1. Solutions

2. Solutions Remember that a solution is a homogeneous mixture
Solubility -the amount of solute that can be dissolved to form a solution.
Solvent – the substance in a solution present in the greatest amount.
Solute – the substance in a solution present in the least amount.
Saturate – a solution that has come to equilibrium. The rate of dissolving is equal to the rate of re-crystalizing.
High solubility – more than .10 mol/L will dissolve (AgNO3)
Low solubility – less than .10 mol/L will dissolve (AgCl)

3. Classification of solutions
Every combination of solids, liquids, and gases is possible.
This means that there are nine possible kinds of solutions.
Solute Solvent
Liquid
Gas
Solid
Solid
Liquid
Gas
Solid
Liquid
Gas

4. Concentration
Concentration is one way to express the relationship between the solute and solvent in any solution.
Since the mixture is homogeneous, the relative number of particles is the same in a drop as in a liter.
One way to show this relationship is by expressing the number of moles of solute dissolved in one liter of solvent.
That is:

5. Sheet 20/5—1 #1-12 Multiple choice #1-10
Examples 1) What is the concentration of a solution made by dissolving 2.00 moles of NaCl(s) in a volume of 4.00 L of H2O(l) ?
2) How many moles of HCl(aq) are present in 5.00 L of a mol/L solution of hydrochloric acid?
3) What volume of 1.00 mol/L NaOH(aq) can be made from dissolving mol of NaOH(s) in H2O(l)?
Sheet 20/5—1 #1-12
Multiple choice #1-10

6. Classifying solutions - Qualitative
Solutions may be – acidic
- basic
- neutral
How would you test this?
pH paper, indicators, litmus, other?
Solutions may be – electrolytes
- non electrolytes
conductivity meter, conductivity apparatus, other?

7. Dilution Problems
In dilution problems, only the amount of solvent changes.
Therefore, the number of moles of solute in the final solution is the same as the number of moles of solute in the original solution.
Examples
1) What volume of 12.0 mol/L HCl(aq) is needed to produce 2.00 L of mol/L HCl(aq)?

8. 2) What volume of 4. 00 mol/L NaOH(aq) is needed to produce 1
2) What volume of 4.00 mol/L NaOH(aq) is needed to produce 1.00 L of mol/L NaOH(aq)?
3) What is the maximum volume of mol/L HNO3(aq) that could be prepared from 45.0 mL of 8.00 mol/L HNO3(aq) ?
4) What is the final concentration of a solution where 25.0 mL of 18.0 mol/L H2SO4(aq) is added to mL of H2O(l)?
Do Sheet 20/ #5 – 9, then do 1 - 4

9. Addition Problems
We can use the same method to solve addition problems. In this type of problem, one solution is diluted with a lower concentration of the same kind of solution.
Examples:
1) L of a 6.00 mol/L HCl (aq) solution is mixed with L of a 5.00 mol/L HCl (aq) solution. What is the [HCl] in the final mixture?
300 mL of 5.00 mol/L NaOH (aq) is combined with 200 mL of an unknown concentration of NaOH (aq). The concentration of the final mixture is 6.00 mol/L. What is the concentration of the unknown solution?

10. Sheet 20/5--6 #5 - 10

11. How to make a standard solution This is a solution of a known and exact concentration
Procedure
Calculate the mass of solute required and weigh out this amount on a scale.
Dissolve the solute in a small amount of water.
Top up the water to the final desired volume.
Stopper and mix well.
Example: Mix up 2.00 L of a mol/L NaOH(aq) solution.

12. Acids and Bases
Svante Arrhenius was the first to do extensive study with acids and bases.
ACIDS
HCl(aq) → H+(aq) + Cl –(aq)
HNO3(aq) → H+(aq) + NO3-(aq)
HI(aq) → H+(aq) + I-(aq)
Acids – substances that give up a hydrogen ion in solution.

13. Acids are substances that donate hydrogen ions in solution
Svante Arrhenius

14. Bases NaOH(s) → Na+(aq) + OH–(aq) Ba(OH)2(s) → Ba2+(aq) + 2 OH-(aq)
Al(OH)3(s) → Al3+ (aq) OH-(aq)
Bases – substances that give up hydroxide ions in solution.

15. Bases are substances that donate hydroxide ions in solution
Svante Arrhenius

16. The Hydronium Ion
It is believed that the positively charged hydrogen ion is attracted to and sticks on to the negative end of the very polar water molecule.
This forms a new species called the Hydronium Ion.
H+ + H2O (l)  H3O+ (aq)

17. Ionization of Water 2 H2O (l)  H3O+ (aq) + OH– (aq)
Pure water always contains some ions.
This is the result of the ionization of some (very few) of the water molecules.
That is:
2 H2O (l)  H3O+ (aq) + OH– (aq)
The concentration of the H+ in pure water is very low…about 1.0 x mol/L … and…
The concentration of OH- in pure water is the same… x mol/L.

18. In a solution: If the [H+] > [OH-] the solution is acidic.
If the [OH-] > [H+] the solution is basic.
If the [H+] = [OH-] the solution is neutral.

