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Investigative Analytical Science > Titrations > Flashcards

Titrations Flashcards

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1
Q

What is Volumetric Analysis?

A

A method of analysis based on titration.

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2
Q

What does titration determine?

A

Amount of a particular substance A.

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3
Q

How does titration determine the amount of a particular substance A?

A

By adding a measured volume of a solution with a known concentration of B until the reaction is complete.

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4
Q

What plays a role in titration?

A

Volume measurements.

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5
Q

What is the key of titration?

A

Accurate measurement of volume.

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6
Q

What happens in titration?

A

Solution (titrant) is added from a burette to a solution in a flask (titrand) until it is shown that tritant reacted stoichiometrically with the titrand.

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7
Q

What do we need in titration?

A

A properly balanced equation.

Understanding of equations’ stoichiometry.

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8
Q

What are the standards?

A

Reagents of accurately known concentration (x+-y) units, used in volumetric analysis.

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9
Q

What happens in Primary standard?

A

Substance with sufficient purity can make a standard solution by weighing its quantity –> dissolving it –> diluting to known solution volume.

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10
Q

What happens in a secondary standard?

A

Solution with found concentration is compared against a primary standard to find the accurate concentration.

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11
Q

Where is titration carried out?

A

In a conical flask with liquid/dissolved sample.

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12
Q

How is titrant solution delivered?

A

Volumetrically.
Slowly.
With a burette.
Shaking to reaction flask.

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13
Q

How is the titrant delivery called?

A

Titration.

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14
Q

When is the titration complete?

A

When the equivalent titrant is added with the whole analyte, based on the equation.

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15
Q

How is the titrant completion called?

A

Equivalent point.

End point.

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16
Q

Are the equivalent and end points the same?

A

No.

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17
Q

What does the equivalent point tell us?

A

The volume of titrant needed to reach equivalent point.

Moles of titrant are used by the analyte later.

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18
Q

How is titration classified?

A

Based on reaction type used.

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19
Q

What happens in an Acid-base titration?

A

Acidic/basic titrant reacts with acidic/basic analyte.

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20
Q

What happens in redox titrations?

A

Titrant is oxidizing/reducing agent.

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21
Q

How do we work out a balanced redox equation?

A

Examine half reactions.

Balance them with electrons required to transferred.

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22
Q

How do we balance redox titrations?

A

Use 2 reactants –> end point –> check equivalent of titrant to titrand.

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23
Q

What happens in precipitation titrations?

A

Analyte and titrant react –> form precipitate.

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24
Q

What must the primary standard be?

A

100% pure.
Known purity.
Stable at drying oven temperatures.
Not hygroscopic = not absorb water when exposed to laboratory air.

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25
Q

What are some substances that absorb water when exposed to laboratory air and should not be used in as primary standards?

A

HCL.
Phosphorous pentoxide.
NaOH.

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26
Q

What must the reaction where primary standard takes place be?

A

Quantitative.

Fast.

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27
Q

Why must the reaction of primary standard be fast?

A

Small volumes are used.

No delay needed.

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28
Q

What kind of a formula weight must primary standards have?

A

A high formula weight.

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29
Q

Why must a primary standard have a high formula weight?

A

To not reduce significant figures in calculated result.

Give a reasonable amount.

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30
Q

What is commonly used for standardising acids like HCL?

A

Sodium carbonate.

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31
Q

What is used for bases like NaOH?

A

Potassium hydrogen phthalate.

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32
Q

When can chemicals which do not meet primary standard requirement be used as a standard?

A

Only after standardisation with a primary standard.

When accurate measurement of titrant will take place.

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33
Q

What is a secondary standard?

A

A compound with a purity from chemical analysis.

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34
Q

How is the secondary standard served?

A

As the reference material for a titrimetric analysis method.

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35
Q

What are the characteristics of an ideal standard solution for a titrimetric method?

A

Sufficiently stable to determine its concentration once.

React rapidly with analyte to minimise time required between reagent additions.

React completely with analyte to realise end points.

Undergo a selective reaction with analyte described by a balanced equation.

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36
Q

What are the common titrations reagents?

A

HCL.

NaOH.

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37
Q

What is the characteristic of HCL and NaOH as titration reagents?

A

They stay stable for long time.

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38
Q

What is HCL?

A

A gas.

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39
Q

What is the solution of HCL in water?

A

36%.

11.8M.

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40
Q

Why we cannot make an accurate solution of HCL?

A

Needs to be titrated against primary standard.

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41
Q

How is NaOH appearing?

A

As pellets.

Small beads.

