Topic 1: Atoms, Molecules, Stoichiometry and Redox Reactions
1. Relative atomic mass: Average mass of one atom (an isotope in its relative abundance) in
an element relative to 1/12 the mass of a 12-Carbon
2. Relative isotopic mass: Mass of one atom of an isotope of an element relative to 1/12 the
mass of a 12-Carbon
3. Relative molecular mass: Average mass of one molecule of a substance relative to 1/12
the mass of a 12-Carbon
4. Relative formula mass: Average mass of one formula unit of a substance relative to 1/12
the mass of a 12-Carbon
5. Isotopic Abundance: The abundance of each isotope in a mixture
6. One mole of a substance contains the same number of particles as there are in 12g of the
12-Carbon isotope
7. Avogadro’s Constant: L= 6.02 x 1023mol-1
8. Molar mass: The mass of one mole of a substance
9. Stoichiometry: The study of quantitative aspects of chemical formulae and reactions. It
involves both the determination of chemical formulae as well as calculations using
balanced chemical equations based on the central idea of the mole
10. Empirical Formula: Simplest formula showing the ration of the atoms of the different
elements in the compound
11. Avogadro’s Law: Equal volume of gases, under the same temperature and pressure,
contain the same number of molecules
12. Redox reaction: A reaction that involves Reduction and Oxidation simultaneously
13. Reduction: A process whereby a substance gains electrons, resulting in a decrease in
oxidation number
14. Oxidation: A process whereby a substance loses electrons, resulting in an increase in
oxidation number
15. A redox reaction is a reaction in which the same reactant is both oxidised and reduced
16. Disproportionation: A redox reaction in which the same reactant is both oxidised and
reduced
17. Aqueous metathesis: Double displacement reaction
18. Eocell = EoR - EoO
Common Oxidising Agents
Oxidant
MnO4- (purple)
Cr2O72- (orange)
I2 (brown in aq solution)
H2O2 (colourless)
Reaction Medium
Acidic
Alkaline
Acidic
Neutral
Acidic
Product
Mn2+ (colourless)
MnO2 (brown solid/ppt)
Cr3+ (green)
I- (colourless)
H2O
Common Reducing Agents
Reductant
Fe2+ (pale green)
I- (colourless)
S2O32- (colourless)
H2O2 (colourless)
C2O42- (colourless)
Reaction Medium
Acidic
Acidic/Alkaline/Neutral
Neutral
Acidic
Acidic
Product
Fe3+ (yellow)
I2 (brown in aq solution)
S4O62- (colourless)
O2
CO2
Topic 2: Atomic Structure and Physical Properties
1. Angle of deflection ∝
𝒄𝒉𝒂𝒓𝒈𝒆
𝒎𝒂𝒔𝒔
2. Atomic number (Z): The number of protons in the nucleus of each atom of an element
3. Mass number (A): Total number of neutrons and protons present in the nucleus if an atom
of an element
4. Isotopes: Elements of the same atomic number but different number of neutrons. They
have the same electronic configuration and chemical properties but different relative
isotopic masses and physical properties
5. Isoelectronic: Same number of electrons
6. Isotonic: Same number of neutrons
7. Isotopic: Same number of protons
8. Principal Quantum Number: describes the main energy level of an electron and the size of
an atomic orbital
9. Subshells: subdivision of each principal quantum shell
10. Atomic Orbital: region of space with ≥90% probability of finding an electron
11. Aufban’s Principle: Electrons in their ground states occupy orbitals in order of energy
levels. The orbital with the lowest energy is always filled first
12. Hund’s Rule of Multiplicity: When filling subshells that contain more than one orbital with
the same energy level, each orbital must be singly occupied before electrons are paired
13. Pauli’s Exclusion Principle: An orbital cannot hold more than two electrons and the two
electrons sharing the same orbital must have opposite spins
14. Effective nuclear charge: Zeff is the net nuclear charge experienced by an outer electron
(Zeff ≈ Z - S)
15. Shielding effect: Partial offsetting of the charge of the nucleus by inner-shell electrons
16. Electronegativity: The relative ability of an atom in a molecule to attract the shared
electrons in a covalent bond
17. 1st Ionisation Energy: The energy required to remove 1 mole of electrons from 1 mole of
gaseous atoms to form 1 mole of unipositively charged gaseous ions
Topic 3: Chemical Bonding
1. Ionic Bond: The electrostatic forces of attraction between oppositely charged ions in an
ionic compound
2. Lattice Energy: The result of electrostatic interactions between the oppositely charged ions
in an ionic compound/ The energy released when 1 mole of a pure ionic solid is formed
from its constituent gaseous ions
𝒒+ 𝒒−
3. Lattice Energy ∝ |r+ +r−|
