S.1

Classifying Matter

Mixtures

• mixtures contain more than one element or compound in no fixed ratio

• air is described as a mixture of gases because the seperate components are interspersed with each other but not chemically combined

  • nitrogen and oxygen have the same properties as they do in pure samples

• homogenous mixture → uniform composition and properties

  • the inter-particle attraction within the different components must be similar in nature to those between the components in the mixture

• heterogenous mixture → non-uniform composition and its properties are not the same throughout

  • interactions between components are different in nature

Solutions

• Solute is dissolved in the solvent

• Filtration

  • solid is seperated from a liquid or a gas using a membrane

  • solid is collected on the membrane as residue, filtrate containing solution passes through

• Distillation

  • seperates solvent from solute

  • solvent has lower boiling point, so it is evaporated and collected as a gas then condensed

• Paper Chromatography

  • small spots of solution containing the samples are placed on the baseline

  • paper is suspended to ensure it is saturated

  • different components have different water affinities so they seperate as the solvent moves up the paper

Kinetic Molecular Theory

• states of matter are characterized by the different energies of the particles

• temperature is directly related to the average kinetic energy of the particles

• solid → gas = sublimation

• gas → solid = deposition

• b-c and d-e: energy is added to break intermolecular bonds, not increase kinetic energy as the average kinetic energy of the particles is sufficient enough already to leave

• Celsius = Kelvin - 273.15, Kelvin = Celsius + 273.15

DIffusion

• process by which particles of a substance spread out more evenly

  • occurs as a result of their random movements

• particles with smaller mass diffuse more quickly

  • Ek=1/2 mv²

The Atom

• Dalton’s model of the atom

  • all matter is composed of tiny indivisible particles called atoms

  • atoms cannot be created or destroyed

  • atoms of the same element are alike in every way

  • atoms of different elements are different

  • atoms can combine together in small numbers to form molecules

• Electrons

  • negative charged particle

• Protons

  • positively charged particle

• Neutron

  • neutral charge, not positive or negative

• Bohr model of the atom

• Atomic/mass number

  • Atomic number: number of protons in the atom

    • no. of protons = no. of electrons if element has no overall charge

  • Mass number: number of protons + number of neutrons

    • electron mass is negligible, so this is regarded as the mass

• Ions

  • positive ion - loses an electron - cation

  • negative ion - gains an electron - anion

• Isotopes

  • different number of neutrons than the periodic table

  • same chemical properties

  • too many or too few neutrons causes radioactivity, making radioisotopes

Mass Spectra

• determines the relative atomic masses of elements from their isotopic composition

• mass spectrometer - measures mass and abundance of isotopes

  • amount of deflection is inversely proportional to mass/charge

  • fragmentation pattern provides insight on the structure of the compound

• Relative atomic mass of an element

  • masses of all elements are in the range 10^-24 - 10^-22, so relative atomic mass is used

• mass spectrum produced by the mass spectrometer using deflection patterns

Emission Spectra

• different elements give out light of distinctive colours

• electromagnetic radiation is emitted in different forms of differing energies

  • all travel at the same speed, but have differing wavelengths (v=λf)

    • c=λf (c=speed of light; 3×10^8) (λ=wavelength) (f=frequency)

• an emission spectrum is produced when an atom moves from a higher to a lower energy level

  • when electromagnetic radiation passes through a collection of atoms, some of the radiation is absorbed and used to excite the atoms from a lower energy level to a higher energy level

  • spectrometer is used to analyze the transmitted radiation relative to the incident radiation and an absorption spectrum is produced

• one packet of energy, a photon, is released for each electron transition

  • change in energy = energy of the photon emitted

  • energy of the photon = hf (h=Planck’s constant, f=frequency)

S2

S2.1: The Ionic Model

• giving and taking of electrons to form electrostatic attraction

• positive ion → cation, negative ion → anion

Transition Metals

• When the ionization energies for the 1st, 2nd, 3rd, etc. are close to each other, the element will be able to lose those valence electrons easily

  • results in variable oxidation states depending on the number of valence electrons in the d and s shells

Compound Names

• polyatomic ions

  • nitrate: NO3-

  • sulfate: SO4²-

  • phosphate: PO4³-

  • hydroxide: OH-

  • hydrocarbonate: HCO3-

  • carbonate: CO3²-

  • ammonium: NH4+

Ionic structures and properties

• giant ionic lattice structures

  • made up of very strong electrostatic forces of attraction

    • ionic bonds or lattice enthalpy

• physical properties depend on lattice structures

Lattice Enthalpy (LE)

• energy needed to seperate into constituent ions (hence △H>0)

• △H is a negative enthalpy change

• LE values are a function of the ionic radii and charge

• △H = (Knm)/(Rm+ + Rx-)

  • K - constant

  • n, m - magnitude of charges

  • Rm+, Rx- - ionic radii

• increase in ionic charge = increase in ionic attraction between ions = increase in lattice enthalpy

• increase in ionic radius of on of the ions = decreased attraction between ions = decreased lattice enthalpy

• lattice enthalpy is greater for ions with a larger charge density as they have a small radius and are highly charged

Properties of ionic compounds

• high melting + boiling point

  • due to high LE

• generally soluble in water (polar) but not in non-polar liquids

• good electrical conductivity in liquid/aqueous states

• generally brittle

  • due to crystalline structure

• low volatitilty

  • volatility - tendency of a substance to vaporize

Ionic character and electronegativity

• ionic character is determined by the formula:

