA

• Relative Uncertainty is the ratio comparing the size of the absolute uncertainty.

  • Relative uncertainty = ΔA / A

• All experimental results should be reported in the form:

  • Experimental result = (A ± ΔA) units

Percentage Error

11.2 Graphical Techniques

• Graphical techniques: an effective means of communicating the effect of an independent variable on a dependent variable, and can lead to the determination of physical quantities.

• Sketched graphs have labeled but unscaled axes.

  • They are used to show qualitative trends, such as variables that are proportional or inversely proportional.

  • Units generally would not need to be shown on a sketch, only the variables.

• Drawn graphs have labeled and scaled axes.

  • They are based on quantitative measurements.

  • Drawn graphs always display the appropriate units for variables.

• Dependence is considered any statistical relationship between two sets of data or between two random variables.

  • In a graph of Y versus X, the independent variable (that is, the cause) is plotted on the x-axis, and the dependent variable (the effect) is plotted on the y-axis.

• Correlation can be described as a statistical measure and technique that indicates the degree and direction of the relationship between two sets of variables.

  • A positive correlation is where the two variables increase or decrease in parallel to one another.

  • A negative correlation is one in which one variable increases while the second variable decreases or vice versa.

• Correlations can be deduced from the correlation coefficient, represented by the symbol, r.

  • This coefficient is a measure of the strength of the relationship between two variables. Data are often represented by scatter plots that show the scatter of various points on a graph.

• r = +1, is indicative of a perfect positive linear relationship (all points lie on a straight line)

• r = 0, no linear relationship exists (there is a complete scatter of points)

• r = -1, is indicative of a perfect negative linear relationship

Slope (m): the tangent of the angle θ, that the line makes with the positive direction of the x-axis.

Intercept (c): points where the line cuts the y-axis at x=0

• Can be found by→

  • Extrapolation

  • Equation of a line [y = mx + b]

Best Fit Line: a straight line that minimizes the distance between it and some data.

11.3 Spectroscopic Identification of Organic Compounds

Degree of unsaturation or index of hydrogen deficiency (IHD)

• Used to determine the number of rings or double bonds from a molecular formula

  • Double bond = 1 IHD

  • Triple bond = 2 IHD

  • Ring = 1 IHD

  • Aromatic = 4 IHD (3 double bond + 1 ring)

IHD from Structure

IHD from Molecular formula

• For the generic formula C_cH_hN_nO_oX_x:

  • IHD = (0.5)(2c + 2 – h – x + n)

Ex:

Electromagnetic Spectrum

• Various regions of EMS are the basis of different types of spectroscopy

The energy of electromagnetic radiation, E, is related to the frequency v of the radiation by Planck equation:

h = Plancks constant = 6.63 x 10 ^-34 J s

E = energy of radiation (measured in J)

v = frequency of radiation (measured in Hz)

c = speed o light = 3.00 x 10⁸ m s^-1

λ = wavelength (measured in m)

Techniques used to identify structures of substances

• X-rays – as their energy is high, these cause electrons to be removed from the inner energy levels of atoms. Direction patterns can lead to information such as the bond distances and bond angles in a structure and form the basis of X-ray crystallography.

• Visible and UV light–electronic transitions and hence this type of spectroscopy gives information about the electronic energy levels in an atom or molecule. This is the basis of UV-vis spectroscopy.

• Infrared Radiation – causes certain bonds in a molecule to vibrate (for example, stretch and bend) and as such provides information on the functional groups present. This is the basis of IR spectroscopy.

• Microwaves – cause molecular rotations and can give information on bond lengths.

• Radiowaves – can cause nuclear transitions in a strong magnetic field because radiowaves can be absorbed by certain nuclei, which causes their spin states to change. Nuclear magnetic resonance (NMR) spectroscopy is based on this and information on different chemical environments of atoms can be deduced, which leads to information on the connectivity of the atoms present in a molecule.

Spectroscopy

Infrared Spectroscopy

IR radiation does not have sufficient energy to result in electronic transitions but can cause molecular vibrations, which result in the vibration of certain groups of molecules about their bonds. The basis of IR spectroscopy is the spring model.

In the spring model, every covalent bond is considered as a spring. Such a spring can be stretched (both symmetrically and asymmetrically), bent, or twisted, giving rise to a distortion.

Hookes law: The force required to cause the vibration.

k = spring constant

Atom’s Vibration

• Lighter atoms vibrate at higher frequencies (v).

• Heavier atoms vibrate at lower frequencies (v).

• Multiple bonds (e.g., double and triple bonds) follow the same trend as lighter atoms.

• Different molecules absorb at different frequencies due to variations in bond enthalpy.

• IR absorptions are typically reported in wavenumber (1/λ), measured in cm^-1.

Proton Nuclear Magnetic Resonance (¹H NMR) Spectroscopy

• Provides information about the different chemical environments of hydrogen atoms in a molecule.

• Based on the principle that the nuclei of hydrogen atoms have two possible spin states and act like tiny magnets.

