mass

Chemistry Notes 

1.1 The Nature of Chemistry  

Empirical Knowledge - Knowledge that comes from investigation and observations.  

Theoretical knowledge - Knowledge that explains scientific observations. 

Matter - Matter is defined as anything that has mass and occupies space.  

Theory - An explanation or model based on observation, experimentation, and reasoning.   

1.2 Atomic Structure 

Chemistry - The study of the composition and properties of matter 

Clues to indicate the occurrence of Chemical Change 

• Difficult to reverse 

• Gas produced 

• Energy released/absorbed 

  • Thermal 

  • Light 

• Precipitation forms  

• Color change 

Atoms - Atoms in greek means “uncuttable” or “indivisible”  

The History of the Atom 

• Democritus 

  • Proposed that matter is made up of individual particles called atoms.  

  • Atoms can vary in size, are constantly moving, and separated by empty space. 

• Aristotle 

  • Proposed the four elements model of matter

Note: Neither of the above ideas were based on experimental observations 

Alchemy and Alchemists 

• Searched for the “elixir of life” (to bring immortality), and the philosopher’s stone (to turn common metals into Au) 

• Developed alloys, labware, handling procedures for chemicals and various procedures. 

• The word that is hiding in Alchemist is “chemist”.  

Note: Knowledge acquired was based on experimental observation. (this is the big difference from Democritus and Aristotle) 

John Dalton 

• Matter is made up of tiny, indivisible particles called atoms

• All atoms of an element are identical 

• Atoms of different elements are different (P and C properties)  

• Atoms can rearrange to form new substances in chemical reactions, but they are never created or destroyed.   

• Came up with the Billiard Ball Model as the representation of the atom.  

The periodic table of the elements is also known as the periodic table of atoms 

Atoms and elements are not technically the same. 

Atoms are like alphabets, they are used to create different compounds, etc.  

J.J. Thomson  

• Discovered the electron, e^- (negatively charged subatomic particle). 

• The e- and the positive charge of the atom is evenly distributed.  

• Featureless vs Subatomic particles.  

• Raisin Bun/Plum Pudding Model 

• J.J. Thomson came up with the plum pudding model also known as the blueberry muffin model.  

Ernest Rutherford

• Quote: “It was almost as credible as if you fired a 15” shell at a piece of tissue paper and it came back and hit you” 

• Discovered the proton, p⁺ (positively charged subatomic particle) 

• Said that the protons occupy the center (nucleus) of the atom and the electrons move around them held by electrical attraction. (This means that the majority of the space of atoms is empty)  

• He performed the famous gold foil experiment. 

• Helium nucleus also known as the alpha particle 

• An alpha particle consist of two protons and two neutrons

James Chadwick  

• Discovered the neutron, n⁰ (unchanged subatomic particle) 

• Electrons, Protons and Neutrons are collectively known as nucleons  

• The presence of neutrons helped to explain the overall mass of the atom. 

Neils Bohr 

• Proposed the organization and behavior of electrons 

• Said that electrons revolve around the nucleus in orbits of fixed energy called energy levels (electrons further away have some energy and vice versa) like planets around the sun.  

• He said that the space between the nucleus and the neutrons is empty space. 

• Said that electrons can “jump” to a higher energy level (excited state) if they absorb energy and can return back to their original energy level (ground state) if they lose the energy that they absorbed 

• Came up with the planetary Model  

• First Valence is 2, Second Valance is 8, and the Third Valance is 18 

• 2n², “n” represents the energy level “orbit”.  

• Electrons can transition from their ground state to their excited state.  

• You can use thermal energy to excite electrons and in some cases light will be produced.  

• Opposite charges will attract and like charges will repel each other

• The Quantum mechanical model is the most current model of the atom - post (after) Neil Bor

Subatomic Particles and Isotopes 

• “Nucleons”

• Compose nucleons 

  • Protons p⁺

  • Neutrons n⁰

    • Not all atoms of the same element have the same number of n⁰ 

    • i.e: there are different isotopes of the same element 

  • Electrons e^-

    • Arranged in energy levels 

• Atomic Number, represented as Z - the number of protons p⁺

  • Z = p⁺ ( = e^- for neutral atom)  

• Mass Number, represented as A - the sum of protons p⁺ and neutrons n⁰

  • A = p⁺ (or Z) + n⁰

• Chemically atoms do not share their electrons willingly and are chemically stable

• Nuclear stability means that the atom will not decompose.  

