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Rate of a reaction
proportional to the concentration of a particular reactant raised to the power of it's order.
K
Proportionality/Rate constant
Finding orders of reaction
Only through experiments
Concentration-time graphs plotting
Gas collection, mass loss and colour change
Monitoring rate with a colorimeter
Measures light absorbed by solution. Select a filter with complimentary colours; zero colorimeter with water; plot a calibration curve with known concentrations; take absorbance readings at timed intervals; compare to curve.
Concentration-time graphs - zero order
Gradient = k
Half-life
time taken for half of a reactant to be used up
Concentration-time graphs - first order
Constant half-life. Exponential decay
Finding K on C-T graphs, first order
substitute values of concentration and rate from tangent into rate equation; K=ln2/half-life
Rate-concentration graph - zero order
y-intercept = k
Rate-concentration graph - first order
Gradient = k
Rate-concentration graph - second order
k = gradient of rate-concentration^2 graph
Clock reaction
Calculates initial rate. Time, t, is measured for a visual change. Initial rate is proportional to 1/t
Iodine clock reaction
Formation of iodine, t=time until first appearance of iodine colour. Starch often added to form a complex - dark blue-black colour appears when Na2S2O3 (Sodium thiosulfate) is used up.
Assumptions on average rate of reaction
Constant, and same as initial rate (accurate providing less than 15% of the reaction has taken place)
Why is it expected that reactions happen in a series of steps
Unlikely that multiple species collide simultaneously
The slowest step
rate determining step
The rate-determining step
Only includes species on the rate equation, orders of which match the number of species in the rds
Increase in temperature
increase in rate and rate constant. Shifts boltzmann distribution to the right, increasing the proportion of particles that exceed the activation energy.
The Arrhenius Equation
k=Ae^(-Ea/RT)
lnk = lnA - Ea/RT
Gradient = -Ex/R ; y-intercept = Ln A
Heterogeneous equilibria
Species of different states. Only gases and solutions included in Kc as solids/liquids have constant concentrations.
Kp
equilibrium constant in terms of partial pressures of gases, which are proportional to it's concentration.
Mole fraction of A
moles of A/total moles ; Volume of A/total volume
Sum of the mole fractions
= 1
Partial pressure of A
mole fraction of A x total pressure
Kp expression
includes ONLY gases
Application Kp
Endurance athletes train at high altitude to produce haemoglobin with extra oxygen containing abilities
K is only affected by
temperature
Exothermic reaction, increase temperature
Kc decreases (decrease in product yield) ; Kp: Partial pressures of reactants increase, shift to left, lower Kp value
Endothermic reaction, increase temperature
Kc increases (increase product yield) ; Partial pressures of products increase, shift to right, higher Kp value
Change in concentration or pressure
does not affect K
If total pressure doubles
so does each partial pressure
Increase pressure, Kc
Kc doesn't change. Ratio of expression now less/more than Kc value. Concentrations of reactants/products increase/decrease to return system to equilibrium and restore ratio of Kc
Conjugate acid-base pair
Two species that can be interconverted by transfer of a proton
pH scale
Measures concentration of H+ ions.
pH =
-log[H+]
[H+] =
10^-pH
Change in 1 pH unit
10 fold change in H+
In Strong acids
[H+] = [HA]
Ka
acid dissociation constant
Ka =
[H+][A-]/[HA]
The stronger the acid
the LARGER the Ka value and the SMALLER the pKa value
The weaker the acid
the SMALLER the Ka value and the LARGER the pKa value
Approximation 1 on weak acids
HA dissociates to produce equal ammounts of H+ and A- ; [H+] = [A-]
Approximation 2 on weak acids
Dissociation is small, so [HA]eqm = [HA]start
Ka for weak acids
[H+]^2[HA]
Ka determined experimentally
Prepare standard of weak acid at known concentration, measure pH with a pH metre
Limitations of approximation 1 of weak acid
If ph>6, then [H+] from dissociation of water will be significant compared to acid. So assumption breaks down with very weak/dilute acids
Limitations of approximation 2 of weak acid
Breaks down when [H+] is significant. , so doesnt work for stronger weak acids with Ka>10^-2 and very dilute solutions.
