Titrations Flashcards
(282 cards)
1
Q
What is Volumetric Analysis?
A
A method of analysis based on titration.
2
Q
What does titration determine?
A
Amount of a particular substance A.
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.
4
Q
What plays a role in titration?
A
Volume measurements.
5
Q
What is the key of titration?
A
Accurate measurement of volume.
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.
7
Q
What do we need in titration?
A
A properly balanced equation.
Understanding of equations’ stoichiometry.
8
Q
What are the standards?
A
Reagents of accurately known concentration (x+-y) units, used in volumetric analysis.
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.
10
Q
What happens in a secondary standard?
A
Solution with found concentration is compared against a primary standard to find the accurate concentration.
11
Q
Where is titration carried out?
A
In a conical flask with liquid/dissolved sample.
12
Q
How is titrant solution delivered?
A
Volumetrically.
Slowly.
With a burette.
Shaking to reaction flask.
13
Q
How is the titrant delivery called?
A
Titration.
14
Q
When is the titration complete?
A
When the equivalent titrant is added with the whole analyte, based on the equation.
15
Q
How is the titrant completion called?
A
Equivalent point.
End point.
16
Q
Are the equivalent and end points the same?
A
No.
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.
18
Q
How is titration classified?
A
Based on reaction type used.
19
Q
What happens in an Acid-base titration?
A
Acidic/basic titrant reacts with acidic/basic analyte.
20
Q
What happens in redox titrations?
A
Titrant is oxidizing/reducing agent.
21
Q
How do we work out a balanced redox equation?
A
Examine half reactions.
Balance them with electrons required to transferred.
22
Q
How do we balance redox titrations?
A
Use 2 reactants –> end point –> check equivalent of titrant to titrand.
23
Q
What happens in precipitation titrations?
A
Analyte and titrant react –> form precipitate.
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.
25
What are some substances that absorb water when exposed to laboratory air and should not be used in as primary standards?
HCL.
Phosphorous pentoxide.
NaOH.
26
What must the reaction where primary standard takes place be?
Quantitative.
| Fast.
27
Why must the reaction of primary standard be fast?
Small volumes are used.
| No delay needed.
28
What kind of a formula weight must primary standards have?
A high formula weight.
29
Why must a primary standard have a high formula weight?
To not reduce significant figures in calculated result.
| Give a reasonable amount.
30
What is commonly used for standardising acids like HCL?
Sodium carbonate.
31
What is used for bases like NaOH?
Potassium hydrogen phthalate.
32
When can chemicals which do not meet primary standard requirement be used as a standard?
Only after standardisation with a primary standard.
| When accurate measurement of titrant will take place.
33
What is a secondary standard?
A compound with a purity from chemical analysis.
34
How is the secondary standard served?
As the reference material for a titrimetric analysis method.
35
What are the characteristics of an ideal standard solution for a titrimetric method?
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.
36
What are the common titrations reagents?
HCL.
| NaOH.
37
What is the characteristic of HCL and NaOH as titration reagents?
They stay stable for long time.
38
What is HCL?
A gas.
39
What is the solution of HCL in water?
36%.
| 11.8M.
40
Why we cannot make an accurate solution of HCL?
Needs to be titrated against primary standard.
41
How is NaOH appearing?
As pellets.
| Small beads.
42
What are the characteristics of NaOH?
Very hydroscopic.
| Very difficult to weigh accurately.
43
What does any solution of NaOH prepared need to be?
Titrated against primary standard.
44
Why does any solution prepared of NaOH need to be titrated against primary standard?
To find the accurately concentration of it.
45
What colour will the acid/base solutions be in titration?
Colourless.
46
What do we use to visualise the equivalent point?
An indicator.
47
What is an indicator?
A chemical substance.
48
What does an indicator have?
One colour in acidic solution.
| Different colour in basic solution.
49
What do we use to measure the end point?
The sudden colour change.
50
What is the indication point and end point to one another?
Close to each other.
51
What happens when an acid is in burette and base in flask?
Indicator does not change colour at equivalent point.
52
Why will be no difference in the acid volume added whichever indicator we choose?