  • %ionic character = △Xp/3.2

    • △Xp=electronegativity difference

• bonding continuum

  • Ionic → △Xp>1.8

  • Polar Covalent → 0<△Xp<1.8

  • Covalent → △Xp<1.8

S2.2: The Covalent Model

• atoms sharing electrons

Octet Rule

• tendency of atoms to gain a valence shell consisting of 8 electrons

  • pairs of electrons not involved in the bond are called lone pairs

• ability of two atoms to form a covalent bond is due to similar strength with which they attract valence electrons

• exceptions to the rule

  • applies to small atoms with less than 8 electrons

    • forms an incomplete octet

    • ex: BeCl2, BF3

Bond Strength

• bond length → distance between 2 bonded nuclei

• bond strength = bond enthalpy → energy required to break the bond

  • as bond length increases, bond enthalpy decreases

Coordination Bond

• bond that is formed by both electrons in pair originating from the same atom

  • ex: H3O+

Valence Shell Electron Pair Repulsion (VSEPR)

• because electron pairs in the same valence shell carry the same charge, they repel each other and so spread themselves as far as possible

• electron pair = electron domain

  • all electron locations in valence shell, all single, double, triple = 1 electron domain

• repulsion applies to electron domains (can be single/double/triple or non-bonding pairs)

• total number of electron domains around central atom determines geometrical arrangement of electron domains

• shape of molecules is determined by angles between bonded atoms

• lone pairs and multiple bonds cause slightly more repulsion

  • lone pairs have a higher concentration of charge (not shared)

  • multiple bonds have a higher concentration of charge (more electrons)

  • non-bonding/lone pair > multiple bond > single bond

Structure

• two electron domains

  • linear shape (bond angle of 180°)

• three electron domains

  • triangular planar (bond angle of 120°) → all single bonds

  • bent/V-shaped (120°, 121°, 118°) → one double bond

  • bent/V-shaped (117° bond angle) → one double bond, one lone pair

• four electron domains

  • tetrahedral (109.5°) → all single bonds

  • trigonal pyramid (107°) → one lone pair

  • bent/V-shaped (104.5°) → two lone pairs

Bond Polarity

• polar bonds - differing electronegativities

  • different pulling strength for electrons

• the more electronegative atom exerts a greater pulling power on the shared electrons → gains more “possesion”

• bond dipole - two partially seperated opposite electric charges

  • the more electronegative atom becomes partially negative

  • the less electronegative atom becomes partially positive

  • increased electronegativity difference = increased bond polarity

• pure covalent bonds → zero electronegativity difference

• polar bonds introduce some ionic nature to the covalent bonds

• Pure Covalent

  • equal sharing of electrons

  • discrete molecules

• Polar Covalent

  • partial/unequal sharing/transfer of electrons

• Ionic

  • complete transfer of electrons

  • lattice of oppositely charged ions

Molecular Polarity

• depends on polar bonds it contains

• depends on molecular geometry

• non-polar

  • dipoles can cancel out, creating non-polar molecules

• polar

  • if the molecule contains bonds of different polarity or the bonds are not symmetrically arranged, dipoles will not cancel out

  • creates a net dipole (turning force in electric field)

• IR active

  • happens only when an overall dipole moment related to the position and vibration of its bonds is found

Covalent Network Structures

• discrete molecules → finite amount of atoms

• covalent networks → no finite number of atoms

  • single repeating pattern of covalent bonds

• allotropes → different bonding/structural patterns of the same element in the same physical state, causing different chemical and physical properties

Allotropes of Carbon

• Diamond

  • structure: sp3 hybridied and covalently bonded to 4 others tetrahedrally arranged in a regular repetitive pattern (angle → 109.5°)

  • non-conductor of electicity, all electrons are bonded

  • very efficient thermal conducter, better than metals

  • highly transparent, lustrous crystal

  • hardest natural substance, brittle, very high melting point

  • uses: polished for jewelry, tools and machinery for grinding/cutting glass

• Graphite

  • structure: sp2 hybridized and covalently bonded to 3 others, forming hexagons in parallel angles with bond angles of 120° (weak London dispersion forces)

  • good electrical conductor; contains one delocalized electron per atom

  • not a good thermal conductor, unless the heat conducts parallel to crystal layers

  • non-lustrous, grey crystalline solid

  • soft and slippery due to sliding layers, brittle, very high melting point, stable mostly

  • uses: dry lubricant, pencils, electrode rods in electrolysis

• Graphene

  • sp2 hybridized and covalently bonded to 3 others, forming hexagons with bond angles of 120° (single layer → 2D only) honeycomb/chicken wire

  • very good electrical conductor, one delocalized electron per atom

  • best thermal conductivity known

  • almost completely transperent

  • thickness of just one atom (2D) → thinnest material to ever exist, 100x stronger than steel (strongest), very flexible, very high melting point

  • uses: transmission electron microscopy (TEM) grids, photovoltaic cells, touchscreens, high performance electronic devices, etc.

