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CHEMISTRY DEFINITIONS
Topic 1: Atoms, Molecules, Stoichiometry and Redox Reactions
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Relative isotopic mass: Mass of one atom of an isotope of an element relative to 1/12 the
mass of a 12-Carbon
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Relative formula mass: Average mass of one formula unit of a substance relative to 1/12
the mass of a 12-Carbon
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One mole of a substance contains the same number of particles as there are in 12g of the
12-Carbon isotope
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02 x 1023mol-1
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Stoichiometry: The study of quantitative aspects of chemical formulae and reactions
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Empirical Formula: Simplest formula showing the ration of the atoms of the different
elements in the compound
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Redox reaction: A reaction that involves Reduction and Oxidation simultaneously
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Oxidation: A process whereby a substance loses electrons, resulting in an increase in
oxidation number
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Disproportionation: A redox reaction in which the same reactant is both oxidised and
reduced
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Eocell = EoR - EoO
Common Oxidising Agents
Oxidant
MnO4- (purple)
Cr2O72- (orange)
I2 (brown in aq solution)

Reaction Medium
Acidic
Alkaline
Acidic
Neutral

Product
Mn2+ (colourless)
MnO2 (brown solid/ppt)
Cr3+ (green)
I- (colourless)

H2O2 (colourless)

Acidic

H2O

Reaction Medium
Acidic
Acidic/Alkaline/Neutral
Neutral
Acidic
Acidic

Product
Fe3+ (yellow)
I2 (brown in aq solution)
S4O62- (colourless)
O2
CO2

Common Reducing Agents
Reductant
Fe2+ (pale green)
I- (colourless)
S2O32- (colourless)
H2O2 (colourless)
C2O42- (colourless)

Topic 2: Atomic Structure and Physical Properties
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Atomic number (Z): The number of protons in the nucleus of each atom of an element
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Isotopes: Elements of the same atomic number but different number of neutrons
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Isoelectronic: Same number of electrons
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Isotopic: Same number of protons
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Subshells: subdivision of each principal quantum shell
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Aufban’s Principle: Electrons in their ground states occupy orbitals in order of energy
levels
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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
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Effective nuclear charge: Zeff is the net nuclear charge experienced by an outer electron
(Zeff ≈ Z - S)
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Electronegativity: The relative ability of an atom in a molecule to attract the shared
electrons in a covalent bond
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Ionic Bond: The electrostatic forces of attraction between oppositely charged ions in an
ionic compound
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Lattice Energy ∝ |r+ +r−|
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Covalent Bond: The electrostatic forces of attraction between the nuclei of atoms and
their shared pair of electrons
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Bond Energy: The energy required to break a particular bond in one mole of a gaseous
substance
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Pi (π) Bond: “Side-on” overlap of two atomic orbitals
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Electron groups around a central atom locate themselves as far away from each
other as possible to minimise electron repulsion
b
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Repulsion between EPs depend on the relative electronegativity of central and
bonding atoms
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Ionic Bond with covalent character: due to the distortion of the electron cloud by cation of
high charge density
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Covalent Bond: Equal sharing of electrons due to no difference in electronegativity
between 2 atoms
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cation is large and anion is small
b
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Bond polarity: separation of positive and negative charges
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Polar Bond: unequal sharing of electrons due to difference n electronegativity of bonding
atoms
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Permanent dipole-permanent dipole interactions: electrostatic forces of attraction
between polar molecules
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Ionic solute – Polar solvent (Ion-dipole interactions)
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Non-polar solute – non-polar solvent (dispersion forces)
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Boyle’s Law: V ∝ 𝑝 (Ceteris paribus)
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Avogadro’s Law: V ∝ n (Ceteris paribus)
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Ideal gas: a hypothetical gas whose pressure-volume-temperature behaviour can be
completely accounted for by the ideal gas equation pV = nRT
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31 JK-1mol-1
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Partial Pressure: Pressure exerted by a particular gas on the sides of a container in a
mixture of gases
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Kinetic Theory of Gases (assumptions applied to an ideal gas):
a
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The forces of attraction and repulsion between gas particles are negligible ie
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Gas particles collide elastically ie
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Gas particles are in constant random motion
e
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At high pressure:

a
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The molecules are close together, thus intermolecular forces are significant
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Real gases behave most ideally at low pressures and high temperatures
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Large molecules – large electron cloud result in stronger intermolecular forces
b
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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
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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

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Heat capacity (C): The amount of heat required to raise an object’s temperature by 1K

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c of water is 4
...


q = mcΔT

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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

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Enthalpy change of atomization (ΔHatomѳ): Energy required when 1 mole of gaseous
atoms are formed the element at 1 atm and 298K
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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
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LE = ΔHhyd - ΔHsol
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ΔHrѳ = ∑ 𝒎𝚫𝐇𝐟 ѳ (𝐩𝐫𝐨𝐝𝐮𝐜𝐭𝐬) − ∑ 𝒏𝚫𝐇𝐟 ѳ (𝒓𝒆𝒂𝒄𝒕𝒂𝒏𝒕𝒔)
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Bond energy: Average energy required to break 1 mole of covalent bond in the gaseous
state
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Bond order
b
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Bond polarity
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Bond Haber Cycle: An extension of Hess’s Law to ionic compounds
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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
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ΔS>0: greater disorder/more no
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ΔS<0: more ordered state/less no
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Factors affecting entropy of a system:
a
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Change in phase
c
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Expansion of a gas
e
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Dissolution of an ionic solution
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ΔG<0: Reaction is feasible and can take place spontaneously (reaction is exergonic)
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There is no net change (during melting and boiling)
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Collision Theory: For effective collisions to occur, reactant molecules must:
a
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Collide with a certain minimum energy called activation energy, Ea
c
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Activation Energy, Ea: The minimum energy which the reacting particles must possess in
order to overcome the energy barrier before the formation of products

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Rate of reaction: Change in concentration of reactant or product per unit time

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For any reaction: aA + bB → cC +dD, Rate = −

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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]

𝑎 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
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Rate equation: An experimentally determined equation that relates the rate of reaction
to the concentration of the reactants raised to appropriate powers
...
Rate = k[A]m[B]n
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Rate constant: affected by temperature and presence of catalyst only (k = Ae-Ea/RT)
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Half-life of a reaction: the time taken for the concentration of a reactant to decrease to
half of its initial value
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First-order Reactions: Rate = k[A]1 = k[A], t½ =
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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

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Reaction intermediate: a chemical species that is produced in an elementary step and
consumed in another
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Molecularity: the number of reactant particles taking part in the reaction in that
elementary step
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The slowest step is the rate-determining step



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