Because the graph is steep/sharp.
53
What would make sense even if there is no difference no matter the indicator?
Titrate best possible colour with each indicator.
54
How is the graph of acid in burette and base in flask?
Equivalent point is in between phenolopthalene and methyl orange.
Colour change is roughly at pH7.
55
How does a general indicator like Hind behave?
Like a weak acid.
56
What happens to Hind and Ind- when Hind behaves like a weak acid?
Hind has one colour.
| Ind- has different colour.
57
When do Hind and Ind have the same concentration?
At a particular point.
58
What is the equation of Hind and Ind?
Hind --> H+ + Ind-.
59
What is the pH equation when Hind and Ind are equal?
pH = pKind.
| Unique for each indicator.
60
What does the pH direction change tell us?
Colour change on curve.
| pH range of an indicator.
61
What will the indicators have at equilibrium/end point?
Sae colour.
62
What happens to the colour of indicator when pH range ends?
It fades.
63
What is known for phenolphthalene?
pKind has higher pH than equivalent point.
| Smaller titration volume.
64
What is known about methyl orange?
It is at a lower pH than equivalent point.
| Larger titration volume.
65
Where is the equivalent point of a titration?
Where 2 substances are mixed in equation proportions accurately.
66
What indicator do we need to choose?
The one that changes colour as close as possible to the equivalent point.
67
What does vary in each titration?
Equivalent point.
68
What does colour change as close to equivalent point as possible mean for an indicator?
Best match between equivalent point and titration point.
69
Which indicator would we choose in real world?
One with pH range on steepest curve part.
70
What is it almost impossible to measure in a titration with the best match between end point and equivalent point?
Difference.
71
What happens in the graph of strong acid vs weak base in flask with using phenolphthaline and methyl orange as indicators?
```
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.
72
What happens in the graph of weak acid vs strong base in flask when using methyl orange and phenolphthalene?
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.
73
What do titration graph show us?
Indicator choice is important.
74
What happens in a graph of weak acid vs weak base in flask?
Neither indicator is a use.
Phenolopthalene finishes changing before equivalent point.
Methyl orange does not change colour at all.
End point never reached.
75
What is the secret in indicator choice?
Choose indicator with a colour change close to pH 7.
76
What will an indicator with a colour change between 6-8 do?
Change very near/at equivalent point.
77
Which two indicators are the best choices?
Bromothymol blue: pKind=7.3.
| Phenol red: pKind=7.4.
78
Why is bromothymol blue better?
Good pKind.
End point and equivalent point will be very close.
Colour change: deep blue to bright yellow = easy to observe.
79
What will we use as our primary standard if we want to find the standard HCl solution?
Sodium carbonate.
80
What is the first and most obvious way to start in titrimetric analysis?
A balanced equation.
81
Why the best way to start in titrimetric analysis is a balanced equation?
Because it can give us the stoichiometric relationship between the reactants.
82
How we use the primary standard after we prepare it?
To standardise an acid solution.
83
What do we get when we prepare a primary standard of HCl 0.1M?
2HCl + Na2CO3 --> 2NaCl + CO + H2O.
84
What is the ratio between 2HCl and Na2CO3 in the primary standard?
2:1.
85
What do we have to consider if we want to standardise HCl with a sodium carbonate solution?
1. Want titration with sensible values.
2. Remember stoichiometry 2:1 = HCl: carbonate.
3. Choose an indicator.
86
What do we mean when we say 'titration with sensible values' ?
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.
87
What do we mean when we say 'remember stoichiometry 2:1 HCl: carbonate'?
```
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.
```
88
What do we mean when we say 'choose an indicator'?
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.
89
What do we need to make a standard sodium carbonate solution?
25mL for titration.
| Repetitions for harmonic results.
90
How many repetitions should we do for the titration with 25mL?
3 x 25mL = 75mL.
91
What should we do before use the pipette?
Rinse it.
92
If [HCL] = 0.1M and then [carbonate] = 0.05M, how many grams of sodium carbonate should we use?