4. Metallic Bond: The electrostatic forces of attraction between metal cations and a sea of
delocalised electrons in a metal
5. Covalent Bond: The electrostatic forces of attraction between the nuclei of atoms and
their shared pair of electrons
6. Bond Length (Covalent Bond): The distance between the nuclei of the 2 bonded atoms
when their atomic orbitals overlap to form molecular orbitals
7. Bond Energy: The energy required to break a particular bond in one mole of a gaseous
substance
8. Sigma (σ) Bond: “Head-on” overlap of two atomic orbitals
9. Pi (π) Bond: “Side-on” overlap of two atomic orbitals
10. Valence Shell Electron Pair Repulsion (VSEPR) Theory states:
a. Electron groups around a central atom locate themselves as far away from each
other as possible to minimise electron repulsion
b. LP-LP repulsion > LP-BP repulsion > BP-BP repulsion (as electron density of LP
resides closer to the atom as compared to electron density of BP)
c. Repulsion between EPs depend on the relative electronegativity of central and
bonding atoms
11. Pure Ionic Bond: Complete transfer of electrons, resulting from a large difference in
electronegativity between the 2 atoms
12. Ionic Bond with covalent character: due to the distortion of the electron cloud by cation of
high charge density
13. Polar covalent Bond: Unequal sharing of electrons due to some electronegativity
difference between 2 atoms
14. Covalent Bond: Equal sharing of electrons due to no difference in electronegativity
between 2 atoms
15. Fajan’s Rules: Bonding tends to be ionic if
a. cation is large and anion is small
b. charges on both cation and anion are small
16. Bond polarity: separation of positive and negative charges
17. Non-polar Bond: bonding electrons shared equally due to no differences in
electronegativity
18. Polar Bond: unequal sharing of electrons due to difference n electronegativity of bonding
atoms
19. Dispersion Forces (London forces or instantaneous dipole-induced dipole interactions):
caused by temporary fluctuations in the electron density of an atom or non-polar molecule
20. Permanent dipole-permanent dipole interactions: electrostatic forces of attraction
between polar molecules
21. Hydrogen bonding: occurs between molecules with a H atom bonded to a small, highly
electronegative atom with LP (F, O and N only)
22. Ionic solute – Polar solvent (Ion-dipole interactions)
23. Polar solute – Polar solvent (H bonding and pd-pd interactions)
24. Non-polar solute – non-polar solvent (dispersion forces)
25. Non-polar solute – Polar solvent (immiscible)
Topic 4: The Gaseous State
1
1. Boyle’s Law: V ∝ 𝑝 (Ceteris paribus)
2. Charles’ Law: V ∝ T (Ceteris paribus)
3. Avogadro’s Law: V ∝ n (Ceteris paribus)
4. Ideal gas equations: pV = nRT
5. Ideal gas: a hypothetical gas whose pressure-volume-temperature behaviour can be
completely accounted for by the ideal gas equation pV = nRT
6. R: the proportionality constant/molar gas constant = 8.31 JK-1mol-1
7. Combined gas equation:
𝑝1 𝑉1
𝑇1
=
𝑝2 𝑉2
𝑇2
8. Partial Pressure: Pressure exerted by a particular gas on the sides of a container in a
mixture of gases
9. Dalton’s Law of partial pressure: The total pressure of a mixture of gases is the sum of the
partial pressures of the constituent gases (pTotal = pA + pB +pC)
10. Kinetic Theory of Gases (assumptions applied to an ideal gas):
a. The gas particles occupy negligible volume as compared to the volume of its
container
b. The forces of attraction and repulsion between gas particles are negligible ie. There
are negligible intermolecular forces
c. Gas particles collide elastically ie. There is no loss in kinetic energy after collisions
d. Gas particles are in constant random motion
e. The average kinetic energy of the gas particles is directly proportional to the
absolute temperature
11. At high pressure:
a. There are many gas particles per unit volume; the volume of particles cannot be
considered negligible
b. The molecules are close together, thus intermolecular forces are significant
12. At low temperature: the gas particles have low speeds and average kinetic energy and are
less able to overcome intermolecular forces, which are thus significant
13. Real gases behave most ideally at low pressures and high temperatures
14. Potential deviants:
a. Large molecules – large electron cloud result in stronger intermolecular forces
b. Polar molecules – stronger intermolecular forces of attraction such as pd-pd
interactions or H bonding
15. Maxwell-Boltmann Distribution Curve: As the temperature of the gas increases, the
average speed and kinetic energy of the gas molecules increases
Topic 5: Chemical Energetics and Thermodynamics
1.