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42
Q

What are the characteristics of NaOH?

A

Very hydroscopic.

Very difficult to weigh accurately.

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43
Q

What does any solution of NaOH prepared need to be?

A

Titrated against primary standard.

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44
Q

Why does any solution prepared of NaOH need to be titrated against primary standard?

A

To find the accurately concentration of it.

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45
Q

What colour will the acid/base solutions be in titration?

A

Colourless.

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46
Q

What do we use to visualise the equivalent point?

A

An indicator.

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47
Q

What is an indicator?

A

A chemical substance.

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48
Q

What does an indicator have?

A

One colour in acidic solution.

Different colour in basic solution.

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49
Q

What do we use to measure the end point?

A

The sudden colour change.

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50
Q

What is the indication point and end point to one another?

A

Close to each other.

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51
Q

What happens when an acid is in burette and base in flask?

A

Indicator does not change colour at equivalent point.

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52
Q

Why will be no difference in the acid volume added whichever indicator we choose?

A

Because the graph is steep/sharp.

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53
Q

What would make sense even if there is no difference no matter the indicator?

A

Titrate best possible colour with each indicator.

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54
Q

How is the graph of acid in burette and base in flask?

A

Equivalent point is in between phenolopthalene and methyl orange.
Colour change is roughly at pH7.

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55
Q

How does a general indicator like Hind behave?

A

Like a weak acid.

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56
Q

What happens to Hind and Ind- when Hind behaves like a weak acid?

A

Hind has one colour.

Ind- has different colour.

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57
Q

When do Hind and Ind have the same concentration?

A

At a particular point.

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58
Q

What is the equation of Hind and Ind?

A

Hind –> H+ + Ind-.

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59
Q

What is the pH equation when Hind and Ind are equal?

A

pH = pKind.

Unique for each indicator.

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60
Q

What does the pH direction change tell us?

A

Colour change on curve.

pH range of an indicator.

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61
Q

What will the indicators have at equilibrium/end point?

A

Sae colour.

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62
Q

What happens to the colour of indicator when pH range ends?

A

It fades.

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63
Q

What is known for phenolphthalene?

A

pKind has higher pH than equivalent point.

Smaller titration volume.

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64
Q

What is known about methyl orange?

A

It is at a lower pH than equivalent point.

Larger titration volume.

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65
Q

Where is the equivalent point of a titration?

A

Where 2 substances are mixed in equation proportions accurately.

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66
Q

What indicator do we need to choose?

A

The one that changes colour as close as possible to the equivalent point.

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67
Q

What does vary in each titration?

A

Equivalent point.

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68
Q

What does colour change as close to equivalent point as possible mean for an indicator?

A

Best match between equivalent point and titration point.

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69
Q

Which indicator would we choose in real world?

A

One with pH range on steepest curve part.

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70
Q

What is it almost impossible to measure in a titration with the best match between end point and equivalent point?

A

Difference.

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71
Q

What happens in the graph of strong acid vs weak base in flask with using phenolphthaline and methyl orange as indicators?

A 
Phenolphthalene:
End point far from equivalent point.
Incomplete titration.
True titrant amount needed to fully react with analyte unknown.
Big difference.

Methyl orange:
Best end point and equivalence point difference.
Very small difference.
Effectively same.

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72
Q

What happens in the graph of weak acid vs strong base in flask when using methyl orange and phenolphthalene?

A

Methyl orange:
Never get to end point.
pH never gets low enough to reach indicators range.
Never colour change.

Phenolopthalenine:
Best difference between end point and equivalence point.
Very small difference.
Effectively same.

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73
Q

What do titration graph show us?

A

Indicator choice is important.

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74
Q

What happens in a graph of weak acid vs weak base in flask?

A

Neither indicator is a use.
Phenolopthalene finishes changing before equivalent point.
Methyl orange does not change colour at all.
End point never reached.

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75
Q

What is the secret in indicator choice?

A

Choose indicator with a colour change close to pH 7.

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76
Q

What will an indicator with a colour change between 6-8 do?

A

Change very near/at equivalent point.

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77
Q

Which two indicators are the best choices?

A

Bromothymol blue: pKind=7.3.

Phenol red: pKind=7.4.

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78
Q

Why is bromothymol blue better?

A

Good pKind.
End point and equivalent point will be very close.
Colour change: deep blue to bright yellow = easy to observe.

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79
Q

What will we use as our primary standard if we want to find the standard HCl solution?

A

Sodium carbonate.

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80
Q

What is the first and most obvious way to start in titrimetric analysis?