Sodium carbonate FM = 105.9g/mole.
| 0.05 moles x 105.9g/mole = 5.295 g for 100 mL solution.
93
What if we weigh a bit less than 0.5295g/mole?
Maybe 0.5255g/mole?
0.5255g / 105.9g/mole = 0.0049 moles / 0.1 L = 0.049M solution.
Not exact measurement, but close.
94
Why should we make these calculations?
To know exactly molarity.
Accurate weight.
Measure concentration.
95
How do we start titration after we do calculations of solutions will be used?
Pipette 25.0mL carbonate in conical flask.
Add 2 drops indicator.
Indicator = bromothymol blue.
96
What will we do after we prepare carbonate solution?
Prepare HCl solution.
97
How will we prepare HCl solution?
Fill burette with HCl.
1mL carbonate in test-tube.
HCl in another test-tube.
Add 1 indicator drop to each.
98
Why should we add 1 drop indicator to each test-tube of carbonate and HCl solution?
To recognise colours and change at end point.
99
After titration what should we do?
Repeat process one more time at least.
100
Why should we repeat process one more time at least?
To get > 0.1mL apart results.
101
What are our values in the end?
Harmonic = concordant.
102
What is the average volume in the end?
26.35mL.
103
Do we use units for titration?
No.
104
```
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.
?
```
```
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.
```
105
What can we find with our standardised HCl solution?
Molarity of unknown sodium hydroxide solution.
| Think is = 0.1M
106
How can we find molarity of unknown sodium hydroxide solution?
Pipette 25.0mL hydroxide solution in conical flask.
Use any indicator --> bromothymol blue.
Titrate standardised acid 3 times.
107
Why can we use any indicator to find molarity of unknown sodium hydroxide solution?
It is a strong acid HCl vs strong base NaOH.
108
What do we do if we see data in 1st titration are over end point?
1: 25.15mL
2: 24.6mL
3: 24.6mL.
Reject them.
109
What is the average volume from 2 titrations?
2: 24.6mL.
3: 24.6mL.
24.60ml.
110
What is the average amount we find from titrations?
End point which matches equivalent point.
111
How is the titration called when the 3 titrations are different to each other?
Rough titration.
112
```
How can we find Mnaoh then from our data?
Mhcl = 0.093M.
Vhcl = 24.60mL.
Mnaoh = ?
Vnaoh = 25.0mL.
```
```
HCL + NaOH --> NaCl + H2O.
Stoichiometry 1:1.
1 / =
(Mhcl x Vhcl) / (Mnaoh x Vnaoh)
Mhaoh = (0.093M x 24.6mL) / 25.0mL = 0.091M.
```
113
What is a different way to measure the equivalent point?
pH titration curves.
114
How can we use pH titration curves to measure the equivalent point in a different way?
Collect data.
pH.
Volume added in burette.
Intervals during titration.
115
What is different in pH titration curves in measuring equivalent point?
Do not stop at end point.
116
Why do not we stop at end point in pH titration curves?
We are not sure where it is.
117
What do we do if we do not stop at end point in pH titration curves?
Continue to use data.
| Plot graph.
118
How is the graph in pH titration curves?
pH on y-axis.
Volume added on x-axis.
No indicator.
119
On what do the graphs depend in pH titration curves?
On relative strength point of weak and acid.
120
What can we do in a graph of a titration of strong acid (base) and strong base?
Add larger volumes near start of titration when pH rises slowly.
121
Why can we add larger volumes near start of titration in strong acid v strong base?
Because there is still acid.
122
When does pH starts to rise more rapidly?
When acid is consumed.
123
What do we need to do when pH starts rising rapidly?
Collect more points.
| Ass smaller volumes.
124
Until when do we have to add smaller volumes to get more points?
Until rise rate in pH slows.
125
What can we add after rising of Ph slows?
Larger volumes.
126
What does the adding of larger volumes give us?
More data points in either side of equivalent point.
127
Where do the data points in either side of equivalent point in graph help us?
To find equivalent point easier.
128
What happens in a weak acid (flask) v strong base graph of titration?
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
Why pH slows down when buffer is added?