• The position of the spin relative to the standard spin (tetramethylsilane, TMS) is called the chemical shift.

• Gives information about the relative number of hydrogens in each environment, presented as a ratio.

• The area under the curve in a spectrum corresponds to the number of hydrogens in that environment.

Mass Spectroscopy

• Provides additional information about functional groups in a molecule.

• When a gaseous molecule is ionized, a molecular ion (M⁺) is formed.

• The fragmentation pattern in a mass spectrum shows the masses of fragments after functional groups are removed.

• The molecular ion peak in a mass spectrum represents the molecular mass of the compound.

Absolute and Relative Uncertainty

• Absolute Uncertainty is the margin of uncertainty associated with the result of a measurement.

  • Its symbol is given by ΔA

• Relative Uncertainty is the ratio comparing the size of the absolute uncertainty.

  • Relative uncertainty = ΔA / A

• All experimental results should be reported in the form:

  • Experimental result = (A ± ΔA) units

Percentage Error

11.2 Graphical Techniques

• Graphical techniques: an effective means of communicating the effect of an independent variable on a dependent variable, and can lead to the determination of physical quantities.

• Sketched graphs have labeled but unscaled axes.

  • They are used to show qualitative trends, such as variables that are proportional or inversely proportional.

  • Units generally would not need to be shown on a sketch, only the variables.

• Drawn graphs have labeled and scaled axes.

  • They are based on quantitative measurements.

  • Drawn graphs always display the appropriate units for variables.

• Dependence is considered any statistical relationship between two sets of data or between two random variables.

  • In a graph of Y versus X, the independent variable (that is, the cause) is plotted on the x-axis, and the dependent variable (the effect) is plotted on the y-axis.

• Correlation can be described as a statistical measure and technique that indicates the degree and direction of the relationship between two sets of variables.

  • A positive correlation is where the two variables increase or decrease in parallel to one another.

  • A negative correlation is one in which one variable increases while the second variable decreases or vice versa.

• Correlations can be deduced from the correlation coefficient, represented by the symbol, r.

  • This coefficient is a measure of the strength of the relationship between two variables. Data are often represented by scatter plots that show the scatter of various points on a graph.

• r = +1, is indicative of a perfect positive linear relationship (all points lie on a straight line)

• r = 0, no linear relationship exists (there is a complete scatter of points)

• r = -1, is indicative of a perfect negative linear relationship

Slope (m): the tangent of the angle θ, that the line makes with the positive direction of the x-axis.

Intercept (c): points where the line cuts the y-axis at x=0

• Can be found by→

  • Extrapolation

  • Equation of a line [y = mx + b]

Best Fit Line: a straight line that minimizes the distance between it and some data.

11.3 Spectroscopic Identification of Organic Compounds

Degree of unsaturation or index of hydrogen deficiency (IHD)

• Used to determine the number of rings or double bonds from a molecular formula

  • Double bond = 1 IHD

  • Triple bond = 2 IHD

  • Ring = 1 IHD

  • Aromatic = 4 IHD (3 double bond + 1 ring)

IHD from Structure

IHD from Molecular formula

• For the generic formula C_cH_hN_nO_oX_x:

  • IHD = (0.5)(2c + 2 – h – x + n)

Ex:

Electromagnetic Spectrum

• Various regions of EMS are the basis of different types of spectroscopy

The energy of electromagnetic radiation, E, is related to the frequency v of the radiation by Planck equation:

Unit 12: Atomic Structure 

The Nuclear Atom

Isotopes

• Isotopes: atoms of the same element that have the same number of protons in the nucleus but a different number of neutrons

  • -differnt mass number

• Ex: iodine-131, cobalt-60, lutetium-177

• Stable element and isotopes have same chemical properties because same number of electrons

Calculating Relative Atomic Mass

Mass Spectrometer: used to determine the relative atomic masses of elements and is used to determine the structure of organic compounds

Steps to Calculate Relative Atomic Mass:

1. The sample being studied would be vaporized to form a gas

2. It is bombarded with high-energy electrons, producing positive ions (+1)

3. The positive ions are accelerated in an electric field

4. The positive ions are deflected in a magnetic field depending on the mass ot charge ratio (m/z)

5. The ions with a higher m/z are deflected less in the magnetic field

6. The positive ions reach the detector, where they produce a mess spectrum

Electrons in Atom

Atomic Orbitals

Energy Levels and Sublevels

• Bohr Model → electrons exist in energy levels

• Principle Energy Levels: assigned numbers with n = 1 being the closest to the nucleus and of lowest energy, with the higher numbers being further from the nucleus

• The main energy levels ar split into sub-levels: s,p,d,f

  • Atomic Orbital: a region of space where there is a high possibility of finding an electron