Activation Energy: The minimum energy which the reacting particles must possess in
order to overcome the energy barrier before the formation of products
2.
Enthalpy change of reaction (ΔHrѳ): Enthalpy change when molar quantities of reactants
as specified by the chemical equation react to form products at 1 atm and 298K
3.
Enthalpy change of combustion (ΔHcѳ): Heat evolved when 1 mole of a substance is
completely burnt in excess oxygen at 1 atm and 298K
4.
Heat capacity (C): The amount of heat required to raise an object’s temperature by 1K
5.
Specific eat capacity (c): The amount of heat required to raise the temperature of 1g of
the substance by 1K
6.
c of water is 4.18 Jg-1K-1
7.
q = mcΔT
8.
Hess’s Law of Constant Heat Summation: The enthalpy change of a particular reaction is
determined only by the initial and final states of the system regardless of the pathway
taken
9.
Enthalpy change of formation (ΔHfѳ): Enthalpy change when 1 mole of a substance is
formed from its constituent elements in their standard states at 1 atm and 298K
10. Enthalpy change of neutralisation (ΔHneutѳ): Heat evolved when 1 mole of H2O is formed
in the neutralisation reaction between an acid and a base at 1 atm and 298K
11. Enthalpy change of atomization (ΔHatomѳ): Energy required when 1 mole of gaseous
atoms are formed the element at 1 atm and 298K
12. Enthalpy change of hydration (ΔHhydѳ): Heat evolved when 1 mole of free gaseous ions is
surrounded by water molecules at 1 atm and 298K
13. Enthalpy change of solution (ΔHsolѳ): Enthalpy change when 1 mole of solute is
completely dissolved in a solvent to form an infinitely dilute solution at 1 atm and 298K
14. Lattice Energy: Heat evolved when 1 mole of pure solid is formed from its constituent
gaseous ions
15. LE = ΔHhyd - ΔHsol
16. ΔHrѳ = ∑ 𝒏𝚫𝐇𝐜 ѳ (𝒓𝒆𝒂𝒄𝒕𝒂𝒏𝒕𝒔) − ∑ 𝒎𝚫𝐇𝐜 ѳ (𝒑𝒓𝒐𝒅𝒖𝒄𝒕𝒔)
17. ΔHrѳ = ∑ 𝒎𝚫𝐇𝐟 ѳ (𝐩𝐫𝐨𝐝𝐮𝐜𝐭𝐬) − ∑ 𝒏𝚫𝐇𝐟 ѳ (𝒓𝒆𝒂𝒄𝒕𝒂𝒏𝒕𝒔)
18. Bond dissociation energy: The energy required to break 1 mole of a specified covalent
bond in a specified compound in the gaseous state
19. Bond energy: Average energy required to break 1 mole of covalent bond in the gaseous
state
20. Factors affecting bond energy:
a. Bond order
b. Effectiveness of bond overlap
c. Bond polarity
21. 1st Electron Affinity: The enthalpy change when 1 mole of electrons is added to 1 mole of
gaseous atoms to form 1 mole of singly charged gaseous anions
22. Bond Haber Cycle: An extension of Hess’s Law to ionic compounds. It can be used to
determine the LE of ionic compounds
23. Entropy: S is a thermodynamic quantity related to the number of ways the energy of a
system can be dispersed through the motions of its particles
24. 2nd Law of Thermodynamics: The total entropy of the universe always tends to increase
25. ΔS>0: greater disorder/more no. of ways energy can be dispersed in a system
26. ΔS<0: more ordered state/less no. of ways energy can be dispersed in a system
27. Factors affecting entropy of a system:
a. Change in temperature
b. Change in phase
c. Mixing of particles
d. Expansion of a gas
e. Change in number of particles
f. Dissolution of an ionic solution
28. Gibbs free energy: ΔG = ΔH – TΔS
29. ΔG<0: Reaction is feasible and can take place spontaneously (reaction is exergonic)
30. ΔG=0: Reaction is at equilibrium. There is no net change (during melting and boiling)
31. ΔG>0: Reaction is not feasible and cannot take place spontaneously (reaction is
endergonic, and spontaneous if in reverse direction)
ΔH<0
ΔS<0
ΔS>0
Feasible at low
Feasible at all
temperatures
temperatures
ΔH>0
Not feasible
Feasible at high
temperatures
Topic 6: Reaction Kinetics
1.