A

A balanced equation.

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81
Q

Why the best way to start in titrimetric analysis is a balanced equation?

A

Because it can give us the stoichiometric relationship between the reactants.

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82
Q

How we use the primary standard after we prepare it?

A

To standardise an acid solution.

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83
Q

What do we get when we prepare a primary standard of HCl 0.1M?

A

2HCl + Na2CO3 –> 2NaCl + CO + H2O.

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84
Q

What is the ratio between 2HCl and Na2CO3 in the primary standard?

A

2:1.

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85
Q

What do we have to consider if we want to standardise HCl with a sodium carbonate solution?

A
  1. Want titration with sensible values.
  2. Remember stoichiometry 2:1 = HCl: carbonate.
  3. Choose an indicator.
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86
Q

What do we mean when we say ‘titration with sensible values’ ?

A
  1. Not a strong sodium carbonate solution to need a lot of HCl mL’s to reach point.
  2. Not a weak sodium carbonate solution to make titration very very small.
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87
Q

What do we mean when we say ‘remember stoichiometry 2:1 HCl: carbonate’?

A 
Want:
[carbonate] = 1/2 of [HCl].
If:
[HCl] = 0.1M --> 
[carbonate] = 0.050M.
Pipette:
25.0mL 0.050M carbonate in conical flask -->
Titration volume = HCl 25mL.
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88
Q

What do we mean when we say ‘choose an indicator’?

A

This is:
Titration of weak base with strong acid.
Equivalent point:
ideal, clear colour change at end point.
Equivalent point:
Close to end point –> not identify difference.

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89
Q

What do we need to make a standard sodium carbonate solution?

A

25mL for titration.

Repetitions for harmonic results.

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90
Q

How many repetitions should we do for the titration with 25mL?

A

3 x 25mL = 75mL.

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91
Q

What should we do before use the pipette?

A

Rinse it.

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92
Q

If [HCL] = 0.1M and then [carbonate] = 0.05M, how many grams of sodium carbonate should we use?

A

Sodium carbonate FM = 105.9g/mole.

0.05 moles x 105.9g/mole = 5.295 g for 100 mL solution.

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93
Q

What if we weigh a bit less than 0.5295g/mole?

A

Maybe 0.5255g/mole?
0.5255g / 105.9g/mole = 0.0049 moles / 0.1 L = 0.049M solution.
Not exact measurement, but close.

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94
Q

Why should we make these calculations?

A

To know exactly molarity.
Accurate weight.
Measure concentration.

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95
Q

How do we start titration after we do calculations of solutions will be used?

A

Pipette 25.0mL carbonate in conical flask.
Add 2 drops indicator.
Indicator = bromothymol blue.

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96
Q

What will we do after we prepare carbonate solution?

A

Prepare HCl solution.

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97
Q

How will we prepare HCl solution?

A

Fill burette with HCl.
1mL carbonate in test-tube.
HCl in another test-tube.
Add 1 indicator drop to each.

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98
Q

Why should we add 1 drop indicator to each test-tube of carbonate and HCl solution?

A

To recognise colours and change at end point.

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99
Q

After titration what should we do?

A

Repeat process one more time at least.

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100
Q

Why should we repeat process one more time at least?

A

To get > 0.1mL apart results.

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101
Q

What are our values in the end?

A

Harmonic = concordant.

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102
Q

What is the average volume in the end?

A

26.35mL.

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103
Q

Do we use units for titration?

A

No.

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104
Q 
How will we continue calculations for titration to find Mhcl if we know:
Balanced equation: 
2HCL + Na2CO3 --> 2NaCL + CO2 + H2O.
Stoichiometry = 2:1 HCl : carbonate.
Vhcl = 26.35mL.
Mcarb = 0.049M.
Vcarb = 25.0mL.
?
A 
Use equation:
2HCL + Na2CO3 --> 2NaCl + CO2 + H2O 
2/1 = (Mhcl x Vhcl) / (Mcarb x Vcarb)
Mhcl = (Mcarb x Vcarb) x 2 / (Vhcl x 1)
Mhcl = (0.049M x 25.0mL) x 2 / 26.35mL = 0.093M.
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105
Q

What can we find with our standardised HCl solution?

A

Molarity of unknown sodium hydroxide solution.

Think is = 0.1M

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106
Q

How can we find molarity of unknown sodium hydroxide solution?

A

Pipette 25.0mL hydroxide solution in conical flask.
Use any indicator –> bromothymol blue.
Titrate standardised acid 3 times.