• There are 118 different types of atoms.  

• There are 92 naturally occurring atoms; the rest are synthetic.  

• The greater the sum of the protons and electrons the greater the mass. 

• Z = to the number of protons that you have (p⁺)

• A = to the sum of protons and neutrons (p⁺ + n⁰) or (Z + n⁰)

Communicating Isotopes 

• Elements Name/Symbol - A 

  • e.g. Magnesium - 24

  • Mg - 24 

• Standard Atomic Notation 

  • e.g. ²⁴₁₂MG

• Example: Determine each of the following for Oxygen-18: 

  • A = 18 

  • Z = 8

  • p⁺ = 8

  • n⁰ = 10

  • e^- = 8 

  • SAN: ¹⁸₈O

Question: An Isotope has a mass number of 198. Determine the number of protons it has if the number of neutrons is 1.54 times the number of protons.  

1.5 Periodic Table and Periodic Law 

Group - A column of elements in the periodic table; sometimes referred to as a family.  

Period - A row in the periodic table. 

Period Law - A rule, developed from many observations, stating that when the elements are arranged in order of increasing atomic number, their properties show a periodic recurrence and gradual change.  

Chemical properties of Halogens are very similar while physical properties are not.  

Main group elements (groups 1,2,13,14,15,16,17,18)

(representative) 

Lewis Symbol - A representation of an element consisting of the chemical symbol and dots to represent the valence electrons; electron dot diagram.  

Octet rule - A generalization stating that when atoms combine, they tend to achieve 8 valence electrons. 

Ion - A charged entity formed when an atom gains or loses one or more electrons. 

Cation - A positively charged ion formed by the removal of one or more electrons from the valence shell of a neutral atom.  

• If you are in group one (from lithium to francium), more often than not you will make a positive one common ion charge. They will all lose that one valence electron 

• Hydrogen can make a negative common ion charge.  

• Group two, will have 2 positive ion charge 

• Boron will never make ions and if it does it will loose three 

• Aluminum will make three plus 

• Tin can lose up to four but it can be multivalent and only loose 2.  

• Group 16 will gain 2 

• Group 17 will gain 1 

1.3 Ions and the Octet Rule 

• Ion - charged atom (monatomic) or molecule (polyatomic ion)   

• Charged means that the (proton ion does not equal the electron ion)  

• An ion can be positively charged or negatively (anion) charged 

• Example: Classify  the following ions using the terms: monatomic, polyatomic, cation, anion 

• a) Mg²⁺ - monatomic, cation 

• b) CO₃^2– - polyatomic, anion 

• c) O₂^2– - polyatomic, anion (also known as peroxide ion)(could be called an diatomic ion because of the 2)        

• d) O^2– - monatomic, anion 

• Monatomic cations (metal ions)

  • Could be single valent or multivalent 

    • Na⁺

    • Mg²⁺ 

    • Fe²⁺ or Fe³⁺ or Ion(II) and Iron(III)

    • Cu⁺ or Cu²⁺ or copper(I) and copper(II)

• Na⁺, Mg²⁺, O^2–, F^– are known as isoelectronic 

• Isoelectronic 

  • With each other and with Ne  

• Polyatomic ions - 

1.7 Periodic Trends in Atomic Properties

• When going from left to right on the periodic table the size of the elements goes up.  

• Electrons can be moved in succession, more than one electron can be removed from an atom. 

• It is much more difficult to take a negative charge.  

• The I.E. for Li is 520 kJ/mol 

• Li_(g) + 520 KJ/mol  → Li⁺_(g) + e^- 

• X_(g) + I^stI.E. → X⁺_(g) + e^-  

• X⁺_(g) + 2^ndI.E. → X²⁺ + e^-  

• X²⁺ + 3^rdI.E. → X³⁺_(g) + e^- 

• Electron Affinity (E.A.): 

  • The energy charge that occurs when an outermost e^– is added to an atom in the gaseous state. 

  • If an atom gets less active (less energy) then there is a positive electron affinity. 

  • If the atoms get more active (more active) then there is a negative electron affinity. 

• The addition of an electron to an atom results in the formation of a negative ion (anion). If the anion is stable, energy is released. The quantitative value for a stable formation can be either positive or negative depending on which convention is used but like I.E. It is stated for 1 mole of atoms. Our text uses positive values to indicate stable e^– additions (high positive value…high E.A. and vice versa).  

