Solutions and Dilutions Flashcards
(118 cards)
1
Q
What do most of the analytical methods need to be in solution?
A
An analyte.
2
Q
What is usually the concentration for qualitative analysis?
A
Not critical.
3
Q
What do reagents not need to be?
A
Accurate.
4
Q
What is it important in quantitative work?
A
Reagents to be accurate.
5
Q
What do we need to know in quantitative work?
A
Volumes.
Concentrations.
Accuracy.
6
Q
How do we calculate quantity of analyte?
A
Quantity of analyte = total amount of material x concentration of analyte.
7
Q
What happens to the amount of analyte when we double the amount of material?
A
It doubles.
8
Q
What happens to the amount of analyte when we double the concentration?
A
It doubles.
9
Q
What can the units be for analyte’s quantity?
A
Grams.
Moles.
10
Q
How do we measure moles?
A
Moles = Molarity x volume.
11
Q
How do we measure grams?
A
Mass x percentage.
12
Q
How do we measure solute’s volume?
A
Total volume x percentage by volume.
13
Q
How can we measure how much alcohol beer has if we know 500mL of it at 4.2% abv?
A
500 x (4.2/100)= 21mL.
14
Q
How much alcohol does alcopop have if 300mL of it at 5.0%?
A
300 x (5.0/100)= 15mL.
15
Q
How much alcohol does wine have if 175mL at 12%?
A
175 x (12/100)= 21mL.
16
Q
How much alcohol does whisky have if 25mL at 40%?
A
25 x (40/100)= 10mL.
17
Q
What is the rule of solubility?
A
Like dissolves like.
18
Q
What does ‘Like dissolves like’ mean?
A
Polar solvent dissolves polar molecule.
Non-polar solvent dissolves non-polar analyte.
19
Q
Why do we have to choose purposeful properties for solvents?
A
Because we use them for analysis.
20
Q
What do we check properties for to get a good solvent?
A
- Good solubility for analyte.
- Cheap to be diluted.
- Moderate volatility to have little evaporation from the solution.
- Safe from toxicity and flammability.
- Useful.
21
Q
With what factor does solubility vary?
A
Temperature.
Heat it up.
Let it cool.
22
Q
For what is water a good solvent?
A
Ionic materials.
Polar compounds.
23
Q
Why could water not be suitable for solubility?
A
It might react.
24
Q
What other solvents can we use if water is not suitable?
A
- Acid, hydrochloric, nitric for inorganic materials.
- Concentrated acid, hot acid, alkali.
- Specialised solvents: mixed acids, complexing agents: aqua regia for gold, HF (nasty), perchloric acid.
25
What is the disadvantage of many of the complexing agents?
They are hazardous.
26
What is the aqua regia?
'Royal water' mix of nitric and hydrochloric acids.
27
What can aqua regia dissolve?
Gold.
| Platinum.
28
What can work if all else fail?
1. Fusions with flux (borax).
2. Polar organic solvents: diethyl ether, alkanals, alkanoles, acetone, alkanols. ethanol, hexanol, chlorinated hydrocarbons.
3. Non-polar solvents: paraffins, heptane.
29
What do we have to check about polar organic solvents?
Solvent's nature.
30
Why should we check solvent's nature?
To not react unwanted.
31
For what are pentane and heptane good solvents?
Non polar material: oils, fats, waxes.
32
Where can we find solubility information?
On Certificates of Analysis.
| Seller websites.
33
What do we aim for in solubility?
Pu analyte in solution --> dilute it --> put it in analytical processes.
Best soluble substance.
34
Where can help heat in solubility?
In dissolving analyte.
35
Why should we be careful with heat in solubility?
Because it does not recrystalise on cooling.
Need to dilute it further on our own.
36
What do we need in solubility?
Something that will leave less dilution after its use.
37
What is the Ultra-Sonic?
Ultrasound from 20-400kHz.
38
What happens in cavitation?
Waves in ultra-sound --> expand --> --> contract liquid.
If gas --> makes bubbles.
39
Where does solubilisation take place?
On material's surface.
40
What do we create when we break something to smaller pieces?
Larger surface area.
| More solubilisation.
41
What do we need in a chemical reaction to use it and find how much analyte we have?