Buffers resist pH changes.
| Add small amounts of base/acid.
130
What can we work out of pH titration curves?
Where equivalent point is.
131
How can we find the equivalent point on a pH titration curve?
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
Why can estimation of equivalence point in pH titration curves is tricky?
It is a matter of judgement.
| But manageable.
133
How is it better to estimate equivalence point from a pH titration curve and why?
In pencil, because it is easier to correct.
134
How are all these examples of finding equivalent point called?
Direct titration.
135
Why sometimes finding equivalent point it is not simple?
1. might have a strong base and weak acid or vice versa.
| 2. not suitable indicator --> cannot see end point.
136
What is a very weak acid?
Benzoic acid.
137
Which is an alternative technique to find equivalent point in titrations?
Back titration.
138
What happens in a back titration?
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
What can we find in back titrations if we know how much we added and excess from first trial?
Original concentration we use.
140
What happens if we add strong base with weak acid?
All acid is consumed.
| Some base is left.
141
How do we call the left over of the base in a titration?
Excess.
142
Why is it a little tricky to measure weak acid by direct titration?
it is weak.
143
What do we do if we have a weak acid?
1. Add excess strong base (NaOH).
| 2. 'back titrate' unreacted NaOH with strong acid (HCl).
144
What is the equation of strong base and strong acid?
NaOH + AcOH --> AcONa + H2O.
145
What is the ration in
| NaOH + AcONa + H2O?
1:1.
146
What do we do after we find the equation of strong base with strong acid?
Think of acid's molarity.
147
Why should we think about acid's molarity?
To have an idea of how much NaOH to use and how much to left behind.
148
How much do we think acid's molarity is?
0.03-0.05M.
149
How much will the moles of acetic acid be if we take 25.0mL?
(0. 03M/1000) x 25mL = 7.5 X 10-4 moles.
| (0. 05M/1000) x 25mL = 1.25 x 10-3 moles.
150
What do we need after we cunt acid's molarity with 25mL?
Excess of NaOH.
151
What do we to find NaOH molarity?
Add 25mL of 0.091M.
| (0.091m/1000) x 25mL = 2.275 X 10-3 moles.
152
What have we done with NaOH molarity?
Added twice NaOH moles (thought).
153
What do we realise from acid's and base's molarities in the end?
We have much more NaOH than acetic acid.
| Will be amount of NaOH unreacted by acetic AcOH.
154
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.
?
How much is left over when titrating it with HCL (0.093M).
155
What is the equation of NaOH titrating with HCl?
HCl + NaOH --> NaCl + H2O.
156
What is the ratio of NaOH and HCl titration?
1:1.
157
What is the equation of NaOH with HCl?
Strong acid v strong base reaction.
158
Which indicator is good for a strong acid v strong base reaction?
Any.
159
Which is an indicator we could use for a strong base v strong acid reaction?
Phenolphthalein.
160
What happens in colour changing when we use phenolphthalein as an indicator?
Pink --> colourless.
161
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?
Colourless --> pink.
162
How many times and what results do we get if we do this back titration?
```
3 times.
Volumes:
10.4mL.
10.2mL.
10.2mL.
Average = 10.16mL.
```
163
What is the number of HCl moles then?
(0.093M/1000) x 10.16mL = 9.45 x 10-4 mol.
164
With what is the number of HCl moles same?
With unreacted NaOH moles number.
165
What is the equation and calculations of moles in a back titration then?
Total moles NaOH - moles NaOH used in titration = moles NaOH used by acetic acid.
2. 275 x 10-3 moles - 9.45 x 10-4 moles =
1. 33 x 10-3 moles NaOH used by acetic acid.
166
What is then the moles of acid based on 1:1 NaOH:AcOH stoichiometry?
1.33 X 10-3 moles.
167
How much is the acid molariy?
Contained in 25.0mL --> 0.025L.
Molarity of acid = 1.33 x 10-3 moles / 0.025L =
0.053M.
168
On what is a redox titration based?
On an oxidation-reduction reaction.
169
Hoe does the oxidation-reduction reaction work?
In pairs.
| One is oxidised, one is reduced.