Collision Theory: For effective collisions to occur, reactant molecules must:
a. Collide with each other
b. Collide with a certain minimum energy called activation energy, Ea
c. Collide with the correct orientation
2.
Activation Energy, Ea: The minimum energy which the reacting particles must possess in
order to overcome the energy barrier before the formation of products
3.
Transition State/ Activated Complex: the arrangement of atomic nuclei and bonding
electrons at the maximum potential energy
4.
Rate of reaction: Change in concentration of reactant or product per unit time
5.
Rate = −
6.
For any reaction: aA + bB → cC +dD, Rate = − 𝑎
7.
Average rate of reaction = −
8.
Instantaneous rate of reaction: the reaction rate at a specified time = gradient of the
𝒅[𝒓𝒆𝒂𝒄𝒕𝒂𝒏𝒕]
𝒅𝒕
𝒐𝒓
𝒅[𝒑𝒓𝒐𝒅𝒖𝒄𝒕]
𝒅𝒕
1 𝑑[𝐴]
= −
1 𝑑[𝐵]
𝑑𝑡
𝑏 𝑑𝑡
𝒇𝒊𝒏𝒂𝒍 [𝒓𝒆𝒂𝒄𝒕𝒂𝒏𝒕]− 𝒊𝒏𝒊𝒕𝒊𝒂𝒍 [𝒓𝒆𝒂𝒄𝒕𝒂𝒏𝒕]
=
1 𝑑[𝐶]
𝑐 𝑑𝑡
=
1 𝑑[𝐷]
𝑑 𝑑𝑡
∆𝒕
tangent of a concentration-time graph at the specified time
9.
Initial rate of reaction: the reaction rate at the start of a reaction when an infinitesimally
small amount of the reactant has reacted = gradient of the tangent of a concentrationtime graph at t=0
10. Rate equation: An experimentally determined equation that relates the rate of reaction
to the concentration of the reactants raised to appropriate powers.
11. Rate = k[A]m[B]n
12. Overall order of reaction = order of reaction wrt A + order of reaction wrt B
13. Rate constant: affected by temperature and presence of catalyst only (k = Ae -Ea/RT)
14. Order of reaction: the power of the concentration terms in the experimentally
determined rate equation
15. Half-life of a reaction: the time taken for the concentration of a reactant to decrease to
half of its initial value
16. Zero-order Reactions: Rate = k[A]0 = k, t½ =
[𝐴]0
2𝑘
17. First-order Reactions: Rate = k[A]1 = k[A], t½ =
𝑙𝑛2
𝑘
1
18. Second-order Reactions: Rate = k[A]2, t½ = 𝑘[𝐴]𝑜
19. Pseudo-zero order Reactions: A particular reactant is present in large excess with respect
to other reactants, thus its concentration with not change significantly during the reaction,
causing the rate to appear to be independent of this reactant concentration
20. A reaction mechanism: A sequence of steps through which reactants react to form
products
21. Reaction intermediate: a chemical species that is produced in an elementary step and
consumed in another
22. Elementary step: a step which cannot be broken down into simpler steps
23. Molecularity: the number of reactant particles taking part in the reaction in that
elementary step
24. Molecularity of an elementary step = overall order of reaction
25. The slowest step is the rate-determining step