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107
Q

Why can we use any indicator to find molarity of unknown sodium hydroxide solution?

A

It is a strong acid HCl vs strong base NaOH.

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108
Q

What do we do if we see data in 1st titration are over end point?

1: 25.15mL
2: 24.6mL
3: 24.6mL.

A

Reject them.

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109
Q

What is the average volume from 2 titrations?

2: 24.6mL.
3: 24.6mL.

A

24.60ml.

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110
Q

What is the average amount we find from titrations?

A

End point which matches equivalent point.

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111
Q

How is the titration called when the 3 titrations are different to each other?

A

Rough titration.

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112
Q 
How can we find Mnaoh then from our data?
Mhcl = 0.093M.
Vhcl = 24.60mL.
Mnaoh = ?
Vnaoh = 25.0mL.
A 
HCL + NaOH --> NaCl + H2O.
Stoichiometry 1:1.
1 / =
(Mhcl x Vhcl) / (Mnaoh x Vnaoh)
Mhaoh = (0.093M x 24.6mL) / 25.0mL = 0.091M.
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113
Q

What is a different way to measure the equivalent point?

A

pH titration curves.

114
Q

How can we use pH titration curves to measure the equivalent point in a different way?

A

Collect data.
pH.
Volume added in burette.
Intervals during titration.

115
Q

What is different in pH titration curves in measuring equivalent point?

A

Do not stop at end point.

116
Q

Why do not we stop at end point in pH titration curves?

A

We are not sure where it is.

117
Q

What do we do if we do not stop at end point in pH titration curves?

A

Continue to use data.

Plot graph.

118
Q

How is the graph in pH titration curves?

A

pH on y-axis.
Volume added on x-axis.
No indicator.

119
Q

On what do the graphs depend in pH titration curves?

A

On relative strength point of weak and acid.

120
Q

What can we do in a graph of a titration of strong acid (base) and strong base?

A

Add larger volumes near start of titration when pH rises slowly.

121
Q

Why can we add larger volumes near start of titration in strong acid v strong base?

A

Because there is still acid.

122
Q

When does pH starts to rise more rapidly?

A

When acid is consumed.

123
Q

What do we need to do when pH starts rising rapidly?

A

Collect more points.

Ass smaller volumes.

124
Q

Until when do we have to add smaller volumes to get more points?

A

Until rise rate in pH slows.

125
Q

What can we add after rising of Ph slows?

A

Larger volumes.

126
Q

What does the adding of larger volumes give us?

A

More data points in either side of equivalent point.

127
Q

Where do the data points in either side of equivalent point in graph help us?

A

To find equivalent point easier.

128
Q

What happens in a weak acid (flask) v strong base graph of titration?

A
  1. Start of graph –> rapid rise in pH as base is added.
  2. Buffer region = buffer with acetic acid and sodium acetate produced –> pH slows down.
  3. equivalent point approached –> Ph rise rapidly.
  4. 2nd buffering region = after equivalent point –> pH rate decreases, no acid left, more basic solution.
  5. Solution is very basic –> pH does not increase.
129
Q

Why pH slows down when buffer is added?

A

Buffers resist pH changes.

Add small amounts of base/acid.

130
Q

What can we work out of pH titration curves?

A

Where equivalent point is.

131
Q

How can we find the equivalent point on a pH titration curve?

A
  1. Make lines through data above and below steepest part of curve.
  2. equivalence point = volume at mid point between 2 lines on curve.
  3. estimate values.
132
Q

Why can estimation of equivalence point in pH titration curves is tricky?

A

It is a matter of judgement.

But manageable.

133
Q

How is it better to estimate equivalence point from a pH titration curve and why?

A

In pencil, because it is easier to correct.

134
Q

How are all these examples of finding equivalent point called?

A

Direct titration.

135
Q

Why sometimes finding equivalent point it is not simple?

A
  1. might have a strong base and weak acid or vice versa.

2. not suitable indicator –> cannot see end point.

136
Q

What is a very weak acid?

A

Benzoic acid.

137
Q

Which is an alternative technique to find equivalent point in titrations?

A

Back titration.

138
Q

What happens in a back titration?

A
  1. Intermediate indicator added in excess.
  2. excess of reactant determined by titration of unreacted intermediate with suitable titrant.
  3. find how much left from first titration.
139
Q

What can we find in back titrations if we know how much we added and excess from first trial?

A

Original concentration we use.

140
Q

What happens if we add strong base with weak acid?

A

All acid is consumed.

Some base is left.

141
Q

How do we call the left over of the base in a titration?

A

Excess.

