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A Level Chemistry > Module 5 Section 1: Rates, Equilibrium and pH > Flashcards

Module 5 Section 1: Rates, Equilibrium and pH Flashcards

1
Q

Expression for equilibrium constant

A

Remember [ ]

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2
Q

Calculate Kc (with units) for the reaction:

A

Kc =

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3
Q

What is a homogeneous mixture

A

All the reactants and products are put into the expression for the equilibrium constant

All reactants and products are in the same state

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4
Q

What is a heterogeneous mixture

A

Only gases and aqueous substances go into the expression for the equilibrium constant (solids and liquids are left out)
However, if the solvent isn’t aqueous you must include any water in the Kc expression

Reactants and products are different states

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5
Q

Expression for the equilibrium constant for:

A

Reactants and products are a mixture of aqueous and solid (heterogeneous) so only aqueous substances are included in Kc

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6
Q

Find the units when (mol dm-3)^2 / (mol dm-3)^4

A

(mol dm-3)^2 / (mol dm-3)^4 = 1 / (mol dm-3)^2 = 1 / mol^2 dm^-6 = mol^-2 dm^6

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7
Q

Equation: A(g) ⇌ 2B(g)
4.00 moles of A was placed in a 20.0dm3 container and heated to 320K until equilibrium had been established
The equilibrium mixture was found to contain 1.50 moles of A
Calculate Kc at this temperature

A

A(g) ⇌ 2B(g) 1 : 2

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8
Q

What does the order of reaction tell you

A

With respect to a particular reactant, the order of reaction says how the reactants concentration affects the rate

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9
Q

What does the zero order mean

A

If you double the reactant’s concentration, the rate stays the same
Concentration does not influence rate
Increases by 2^0 = 1 (stays same)

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10
Q

What does the first order mean

A

If you double the reactant’s concentration, the rate also doubles
Increases by 2^1 = 2 (doubles)

Tripling the concentration triples the rate

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11
Q

What does the second order mean

A

If you double the reactant’s concentration, the rate multiplies by 4
Increases by 2^2 (multiplies by 4)

Tripling the rate increases the rate by a factor of 3^2

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12
Q

What is the overall order

A

The sum of the orders of all the different reactants

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13
Q

What is the only way to find out the orders of reaction

A

Can only be found from experiments
Not from chemical equations

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14
Q

Methods to work out the order of reaction for a specific reactant in a reaction

A

Continuously monitor the change in concentration of the chosen reactant against time to construct a rate-concentration graph
Or
Use an initial rates method to find out how the initial rate changes as you vary the concentration of the chosen reactant

For each method, the concentrations of other reactants must be in excess so the change in rate is only due to changing the conc of the chosen reactant

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15
Q

How to construct a rate-concentration graph from a concentration-time graph

A

Find the gradient (represents the rate) at various points along the conc-time graph
This gives a set of points for the rate-conc graph
Plot points and join them up with line or smooth curve
Shape of the graph shows the order

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16
Q

What the different orders look like on a concentration-time graph

A

Zero order: horizontal line to show that changing the concentration doesn’t change the rate

First order: straight line through origin, shows the rate is proportional to [X]

Second order: curve meaning that the rate will be proportional to [X]^2

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17
Q

How to work out the order of reactions for reactants using experimental data

A

Look at what happens when the conc doubles for a specific reactant and see how this affects the rate
If rate stays the same, order is zero
If rate doubles, order is first
If rate quadruples, order is second

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18
Q

What do rate equations mean

A

They tell us how the rate is affected by the concentrations of reactants

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19
Q

Rate equation with units

A

k: rate constant, bigger it is the faster the reaction
m, n: the orders of the reaction with respect to reactants A and B
m says how the conc of A affects the rate (same for n and B)
Units of rate are mol dm-3 s-1

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20
Q

Find rate equation for reaction:

A

Rate equation: k[CH3COCH3]^1 [H+]^1 [I2]^0
[X]^1 can be written as [X] and [X]^0 = 1 so can be left out
Simplified: k[CH3COCH3] [H+]

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21
Q

Steps to work out units for a rate equation

A

1.Rearrange the equation to make k the subject
2.Substitute units into the expression for k
3.Cancel common units and show the final units on a single line

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22
Q

How to find rate constant from rate-concentration graph of first order reactions

A

If the overall reaction is first order the rate constant is equal to the gradient of the rate-concentration graph of that reactant