Kw
The ionic product of water, 1x10^-14
Kw =
[H+][OH-]
Change in temperature effect Kw
Kw varies
Water is neutral so
[H+] = [OH-]
For pH values with whole numbers
indices for [H+] and [OH-] add up to -14
A strong base
an alkali that completely dissociates it OH- in solution.
pH of Strong base
[OH-] = [NaOH], and Kw
Ammonia is a weak base
Buffer Solution
A system that minimises pH changes on addition of small amounts of an acid or a base
Preparing buffer solutions from weak acid and it's salt
Solution of a weak acid and it's salt. HA only partially dissociates a small amount - source of HA. Salt completely dissolves, provides A- ions.
Preparing buffer solutions from partial neutralisation
Add an aqueous solution of an alkali to an excess of weak acid. Weak acid is partially neutralised, forming conjugate base and some HA left over.
Conjugate base of buffer solution removes added acid
[H+] increases, reacts with A- ions, equilibrium shifts to the left, removing most H+ ions.
Weak acid of buffer solution removes added alkali
[OH-] increases, reacts with H+ ions, HA dissociates, shifting equilibrium to the right to restore H+ ions.
Buffer most effective when
[HA]=[A-] ; pH of buffer is same as pKa of HA
Fine tune pH of buffer solution
ratio of concentrations of the weak acid and it's base can be adjusted
pH buffer solution
[H+] = Ka x [HA]/[A-]
Ideal pH for blood plasma
7.35-7.45
Buffer system in blood
Carbonic acid-hydrogencarbonate
Carbonic acid-hydrogencarbonate
Body prevents build up of H2CO3
Converting it into CO2 which is then exhaled by lungs
pH titration curve - equivalence point
the volume of one solution that reacts exactly with the volume of the other. Amounts use match stoichiometry of the reaction.
acid-base indicator
a weak acid that has a distinctly different colour from its conjugate base.
Methyl orange
Red in acid, yellow in alkali
Added Alkali
shift right
Added Acid
shift left
At the end point
[HA] = [A-], Ka = [H+], pKa = pH
Indicator must have a colour change
which coincides with vertical section of curve. Ideally, end point = equivalence point
weak acid - weak base reaction
No indicator is suitable
Standard enthalpy change of solution
one mole of a solute dissolves in a water under standard conditions
Lattice Enthalpy
The enthalpy change that accompanies the formation of one mole of an ionic compound from its gaseous ions under standard conditions
standard enthalpy of formation
The change in Enthalpy that accompanies the formation of 1 mole of a compound from its elements with all substances in their standard states.
Standard enthalpy change of atomisation
the enthalpy change when one mole of gaseous atoms is formed from an element under standard conditions
First ionisation energy
The energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions
First electron affinity
The enthalpy change that takes place when one electron is added to each atom in one mole of gaseous atoms to form one mole of gaseous 1- ions
Enthalpy change of Hydration
The enthalpy change that accompanies the dissolving of gaseous ions in water to form one mole of aqueous ions
Sol H
q/n
Why are first electron affinities exothermic?
Because electrons are attracted to nucleus
Why are second electron affinities endothermic?
Because the ion is already negative so repels the electron. So energy is needed to force it in
Effect of ionic size on Lattice Enthalpy
ionic radius increases, attraction decreases, less negative, melting point decreases
Effect of ionic charge on Lattice Enthalpy
ionic charge increases, attraction increases, more negative, melting point increases
Effect of ionic size on Hydration
ionic radius increases, attraction between ion and water decreases, less negative
Effect of ionic charge on Hydration
ionic charge increasing, attraction increases, more negative
Soluble
attraction between ions in ionic lattice needs to be broken
If sum of hydration is larger than lattice enthalpy
then enthalpy change of solution is exothermic and should dissolve.
Enthalpy change of solution doesn't take into account
temperature and entropy
Entropy
dispersal of energy within the chemicals making up the chemical system
Increasing entropy
Solids -> Liquids -> Gases
If a system changes to become more disordered
entropy change will be positive
Feasibility
describes whether a reaction is able to happen, energetically
Free energy
Overall change in energy during a chemical reaction.
Note 
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