142
Q

Why is it a little tricky to measure weak acid by direct titration?

A

it is weak.

143
Q

What do we do if we have a weak acid?

A
  1. Add excess strong base (NaOH).

2. ‘back titrate’ unreacted NaOH with strong acid (HCl).

144
Q

What is the equation of strong base and strong acid?

A

NaOH + AcOH –> AcONa + H2O.

145
Q

What is the ration in

NaOH + AcONa + H2O?

A

1:1.

146
Q

What do we do after we find the equation of strong base with strong acid?

A

Think of acid’s molarity.

147
Q

Why should we think about acid’s molarity?

A

To have an idea of how much NaOH to use and how much to left behind.

148
Q

How much do we think acid’s molarity is?

A

0.03-0.05M.

149
Q

How much will the moles of acetic acid be if we take 25.0mL?

A

(0. 03M/1000) x 25mL = 7.5 X 10-4 moles.

(0. 05M/1000) x 25mL = 1.25 x 10-3 moles.

150
Q

What do we need after we cunt acid’s molarity with 25mL?

A

Excess of NaOH.

151
Q

What do we to find NaOH molarity?

A

Add 25mL of 0.091M.

(0.091m/1000) x 25mL = 2.275 X 10-3 moles.

152
Q

What have we done with NaOH molarity?

A

Added twice NaOH moles (thought).

153
Q

What do we realise from acid’s and base’s molarities in the end?

A

We have much more NaOH than acetic acid.

Will be amount of NaOH unreacted by acetic AcOH.

154
Q

What can we find if we know we have added:
2.275 x 10-3 moles NaOH
used:
7.5 x 10-4 moles - 1.25 x 10-3 moles of NaOH –> react with acid.
?

A

How much is left over when titrating it with HCL (0.093M).

155
Q

What is the equation of NaOH titrating with HCl?

A

HCl + NaOH –> NaCl + H2O.

156
Q

What is the ratio of NaOH and HCl titration?

A

1:1.

157
Q

What is the equation of NaOH with HCl?

A

Strong acid v strong base reaction.

158
Q

Which indicator is good for a strong acid v strong base reaction?

A

Any.

159
Q

Which is an indicator we could use for a strong base v strong acid reaction?

A

Phenolphthalein.

160
Q

What happens in colour changing when we use phenolphthalein as an indicator?

A

Pink –> colourless.

161
Q

What happens in colour changing when reaching end point and equivalent point if we use HCl as a titrant and excess NaOH as a titrand?

A

Colourless –> pink.

162
Q

How many times and what results do we get if we do this back titration?

A 
3 times.
Volumes:
10.4mL.
10.2mL.
10.2mL.
Average = 10.16mL.
163
Q

What is the number of HCl moles then?

A

(0.093M/1000) x 10.16mL = 9.45 x 10-4 mol.

164
Q

With what is the number of HCl moles same?

A

With unreacted NaOH moles number.

165
Q

What is the equation and calculations of moles in a back titration then?

A

Total moles NaOH - moles NaOH used in titration = moles NaOH used by acetic acid.

  1. 275 x 10-3 moles - 9.45 x 10-4 moles =
  2. 33 x 10-3 moles NaOH used by acetic acid.
166
Q

What is then the moles of acid based on 1:1 NaOH:AcOH stoichiometry?

A

1.33 X 10-3 moles.

167
Q

How much is the acid molariy?

A

Contained in 25.0mL –> 0.025L.
Molarity of acid = 1.33 x 10-3 moles / 0.025L =
0.053M.

168
Q

On what is a redox titration based?

A

On an oxidation-reduction reaction.

169
Q

Hoe does the oxidation-reduction reaction work?

A

In pairs.

One is oxidised, one is reduced.

170
Q

Where does the oxidation-reduction reaction occur?

A

Between titrant and analyte.

171
Q

What can we identify with the oxidation-reduction reaction?

A

Unknown analytes’ concentration.

172
Q

How can we identify unknown analytes’ concentration from an oxidation-reduction reaction?

A

By knowing oxidant’s concentration –> measuring reductant’s concentration.

173
Q

What do we monitor in the oxidation-reduction titrations?

A

The reaction potential = voltage.

174
Q

What we do not monitor in the oxidation-reduction titrations?

A

Concentration of reacting species.

175
Q

Why we do not monitor concentration of reacting species in oxidation-reduced titrations?

A

Because they are both colourless.

176
Q

What do we do after we monitor the voltage in the oxidation-reduction titrations?

A

Get data –> plot a graph like pH graph titration.