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23
Q

What does [A] represent

A

Concentration of A

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24
Q

What do colorimeters do

A

Measure the absorbance of a particular wavelength of light by a solution

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25
Q

How to measure the rate of reaction using a colorimeter

A

Can be set to measure the absorbance of a wavelength that is absorbed by one of the reactants but not by the products
The change in absorbance over the course of the reaction can be used to measure the rate

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26
Q

How to use a colorimeter to produce concentration time graph

A
  1. Produce standard solutions of known concentrations
  2. Select correct filter
  3. Zero colorimeter with water
  4. Plot calibration curve (from absorbance readings of standard solutions)
  5. Carry out reaction of unkown solution and take absorbance readings at regular intervals
  6. Use calibration curve to measure the concentration of iodine at each absorbance reading
  7. Plot concentration time graph
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27
Q

How to plot and work out rate of reaction from a concentration time graph

A

You can plot a concentration-time graph by repeatedly taking measurements during a reaction
The rate at any point in the reaction is given by the gradient at that point on the graph
If the graph is a curve, the draw a tangent to the curve and find the gradient of that
Sign in front of gradient (+or-) doesn’t matter, take positive value

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28
Q

How to find rate from the gradient

A

Gradient = change in y/change in x

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29
Q

Shape of order of reaction for concentration-time graphs

A

Zero order: negative straight line with unchanging gradient
First order: downward curve with a decreasing gradient over time
Second order: downwards curve which is steeper at the start, but falling off more slowly

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30
Q

What is the half-life for a reactant

A

This is the time taken for half of the reactant to be used up

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31
Q

Half life for first order of reaction

A

First order: this is independent of the concentration
So each half life will be the same length every time the reactant concentration halves

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32
Q

Equation to find the rate constant for a first order reaction using the half life

A

k = ln2/t 1/2

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33
Q

Find concs

A

Answers

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34
Q

Find Kc

A
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35
Q

0.5 mol B and 0.3 mol C mixed in 10dm3
At equilibrium there was found to be 0.1 mol of A present
Calculate Kc

A
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36
Q

What is total pressure

A

The total pressure of a gas mixture is the sum of all the partial pressures of individual gases

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37
Q

Equation for mole fraction of a gas in a mixture

A

Number of moles of gas / total number of moles of gas in the mixture

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38
Q

Equation for partial pressure

A

Mole fraction of gas x total pressure of the mixture

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39
Q

Find partial pressure

A

Answer

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40
Q

Find partial pressure
Use mole fraction equation

A

Ans

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41
Q

How to find Kp

A
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42
Q

Calculate Kp for decomposition of PCl5

A
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43
Q

Calculate partial pressure of monomer in the equilibrium and find total pressure exerted by equilibrium mixture

A
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44
Q

What to do when calculating kP expressions for heterogeneous equilibrium

A

You do not include solids or liquids, only gases

With: NH4HS (s) ⇌ NH3 (g) + H2S (g)
Expression for kP is kP = p(NH3) p(H2S)
(No bottom line as reactant is solid)

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45
Q

Draw each concentration-time graph for different orders

A
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46
Q

Draw each rate-concentration graph for different order

A
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47
Q

Work out rate constant:

A
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48
Q

What is the initial rate of a reaction and how do you find it on a graph

A

This is the rate right at the start of the reaction
Calculate the gradient of the tangent at time = 0

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49
Q

How to show that the value of initial rates depend on the reactant concentration

A

Carry out reaction, continuously monitoring one reactant
Use this to draw concentration-time graph
Repeat experiment using different initial concentration of the reaction
Keep the concentrations of other reactants the same, draw another concentration time graph
Use the graphs to calculate the initial rate for each experiment
Repeat process for each reactant (different reactants may affect the rate differently)

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50
Q

How to use clock reactions to measure initial rates

A

In a clock reaction, you measure how the time taken for a set amount of product to form changes as you vary the concentration of one of the reactants
There is usually an easily observable endpoint, such as colour change, to tell you when the desired amount of product has formed
The quicker the clock reaction finishes, the faster the initial rate of the reaction

(Known conc of reactant, rate stays the same, so time taken for end point to finish can be used in change in conc/time to find initial rate (rate stays same)) ask Pritchard?