• Fullerene (C60)

  • sp2 hybridized, bonded in a sphere of 60 carbon atoms, consisting of 12 pentagons and 20 hexagons (closed spherical cage)

  • poor conductors of electricity, delocalized electron has little movement

  • very low thermal conductivity

  • black powder

  • very light and strong, reacts with potassium (K) to make superconducting crystalline material, low melting point

  • uses: lubricants, medical, industrial devices for binding specific target molecules; related forms are used to make nanotubes/nanobuds used as capacitators in the electronics industry, and catalysts

London Dispersion Forces

• non-polar molecules do not have a permanent dipole

  • electrons behave somewhat like clouds of negative charge, density of the cloud could be greater over one atom at any moment

  • when there is a differing density, the bond will have seperation of charge, creating a weak dipole (temporary/instantaneous dipole)

• creates weak forces of attraction that occur between opposite ends of two temporary dipoles

  • weakest form of intermolecular force

  • strength increases with molecular size (greater number of electrons)

• LDF is the only force that exists for non-polar molecules

• also exists in polar molecules, but is often overlooked for stronger forces

Dipole-dipole attraction

• polar molecules have permanent seperation of charge (electronegativity difference)

  • known as a permanent dipole

  • opposite charges on neighbouring molecules attracting each other

• strength varies on distance and relative orientation of the dipoles

Dipole-induced dipole attraction

• occurs in mixtures with both polar and non-polar molecules

• the permanent dipole from a polar molecule creates a temporary seperation of charge in the non-polar

• act in addition to LDF (non-polar) and dipole-dipole attraction (polar)

• van der Waal’s force: all 3 forces added together

Hydrogen Bonding

• when a molecule contains hydrogen covalently bonded to fluorine, oxygen, or nitrogen (electronegative atoms)

  • particular case of dipole-dipole attraction

• the large electronegativity difference between hydrogen and the bonded atoms results in the electron pair being pulled away from hydrogen

• hydrogen now exerts a strong attractive force on a lone pair in the electronegative atom due to its small size and the absence of other electrons to shield the nucleus

• strongest form of intermolecular force

Melting and Boiling Point

• changing state = breaking intermolecular forces

• covalent substances generally have lower MP and BP than ionic substances

  • relatively weak intermolecular forces < electrostatic attraction

  • covalent substances are generally liquid/gas at room temperature

• strength of intermolecular forces increase with molecular size and extent of polarity

Solubility

• non-polar substances are generally dissolvable in non-polar solvents by formation of LDF between solute and solvent

• polar covalent compounds can generally dissolve in water (highly polar H2O) through dipole interactions and hydrogen bonding

• solubility of polar compounds is reduced in larger molecules

  • polar bonds only a small part of the structure

  • non-polar parts reduce solubility

• inability of non-polar groups to associate with water means non-polar substances do not dissolve well in water

• polar substances have low solubility in non-polar solvents

  • they remain together due to dipole-dipole attractions

• giant molecular are generally insoluble in all solvents

  • too much energy required to break the strong covalent bonds

Electrical Conductivity

• covalent compounds do not contain ions, so they cannot conduct electricity in the solid or liquid state

  • some polar covalent molecules (when they can ionize) will conduct electricity

Resonance Structures

• delocalization - tendency to be shared between more than bonding position

  • delocalized electrons spread out → greater stability for molecule/ion

• delocalization occurs when there is more than one position for a double bond within a molecule

  • two equally valid positions for a double bond

  • expected is 1 single and 1 double bond, in reality is is 2 equal bonds, intermediate in length and strength

• resonance - molecule is a combination of two Lewis formulas

  • electrons from the double bond delocalize and spread themselves equally between both possible bonding positions

    • shown with a dotted line

  • known as a resonance hybrid

• resonance influences bond strengths/lengths which in turn can influence reactivity

Benzene (C6H6)

• six carbon atoms arranged in a hexagonal ring, each bonded to a hydrogen atom in a triangular planar arrangement with bond angle 120°

• true form of benzene is the resonance hybrid

  • circle represents equally spread delocalized electrons

• 1-1 ratio of carbon to hydrogen indicates a high degree of unsaturation, greater than that of alkenes or alkynes

  • does not show characteristic properties

  • no isomers, reluctant to undergo additional reactions

Expanded Octet

• when the central atom is period 3 or lower, sometimes there are more than 8 electrons around the central atom

• d orbitals available in the valence shell have energy values similar to those of the p orbitals

  • promotion of electrons (3p→3d) allows additional electron pairs to form

• causes some elements to expand their octets (5-6 electron domains)

Species with five electron domains

• triangular bipyramidal shape → angles of 90° and 120°

• if one or more domains are non-bonding electrons, they will repel the most

  • one lone pair gives an unsymmetrical tetrahedron or see-saw shape (bond angles <120° and <90°)

• to minimize additional repulsion and bonding domains, the lone pair must be located in an equatorial position (horizontal plane around central atom) instead of an axial osition (above/below horizontal plane)

  • two lone pairs give a T-shaped structure (bond angles <90°)

  • three lone pairs give a linear shape (bond angle 180°)

Species with six electron domains

• octahedral shape with angles of 90°

  • no lone pairs of electrons → symmetrical octahedral shape

  • one lone pair → square pyramidal shape (bond angles slightly less than 90°)

  • two lone pairs → square planar shape (bond angles of 90°)

    • maximizes distance apart by arranging pairs on opposite sides

Formal Charge

• formal charge used to predict a preferred Lewis formula

• treats covalent bonds as if they were purely covalent with equal electron distribution

• FC = V - (1/2 B + L)

  • V = valence, B = bonding, L = lone (number of electrons)

• low FC means less charge transfer has taken place in forming a structure from its atoms

  • generally means most stable → preferred structure

• sum of formal charges for a species must be equal to the charge

Sigma Bond

• when two atomic orbitals combine head-on along the bond axis (imaginary line)