170
Where does the oxidation-reduction reaction occur?
Between titrant and analyte.
171
What can we identify with the oxidation-reduction reaction?
Unknown analytes' concentration.
172
How can we identify unknown analytes' concentration from an oxidation-reduction reaction?
By knowing oxidant's concentration --> measuring reductant's concentration.
173
What do we monitor in the oxidation-reduction titrations?
The reaction potential = voltage.
174
What we do not monitor in the oxidation-reduction titrations?
Concentration of reacting species.
175
Why we do not monitor concentration of reacting species in oxidation-reduced titrations?
Because they are both colourless.
176
What do we do after we monitor the voltage in the oxidation-reduction titrations?
Get data --> plot a graph like pH graph titration.
177
How is the graph of oxidation-reduction titrations?
Volume of titrant from burette v voltage.
Curve.
Voltage rising --> achieve equivalent point.
178
When is the voltage graph of oxidation-reduction titrations useful?
When there is no other way to observe equivalence point = end point.
179
What happens in oxidation-reduction titrations?
Electrons are transferred during redox reaction.
180
How is potassium permanganate characterised?
Self indicating.
181
What does it mean that potassium permanganate is self indicating?
Its colour changes during oxidation.
Its colour stays purple during reduction.
--> It shows equivalent point.
182
How is the colour of potassium permanganate characterised and why?
Deep purple.
| Due to permanganate ion MnO-4.
183
What agent is permanganate?
Strong oxidising agent.
184
What happens to permanganate when it reacts with a reducing agent Mn (+7)?
It is reduced to Mn (+2).
| MnO-4 (purple) --> Mn2+ (colourless).
185
Where does the permanganate reacting with reducing agent is useful?
When reducing agents are:
Fe (2).
Oxalate anion.
186
What happens to oxalate anion?
Carbon is oxidised from C (+3) to C (+4).
| C2O4 2- (colourless) --> CO2 (a gas).
187
What must we know before a titration can be performed?
The stoichiometric relationship between oxidant and reductant = balanced equation.
188
What is the reaction between MnO4- + C2O4 2-?
MnO4- + C2O4 2- --> Mn2+ + CO2.
189
What do we have to do the reaction to have a balanced redox equation?
Split it into a reduction reaction and an oxidation reaction.
1/2 reactions.
190
What are the 1/2 reactions of MnO4- + C2O4 2- --> Mn2+ + CO2 to get a balanced redox equation?
MnO4- + 8H+ + 5e- --> Mn2+ + 4H2O.
C2O4 2- --> 2CO2 + 2e-.
191
What do we do to the 1/2 equations?
Multiply first by 2.
| Multiply second by 5.
192
Why do we balance 1/2 equations?
To balance electron transfer.
193
What is the balanced redox equation in the end?
2MnO4 - + 5C2O4 2- + 16H+ --> 2Mn2+ + 8H2O + 10CO2.
194
What is the stoichiometry of the balanced redox equation between MnO4- and C2O4 2-?
2:5.
195
What do we use as primary standard to start the oxidation-reduction titration?
Sodium oxalate.
| Or oxalic acid.
196
What can oxalic acid give us?
Oxalate.
197
What do we have to do when we want to start the oxidation-reduction titration and use oxalate as primary standard?
Prepare 0.05M oxalate solution.
198
Why do we prepare 0.05M oxalate solution?
To standardise 0.02M KMno4 solution.
199
Why do we need to standardise permanganate?
Because it is not a primary standard.
| We do not know exact grams we need to prepare solution.
200
What will we do with the 0.05M oxalate solution we prepared?
Place 25.0mL in a conical flask.
201
What is a requirement in the oxidation-reduction titrations?
H+.
202
Why do we need H+ for the oxidation-reduction titrations?
To make the solution acidic.
203
What do we add after we put 25.0mL oxalate solution in a conical flask?
25mL strong acid solution = 1M H2SO4.
204
How is the reaction of oxalate and H2SO4 characterised?
Slow.
205
What do we have to do to the oxalate + H2SO4 reaction as it is slow?