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51
Q

Assumptions for clock reactions

A

The concentration of each reactant doesn’t change significantly over the time period of your clock reaction
The temperature stays constant
When the endpoint is seen, the reaction has not proceeded too far

If these assumptions are reasonable for your experiment, you can assume that the rate of reaction stay constant during the time period of your measurement, so the rate of your clock reaction will be a good estimate for the initial rate of you reaction

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52
Q

How does the iodine clock reaction work

A

Small amount of sodium thiosulfate solution and starch are added to an excess of hydrogen peroxide and iodide ions in acid solution
Starch is used as an indicator as it turns blue black in the presence of iodine
The sodium thiosulfate that is added to the added to the reaction mixture reacts instantaneously with any iodine that forms
To begin, all the iodine that forms in the first reaction is used up straight away in the second reaction
But once all the sodium thiosulfate is used up, any more iodine that forms will stay in solution, so the starch indicator will suddenly turn the solution blue-black, this is the end of the clock reaction
Varying iodide or hydrogen peroxide concentration while keeping the other constant will give different times for the colour change

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53
Q

Two reactions of iodine clock reaction

A
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54
Q

What is a rate determining step

A

Reaction mechanisms can have one step or a series of steps
In series of steps, each step can have a different rate
The overall is decided by the step with the slowest rate
This is the rate determining step

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55
Q

How to pick the reactants from the chemical equation are involved in the rate determining step

A

If a reactant appears in the rate equation, it must affect the rate
So this reactant, or something derived from, must be in the rate determining step

If a reactant doesn’t appear in the rate equation, then it isn’t involved in the rate determining step (and neither is anything derived from it)

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56
Q

Important points about rate determining steps and mechanisms

A

The RDS doesn’t have to be the first step in a mechanism
The reaction mechanism can’t usually be predicted from just the chemical equation

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57
Q

How to tell how many molecules are involved in the rate determining step from the order

A

The order of a reaction with respect to a reactant shows the number of molecules of that reactant which are involved in the rate determining step

So, if a reaction’s second order with respect to X, there’ll be two molecules of X in the rate determining step

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58
Q

Predict rate equation from this mechanism:

A
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59
Q

Predict the mechanism from the rate equation for the reaction:
(CH3)3CBr + OH- -> (CH3)3COH + Br-
Rate equation: k[(CH3)3CBr]

A
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60
Q

Possible exceptions to suggesting a mechanism for this reaction

A
  1. From the chemical equation, it looks like 2 N2O5 molecules react with each other so the reaction may be predicted as second order with respect to N2O5
  2. However the rate equation shows that the reaction is first order
  3. There’s only 1 molecule of N2O5 in the RDS
    So there’s only one mechanism
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61
Q

What needs to happen for a reaction to occur

A

Particles need to:
Collide with eachother
Have enough energy to react (exceed activation energy)
Correct orientation

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62
Q

How does increasing the temperature increase the rate in terms of particles

A

Gives particles more kinetic energy so they speed up and collide more often
More reactant particles have the required activation energy so there are more successful collisions

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63
Q

How does increasing the temperature increase the rate

A

According to the rate equation, reaction rate depends only on the rate constant and reactant concentrations
So changing the temperature must change the rate constant

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64
Q

How does temperature increasing the rate constant result in a higher rate of reaction

A

The rate constant applies to a particular reaction at a certain temperature
At a higher temperature, the reaction will have a higher rate constant
(Higher rate constant, faster the rate)

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65
Q

What is the Arrhenius equation

A

This links the rate constant (k) with activation energy (Ea) and temperature (T)

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66
Q

How does activation energy and temperature effect k in the Arrhenius equation

A

As the activation energy increases, k gets smaller
Means that a large Ea will have a slow rate (as less successful collisions)
Shows that as the temperature rises, k increases

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67
Q

Putting Arrhenius equation into logarithmic form

A
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68
Q

How to create Arrhenius plot from logarithmic form of Arrhenius equation

A

Plot lnk against 1/T
This will produce a graph with a gradient of -Ea/R and a y intercept of lnA

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69
Q
A
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70
Q
A
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71
Q

What is Le Chatelier’s principle

A

If there’s a change in concentration, pressure or temperature, the equilibrium will move to help counteract the change

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72
Q

What does the size of the Kc say about the position

A

The larger the value of Kc, the further to the right the equilibrium lies and the more products there will be relative to reactants
The smaller the value of Kc, the further to the left the equilibrium lies and the more reactants there will be relative to products

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73
Q
A
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74
Q

General rule for what happens to an equilibrium when you change the temperature

A

If changing the temperature causes less product to form, the equilibrium moves to the left, and the equilibrium constant decreases
If changing the temperature causes more product to form, the equilibrium moves to the right, and the equilibrium constant increases