  • overlap of s, p, and hybrid orbitals in different combinations

  • always the bond that forms in a single covalent bond

• electron density is concentrated between the nuclei of the bonded atoms

Pi Bond

• when two p orbitals collide laterally (sideways-on)

• electron density is concentrated above and below the plane of the bond axis

• only forms within double and triple bonds

• weaker than sigma bonds as electron density is further from nucleus

Sigma and Pi Bonds

• s+s → sigma

• s+p → sigma

• p+p (head-on) → sigma

• hybrid + s → sigma

• hybrid + hybrid → sigma

• p+p (laterally) → pi

• single bond - 1 sigma

• double bond - 1 sigma + 1 pi

• triple bond - 1 sigma + 2 pi

Hybridization

• formation of covalent bonds often starts with excitation of the atoms

  • amount of energy put in to achieve this is more than compensated by the extra energy released on forming bonds

• if different orbitals are used in forming covalent bonds, unequal bonds are expected

  • instead, unequal atomic orbitals within an atom mix to form new hybrid atomic orbitals which are identical but different from the original bonds

• hybrid orbirtals have different energies, shapes, and orientation in space from their parent orbitals

  • allows them to form stronger bonds by allowing for greater overlap

• sp³ orbitals

  • 1 s orbital and 2 p orbitals produce 4 sp³ orbitals

  • shape and energy have properties of s and p, but more like p than s

• sp² orbitals

  • 1 s orbital and 2p orbitals produce 3 sp² orbitals

• sp orbitals

  • 1 s orbital and 1 p orbital produce 2 sp orbitals

Carbon → Hybridization

• C: atomic number = 6 (1s²2s²2px12py1)

  • forms 4 covalent bonds, but only has two singly occupied bonding electrons

• excitation occurs (2s → 2p) to change from ground state

• sp³ hybridization

  • orbitals orientate themselves at 109.5°, forming a tetrahedron

  • each hybrid orbital overlaps with an atomic orbital → 4 sigma bonds

• sp² hybridization

  • when carbon forms a double bond

  • orientate themselves at 120°, forming a triangular planar

  • each hybrid orbital overlaps with a neigbouring atomic orbital → 3 sigma bonds

  • as the 2 carbon atoms approach each other, the p orbitals in each atom that did not hybridize overlap sideways

    • forms a pi bond

    • double bond (C2) → 1 sigma, 1 pi

    • characteristic lobes of electron density above and below the bond axis

• sp hybridization

  • orientate themselves at 180°, giving a linear shape

  • overlap of the two hybrid orbitals with other atomic orbitals → 2 sigma bonds

  • when carbon forms a triple bond

  • C2H2

    • each carbon atom has 2 unhybridized p orbitals that are orientated 90° to each other

      • combines to form 2 pi bonds

    • four lobes of electron density turns into a cylinder of negative charge around the atom, making the molecule susceptible to electrophilic reactants (attracted to electron-dense regions)

Hybridization and Molecular Geometry

• tetrahedral → sp³

• triangular planar → sp²

• linear → sp

• lone pairs can also be used in hybridization

  • non-bonding pairs can also hybridize

    • ex: NH3 → lone pair in N resides in the sp³ orbital

Hybridization and Benzene (C6H6)

• each of the 6 carbon atoms are sp² hybridized

  • forms 3 sigma bonds (120°) → planar shape

  • leaves the unhybridized p electron on each carbon atom

    • dumbbell shape perpendicular to the plane of the ring

    • do not form pi bonds but effectively overlap in both directions

    • spreads themselves evenly to be shared by all 6 carbon atoms

  • forms a delocalized pi electron cloud

    • electron density is concentrated in 2 donut-shaped rings above and below the plane

2.3: The Metallic Model

Metallic Bonding

• metals: low ionization energies so they react by losing valence electrons forming a positive ion

  • metallic character: loss of control over outer shell electrons

• when there is no other element present to accept the electrons and form an ionic compound, the outer electrons are held loosely by the nucleus so they ‘wander off’

  • delocalized electrons

• metal atoms form a regular lattice structure through which electrons move freely

  • metallic bonding: force of electrostatic attraction between lattice of cations and delocalized electrons

Uses of metals

• Good electrical conductivity

  • because of highly mobile delocalized electrons

    • used for electrical circuits (copper)

• Good thermal conductivity

  • because of delocalized electrons and closely packed ions

    • used for pots and pans for cooking

• Malleable (can be shaped under pressure)

  • because of the lack of direction in the movement of delocalized electrons

    • used for machinery

• Ductile (can be drawn out into threads)

  • because the metallic bond remaining intact while formation changes

    • used for electric wires and cables

• High melting points

  • because of strong electrostatic forces

    • used for high-speed tools

• Shiny, lustrous appearance

  • because delocalized electrons in metal crystal structure reflect light

    • used for jewelry

• non-directional nature of metallic bonding allows metals to mix with other metals or non-metals in the molten state

  • resulting mixture is an alloy

    • enhances properties of the metallic structure

Metallic bond strength

• determined by

  • number of delocalized electrons

  • charge of the cation

  • radius of the cation

• the greater the electron density and the smaller the cation, the greater the electrostatic attraction

• Across a period

  • increasing melting point

    • greater attraction between ions and delocalized electrons

  • lower degree of reactivity

• Down a group

  • decreasing melting point

    • weaker attraction between ions and delocalized electrons

  • higher degree of reactivity

Transition elements

• elements with an incomplete d-sublevel OR elements that can give rise to cations with an incomplete d-sublevel

