5.1 Rates, equilibrium and pH Flashcards
rate of reaction definition
change in concentration (of product or reactants) over time
order definition
power to which a concentration (of a reactant) is raised
overall order definition
sum of powers in the rate equation
rate constant definition
probability constant
half-life definition
time taken for concentration of a reactant to halve
rate-determining step definition
slowest step in a reaction mechanism
zero order
concentration of that reagent doesn’t affect the rate
first order
concentration of that reagent is proportional to the rate of reaction
second order
rate of reaction is proportional to the square of concentration of this reagent
rate constant relationship with constant half-life
kt1/2 = ln2 = 0.693
t1/2 = constant half-life
only for first order reactions where half-life is constant
plotting concentration against time
0 order = linear, negative gradient (decreasing half-life)
1st order = curved, negative gradient becomes less steep (constant half-life)
2nd order = curved, steeper but becomes less steep faster (increasing half-life)
plotting rate against concentration
0 order = flat horizontal line
1st order = linear line
2nd order = exponentially increased curved line
how to determine number of and what molecules are in rate determining step
0 order = not involved
1st order = 1 molecule of this reactant
2nd order = 2 molecules of this reactant
etc.
what number of different reactants in rate determining step means
1 = decomposition 2 = collision
why Arrehenius is important
standard rate equation doesn’t take into account temperature
Arrehenius units
T = temperature (Kelvin) R = gas constant (8.314 J K^-1 mol) EA = activation energy (J mol^-1) A = pre-exponential factor, takes into account number of molecules that exceed activation energy k = rate constant
why Arrehenius is important
standard rate equation doesn’t take into account temperature
Arrehenius units
T = temperature (Kelvin) R = gas constant (8.314 J K^-1 mol) EA = activation energy (J mol^-1) A = pre-exponential factor, takes into account number of molecules that exceed activation energy k = rate constant
mole fraction formula
mole fraction = number of moles/total number of moles of gas
factors affecting Kc
only temperature
partial pressure formula
partial pressure = mole fraction x total pressure
formula of Kp
exact same as formula of Kc but with partial pressures
only gaseous reactants/products
mole fraction formula
mole fraction = number of moles/total number of moles of gas
mole fraction formula
mole fraction = number of moles/total number of moles of gas
ph formula
pH = -log[H+]
pH = negative log of concentration of H+
can only work directly with strong acid
pH scale
logarithmic
pH1 is 10 times stronger than pH2
phenolphthalein
colourless (0-8) to pink (9-14)
methyl orange
red/orange (0-3) to yellow (4-14)
bromophenol blue,
yellow (0-3) to blue (4-14)
conjugate acid-base pair definition
conjugate acid base pairs differ by the presence or absence of a transferable proton
amphoteric definition
can act as proton donor or proton acceptor
monobasic definition
donates 1 proton
dibasic definition
donates 2 protons
tribasic definition
donates 3 protons
why Kc is always low during acid dissociation
[H2O] will always be close to 55 mol dm^-3
much higher in comparison to [H3O^+] and [A^-]
Ka formula
Ka = [H^+][A^-]/ [HA]
[H2O] removed as it remains near constant
ionic product of water expression
Kw = [H^+] [OH^-]
Brønsted-Lowry acid definition
proton donor
Brønsted-Lowry base definition
proton acceptor
calculating pH of weak acids
write expression of Ka
assume [H^+] = [A^-]
calculate [H^+] then pH
why [HA] at start = [HA] at equil. can be assumed during weak acid dissociation
dissociation is very small
any decrease in [HA] is negligible
why [H^+] at equil. = [A^-] at equil. can be assumed during weak acid dissociation
even though some water does dissociate, it is very small compared to [H^+]
buffer solution definition
system that minimises pH changes on addition of small amounts of an acid or a base
composition of buffer solutions
a weak acid and a salt of that acid
e.g. ethanoic acid and sodium ethanoate
what happens to equilibria when acid is added example
- CH3COOH <=> CH3COO- + H+
- CH3COONa <=> CH3COO- + Na+
[H+] increases
position of equilibrium 1 shifts left
H+ ions react with ethanoate ions, removing H+
minimal change in pH
what happens to equilibrium when alkali is added example
- CH3COOH <=> CH3COO- + H+
- CH3COONa <=> CH3COO- + Na+
[OH-] increases
reacts and removes H+, decreasing [H+]
pos. of equilibrium shifts right
more ethanoic acid dissociates, replacing missing protons
assumptions when calculating pH of buffers
salt has fully dissociated
weak acid has not dissociated
factors of pH of buffer
Ka of weak acid
concentration ratio of weak acid:conjugate base or salt
titration curves
shows how pH of solution changes when different volumes of different acids and bases are neutralised
strong acid strong base titration curve shape
high starting pH
very gradual change until equivalence point
very sudden drop
gradual change until pH of acid
strong acid weak base titration curve
pH not as high slightly steeper curve until equivalence point sharp drop (shorter) gradual change until pH of acid
weak acid strong base titration curve
gradual change until equivalence point sharp drop (shorter) slightly steeper change to pH of acid
weak acid weak base titration curve
slightly steeper change until equivalence point
less steep drop
steeper change until pH of acid
importance of titration curves to indicators
helps choose appropriate indicator to show equivalence point
Haber process conditions and why
N2 (g) + 3H2 (g) ⇌ 2NH3(g) forward reaction is exothermic, reverse is endothermic medium temperature (400-450°C) as high temp. increases rate of reaction but decreases yield high pressure (200/250 atm.) to increase rate and increase yield iron catalyst (increase rate of forward and backward reaction)
