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A-Level Chemistry > 5.1 Rates, equilibrium and pH > Flashcards

5.1 Rates, equilibrium and pH Flashcards

1
Q

rate of reaction definition

A

change in concentration (of product or reactants) over time

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2
Q

order definition

A

power to which a concentration (of a reactant) is raised

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3
Q

overall order definition

A

sum of powers in the rate equation

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4
Q

rate constant definition

A

probability constant

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5
Q

half-life definition

A

time taken for concentration of a reactant to halve

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6
Q

rate-determining step definition

A

slowest step in a reaction mechanism

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7
Q

zero order

A

concentration of that reagent doesn’t affect the rate

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8
Q

first order

A

concentration of that reagent is proportional to the rate of reaction

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9
Q

second order

A

rate of reaction is proportional to the square of concentration of this reagent

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10
Q

rate constant relationship with constant half-life

A

kt1/2 = ln2 = 0.693
t1/2 = constant half-life
only for first order reactions where half-life is constant

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11
Q

plotting concentration against time

A

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)

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12
Q

plotting rate against concentration

A

0 order = flat horizontal line
1st order = linear line
2nd order = exponentially increased curved line

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13
Q

how to determine number of and what molecules are in rate determining step

A

0 order = not involved
1st order = 1 molecule of this reactant
2nd order = 2 molecules of this reactant
etc.

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14
Q

what number of different reactants in rate determining step means

A 
1 = decomposition
2 = collision
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15
Q

why Arrehenius is important

A

standard rate equation doesn’t take into account temperature

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16
Q

Arrehenius units

A 
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
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17
Q

why Arrehenius is important

A

standard rate equation doesn’t take into account temperature

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18
Q

Arrehenius units

A 
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
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19
Q

mole fraction formula

A

mole fraction = number of moles/total number of moles of gas

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20
Q

factors affecting Kc

A

only temperature

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21
Q

partial pressure formula

A

partial pressure = mole fraction x total pressure

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22
Q

formula of Kp

A

exact same as formula of Kc but with partial pressures

only gaseous reactants/products

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23
Q

mole fraction formula

A

mole fraction = number of moles/total number of moles of gas

24
Q

mole fraction formula

A

mole fraction = number of moles/total number of moles of gas

25
ph formula
pH = -log[H+] pH = negative log of concentration of H+ can only work directly with strong acid
26
pH scale
logarithmic | pH1 is 10 times stronger than pH2
27
phenolphthalein
colourless (0-8) to pink (9-14)
28
methyl orange
red/orange (0-3) to yellow (4-14)
29
bromophenol blue,
yellow (0-3) to blue (4-14)
30
conjugate acid-base pair definition
conjugate acid base pairs differ by the presence or absence of a transferable proton
31
amphoteric definition
can act as proton donor or proton acceptor
32
monobasic definition
donates 1 proton
33
dibasic definition
donates 2 protons
34
tribasic definition
donates 3 protons
35
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^-]
36
Ka formula
Ka = [H^+][A^-]/ [HA] | [H2O] removed as it remains near constant
37
ionic product of water expression
Kw = [H^+] [OH^-]
38
Brønsted-Lowry acid definition
proton donor
39
Brønsted-Lowry base definition
proton acceptor
40
calculating pH of weak acids
write expression of Ka assume [H^+] = [A^-] calculate [H^+] then pH
41
why [HA] at start = [HA] at equil. can be assumed during weak acid dissociation
dissociation is very small | any decrease in [HA] is negligible
42
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^+]
43
buffer solution definition
system that minimises pH changes on addition of small amounts of an acid or a base
44
composition of buffer solutions
a weak acid and a salt of that acid | e.g. ethanoic acid and sodium ethanoate
45
what happens to equilibria when acid is added example
1. CH3COOH <=­> CH3COO-­ + H+ 2. CH3COONa <=­> CH3COO­- + Na+ [H+] increases position of equilibrium 1 shifts left H+ ions react with ethanoate ions, removing H+ minimal change in pH
46
what happens to equilibrium when alkali is added example
1. CH3COOH <=> CH3COO-­ + H+ 2. CH3COONa <=> CH3COO-­ + Na+ [OH-] increases reacts and removes H+, decreasing [H+] pos. of equilibrium shifts right more ethanoic acid dissociates, replacing missing protons
47
assumptions when calculating pH of buffers
salt has fully dissociated | weak acid has not dissociated
48
factors of pH of buffer
Ka of weak acid | concentration ratio of weak acid:conjugate base or salt
49
titration curves
shows how pH of solution changes when different volumes of different acids and bases are neutralised
50
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
51
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 ```
52
weak acid strong base titration curve
``` gradual change until equivalence point sharp drop (shorter) slightly steeper change to pH of acid ```
53
weak acid weak base titration curve
slightly steeper change until equivalence point less steep drop steeper change until pH of acid
54
importance of titration curves to indicators
helps choose appropriate indicator to show equivalence point
55
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) ```

A-Level Chemistry (12 decks)

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