• proximity in energy between outer occupied sublevels enables them to delocalize large amounts of d-electrons to form metallic bonds

Transition elements: High melting point

• metals have a large amount of delocalized electrons and a large positive charge on the metal cations which leads to strong metallic bonding → high melting points

• transition metal trends are less evident due to ability of transition elements to delocalize large numbers of electrons and the similar ionic radii

• difficult to predict trends accurately compared to the s-block metals

Transition elements: High electrical conductivity

• metals have a large amount of delocalized electrons which increases their conductivity

  • example: copper (Cu) is used in wires

2.4: Models to Materials

Bonding Triangle

• bonding seen as a continuum (ionic, covalent, metallic) → different bonding types are present to different degrees

  • position on triangle determined from electronegativity

• high electronegativity difference = ionic

• low electronegativity difference = covalent or metallic

• intermediate electronegativity difference = polar covalent

Composite Materials

• mixture between two or more different materials

  • materials have seperate phases (different positions on bonding triangle)

• mixture retains properties of individual materials that compose it

  • example: fibreglass, concrete

Alloys

• produced by adding one metal element to another metal (or carbon) in liquid/molten state so the different atoms can mix

  • in solid, ions of the different metals are scattered through the lattice

    • forms a structure of uniform composition

• metallic bonds are present → delocalized electrons bind the lattice

  • possible due to the non-directional nature of the delocalized electrons and accomodation of the lattice to different sizes of ions

• alloys have properties distinct of component elements (different packing of cations in lattice)

• pure metal → regular arrangement of atoms

  • interrupted in an alloy by different cations

    • more difficult for atoms to ‘slip over each other’ → stronger

• alloy is stronger, more chemically stable, and more resistant to corrosion

Polymers

• monomers (small molecules) are able to react together to form a linked chain held together by covalent bonds, forming a polymer

• polymers are macromolecules → composed of thousands of atoms and so are relatively large compared with other molecules

• nature/properties of a polymer vary with the monomer, length, and amount of branching

• structure is shown as a repeating unit with open bonds on each end

• natural polymers - found naturally (example: protein, starch, DNA)

• synthetic polymers - human-made and non-biodegradable (example: plastics)

Addition Polymers

• addition reaction - a multiple bond in a molecule breaks and creates new bonding positions

  • alkenes/alkynes have double/triple carbon-carbon bonds respectively so they readily undergo addition reactions

    • they can act as monomers and form addition polymers

• %atom economy = molar mass of desired product / molar mass of all reactants x 100

  • addition polymerization reactions do not generate a by-product so it has a 100% atom economy

Condensation Polymers

• condensation reaction - two functional groups react to form a new covalent bond with the release of a small molecule (H2O, HCl, NH3, etc.)

  • A-OH + H-B → A-B + H2O

• to form condensation polymers, monomers must have functional groups (active ends)

  • allows them to form new covalent bonds with neighbours on both sides

  • the functional groups in neighbouring molecules must be able to react together

Carboxylic acid + alcohol → polyester (ester link)

• when one monomer has two carboxylic acid groups (COOH) and the other has two alcohol groups (OH), an ester link forms between them

  • chain extends in both directions → polyester

Carboxylic acid + amine → polyamide (amide link)

• when one monomer has two carboxylic acid groups (COOH) and another monomer has two amine groups (NH2), an amide link forms between them

  • forms a polymer known as polyamide

S3

3.1: Classification of Elements

(see 3.2 Periodic Trends old syllabus for more detail)

Atomic Radius

• increases down a group and decreases across a period

  • down a group: number of occupied electron levels increase

  • across a period: number of occupied electron levels stay the same but the number of protons increase, increasing the nucleus’ force of attraction to the outer electrons

Ionic Radius

• positive ions are larger than parent atoms due to loss of outer energy level (valence)

• negative ions are smaller than parent atoms due to addition of electrons

  • increased electron repulsion causes electrons to move

• increase in nuclear charge (number of protons) causes ionic radius to decrease

  • increased attraction between outer electrons and nucleus

• ionic radius increases down a group due to increased amount of occupied energy levels

Ionization Energy

• increases across a period and decreases down a group

Electron Affinity

• decreases down a group and increases across a period

Electronegativity

• decreases down a group and increases across a period

  • across → increase in nuclear charge increases attraction between nucleus and bond electrons

  • down → increases distance between nucleus and bond electrons so reduced attraction

Group 1: Alkali metals

• physical properties

  • good conductors of electricity and heat (mobility of outer electrons)

  • low density

  • shiny grey surfaces when freshly cut with a knife

• chemical properties

  • very reactive metals

  • forms ionic compounds with non-metals

  • forms single charged ions (X+)

  • reactivity increases down group (lower IE)

• reaction with water: forms hydrogen and metal hydroxide

  • lithium: floats and reacts slowly (releases hydrogen but keeps shape)

  • sodium: reacts vigorously (heat produced melts the unreacted metal)

  • potassium: reacts more vigorously (heat produced ignites hydrogen)

Group 17: Halogens

• physical properties

  • coloured

  • gradual change from gases (F2, Cl2) to liquid (Br) and solid (I2, At2)

• chemical properties

  • very reactive non-metals

  • reactivity decreases down group (lower attraction)