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75
Q

What does increasing concentration do to the equilibrium constant

A

The value of the equilibrium constant is fixed at a given temperature
So if the concentration of one thing in the equilibrium mixture changes then the concentration of the others must change to keep the value of Kc the same
Kc doesn’t change

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76
Q

What does changing the pressure do to the equilibrium constant

A

Increasing the pressure shifts the equilibrium to the side with fewer gas molecules - reduces the pressure
Decreasing the pressure shifts the equilibrium to the side with more gas molecules - this raises the pressure again
Kp (Kc) stays the same no matter what you do to the pressure

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77
Q

What do concentrations, pressure and temperature do to the values of Kc or Kp

A

Pressure and concentration do not affect the values of Kc or Kp
They do change the amounts of products and reactants present at equilibrium

Changes in temperature not only alter the amounts of product and reactants present at equilibrium but also change the value of equilibrium constant

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78
Q

What do catalysts do to the position of equilibrium and Kc and Kp

A

Catalysts have no effect on the position of equilibrium of the value of Kc/Kp
Can’t increase yield but do mean that equilibrium is approached faster

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79
Q

How to find the half life of a first order reaction

A

Means that the half life of a first order reaction can be read off its concentration-time graph by seeing by seeing how long it takes to halve reactant concentration

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80
Q

Find rate constant from the half life

A
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81
Q

Determine two step mechanism and find RDS
Rate equation = k[NO2]^2

A

Step one is RDS

82
Q

Determine two step mechanism and find RDS

A
83
Q

Solubility of ionic salts

A

Solubles
All Group 1 salts are soluble
All Halides (except for silver halides) are soluble
All nitrates are soluble

Group 2 nitrates are soluble
Group 2 halides soluble
Group 2 sulfate solubility decreases down the group
Group 2 hydroxides increases down the group

84
Q

Process of how to work out (not actual answer)

A

Looking at the mole ratio it is 1:1
Means that the partial pressure of PCl3 will be the same as chlorine
Total pressure is the sum of all partial pressures of gases
Means that you can find the partial pressure of PCl5 by finding difference between 714 and (263+263)

85
Q

Process of how to work out (not actual answer)

A

Write out Initial moles, Change in moles and Equilibrium moles (ICE)
Find moles of PCl5 at equilibrium using mole ratios
Mole fraction of PCl5 x 714
Partial pressure = total pressure x mole fraction

86
Q

What are Brønsted-Lowry acids

A

Bronsted-Lowry acids are proton donors - they release hydrogen ions (H+) when they’re mixed with water

87
Q

How do you normally get H+ ions in water from an acid

A

You never get H+ ions by themselves in water
They’re always combined with H2O to form hydroxonium ions, H3O+

88
Q

What do Brønsted-Lowry bases do when in a solution

A

When they’re in solution, they bind with hydrogen ions from water molecules

89
Q

What do the HA and B mean in acid equilibria

A

HA is any acid
B is a base

90
Q

What are monobasic acids

A

Acids that only have one proton that they can release into solution.
E.g. HCl and HNO3

91
Q

What are dibasic acids

A

Acids that can release two protons into the solution
E.g. H2SO4

92
Q

What are tribasic acids

A

They can release 3 protons into the solution
E.g. phosphoric acid

93
Q

How are conjugate pairs linked in acid equilibrium

A

They are species that are linked by the transfer of a proton
Always opposite sides of the reaction equation

94
Q

What are conjugate bases and conjugate acids

A

Conjugate base: the species that has lost a proton
E.g. in the forwards reaction HCl releases a proton to form the conjugate base Cl-

Conjugate acid: the species that has gain a proton
E.g. in the reverse reaction Cl- accepts a proton to form its conjugate acid HCl

95
Q

How is the equilibrium set up between Brønsted-Lowry acids and bases

A

In the forward reaction, HA acts as an acid as it donates a proton.

In the reverse reaction, A- acts as a base and accepts a proton from the BH+ ion to form HA.