Heat the flask to 80 degrees.
206
What else do we add in the oxalate + H2SO4 reaction?
Small quantity of KMnO4 from burette.
207
What colour does KMnO4 have in the beginning?
Purple.
208
What happens to the purple colour after a while KMnO4 is added to the oxalate + H2SO4 reaction?
It disappears.
209
What else happens except from the colour disappearance, when KMnO4 is added in the oxalate + H2SO4 reaction?
Some Mn2+ are created.
210
How do Mn2+ act?
As a catalyst.
211
What happens to the reaction when produced Mn2+ act as a catalyst?
It speeds up.
| Purple colour stays.
212
When do we recognise colour change and end point = equivalent point in the oxalate + H2SO4 + KMnO4 reaction?
When keep adding rapidly KMnO4 until purple colour just stays visible.
213
What happens if we add permanganate slowly?
Flask cools down.
| Reaction slows down.
214
What do we do when we are unsure if we reached end point?
Warm flask bit more.
Pink colour disperse.
Add permanganate.
215
What shall we do when we are halfway to end point?
Warm again.
Will not work.
Pink colour still here.
216
What happens when oxalate is oxidised?
CO2 produced.
Solution fizzes.
Bubbles away.
217
What do we know from our oxidation-reduction titration so far and what do we need to find?
```
Know:
Mox = 0.050M.
Vox = 25.0mL.
Vmn = 23.8mL.
Want:
Mmn = ?
```
218
What is the equation between oxalate + H2SO4 + KMnO4?
2MnO4- + 5C2O4 2- + 16H+ --> 2Mn2+ + 8H2O + 10CO2.
219
What is the ration of the oxalate + H2SO4 + KMn04 reaction?
5:2.
220
How can we find Mmn?
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
What can we analyse by using KmnO4 standardised solution?
Any oxalate solution.
| Any reductant contained solution = Fe2+.
222
What happens to permanganate over time?
It degrades.
223
Where is permanganate normally kept?
In brown bottle.
| Dark cupboard.
224
Why is permanganate usually kept in a brown bottle?
To reduce light in solution.
| Prolong its usefulness.
225
When does permanganate solution go off?
When Mn4+ is produced --> insoluble precipitate of manganese dioxide produced --> brown bottle gets coloured.
226
How long does permanganate solution last?
For weeks.
227
What is the advantage of permanganate solution lasting for weeks?
It does not need to be standardised often.
228
How can we analyse oxalate solutions by using KnO4 standardised solution?
Oxalate solution has an unknown concentration.
| Carry out 3 titrations.
229
What are the 3 titrations we can do to analyse oxalate concentration?
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
What can be the strong acid we will use in the titration and why?
1M sulphuric solution.
| To have plenty H+ ions.
231
What do we know and what do we need to find?
```
Know:
Mmn = 0.021M.
Vox = 25.0mL.
Vn = 16.8mL.
To find:
Mox = ?
```
232
What is the equation between MnO4- and C2O4 2-?
2MnO4- + 5C2O4 2- + 16H+ --> 2Mn2+ + 8H2O + 10CO2.
233
What is the ratio of the 2MnO4- + 5C2O4 2- equation?
5:2.
234
What calculations do we have to do to find Mox based on the known values?
```
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
Where do all the calculation rely?
On a properly balanced equation.
236
What happens if we get the balanced equation right?
Rest fall into place.
237
What are most ionic compound to some degree?
Soluble.
| Insoluble.
238
How is the degree of solubility of a particular compound expressed?
Ksp = solubility product.
239
What can we get if we combine solutions with different ions?
Insoluble solutions.
240
What is the relations between value and solubility of compound?
Larger volume --> more soluble compound.
241
How soluble are NaCl and AgCl at 20 degrees?
Ksp NaCl = 37.66 --> soluble.
| Ksp AgCl = 1.77 x 10-10 --> insoluble.
242
How solubility can change based on temperature?
When temperature changes so doe s solubility.
243
Where is solubility used?
In precipitation titrations.
244
What do we have to do to find the concentration of a solution that contains CL- ions?