  • form ionic compounds with metals and covalent compounds with non-metals

• displacement reactions

  • the more reactive halogen displaces the less reactive halogen

• halides

  • halogens produce insoluble salts with silver forming precipitates

Period 3 Oxides

• ionic character of period 3 oxides decrease from left to right

  • electronegativity value approaches oxygen, so the difference is less

• ionic oxides

  • dissolve in water to form alkaline solutions

  • reacts with acid to form a salt and water

• non-metallic oxides

  • reacts with water to form acidic solutions

• amphoteric oxides

  • essentially insoluble (does not affect pH when added to water)

  • shows both basic and acidic behaviour

Acid Rain

• produced by non-metal oxides

• sulfur oxides

  • S(s) + O2(g) → SO2(g) sulfur dioxide

  • H2O(l) + SO2(g) → H2SO3(aq) dissolve in rainwater

  • 2SO2(g) + O2(g) → 2SO3(g) sulfur trioxide

  • H2O(l) + SO3(g) → H2SO4(aq) dissolve in rainwater (acid)

• nitrogen oxides

  • N2(g) + O2(g) → 2NO(g) nitrogen monoxide

  • N2(g) + 2O2(g) → 2NO2(g) and 2NO(g) + O2(g) → 2NO2(g) nitrogen dioxide

  • H2O(l) + 2NO(g) → HNO2(aq) + HNO3(aq) dissolve in rainwater

  • 2H2O(l) + 4NO2(g) + O2(g) → 4HNO3(aq) oxidized

Oxidiation States

• oxidation is:

  • addition of oxygen

  • removal of hydrogen

  • electron loss

  • an increase in oxidation state

• rules to assign oxidation states:

  1. atoms in the free (uncombined) element have an oxidation state of zero

  2. in simple ions, the oxidation state is the same as charge of the ion

  3. oxidation states of all atoms in a neutral compound must add up to zero

  4. oxidation states of all atoms in a polyatomic ion must add up to the charge

  5. usual oxidation state for an element in a compound is the one most commonly found

  6. F (fluorine) has oxidation state of -1 all the time (most electronegative)

  7. O (oxygen) has oxidation state of +2 except in peroxides

  8. Cl (chlorine) has oxidation state of -1 except when bonded to more electronegative ions

  9. H (hydrogen) has oxidation state of +1 except when forming ionic hydrides

  10. oxidation state of a transition metal in a complex ion can be found using the charge on the ligands

Transition Metals

• metals in the d-block have similar physical and chemical properties

• zinc is not a transition metal

  • has a full d sublevel in both species

• physical properties

  • high electrical and thermal conductivity

  • high melting point

  • high tensile strength

  • malleable and ductile

• chemical properties

  • forms compounds with more than one oxidation state

  • form a variety of complex ions

  • form coloured compounds

  • acts as catalysts when either elements or compounds

• magnetic properties

  • only found in iron, nickel, and cobalt

    • due to presence of unpaired electron

    • every spinning electron can act as a magnet

Variable Oxidation States

• transition metals display a wide range of oxidation states

• all transition metals show both +2 and +3 oxidation states

• maximum oxidation states increase in steps of +1 and reaches a maximum at manganese then decreases in steps of -1

• oxidation states above +3 generally show covalent character

• compounds with higher oxidation states tend to be oxidizing agents

Coloured Compounds

• transition metal ions in solution have a high charge density

  • attracts water molecules which form coordination bonds with the positive ions

• complex ions are formed when a central ion is surrounded by molecules/ions that possess at least one lone pair of electrons (ligands)

  • number of coordination bonds from the ligands to the central ion is called the coordination number

• colours appear because of a spurt in the d-orbital’s energy levels

  • when light passes through, it excites electrons and increases the energy level for these electrons

  • the ions absorb some colours and reflects the ones opposite it

  • when the energy (light) is absorbed, the d-orbitals split into two levels

3.2 Functional Groups: Classification of Organic Compounds

• empirical formula → simplest whole-number ratio of atoms

• molecular formula → actual number of atoms

• full structural formula → shows every bond and atom at 90°/180°

• condensed structural formula → omits bonds, displays minimal information

• skeletal formula → shorthand representation of a structural formula

• aromatic compounds: molecules which contain a benzene ring

• catenation: ability of carbon to link to itself and form chains/rings

Functional Groups

• atoms/groups of atoms that are present in organic compounds and are responsible for a compound’s physical properties and chemical reactivity

• *halogen atoms are regarded as substituents as they have taken the position of a hydrogen atom

  • IUPAC names example: chloroethane, 2-bromopropane, etc.

• **syllabus does not require knowledge of arenes as compounds but expects you to recognize the phenyl group when it is present in a structure

  • naming is also not required

Functional Groups: chemical reactivity

• reaction pathway → several reactions to produce a target compound

  • product of one reaction is reactant of the next

  • ex: ethane (CH2) → ethanoic acid (CH3COOH)

Amino acids: condensation reaction to form link

• amino acids contain 2 functional groups:

  • amine (-NH2) and carboxylic acid (-COOH)

• amino acids react together via condensation reaction

  • molecule of water eliminated, acid and amino groups form new bond

    • bond is a substituted amide link (peptide bond) forming a dipeptide

• the dipeptide still has a functional group (-NH2, -COOH)

  • can perform a condensation reaction again, forming a tripeptide and eventually a chain of many linked amino acids (polypeptide)