HA and A- are called a conjugate pair
HA is the conjugate acid of A-
A- is the conjugate base of the acid, HA

Similarly, B and BH+ are a conjugate pair.
The base B takes a proton to form BH+
So B is the conjugate base of BH+
BH+ is the conjugate acid of B

96
Q

How does H2O react with acids and bases uniquely to form an equilibrium

A

Water acts as a base and reacts with acids to form a conjugate acid (H3O+)
Water acts as an acid and reacts with bases to form a conjugate base (OH-)

97
Q

How do reactive metals react with acids
Include ionic equation

A

React with acids to release H2 gas
Metal atoms donate electrons to the H+ ions in the acid solution
The metal atoms are oxidised and the H+ ions are reduced

98
Q

How do carbonates react with acids
Include ionic equation

A

Carbonates react with acids to produce carbon dioxide and water

99
Q

How do alkalis react with acids
Include ionic equation

A

Alkalis are bases that release hydroxide ions in water.
They react with acids to form water

100
Q

How do insoluble bases react with acids
Include ionic equation

A

Most insoluble bases are metal oxides.
Like alkalis, they react with acids to form water

101
Q

Lavoisier’s theory of acids

A

Came up with the first theory of acids and bases in 18th century
Knew that sulfuric acid had the formula H2SO4 and nitric acid had the formula HNO3
So he proposed that acids had to have oxygen in them.
Was later shown that acids like hydrochloric acid (HCI) and hydrogen sulfide (H2S) don’t have any oxygen in them

102
Q

Arrhenius’ theory of acids

A

End of the 19th century
Arrhenius suggested acids release protons in aqueous solution, whilst bases release hydroxide ions.
He said when acids and bases react together they always form water and a salt.
Works in many examples but doesn’t work for bases such as ammonia (NH3), which don’t contain any hydroxide ions

103
Q

Brønsted and Lowry’s theory of acids

A

Came up with their definition of acids and bases independently
It’s based on Arrhenius theory, but broadens definition of a base to be a proton acceptor.
Came up with the idea that acids and bases react to form conjugate pairs, rather than a salt and water.
This definition currently explains most of our observations, so is one of the theories we still use today

104
Q

Where is the equilibrium position for the dissociation of strong acids

A

Far to the right to make more ions

105
Q

Identify acids 1 and 2 and bases 1 and 2

A

In the forwards reaction
HCl is an acids because it donates H+
OH- is a base as it accepts H+

In reverse reaction
H2O is an acid as it donates H+
Cl- is a base as it accepts H+

106
Q

Will a strong acid form a stronger or weaker conjugate base

A

Weaker conjugate base
Strong acids dissociate completely giving less H+ so the conjugate base has less basic properties as the reaction goes to completion (reaction goes forward only so base (Cl-) cannot accept H+ as easily)

107
Q

Will a weak acid form a stronger or weaker conjugate base

A

Stronger conjugate base e.g. H2 dissociates to form H+ and H- (strong base)
Weak acids dissociate in an equilibrium so the conjugate bases are able to accept H+ and reform the acid (reaction goes in both directions)

108
Q

What are Brønsted-Lowry bases

A

Bronsted-Lowry bases are proton acceptors.

109
Q

What is pH and how is it measured

A

pH is a measure of how acidic or basic something is
It measures the concentration of hydrogen ions in solution
Concentration of hydrogen ions can vary enormously so a logarithmic scale called the pH scale is used

110
Q

Formula for calculating pH

A
111
Q

Calculate pH

A
112
Q

Formula for find concentration of H+

A
113
Q

Calculate concentration of H+

A
114
Q

How many moles of H+ ions do strong monobasic acids produce in solution

A

Strong monobasic acids such as hydrochloric acid and nitric acid ionise fully in solution.
They’re also monobasic, which means one mole of acid produces one mole of hydrogen ions
So the H+ concentration is the same as the acid concentration
Dibasic acids produce double their moles of H+

115
Q

Find the pH of 0.1 mol dm-3 HCl

A
116
Q

Find pH of 0.05 mol dm-3 HNO3

A
117
Q

How does water form an equilibrium and write out the equilibrium equations

A

Water dissociates into hydroxonium ions and hydroxide ions
To work out the equilibrium constant you must calculate Kw

118
Q

What is Kw

A

Kw is called the ionic product of water
Same as Kc but for water

119
Q

How to calculate Kw for water

A
120
Q

What are the units for Kw

A

Always mol^2 dm-6

121
Q

What happens to Kw when water is not pure and in a solution

A

It doesn’t matter whether water is pure or part of a solution - this equilibrium is always happening, and Kw is always the same at the same temperature

122
Q

How to find the pH of pure water

A

For pure water, there’s a 1:1 ratio of H+ and OH- ions due to dissociation.
This means [H+] = [OH-] and Kw = [H+]^2
So if you know Kw of pure water at a certain temperature, you can calculate [H+] and use this to find the pH