  • S orbital → spherical

  • P orbital → dumbbell shaped

  • D orbital → 4-petal flower

• Pauli Exclusion Principle: two electrons cannot have the same quantum number

  • Number of electrons per main energy level = 2n^2

• Heisenberg’s Uncertainty Principle: It is not possible to know, at th same time, the exact position and momentum of an electron. Instead, only a probability can be stated than an electron will be somewher in a given region of space

Electron Configuration

• Aufbau Principle: electrons fill atomic orbital of lowest energy first

• Within a main energy level, s orbitals are of lower energy than p orbitals and therefore fill first

• Degenerate orbitals: atomic orbials within a sub-level are of equal enrgy

  • 3 p orbitals (2p, 3p, & 4p sub levels)

  • 5 d orbitals (3d sublevel)

• There is an overlap between the 3d and 4s sublevel → 4s is of lower energy and therefore fills first

• Condensed electron Configuration: a shorthand version of writing the electron configuration for atoms or ions using nobel gasses

  • Ex: bromine

• Exceptions of Aufbau Principle

  • Chromium

  • Copper

• Orbital Diagrams: used to represent electrons in atomic orbitals

  • Boxes represent atomic orbitals and arrows represent electron

  • Arrows point opposite direction to represent opposite spins

  • Only two electrons/arrows per box

Line Spectra

• Higher energy → higher frequency → shorter wavelength

• Lower energy → lower frequency → longer wavelength

Continuous, Emission & Absorption Line Spectra

The spectrum shows all wavelengths of visible light

Absorption Line Spectra: produced when electrons absorb energy and transition form lower to higher energy levels

Some visible light wavelengths are missing, and are usually shown as black lines on a colored background

Emmision line spectra: produced when electrons emit enryg and tramisition form higher to lower energy levels

The energy emitted when the electrons make these transitions corresponds to the wavelength or frequency of visible light

When moving to the high-energy end of the spectrum (right ot left), lines get closer together

Violet light, having a shorter wavelength than red light, is of higher energy

As electrons absorb/emit energy, they transition between the energy levels

Hydrogen Emission Spectrum

As voltage passes through a sample of hydrogen, electrons are excited to higher energy levels

Electron Transitions →

N = 1, Ultraviolet radiation

N = 2, visiblelight

N = 3, infrared radiation

The longerthe arrow, the greater the amount of energy emitted

The electron trnisiton form n=5 to n = 1 emits the largest amount of energy

The Nuclear Atom

Isotopes

Isotopes: atoms of the same element that have the same number of protons in the nucleus but a different number of neutrons

-different mass number

Ex: iodine-131, cobalt-60, lutetium-177

Stable element and isotopes have same chemical properties because same number of electrons

Calculating Relative Atomic Mass

Mass Spectrometer: used to determine the relative atomic masses of elements and is used to determine the structure of organic compounds

Steps to Calculate Relative Atomic Mass:

Bohr Model → electrons exist in energy level

Principle Energy Levels: assigned numbers with n = 1 being the closest to the nucleus and of lowest energy, with the higher numbers being further from the nucleus

The main energy levels ar split into sub-levels: s,p,d,f

Atomic Orbital: a region of space where there is a high possibility of finding an electron

Pauli Exclusion Principle: two electrons cannot have the same quantum number

Number of electrons per main energy level = 2n^2

Heisenberg’s Uncertainty Principle: It is not possible to know, at th same time, the exact position and momentum of an electron. Instead, only a probability can be stated than an electron will be somewher in a given region of space

Electron Configuration

3 p orbitals (2p, 3p, & 4p sub levels)

5 d orbitals (3d sublevel)

There is an overlap between the 3d and 4s sublevel → 4s is of lower energy and therefore fills first

Condensed electron Configuration: a shorthand version of writing the electron configuration for atoms or ions using nobel gasses

Ex: bromine

Orbital Diagrams: used to represent electrons in atomic orbitals

Boxes represent atomic orbitals and arrows represent electron

Arrows point opposite direction to represent opposite spins

Only two electrons/arrows per box

Line Spectra

Higher energy → higher frequency → shorter wavelength

Lower energy → lower frequency → longer wavelength

Continuous, Emission & Absorption Line Spectra

The spectrum shows all wavelengths of visible light

Absorption Line Spectra: produced when electrons absorb energy and transition form lower to higher energy levels

Some visible light wavelengths are missing, and are usually shown as black lines on a colored background

Emmision line spectra: produced when electrons emit enryg and tramisition form higher to lower energy levels

The energy emitted when the electrons make these transitions corresponds to the wavelength or frequency of visible light

When moving to the high-energy end of the spectrum (right ot left), lines get closer together

Violet light, having a shorter wavelength than red light, is of higher energy

As electrons absorb/emit energy, they transition between the energy levels

Hydrogen Emission Spectrum

As voltage passes through a sample of hydrogen, electrons are excited to higher energy levels

Electron Transitions →

N = 1, Ultraviolet radiation

N = 2, visiblelight

N = 3, infrared radiation

The longerthe arrow, the greater the amount of energy emitted

The electron trnisiton form n=5 to n = 1 emits the largest amount of energy