• The E.A for Cl is 349 KJ/mol 

• Cl_(g) + e^– → Cl^–_(g) + 349 KJ/mol 

• The electron affinity matches the I.E 

• The higher the ionization energy the higher the electron affinity.  

• The E.A. for Cl is 349 KJ/mol  

• Cl(g) + e^– → Cl^– + 349 KJ/mol 

• The number would be written on the left of the equation if it is unstable.  

• Summery: 

  • As the radius decreases, both I.E. and E.A. increases and vice versa. 

• Arrange the following entities in order of increasing size. 

  • Rb 

  • Ca²⁺

  • Fr 

  • P^3-

  • S^2-

• Correct answer 

  • P^3-

  • S^2-

  • Ca²⁺

  • Rb

  • Fr

2.1 Chemical Bonds & Compounds 

Electrolytes - substance that conducts electric current as a result of a dissociation into positively and negatively charged particles called ions. 

Physical Properties of Ionic Compounds:  

• Crystalline solid at SATP (standard ambient temperature (25^oc) and pressure (100 KPa) ) 

  • Crystalline means being highly ordered entities in their positions 

  • Barometer is used to measure air pressure.  

• hard, but brittle 

  • Hardness is how hard it is to remove one layer of the entities.  

• High melting point 

• Solubility in water varies 

  • Mass of solute per 100 mL H₂O_(L) at 20^oc 

• Can conduct electricity when dissolved in water (electrolight) or when melted.   

• STP - The temperature is 0^oc and the pressure is 101.3 KPa

2.2 Molecular Elements & Compounds 

• A covalent bond forms between two neutral atoms when they share valence e^–. Each e^– that is shared is co-owned by each atom and so shared e^– count towards each atom’s full valence shell.  

• Atoms share 2,4, and 6 valence e^– when forming covalent bonds with each other. 

• H–H

• H:H or H–H

• :O: - this is called a “Lone pair”  

• Remember to learn how to draw the lewis diagrams for shared bonds etc.  

• When two or more atoms bond covalently in such a way that each atom has a full valence shell of electrons, the particles that result are called a molecule. In the solid state, molecules arrange themselves architecturally to produce a molecular crystal lattice (similar to the ionic case but the neighboring molecules are held together by much weaker intermolecular forces of attraction).    

• Many elements exist in molecular form: 

  • H₂

  • O₂

  • N₂ (Nitrogen)

    • Nitrogen is the most abundant substance in the air.

  • Halogens 

    • F₂

    • P₄

    • S₈

  • These elements are called molecular elements. 

• A valid lewis structure must include lone pairs of electrons.  

• The atoms that have the most single electrons will often be the one in the middle. 

• The octet Rule 

• C₂H₆O

  • These are two chemicals are structural isomers 

  • And they have the same chemical formula but different bonding structure 

• Physical Properties of Molecular Elements & Compounds 

• Physical state varies at SATP (solids can be crystalline or amorphous) 

• Softer than ionic compounds but still brittle 

• Low melting and boiling points 

• Solubility in water varies.  

• Non-electrolytes (except for molecular acids/bases such as HCl and NH₃ which ionize when put in water)

  • Most molecular elements and compounds, the atoms stay together when in water.  

  • Molecular acids will show up as being electrolytes as when they are in water they go through an ionization process which helps conduct electricity.  

• The physical properties of molecular substances depend on the ability of neighboring molecules to attract each other.  

Exam last question: 

Attractive force of neighboring atoms

Covalent bonds of carbon ions in diamond.   

2.3 Chemical Bonding & Electronegativity  

First and Foremost: 

• Assume that all atoms will try to share valence e^– when they bond with each other. I know, this might sound odd at first, but it will help you understand this section better.  

Electronegativity  (EN) is the measure of the ability of an atom to attract shared e^- to itself. 

The general trend is that EN increases as the atomic size decreases. 

Thus: EN increases up a group and from left to right across a period.  

Fr has the lowest EN (0.7)

F has the highest EN (4.0)

The difference in electronegativity, ΔEN, between the two atoms that will be bonding can be used to determine whether the bond between the two atoms will be ionic or covalent. 

If the ΔEN is greater than or equal to 1.7, the bind is classified as ionic. 

E.g. For a bond between Na and Cl, the ΔEN = 3.2 - 0.9 = 2.3. The bond is ionic.  

If the ΔEN is less than 1.7, the bond is classified as covalent.  