123
Q

What does changing the concentration of H+ or OH- do to Kw

A

Changing the concentration of [H+] or [OH-] in solution has no effect on the value of Kw as the equilibrium will shift, changing the concentration of the other substances to keep the value of Kw the same

124
Q

What does changing the temperature change the value of Kw

A

Changing the temperature of the solution changes the value of Kw
Because dissociation of water is an endothermic process
E.g. so warming the solution shifts the equilibrium to the right and Kw increases

125
Q

How to find the pH of strong base using Kw

A

NaOH and KOH are strong bases that fully ionise in water
E.g. NaOH —> Na+ + OH-
They donate one mole of OH- ions per mole of base.
This means that the concentration of OH- ions is the same as the concentration of the base.
E.g. for 0.02 mol dm-3 sodium hydroxide solution, [OH-] is also 0.02 mol dm-3
But to work out the pH you need to know [H+] which is linked to [OH-] through the ionic product of water (Kw = [H+][OH-])
If you know Kw and [OH] for a strong aqueous base at a certain temperature, you can work out [H+] and then the pH

126
Q

Find pH of NaOH

A
127
Q
A
128
Q
A
129
Q

How to find rate constant from zero order reaction concentration-time graph

A

Gradient of the line

130
Q

Definition of pH

A
131
Q
A
132
Q
A
133
Q
A
134
Q
A
135
Q
A
136
Q

What does p represent here and what is the pKw of water

A

p means you must do -log(x)
The Kw of water is 10^-14 mol2 dm-6
Meaning pKw = -log(10^-14) = 14

137
Q

Workout pH of water

A

pKw = -log(10^-14) = 14
Because Kw = [H+]x[OH-] = Kw = [H+]^2
So 10^-14 = [H+]^2 ∴ [H+] = 10^-7
So pH of water = -log(10^-7) = pH 7 (neutral)

138
Q

How to work out pOH

A
139
Q

Use pKw method

A
140
Q
A
141
Q
A
142
Q

Why do we use Ka

A

Weak acids don’t dissociate fully in solution so the [H+] isn’t the same as the acid concentration

143
Q

How to find equation for acid dissociation constant

A

For a weak aqueous acid you get the equilibrium: HA ⇌ H+ + A-
As only a small amount of of HA dissociates you can assume that:
[HA]»[H+] so [HA]start = [HA]equilibrium
Applying the equilibrium law you get:

144
Q

Why

A

This means you can assume that all the H+ ions in solution come from the acid, so [H+(aq)]eqm = [A-(aq)]eqm so [H+] x [A-] is the same as [H+]^2
This assumption only works for a weak acid in aqueous solution with nothing else added

145
Q

Units of Ka

A

Mol dm-3

146
Q

Why does the assumptions of Ka not apply to strong acids in solution with nothing else added

A

Stronger acids dissociate more in solution so different between [HA]start and [HA]eqm is more significant so the theory that they are equal no longer applies

147
Q

Rule with pKa and strength of acids

A

The smaller the pKa, the stronger the acid

148
Q
A
149
Q
A
150
Q
A
151
Q
A
152
Q
A
153
Q
A
154
Q
A
155
Q

What does a pH meter do

A

Electronic device that allows you to measure pH

156
Q

How to use a pH meter

A

Make sure to calibrate pH meter before measuring
Place bulb into distilled water and allow reading to settle then adjust it to make it read 7.0
Do the same with solutions of pH 4 and pH 10 (clean probe between each test with distilled water)
Place probe of pH meter in solution

157
Q

What do titrations tell you

A

Titrations let you find out exactly how much alkali is needed to neutralise a quantity of acid.

158
Q

How to find out the quantity of acid needed to neutralise alkali using values obtained from a titration

A

Plot the pH of the titration mixture against the amount of base added as the titration goes on.
The pH of the mixture can be measured using a pH meter and the scale on the burette can be used to see how much base has been added

159
Q

Draw the shape of the pH curve for the titrations of the following solutions:
Strong acid/strong base
Strong acid/ weak base
Weak acid/ strong base
Weak acid/ weak base

A
160
Q

What happens to the curve if you reverse the titration and titrate a base with an acid

A

The shapes remain the same but they’re reversed

161
Q

Why do the pH curve have their shape

A

The initial pH depends on the strength of the acid.
So a strong acid titration will start at a much lower pH than a weak acid.
To start with, addition of small amounts of base have little impact on the pH of the solution.
Most graphs (apart from weak acid/weak base) have a vertical area — this is the equivalence point or end point.
When this is the case, a tiny amount of base quickly causes a big change in pH.