177
Q

How is the graph of oxidation-reduction titrations?

A

Volume of titrant from burette v voltage.
Curve.
Voltage rising –> achieve equivalent point.

178
Q

When is the voltage graph of oxidation-reduction titrations useful?

A

When there is no other way to observe equivalence point = end point.

179
Q

What happens in oxidation-reduction titrations?

A

Electrons are transferred during redox reaction.

180
Q

How is potassium permanganate characterised?

A

Self indicating.

181
Q

What does it mean that potassium permanganate is self indicating?

A

Its colour changes during oxidation.
Its colour stays purple during reduction.
–> It shows equivalent point.

182
Q

How is the colour of potassium permanganate characterised and why?

A

Deep purple.

Due to permanganate ion MnO-4.

183
Q

What agent is permanganate?

A

Strong oxidising agent.

184
Q

What happens to permanganate when it reacts with a reducing agent Mn (+7)?

A

It is reduced to Mn (+2).

MnO-4 (purple) –> Mn2+ (colourless).

185
Q

Where does the permanganate reacting with reducing agent is useful?

A

When reducing agents are:
Fe (2).
Oxalate anion.

186
Q

What happens to oxalate anion?

A

Carbon is oxidised from C (+3) to C (+4).

C2O4 2- (colourless) –> CO2 (a gas).

187
Q

What must we know before a titration can be performed?

A

The stoichiometric relationship between oxidant and reductant = balanced equation.

188
Q

What is the reaction between MnO4- + C2O4 2-?

A

MnO4- + C2O4 2- –> Mn2+ + CO2.

189
Q

What do we have to do the reaction to have a balanced redox equation?

A

Split it into a reduction reaction and an oxidation reaction.
1/2 reactions.

190
Q

What are the 1/2 reactions of MnO4- + C2O4 2- –> Mn2+ + CO2 to get a balanced redox equation?

A

MnO4- + 8H+ + 5e- –> Mn2+ + 4H2O.

C2O4 2- –> 2CO2 + 2e-.

191
Q

What do we do to the 1/2 equations?

A

Multiply first by 2.

Multiply second by 5.

192
Q

Why do we balance 1/2 equations?

A

To balance electron transfer.

193
Q

What is the balanced redox equation in the end?

A

2MnO4 - + 5C2O4 2- + 16H+ –> 2Mn2+ + 8H2O + 10CO2.

194
Q

What is the stoichiometry of the balanced redox equation between MnO4- and C2O4 2-?

A

2:5.

195
Q

What do we use as primary standard to start the oxidation-reduction titration?

A

Sodium oxalate.

Or oxalic acid.

196
Q

What can oxalic acid give us?

A

Oxalate.

197
Q

What do we have to do when we want to start the oxidation-reduction titration and use oxalate as primary standard?

A

Prepare 0.05M oxalate solution.

198
Q

Why do we prepare 0.05M oxalate solution?

A

To standardise 0.02M KMno4 solution.

199
Q

Why do we need to standardise permanganate?

A

Because it is not a primary standard.

We do not know exact grams we need to prepare solution.

200
Q

What will we do with the 0.05M oxalate solution we prepared?

A

Place 25.0mL in a conical flask.

201
Q

What is a requirement in the oxidation-reduction titrations?

A

H+.

202
Q

Why do we need H+ for the oxidation-reduction titrations?

A

To make the solution acidic.

203
Q

What do we add after we put 25.0mL oxalate solution in a conical flask?

A

25mL strong acid solution = 1M H2SO4.

204
Q

How is the reaction of oxalate and H2SO4 characterised?

A

Slow.

205
Q

What do we have to do to the oxalate + H2SO4 reaction as it is slow?

A

Heat the flask to 80 degrees.

206
Q

What else do we add in the oxalate + H2SO4 reaction?

A

Small quantity of KMnO4 from burette.

207
Q

What colour does KMnO4 have in the beginning?

A

Purple.

208
Q

What happens to the purple colour after a while KMnO4 is added to the oxalate + H2SO4 reaction?

A

It disappears.

209
Q

What else happens except from the colour disappearance, when KMnO4 is added in the oxalate + H2SO4 reaction?

A

Some Mn2+ are created.

210
Q

How do Mn2+ act?

A

As a catalyst.

211
Q

What happens to the reaction when produced Mn2+ act as a catalyst?

A

It speeds up.

Purple colour stays.

212
Q

When do we recognise colour change and end point = equivalent point in the oxalate + H2SO4 + KMnO4 reaction?

A

When keep adding rapidly KMnO4 until purple colour just stays visible.