1. Stoichiometric.
2. Fast.
3. Complete as far as possible.
4. Detecting clearly equivalence point.
5. Safe cheap reagents.
6. Pure reagents.
42
Why do we need stoichiometric reactions?
No side reactions.
Reagent reacts completely with solution.
No fault results.
43
Why do we need a completed chemical reaction?
No equilibrium mixture.
44
How can we detect the equivalence point of a chemical reaction clearly?
When it is all reacted.
45
Why do we want safe and cheap reagents in reactions?
To not cause harm.
Not explode.
Not produce gas.
46
Why do we want pure reagents in reactions?
To know exactly how much analyte added.
Confident measures.
Small error.
47
From where can we buy an analyte and be 99% confident it is pure?
Analar.
48
What do we need to know in quantitative analysis?
How much analyte is present.
49
What analysis can we use when the analyte reacts with a reagent in a stoichiometric reaction and it is all reacted?
a Analyte + r Reagent --> p Product.
| a mole + r mole --> p mole.
50
Based on the analyte's analysis equations what we must have if we need r moles of R?
a moles of A.
51
What we must have to produce p moles of P?
a moles of A.
52
What do we know when we know the equations/relationships between analytes, reagents, and products?
Exactly analyte's concentration.
53
What do we need to remember for a reaction?
Molar ratio= 1:1, 1:2, 2:1.
54
2H2 + O2 --> 2H2O
| Molar ratio between O2 and H20?
1 to 2.
55
2H2 + O2 --> 2H2O
| Molar ratio between H2 and H20?
2 to 2.
| 1 to 1.
56
2O3 --> 3O2
| Molar ratio between O3 and O2?
2 to 3.
57
2O3 --> 3O2
| Molar ratio between O2 and O3?
3 to 2.
58
Which are the 3 important equations we must know?
1. n = g/gfm.
2. c = n/v.
3. C1V1 = C2V2.
59
What does the equation n=g/gfm mean?
Number of moles = mass in grams / gram formula mass.
60
What is c?
concentration (moles/L).
61
What can be the units for concentration?
```
moles/L.
%.
ppm.
mg/mL.
ug/mL.
```
62
What must we use in both solutions?
Correct units.
63
What does the equation C1V1 = C2V2 mean?
From concentrated solution to prepare diluted solution.
```
C1 = concentration we put down.
V1 = volume of first solution we will put in second solution.
C2 = second concentration we get once we add the volume 1.
V2 = second volume we add from first volume.
```
64
Calculate the no. of moles present in 50 cm3 (mL) of 0.05 molar HCl
```
n = c x (v/1000)
n = 0.05 x 50/1000
n = 0.0025 moles.
```
65
Calculate the concentration (moles/L) if 0.1 moles is dissolved in 100 cm3(ml) of water
```
c = n/v in litres
c = n/(100/1000)
c = 0.1/0.1
c = 0.01 moles /L.
```
66
How do we make dilutions?
1. Make a standard solution.
2. Transfer a known volume with pipette to a clean flask.
3. Make up to a new volume.
67
What is the equation we follow in dilutions?
C1V1 = C2V2.
68
How can we calculate the new concentration of the dilution?
New concentration (C2) = Old concentration (C1 x (Pipetted volume (V1) / Size of volumetric (V2).
69
What do we have to maintain the dilution equation?
Same units.
70
What do we have to choose in a dilution?
A suitable ratio of pipette volume/standard flask volume.
71
What should we avoid in dilutions?
Using small volumes.
72
What should we use instead in dilutions?
Serial dilutions.
73
What is more accurate between the 2?
(10/100) x (10/100) x (10/100)
or
(1/1000)?
(10/100) x (10/100) x (10/100).
74
Why should we use small dilutions in dilutions?
To make smaller errors.
75
What is the equation we follow in titrations?
aAnalyte + rReagent --> pProduct.
| NaOH + HCl --> H20 + NaCl.
76
What do we have in titrations?
Analyte solution of unknown concentration.
77
What do we do in titrations?
1. Take fixed volume (pipette).
2. React it with reagent of known concentration (standardised).
3. Add unknown, measures amount by burette.
4. Reach equivalence point (end point).
78
How can we determine our equivalence point?
1. Electrochemically by pH meter.
| 2. By indicator.
79
How can we determine the equivalence point by a pH meter?
1. Add reagent to analyte.
2. pH changes gradually.