162
Q

What types of bases do you need to add to acids of different strengths to cause a pH change

A

You need to add more weak base than strong base to a strong acid to cause a pH change, and the change is less pronounced.
On the other hand, you need to add less strong base to a weak acid to see a large change in pH.

163
Q

What does the final pH of a pH curve depend on

A

The final pH depends on the strength of the base
The stronger the base, the higher the final pH

164
Q

How to detect the end point

A

The end point of a titration is when the indicator changes colour
If a suitable indicator is used then the end point should coincide with the equivalence point

165
Q

What is the equivalence point

A

The equivalence point is when the moles of alkali added equals the moles of acid present - but the pH is not always 7 at the equivalence point

166
Q

What substances are indicators

A

Indicators are weak acids where the conjugate pairs HA and A- are different colours

167
Q

What makes an indicator change colour

A

At low pH: HA is the main species present.
At high pH: A- is the main species present.
The pH at which the colour changes varies depending on the indicator
As the pH of the solution changes in a titration, the equilibrium concentrations of the conjugate pairs will also change.
The colour will change depending on whether the indicator is mainly protonated or deprotonated

168
Q

How to pick the right indicator for a titration

A

The indicator must change colour exactly at the end point of titration.
Pick an indicator that changes colour over a narrow pH range that lies entirely on the vertical part of the pH curve.

169
Q

What indicator to use for a strong acid/strong base titration

A

Can use either methyl orange or phenolphthalein

170
Q

What indicator to use for a strong acid/weak base titration

A

Methyl orange
The pH changes rapidly across the range for methyl orange, but not for phenolphthalein

171
Q

What indicator to use for a weak acid/strong base titration

A

Phenolphthalein
The pH changes rapidly over phenolphthalein’s range but not over methyl orange’s.

172
Q

What indicator to use for a weak acid/weak base titration

A

For weak acid/weak base titrations there’s no sharp pH change
There aren’t any indicators you can use in weak acid/weak base titrations
Use a pH meter

173
Q

Methyl orange and phenolphthalein and the pH ranges they cover

A
174
Q

Method to solve calculation

A
175
Q

Characteristics of indicator to use here

A

For this titration, the curve is vertical between pH 8 and pH 11
So a very small amount of base will cause the pH to change from 8 to 11.
So you want an indicator that changes colour somewhere between pH 8 and pH 11.

176
Q

What is a buffer

A

A buffer is a solution that minimises changes in pH when small amounts of acid or base are added

177
Q

Features of how a buffer works

A

A buffer doesn’t stop the pH from changing completely
Only makes the changes very slight
Buffers only work for small amounts of acid or base
If you put too much in they won’t be able to work
You can get acidic buffers and basic buffers, (only need to know about the acidic)

178
Q

What do acidic buffers contain

A

Buffers are made by setting up an equilibrium between a weak acid and its conjugate base

179
Q

What is the pH of acidic buffers

A

Have a pH less than 7

180
Q

Two ways to set up an equilibrium to make acidic buffer

A

Mix a weak acid with the salt of its conjugate base
Mix an excess of weak acid with a strong alkali

181
Q

How to make an acidic buffer by mixing a weak acid with the salt of its conjugate base

A

E.g. ethanoic acid and sodium ethanoate:
The salt fully dissociates into its ions when it dissolves:
CH3COO-Na+ → CH3COO- + Na+
The weak acid only slightly dissociates:
CH3COOH ⇌ H+ + CH3COO-
Equilibrium set up between weak acid and conjugate base to make lots of undissociated weak acid and lots of conjugate base

182
Q

How to make an acidic buffer by mixing an excess of weak acid with a strong alkali

A

E.g. ethanoic acid and sodium hydroxide
All the base reacts with the acid:
CH3COOH + OH- -> CH3COO- + H2O
The weak acid was in excess so there’s still some left in solution once all the base has reacted
This acid slightly dissociates:
CH3COOH ⇌ H+ + CH3COO-
Equilibrium set up between weak acid and conjugate base to make lots of undissociated weak acid and lots of conjugate base

183
Q

What do both methods of making an acidic buffer achieve

A

Equilibrium set up between weak acid and conjugate base to make lots of undissociated weak acid and lots of conjugate base
Equilibrium solution contains:
Lots of undissociated acid (HA)
Lots of acid’s conjugate base (A-)
Enough H+ ions to make the solution acidic