213
Q

What happens if we add permanganate slowly?

A

Flask cools down.

Reaction slows down.

214
Q

What do we do when we are unsure if we reached end point?

A

Warm flask bit more.
Pink colour disperse.
Add permanganate.

215
Q

What shall we do when we are halfway to end point?

A

Warm again.
Will not work.
Pink colour still here.

216
Q

What happens when oxalate is oxidised?

A

CO2 produced.
Solution fizzes.
Bubbles away.

217
Q

What do we know from our oxidation-reduction titration so far and what do we need to find?

A 
Know:
Mox = 0.050M.
Vox = 25.0mL.
Vmn = 23.8mL.
Want:
Mmn = ?
218
Q

What is the equation between oxalate + H2SO4 + KMnO4?

A

2MnO4- + 5C2O4 2- + 16H+ –> 2Mn2+ + 8H2O + 10CO2.

219
Q

What is the ration of the oxalate + H2SO4 + KMn04 reaction?

A

5:2.

220
Q

How can we find Mmn?

A

5/2 = (Mox x Vox ) /(Mmn x Vmn)
Mmn = (2 x (Mox x Vox) )/ (5 x Vmn)
Mmn = (2 x 0.05M x 25mL) / (5 X 23.8mL)
= 0.021M.

221
Q

What can we analyse by using KmnO4 standardised solution?

A

Any oxalate solution.

Any reductant contained solution = Fe2+.

222
Q

What happens to permanganate over time?

A

It degrades.

223
Q

Where is permanganate normally kept?

A

In brown bottle.

Dark cupboard.

224
Q

Why is permanganate usually kept in a brown bottle?

A

To reduce light in solution.

Prolong its usefulness.

225
Q

When does permanganate solution go off?

A

When Mn4+ is produced –> insoluble precipitate of manganese dioxide produced –> brown bottle gets coloured.

226
Q

How long does permanganate solution last?

A

For weeks.

227
Q

What is the advantage of permanganate solution lasting for weeks?

A

It does not need to be standardised often.

228
Q

How can we analyse oxalate solutions by using KnO4 standardised solution?

A

Oxalate solution has an unknown concentration.

Carry out 3 titrations.

229
Q

What are the 3 titrations we can do to analyse oxalate concentration?

A
  1. Prepare KMnO4 to standardise oxalate, add KMnO4 in a flask.
  2. Add strong acid solution in flask.
  3. Add small quantity of oxalate solution from burette.
230
Q

What can be the strong acid we will use in the titration and why?

A

1M sulphuric solution.

To have plenty H+ ions.

231
Q

What do we know and what do we need to find?

A 
Know:
Mmn = 0.021M.
Vox = 25.0mL.
Vn = 16.8mL.
To find:
Mox = ?
232
Q

What is the equation between MnO4- and C2O4 2-?

A

2MnO4- + 5C2O4 2- + 16H+ –> 2Mn2+ + 8H2O + 10CO2.

233
Q

What is the ratio of the 2MnO4- + 5C2O4 2- equation?

A

5:2.

234
Q

What calculations do we have to do to find Mox based on the known values?

A 
5/2 = (Mox x Vox) / (Mmn x Vmn)
Mox = (5 x Mmn x Vmn) / (2 x Vox)
Mox = (5 x 0.021M x 16.8mL) / (2 x 25.0mL)
Mox = 0.0353M.
235
Q

Where do all the calculation rely?

A

On a properly balanced equation.

236
Q

What happens if we get the balanced equation right?

A

Rest fall into place.

237
Q

What are most ionic compound to some degree?

A

Soluble.

Insoluble.

238
Q

How is the degree of solubility of a particular compound expressed?

A

Ksp = solubility product.

239
Q

What can we get if we combine solutions with different ions?

A

Insoluble solutions.

240
Q

What is the relations between value and solubility of compound?

A

Larger volume –> more soluble compound.

241
Q

How soluble are NaCl and AgCl at 20 degrees?

A

Ksp NaCl = 37.66 –> soluble.

Ksp AgCl = 1.77 x 10-10 –> insoluble.

242
Q

How solubility can change based on temperature?

A

When temperature changes so doe s solubility.

243
Q

Where is solubility used?

A

In precipitation titrations.

244
Q

What do we have to do to find the concentration of a solution that contains CL- ions?

A

Titrate with a standardised solution which soluble, contains silver ions.

245
Q

Why can we use silver ions as a standard solution in titration?

A

Because they do not cause any precipitation issues.

246
Q

What is the equation between NaCl and AgCl solutions?