3. At equivalence point pH changes rapidly.
or
1. Add acid to base.
2. pH high before equivalence point.
3. pH drops rapidly when reach equivalent point.
4. Plot a graph.
80
What happens at the equivalence point?
pH changes very rapidly.
81
What do indicators do?
Change colour when ionised/unionised.
82
What are the indicators?
Weak acids/bases.
83
When does pH changes very rapidly when using an indicator?
>2 pH units.
84
What happens when we add an indicator in a basic solution?
NaOH + indicator --> particular colour in basic solution --> reach equivalence point --> colour change in neutral/acid solution.
85
Why do we have to choose carefully an indicator?
To get a colour change on the steep part of the curve.
86
What is the definition of the equivalence point?
The point in a titration when the amount of added standard reagent is equivalent to the amount of analyte.
Stoichiometrically balanced.
87
What is the definition of the end point?
The point in a titration where a physical change occurs that is associated with the condition of chemical equivalence.
88
What is the definition of the indicators?
Added to produce an observable physical change (signalling the end point) at or near the equivalence point.
89
How can we calculate the titration error (Et)?
Et = Vep - Veq.
90
What is Vep?
Actual volume of reagent required to reach end point.
91
What is Veq?
Theoretical volume of reagent required to reach equivalence point.
92
How much is the difference between Vep - Veq if we do titrations correctly?
Zero.
93
What is the aim of titrating?
Having zero difference between Vep-Veq.
| Being precise.
94
What happens in a rough titration?
We find out roughly where the end point is.
| We repeat titrations to get accurate measurements.
95
What does the titration equation mean?
| (𝐶_1×𝑉_1)/𝑛_1 = (𝐶_2×𝑉_2)/𝑛_2
```
C = molar concentration (molarity)
V = volume (L or mL
n = moles in chemical equation - stoichiometry
C1 = reagent's concentration
V1 = how much added from burette
V2 = unknown solution pipetted into flask to titrate
```
96
What is essential in back titrations?
Calculate moles added rather than concentration (moles/L).
97
How can we determine our equivalence point?
1. Electrochemically by a pH meter.
| 2. With an indicator.
98
When do indicators change colour?
When ionised/unionised.
99
In what pH number do indicators change colour?
Over 2 pH units.
| weak acids/bases
100
Where on the curve do we want our colour change to be?
On the steep part of the curve.
101
When is an indicator not suitable for titrations?
When it changes colour beyond end point.
102
Why do we not accept an indicator that changes colour beyond end point?
Because it gives wrong calculations.
103
Why is the indicator choice important?
We can see where pH and colour changes for the titration.
We can see if the colour change matches the experiment.
104
What colour does phenolphthalein have in acidic and basic solutions?
Colourless = acidic.
| Bright pink = basic.
105
Is end point and equivalence point the same?
Yes.
106
Is the pH of the solution difference before and after the end point?
Yes.
107
Where does bromothymol blue change colour?
At pH = 6-7.5.
| On the deepest part of the curve.
108
Where does phenolphthalein change colour?
At pH = 8-9.5.
| On the steepest part of the curve.
109
Is it easy to see pink becoming colourless, iodine - orange becoming colourless in titrations?
No.
110
What do we use when it is hard to identify a colour change in titrations?
Back titrations.
111
What are the rules in back titrations?
```
Analyte = acid.
Reagent = base.
```
112
What is the process in back titrations?
1. Add a known amount of reagent in excess.
2. Titrate to find excess left with different acid with known concentration.
3. Calculate how much base left over in moles.
113
If you dilute 175 mL of a 1.6 M solution of LiCl to 1.0 L, determine the new concentration of the solution.
2.8M
114
You need to make 10.0 L of 1.2 M KNO3. What molarity would the potassium nitrate solution need to be if you were to use only 2.5 L of it?
4.8M.
115
How many milliliters of 5.0 M copper(II) sulfate solution must be added to 160 mL of water to achieve a 0.30 M copper(II) sulfate solution?
10.0mL
116
You need 0.0100 mole of lead (II) chromate (323.2 g/mol). How much should you weigh on the scale?
3.232g
117
Given 6.40 g of HBr (80.91 g/mol). How many moles is this?
0.05moles
118
0.3000 g of Na2CO3 (105.9888 g/mol) were dissolved in 100.0 mL of water. Calculate the molarity.
0.03M