Homologous Series

• organic compounds are classified into ‘families’ of compounds

• successive members of a homologous series always differ by CH2

  • ex: C2H6, C3H8, C4H10 → alkanes

• members of a homologous series can be represented by the same general formula

• members of a homologous series show a trend in physical properties

  • because they differ by CH2, carbon chains get progressively longer

    • so higher B.P. for example

• longer chians = increased London dispersion forces

Functional groups: physical properties

• most volatile → least volatile

  • alkane > halogenoalkane > aldehyde > ketone > alcohol > carboxylic acid

  • London dispersion forces → dipole-dipole interaction → hydrogen bonding

  • increasing strength of intermolecular attraction →

    • increasing boiling point →

• chain length and functional groups affect intermolecular forces

  • polar functional groups = dipole-dopole or hydrogen bonding

IUPAC Naming

1. Identify the longest straight chian of carbon atoms

   • 1=meth-, 2=eth-, 3=prop-, 4=but-, 5=pent-, 6=hex-, etc.

2. Identify the functional group

   • numbered → nuimber has to be the smallest value possible

3. Identify the side chains or substituent groups

   • halogenoalkane (-F, -Cl, -Br, -I), amine (-NH2)

Esters and Ethers

• esters → form when the alkyl group of an alcohol replaces the hydrogen of a carboxylic acid in a condensation reaction

  • R-COOH + R’OH → R-COO-R’ + H2O

  • the stem comes from the parent acid but the alkyl group of the alcohol is the prefix

    • ex: ethanol + ethanoic acid → ethylethanoate

• ethers → 2 alkyl chains linked by an oxygen atom

  • R-O-R’

  • the longer chain will be the stem and retains its alkane name

  • the shorter chain is regarded as a substituent and is given the prefix alkoxy

    • ex: methoxypropane, ethoxyethane

• prefix - stem - suffix

  • prefix - position, number, and name of substituents

  • stem - number of carbon atoms in longest chain

  • suffix - class of compound determined by functional group

Structural Isomers

• same molecular formula but different arrangements of the atoms

• each isomer is a distinct compound

  • unique physical and chemical properties

• the more branching that is present in an isomer, the lower its boiling point

  • reduced surface contact weakens London dispersion forces

• primary carbon → attached to functional group and at least 2 hydrogen atoms

• seconday carbon → attached to functional group, one hydrogen atom, and 2 alkyl groups

• tertiary carbon → attached to functional group and 3 alkyl groups

Stereoisomers

• atoms are attached in the same order but differing in spatial or 3D arrangements → requires 3D representation

  • Isomerism - structural + stereo

    • Stereo - configurational + conformational

      • Configurational - cis-trans + optical

Conformational Isomers (not needed)

• spontaneously interconnect through bond rotations and so cannot be isolated seperately (usually)

  • some conformers of a compound may be more stable than others so are favoured → influences reactivity of the compound

Configurational isomers

• permanent difference in geometry

  • cannot be interconverted and exist as seperate compounds with some distinct properties

Cis-trans isomers

• double-bonded molecules

  • consists of one sigma and one pi bond (pi bond forming by sideways overlap of two p orbitals)

  • free rotation around this double bond is not possible

    • would push p orbitals out of position and pi bond breaks

• when the molecule contains two or more different groups attached to the double bond, these can be arranged to give 2 different isomers

  • cis → same side, trans → opposite side s

• cyclic molecules

  • cycloalkanes contain a ring of carbon atoms that restricts rotation

    • bond angles are strained from the tetrahedral angles in parent alkane

Optical isomers

• chiral - carbon atom attaches to 4 different atoms/groups

  • also known as asymmetric or stereocentre

• the four groups arranged tetrahedrally with bond angles of 109.5° can be arranged in 2 different 3D configurations which are mirror images

  • known as enantiomers → chiral molecules

    • have opposite configurations at each chiral center

• diastereoisomers → have opposite configurations at one or more (but not all) chiral centers

  • not mirror images of each other

Properties of enantiomers

• optical activity → interaction with light

  • when a beam of plane-polarized light passes through a solution of optical isomers, they rotate the plane of polarization

  • optically active → seperate solutions of enantiomers (at the same concentration) rotate plane-polarized light in equal amounts but opposite directions

    • racemic mixture → chiral compound with equal concentration of 2 optical isomers

  • two optical isomers’ rotations cancel out, so racemic mixtures are optically inactive

• reactivity with other chiral molecules

  • when a racemic mixture is reacted with a single enantiomer of another chiral compound, the two components of the mixture (+ and -) react to produce different products

    • products have distinct chemical and physical properties so can be seperated

    • resolution → two enantiomers seperated from a racemic mixture

Mass Spectrometry

• used to find mass of individual atoms and finding relative abundances of different isotopes → also finds relative molecular mass of a compound

Fragmentation patterns

• ionization process → shooting an electron from electron gun then hitting the incident species and removing an electron

  • X(g) + e- → X+(g) + 2e-

    • X is a molecule

  • collision can be so energetic the molecule breaks into different fragments

• fragmentation pattern is used as evidence to find the structure of a compound

  • peak (largest mass/charge) is molecular ion that passed without fragmenting

Infrared Spectroscopy

• frequency of radiation is often measured as number of waves per centimeter (wavenumber)

• radio waves can be absorbed by certain nuclei, reversing their nuclear spin (environment)

  • used in nuclear magnetic resonance (NMR)

• microwaves cause molecules to increase their rotational energy (bond lengths)

• infrared radiation is absorbed by certain bonds causing them to stretch/bend (bonds)

• visible/ultraviolet light can produce electronic transitions (electronic energy levels)