Titrate with a standardised solution which soluble, contains silver ions.
245
Why can we use silver ions as a standard solution in titration?
Because they do not cause any precipitation issues.
246
What is the equation between NaCl and AgCl solutions?
AgNO3(aq) + NaCl(aq) --> AgCl(s) + NaNO3(aq).
247
Why can we no identify the end point pf the NaCl + AgCl reaction?
Both solutions are colourless.
248
How will we identify the end point of NaCl + AgCl reaction?
When chloride is less and less in solution until there is none left.
249
How can we observe the end point in precipitation titrations?
With electrochemical methods.
| With special indicators.
250
What is a special indicator we can use to observe the end point in precipitation titrations?
Mohr method.
251
Which indicator is used in Mohr method to observe the end point in a precipitated titration?
CrO4 2- ion.
252
What colour does CrO4 2-, chromate ion have?
Bright yellow.
253
What happens in a titration when chromate ion is added?
Chromate ion added to chloride solution --> yellow colour --> titration starts.
254
What happens as the titration with chromate ion proceeds?
White/grey silver chloride is formed.
255
What happens at the end point when chromate ion is added to a chloride solution titration?
All chloride is used up.
| Deep red precipitate of silver chromate (Ag2CrO4) is formed.
256
What happens during the titration when chromate ion is added too chloride solution and silver chromate is formed?
AgCl + ag2CrO4 compete.
257
What is the difference between AgCr04 and AgCl?
Ag2CrO4 is more soluble.
| Only forms when all chloride is used up.
258
What is the issue in the titration of using chromate with chloride solution?
Ag2CrO4 is formed but it needs Ag solution.
| End point happens after equivalent point.
259
How can we overcome the issue that end point comes after equivalent point in the titration of using chromate and chlorate?
Add 1 drop chromate indicator.
| Carry out 'blank' titration with no chloride.
260
What will the use of chromate only tell us?
How much to remove from titration volume to get better value equivalent point value.
261
What do we must control to make sure end point wont be different form equivalent point?
pH.
262
What does it form in a basic solution?
AgOH precipitate.
| Chloride.
263
What can precipitate increase if formed?
Titration volume.
264
Why does precipitate increases titration volume when it is formed?
Because precipitate it is formed as well as chloride --> titrating 2 things.
265
What happens to chromate in an acidic solution?
Dimerizes --> becomes dichromate.
Cr2O7 2-.
Froms AgCr2O7.
266
Why will we not realise the end point in an acidic solution when chromate is involved?
Chromate dimerizes.
Ag2Cr2O7 formed.
Ag2Cr2O soluble.
No precipitate.
267
Why must we control pH in a titration?
It is important to observe end point.
268
What does Fajans method in precipitated titrations use?
Species that changes colour when absorbs precipitate.
269
What happens in Fajans method that improves precipitated titrations?
Anionic dye dichlorofluorescein added to titrand's solution in conical flask.
270
What do we realise in the Fajans method before the end point?
Precipitate of AgCl has negative charge.
271
Why does AgCl has negative charge?
It absorbs excess Cl-.
| Dichlorofluorescein carries negative charge --> pushed back by precipitate --> remains in solution.
272
What colour does dichlorofluorescein have?
Greenish-yellow = faint glow.
273
What happens in the Fajans method after the end point?
Precipitate carries positive charge.
274
Why does precipitate carry positive charge after end point in Fajans method?
It absorbs excess Ag+.
| Dichlorofluorescein absorbs to precipitate = attracted negative charge.
275
What happens when the precipitate has a positive charge after the end point?
Changes colour = yellow --> pink.
276
What does the colour of precipitate indicate in Fajans method?
End point.
277
What do the precipitation titrations involve?
Ag ions.
278
How are the precipitation titrations that use Ag ions named?
'Argentometric titrations'.
279
What do Ag ions do in the precipitation titrations when they are used?
Analyse other ions.
280
Why do Ag ions analyse other ions when used in precipitation titrations?
They are soluble.
281
How are the precipitation titrations called?
'Direct titrations'.
282
How are the calculations of precipitation titrations?
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.