19. pH pH = -log [H+] pH – method of expressing the [H+] in solution
– power of Hydrogen (exponential power)
- Scale normally runs from 0 → 14
pH = -log [H+]

20. Converting [H+] to pH
Calculate the pH of a solution that has a [H+] of x 10-2 mol/L Answer pH = 2.000
This is acidic.
If the [H+] is 1.00 x 10– 9, the pH is 9.000
This is basic.
If the [H+] is 1.00 x the pH is 7.000
This represents neutral conditions.

21. neutral
acidic
basic

22. Fill in the following chart
Can you see any patterns here? pH
[H+] (mol/L)
1.25 x 10-5
3.16 x 10-11
6.92 x 10-9
4.89 x 10-2
7.44 x 10-8
7.44 x 10-12
7.44 x 10-3

23. I noticed that!
Me too!
Svante Arrhenius
Beaker

24. Convert the following into Hydrogen ion concentrationss
pH [H+]
4.3
6.21
9.87
3.245
1.6
2.67
12.573

25. Convert the following into Hydrogen ion concentrations
pH [H+]
x 10-5
x 10-7
x 10-10
x 10-4
x 10-2
x 10-3
x 10-13

26. Complete the following tablepH
[H+] mol/L
[OH-] mol/L
1.0 1
0.1 2
0.01 3 4 5 6 7 8 9 10 11 12 13 14

27. Complete the following tablepH
[H+] mol/L
[OH-] mol/L
1.0 1
0.1 2
0.01 3
0.001 4
0.0001 5 6 7 8 9 10 11 12 13 14 -2
100 -1 10

28. Note that these are log numbers
Note that these are log numbers. Each change of 1 pH unit is a 10 fold change in [H+].
Example: A solution with a pH of 4.00 has 1000 times more H+ than a solution with a pH of 7.00

29. Indicators

30. Measuring pH Most accurate method to measure pH is to use a pH meter.
However, certain chemicals change color as pH changes. These are called indicators.
Indicators are less precise than pH meters.
Many indicators do not have a sharp color change as a function of pH.

32. The pH Scale
Measuring pH

33. Bromothymol blue
ACID
pH below 6.0
BASE
pH above 7.6

34. Neutralization
When an acid and a base are reacted, the products will be an ionic salt and water.
The salt present depends on which acid and which base were used in the neutralization.
Example: combine HCl(aq) with NaOH(aq)
HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq)
Acid base water salt (sodium chloride)

35. Neutralization using modified Arrhenius theory could be defined as the reaction between a hydronium ion and a hydroxide ion that produces water.

36. Write net neutralization reactions for:
KOH (aq) + HI (aq) →

38. Dissociation Reactions
Dissociation - means to separate
Ionic compounds and acids will dissociate into ions when they are dissolved in water.
These reactions must be balanced.
Example:
Ca3(PO4)2(s) → 3 Ca2+(aq) PO43-(aq)
Complete Sheet 20/5 - 14

39. Concentration of ions
Combinations problems

40. Qualitative tests
Solution color
Solubility
Flame color

41. titrations

42. Titrations
Volumetric Analysis

43. Titrations
Titration is a physical process that is used to find the concentration of an unknown solution. Titration involves adding one solution (titrant) from a buret to another solution (sample) in an Erlenmeyer flask. (titration flask)

44. (titrant)
indicator is added
known volume
(sample)

45. Titration procedure (quantitative)
Measured volume of unknown acid [ ] is added to a flask.
An appropriate indicator is added.
Measured amount of base of known [ ] is added using a buret.
Continue until solution changes color. This is called the end point.

46. Carry out the
Titration Lab

47. 24.19 mL

49. Procedure Fill the left side buret with acid. (HCl(aq))
Fill the right side buret with base.(NaOH(aq))
Record the initial levels on each buret.
Add about mL of acid to an erlenmeyer flask.
Add 1 drop of indicator.
Add base from the second buret until the expected color change occurs.
Record the final buret levels on each buret.

50. Notes Always keep the acid on the left and the base on the right.
Stop at the first permanent pink.
Don’t forget to add the indicator.
Back titration is o.k.
You may wash the sides of the flask down with distilled water.
Make sure that all glassware is very clean. Remember that you will be marked on accuracy as well.
[NaOH(aq) ] = mol/L

51. Dissociation (ionization) Reactions
When ionic compounds dissolve in water, they separate into ions – one positive and one negative.
Examples:
Ca(NO3)2(s) → Ca2+(aq) + 2 NO3–(aq)
Ca3(PO4)2 → 3 Ca2+(aq) PO43-(aq)
Write dissociation reactions for each of the following:
BaSO4(s) →
Al(OH)3(s) →
PbSO4(s) →
NH4NO3(s) →
Ba2+(aq) + SO4 2-(aq)
Al3+(aq) + 3 OH – (aq)
Pb2+(aq) + SO42-(aq)
NH4+(aq) + NO3–(aq)

52. Net Ionic Equations
Combine solutions of NaCl(aq) and AgNO3(aq)