E.g. For a bond between C and H, the ΔEN = 2.6 - 2.2 = 0.4. The bond is covalent. 

(Na–Cl:) – The dash represents the introduction of the bond between Na and Cl

Na–Cl: ΔEN_Na–Cl  = EN_Cl - EN_Na

                = 3.2 - 0.9

                = 2.3 (Ionic)

C–H: ΔEN_C–H  = EN_c - EN_H

            = 2.6 - 2.2

            = 0.4 (covalent) 

Not all atoms have the same ability to attract e^– and so if two atoms with different EN values bond covalently, the e^– pair or pairs in question will not be shared equally. The atoms with the greater EN value will attract the shared e^– pair or pairs more. For a situation like this, the covalent bond that results is said to have ionic character because the atoms involved adopt partial negative, δ^–, and partial positive, δ⁺, charges. The covalent bond is called a polar covalent bond. The word “polar” conveys the idea of two opposite ends (in this case, two oppositely charged atoms). The atoms are not real ions since e^– are not transferred over, but since one atom attracts at least one new e^– more than the other atom, partial negative and partial positive ionic charges develop.  

• The atom that attracts the ion more will have the symbol δ^– 

• The atom that attracts the ion least will have the symbol  δ⁺ 

Sudo means almost a full ion.  

• Sudo positive ion 

• Sudo negative ion 

If the ΔEN is at least 0.5, but less than 1.7, the bond is classified as polar covalent.  

So in summary, a bond between atoms A and B will be:

• Ionic, if the ΔEN is ≥ 1.7 

• Polar covalent, if 0.5 ≤ ΔEN < 1.7 

• Nonpolar Covalent, if ΔEN < 0.5 

Note: 

• With the exception of the 0.5 cut-off point, I chose to use the ΔEN guidelines from our textbook for classifying polar covalent and ionic bonds. Many other sources state that a bond will be polar covalent if the ΔEN ≤ 1.7 and ionic if the ΔEN > 1.7.   

Question

• Draw a Lewis Structure for the compound that forms between beryllium and chlorine: 

  • ΔEN_Be–Cl = EN_Cl - EN_Be

               = 3.2 - 1.6

   = 1.6 (polar covalent)

Ionic Compounds 

• binary ionic compound = a compound that - consists of ions of only two elements

• polyatomic ionic compound =  a compound that consists of ions of more than two elements

• oxyanion = a negatively charged polyatomic ion that contains oxygen

2.4 Chemical Formulas and Nomenclature 

The valence of an atom 

• This happens when the element wants to change the amount of electrons it has whether it wants to lose, gain or share the electrons.

• The valence of aluminum is 3 because it wants to lose 3 electrons 

• The valence of sodium is 1 because it wants to lose 1 electron 

• The valence of magnesium is 2 because it wants to lose 2 electrons. 

1. Binary Compound (Metal + Non-metal)

• Na₃P - Sodium phosphide 

• MgO - Magnesium oxide 

• K2S - potassium sulfide 

• AlN - Aluminum nitrate 

• Al₂O₃ - Aluminum oxide 

2. Multivalent Metal + Nonmetal

Common Names for Chemicals:

3. Hydrates → a compound that is able to include water in its lattice crystalline structure

• CuSO₄5H₂O

Desiccant - A desiccant is a substance or chemical that absorbs or attracts moisture from the air, causing a state of dryness (desiccation) in its vicinity.

• On Test 

  • Hydrares: Naming and writing chemical formula 

    • Anhydrous - Literally means “no water”. Substances without water are called Anhydrous. 

    • Hydrated form vs dehydrated from → (Anhydrous)

      • To go from Hydrated to Dehydrated 

        • This means to drive off the water of hydration through heating 

      • Formula used 

        • sodium carbonate heptahydrate

        • The used a chemical called natron 

          • One of the compounds of this is anhydrous sodium carbonate (extracts water from the corpse limiting decomposition) 

          • Stuff highlighted in yellow, in the picture below is important

Molecular Compound:

• When dealing with oxygen drop/change the last letter of the prefixes that have an “o” or  “a”. 