A

AgNO3(aq) + NaCl(aq) –> AgCl(s) + NaNO3(aq).

247
Q

Why can we no identify the end point pf the NaCl + AgCl reaction?

A

Both solutions are colourless.

248
Q

How will we identify the end point of NaCl + AgCl reaction?

A

When chloride is less and less in solution until there is none left.

249
Q

How can we observe the end point in precipitation titrations?

A

With electrochemical methods.

With special indicators.

250
Q

What is a special indicator we can use to observe the end point in precipitation titrations?

A

Mohr method.

251
Q

Which indicator is used in Mohr method to observe the end point in a precipitated titration?

A

CrO4 2- ion.

252
Q

What colour does CrO4 2-, chromate ion have?

A

Bright yellow.

253
Q

What happens in a titration when chromate ion is added?

A

Chromate ion added to chloride solution –> yellow colour –> titration starts.

254
Q

What happens as the titration with chromate ion proceeds?

A

White/grey silver chloride is formed.

255
Q

What happens at the end point when chromate ion is added to a chloride solution titration?

A

All chloride is used up.

Deep red precipitate of silver chromate (Ag2CrO4) is formed.

256
Q

What happens during the titration when chromate ion is added too chloride solution and silver chromate is formed?

A

AgCl + ag2CrO4 compete.

257
Q

What is the difference between AgCr04 and AgCl?

A

Ag2CrO4 is more soluble.

Only forms when all chloride is used up.

258
Q

What is the issue in the titration of using chromate with chloride solution?

A

Ag2CrO4 is formed but it needs Ag solution.

End point happens after equivalent point.

259
Q

How can we overcome the issue that end point comes after equivalent point in the titration of using chromate and chlorate?

A

Add 1 drop chromate indicator.

Carry out ‘blank’ titration with no chloride.

260
Q

What will the use of chromate only tell us?

A

How much to remove from titration volume to get better value equivalent point value.

261
Q

What do we must control to make sure end point wont be different form equivalent point?

A

pH.

262
Q

What does it form in a basic solution?

A

AgOH precipitate.

Chloride.

263
Q

What can precipitate increase if formed?

A

Titration volume.

264
Q

Why does precipitate increases titration volume when it is formed?

A

Because precipitate it is formed as well as chloride –> titrating 2 things.

265
Q

What happens to chromate in an acidic solution?

A

Dimerizes –> becomes dichromate.
Cr2O7 2-.
Froms AgCr2O7.

266
Q

Why will we not realise the end point in an acidic solution when chromate is involved?

A

Chromate dimerizes.
Ag2Cr2O7 formed.
Ag2Cr2O soluble.
No precipitate.

267
Q

Why must we control pH in a titration?

A

It is important to observe end point.

268
Q

What does Fajans method in precipitated titrations use?

A

Species that changes colour when absorbs precipitate.

269
Q

What happens in Fajans method that improves precipitated titrations?

A

Anionic dye dichlorofluorescein added to titrand’s solution in conical flask.

270
Q

What do we realise in the Fajans method before the end point?

A

Precipitate of AgCl has negative charge.

271
Q

Why does AgCl has negative charge?

A

It absorbs excess Cl-.

Dichlorofluorescein carries negative charge –> pushed back by precipitate –> remains in solution.

272
Q

What colour does dichlorofluorescein have?

A

Greenish-yellow = faint glow.

273
Q

What happens in the Fajans method after the end point?

A

Precipitate carries positive charge.

274
Q

Why does precipitate carry positive charge after end point in Fajans method?

A

It absorbs excess Ag+.

Dichlorofluorescein absorbs to precipitate = attracted negative charge.

275
Q

What happens when the precipitate has a positive charge after the end point?

A

Changes colour = yellow –> pink.

276
Q

What does the colour of precipitate indicate in Fajans method?

A

End point.

277
Q

What do the precipitation titrations involve?

A

Ag ions.

278
Q

How are the precipitation titrations that use Ag ions named?

A

‘Argentometric titrations’.

279
Q

What do Ag ions do in the precipitation titrations when they are used?

A

Analyse other ions.

280
Q

Why do Ag ions analyse other ions when used in precipitation titrations?

A

They are soluble.

281
Q

How are the precipitation titrations called?

A

‘Direct titrations’.

282
Q

How are the calculations of precipitation titrations?

A

Same with acid/base and redox titrations.
Balanced equation –> stoichiometry –> known silver’s standard solution and volumes of titrand in flask (analyte) and titrant in burette –> do calculations.

Investigative Analytical Science (16 decks)

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