184
Q

How do pH buffers resist the changes in pH

A

The conjugate pair controls the pH of the buffer solution
The conjugate base mops up an excess of H+ (increasing pH to normal)
Conjugate acid releases H+ if there’s too much base around (decreases pH to normal)

185
Q

What happens if a small amount of H+ ions are added to a solution with an acidic buffer in

A

H+ concentration increases.
Most of the extra H+ ions combine with CH3COO- ions to form CH3COOH.
Shifts the equilibrium to the left, reducing the H+ concentration to close to its original value.
So the pH doesn’t change much.
The larger number of CH3COO- ions make sure the buffer can cope with the addition of acid
As [CH3COOH]»[H+] the ratio of [CH3COOH]/[CH3COO-] remains roughly constant

186
Q

What happens if a small amount of alkali is added to a solution with an acidic buffer

A

E.g. NaOH is added so the OH- concentration increases
Most of the extra OH- ions react with H+ ions to form water
This removes H+ ions from the solution.
This causes more CH3COOH to dissociate to form H+ ions (not a problem as there is lots of CH3COOH)
This shifts the equilibrium to the right.
The H+ concentration increases until it’s close to its original value
So pH doesn’t change much
As [CH3COOH]»[H+] the ratio of [CH3COOH]/[CH3COO-] remains roughly constant

187
Q

What are basic buffers

A

Solution that resists changes in pH when small amounts of acid or alkali are added
Have a pH more than 7

188
Q

Ways to make a basic buffer

A

Solutions made from a mixture of weak base and one of its salts
E.g. ammonia and ammonium chloride
Mix excess weak base and a strong acid
E.g. excess ammonia and HCl acid

189
Q

What happens if a small amount of H+ ions are added to a solution with a basic buffer in

A

The added H+ is removed by reaction with OH-
So position of equilibrium shifts to the right and some NH3 reacts to replace OH-
The [NH3] falls slightly and the [NH4+] rises slightly
But as [NH3] and [NH4+]»[OH-] the ratio of [NH3]/[NH4+] remains roughly constant

190
Q

What happens if a small amount of OH- ions are added to a solution with a weak base in it

A

[OH-] increases so position of equilibrium shifts to the left
The added OH- is removed by reaction with NH4+ to form NH3
The [NH3] rises slightly and the [NH4+] falls slightly
But as [NH3] and [NH4+]»[OH-] the ratio of [NH3]/[NH4+] remains roughly constant

191
Q

Assumptions to make when finding pH of an acidic buffer

A

The salt of the conjugate base is fully dissociated, so assume that the equilibrium concentration of A- is the same as the initial concentration of the salt
HA is only slightly dissociated, so assume that its equilibrium concentration is the same as its initial concentration.
The conjugate base doesn’t only come from dissociation of the weak acid so [H+] is not equal to [A-]

192
Q
A
193
Q
A
194
Q

What pH is the blood in our bodies

A

Blood must be kept between pH 7.35 and 7.45

195
Q

How is pH controlled in the blood

A

pH is controlled using a carbonic acid-hydrogen carbonate buffer system.
Two equilibrium reactions occur

196
Q

How does the equilibrium of buffers work in the blood

A

Levels of H2CO3 are controlled by respiration
Breathing out CO2 reduces [CO2] and moves eqm to the right so the level of H2CO3 is reduced
The levels of HCO3 - are controlled by the kidneys with excess being excreted in the urine

197
Q

What happens if pH of blood goes below 7.35

A

Acidosis
Fatigue, shortness of breath, shock, death

198
Q

What happens if pH of blood goes above 7.45

A

Alkalosis
Muscle spasm, light headedness, nausea

199
Q

What happens if H+ ions are added to the blood

A

[H+(aq)] increases.
The equilibrium position shifts to the left
MoreH+(aq) ions react with the conjugate base, HCO3 - (aq)
This removes most of the H+ (aq) ions

200
Q

What happens if OH- ions are added to the blood

A

[OH-(aq)] increases.
The small concentration of H+(aq) ions reacts with the OH-(aq)
H+ + OH- -> H2O
Which shifts the equilibrium position to the right more H2CO3 dissociates to restore most of H+ ions

201
Q
A
202
Q

What happens at different pHs for an indicator to change colour

A

Low pH
H+ reacts with A- so eqm position shifts towards HA
[HA] increases
High pH
OH- reacts with HA/H+ eqm shifts towards A-
[A-] increases
End point contains equal amounts of HA and A-

A Level Chemistry (15 decks)

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