• Never drop “I” → Di and Tri 

• Only drop “O” with oxygen → Mono

• Examples: 

  • CO = Carbon monoxide 

  • CO₂ = Carbon dioxide 

  • N₂O₄ = dinitrogen tetroxide 

  • PCl₃ = Phosphorus trichloride 

  • SF₆ = Sulfur hexafluoride 

Compounds → that will be on the test 

• Binary Compounds (Metal & Nonmetal)

• Multivalent Metal + Nonmetal

• Polyatomic Ionic Compounds

• Hydrates

• Binary Compounds (Nonmetal + Nonmetal)

• Acids

  • Binary Acid (H + Nonmetal)

  • B. H + Polyatomic Ion Without Oxygen

  • 

  • C. Oxyacids or Oxoacids (H + Polyatomic Ion With Oxygen)

• For Binary Acids

  • Add (aq) next to the label 

3.1 Molecular Compounds 

3.3 Polar Bonds and Polar Molecules

• Essentially talks about Polar covalent bonds 1st half. 

• The electron volume is 3 dimensional. 

• Bond polarity: Bond polarity refers to the distribution of electric charge across a chemical bond between two atoms. 

• Molecular polarity: Molecular polarity is a fundamental concept in chemistry that describes the separation of electric charge within a molecule.

• Polar molecule is a molecule in which the uneven distribution of electrons results in a positive charge at one end and a negative charge at the other end. (You can have different molecules and they can be named polar; however, some molecules can be more polar than others.)

  • The measurement that conveys the measure of molecular polarity is called Dipole moment 

• Non-polar molecule is a molecule in which the electrons are equally distributed among the atoms, resulting in no localized charges (t

• The Debye (D) is used to express electric dipole moments of molecules. One Debye is equal to 3.336 10^–30 Coulomb meters.  

• A and C → assume A as Hydrogen and B as Oxygen forming water 

• Hydrogen feels a positive entity as oxygen pulls itso only electron toward it 

• Each hydrogen has the same partial positive charge as Oxygen is pulling on hydrogens with the same force 

Do we have bond polarity?

• Meaning do we have a polar covalent bond?

  • Yes we do! As in both cases the identical bonds in water and both are polar covalent bonds 

  • The Delta EN = 1.2 thus polar covalent bonds 

• Water is also molecularly polar 

Important Term

• Displacement: Displacement is the change in position. 

3.4 Intermolecular Forces

• Intermolecular forces = the attractive forces between molecules 

  • A force between neighboring non-chemically bound entities. 

• Intermolecular forces are generally much weaker than intramolecular forces (ion-dipole forces can be comparably strong), and are responsible for the various physical properties of atomic/molecule

• The only thing two non-polar molecules can do with each other is the London Dispersion Forces (LDF)  (can be attraction or repulsion)

• Dipole-Dipole forces are only for polar molecules exclusively. 

• The Hydrogen bond focuses only on attraction. 

• Substances with stronger intermolecular force networks will:

  • Tend to be solids or liquids at room temperature.  

  • Have higher melting and boiling points. 

  • Be harder as a solid 

  • Have higher surface tension as a liquid.  

  • Be more viscous (although viscosity also depends on the molecular size). 

• Intra - means within. 

• Inter - (between) means that there will be positive or negative relationships between two objects.

  • international = between two countries

  • Interact t

  • Interchange 

  • Intersect 

  • Intercept (when two things collide in a positive or negative aspect)  

  • Intermediate 

  • Interstellar 

  • Intergalactic 

  • Internet 

• Forces → Think of abel pushing(repel) or pulling(attracting). 

• Ionic compounds don’t have intermolecular forces as they don't contain molecules 

• There is space between molecules where gas, solid or liquid.  

• Cohesion vs Adhesion

  • Cohesion - This is when the same molecules can attract each other. 

  • Adhesion - This is when different molecules can attract each other.  

• Intramolecular Force

  • A force relevant to ionic/covalent/metallic bonding. 

• Ionic 

  • Attraction between oppositely charged ions

  • Repulsion between like charged ions

• Covalent 

  • Attraction between two neutral atoms 

  • The nuclei will feel attracted due to the electron glue that is in between. 

  • Repulsion between shared/lone e^– pairs

  • Lone electron pairs can also repel the bond pairs electrons.

• Metallic 

  • Discussed in SCH 4U 

  • Metals pass their valence electrons with each other, this is why they are such good conductors of electricity. 

    • They all technically exist as positive ions. 

• Forces

  • dipole–dipole force an intermolecular force of attraction that forms between the slightly positive end of one polar molecule and the slightly negative end of an adjacent polar molecule

• London dispersion force a weak attractive force acting between all entities, including non-polar molecules and unbonded atoms, caused by the temporary imbalance of electrons within entities