• x-rays are produced when electrons make transitions between inner energy levels

  • produce diffraction patterns (molecular/crystal structure)

Natural frequency of a chemical bond

• chemical bonds are like springs/rulers

  • each bond vibrates and bends at a natural frequency

    • depends on strengths and atom masses

  • light atoms vibrate at higher frequencies (less weight)

  • multiple (stronger) bonds vibrate at higher frequencies

• simple diatomic molecules can only vibrate when the bond stretches

Exciting bonds

• energy needed to excite bonds occur in the infrared (IR) region

• only polar diatomic molecule bonds will interact with IR radiation

  • presence of positive and negative charge allows the electric field component of the IR radiation to excite the vibrational energy

  • change in vibrational energy produces change in dipole moment

    • intensity of absorption depends on polarity

Stretching and Bending

• in a polyatomic molecule (like water), it is more correct to consider the molecule stretching and bending as a whole, rather than considering individual bonds

  • ex: water can vibrate at 3 fundamental frequencies

    • symmetric stretch, asymmetric stretch, symmetric bend

  • each of the modes of vibration results in a change of dipole in the molecule

    • can be detected with IR spectroscopy

• for a symmetrical linear molecule (like carbon dioxide), there are 4 modes of vibration

  • symmetric stretch is IR inactive: no change in dipole moment

    • dipoles of both C=O bonds are equal and opposite throughout the interaction

Greenhouse Gases

• greenhouse effect: solar radiation passes through the atmosphere and warms the surface of the Earth. The surface radiates some of this energy as longer wavelength infrared radiation which is absorbed by greenhouse molecules which makes the air warmer, causing it to radiate heat. Some of this radiation is re-radiated back to the Earth’s surface and some back to space

• the ability of a molecule to absorb infrared radiation depends on the change in dipole moment that occurs as it vibrates

  • some greenhouse gases are much more effective than others in absorbing IR radiation

• global warming potential: amount of infrared radiation that one tonne of a gas would absorb compared to the amount that would be absorbed by one tonne of carbon dioxide

  • depends on effectiveness and atmospheric lifetime of the gas

Wavenumbers

• absorption of certain wavenumbers of IR radiation helps to identify bonds in a molecule

• some bonds can be identified by shapes of their signal

  • ex: O-H bond is broad, C=O is sharp

• hydrogen bonding broadens IR absorption so can be detected

  • ex: O-H in carboxylic acids have broader absorption

• molecules with several bonds can vibrate in many different ways and with different frequencies

  • complex pattern can be used as a ‘fingerprint’ to be matched

  • comparison of spectrum of a sample with a pure compound can be used as a test of purity

Nuclear Magnetic Resonance Spectroscopy (NMR)

• nuclei of atoms with an odd number of nucleons (H, C, F) have a property called nuclear spin and behave like tiny bar magnets

  • when placed in an external magnetic field, these nuclei can exist in two distinct energy levels

    • depending on whether magnetic field is aligned with/opposed to the external magnetic field

    • energy gap between the energy levels is very small and only requires absorption of low-energy radio waves to close the gap between energy levels

• as electrons shield nucleus from full effects of external magnetic field, differences in electron distribution produce different energy seperations between the two spin energy levels

  • nuclei in different chemical environments produce different signals

    • proton = hydroegn because hydrogen has 1 proton

Magnetic Resonance Imaging (MRI)

• application of NMR spectroscopy

  • uses H’s magnetic moment

• with a powerful magnet, radio waves are used to generate an electronic signal that can be decoded to produce images

• useful in diagnosis of living tissue due to hydrogen in water

H NMR Spectroscopy

• NMR provides:

  1. number of signals in the spectrum

  2. position/chemical shifts of each signal

  3. size/integrated area of each signal

  4. splitting pattern observed for each signal

     • gives information on chemical environments and therefore structure

Chemical environments

• hydrogen nuclei (protons) that have the same chemical environment are said to be equivalent as they give the same signal in NMR

  • number of signals observed therefore depends on number of chemical environments

Chemical shifts

• position where a signal appears in NMR spectrum is measured in terms of chemical shift which has units of parts per million (ppm)

• the closer a hydrogen atom is to an electronegative atom, the more pronounced the electron-withdrawing effect and the higher chemical shift observed

  • the high electronegativity effectively pulls electrons away from the hydrogen atoms thus deshielding the hydrogens’ nuclei

  • nuclei are now more susceptible to effects due to external magnetic field

• hydrogen nuclei in particular environments have characteristic chemical shifts

  • found in section 21 of data booklet

Splitting patterns

• individual signals in NMR do not consist of a single peak

  • signals are split/resolved into distinctive patterns

• splitting occurs as the effective magnetic field experienced by particular nuclei is modified by the magnetic field produced by neighbouring protons

  • spin-spin coupling

• the number and intensity of lines produced are easily predicted

  • based on the number of neighbouring hydrogens involved in coupling

• number of lines: n+1 → n=number of hydrogen atoms on the neighbouring carbons

• intensity: pascal’s triangle

• pattern will continue for each additional proton on neighbouring carbons

• protons bonded to the same atom do not interact as they are equivalent and behave as a group

• protons on carbon atoms not adjacent to each other do not generally interact as they are too far apart for their magnetic fields to interact

• alcohol protons (OH) typically do nto engage in spin-spin coupling

  • signals for OH protons are not split and appear as singlets

  • OH protons are not counted when applying the n+1 rule