CHEM 12

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ACID AND BASE UNIT REVIEW

CONCEPTS

1. List the properties of acids (taste sour, conduct electricity, change the colours of indicators, react with certain metals to produce H₂, are neutralized by bases) and bases (taste bitter, conduct electricity, change the colours of indicators, feel slippery, are neutralized by acids) .

2. What ion is produced when an acid reacts with water? ( H₃O⁺ ) A base? ( OH^- )

3. Define: conjugate ( formula that differs by one H⁺ ), Arhenius acid ( ionizes in water to produce H⁺ ) and base ( ionizes in water to produce OH^- ), Bronsted acid (proton donor) and base (proton acceptor), weak acid (an acid that does not completely ionize in water), strong acid (an acid that completely ionizes in water), ionization (the production of anions and cations), equivalence point of a titration (when the moles of H⁺ = the moles OH^- ), transition point of an indicator (write an equation) ( When [HInd]=[Ind^-] ), buffer ( a solution that will maintain a relatively constant pH that consists of a weak acid or base mixed with the salt that contains the congugate) and hydrolysis (the reaction of the cation or anion with water as an acid or a base).

4. Identify the acids or bases in the following equation. Are the reactants or products favoured?

HC0₃^-(base) + HF(stronger acid)<-----> H₂CO₃(weaker acid) + F^-(base)

The products are favoured as the stronger acid is on the left and pushes the reaction to the right.

5. Classify each compound as a strong or weak acid or base; acidic or basic anhydride; acidic, basic, or neutral salt; or buffer system. Write an equation to show how each reacts with water.

NH3 weak base	AlCl3 acidic salt	H2CO3 weak acid	HClO4 strong acid	KCN basic salt	NH4Cl acidic salt	KOH strong base
SO2 acid anhydride	NaF basic salt	HCl strong acid	NaI neutral salt	K2O basic anhydride	NaOH strong base	CO2 acid anhydride

NH3 and NH4Cl buffer

NaCH3COO and CH3COOH buffer

NH₃ + H₂O <----> NH₄ ⁺ + OH^-

Al(H₂O)₆⁺³ + H₂O <--------> Al(H₂O)₅(OH)⁺² + H₃O⁺

H₂CO₃ + H₂O <--------> H₃O⁺ + HCO₃ ^-

HClO₄ + H₂O --------> H₃O⁺ + ClO₄ ^-

CN^- + H₂O <----> HCN + OH^-

NH₄⁺ + H₂0 <-------> NH₃ + H₃O⁺

KOH -----> K⁺ + OH^-

SO₂ + H₂O -----> H₂SO₃

F^- + HOH_(l) <-----> HF + OH^-

HCl + H₂O --------> H₃O⁺ + Cl ^-

NaI --------> Na⁺ + I^-

K₂O + H₂O --------> 2KOH

NaOH --------> Na⁺ + OH^-

CO₂ + H₂O -----> H₂CO₃

NH₃ + H₂O <----> NH₄ ⁺ + OH^-

CH₃COOH + H₂O <--------> H₃O⁺ + CH₃COO ^-

6. H⁺is short for ______. H₃O⁺

7. Determine the conjugates for each of the bases.

CN^-HCN

NH3 NH4 ^+

F^-HF

OH^-H2O

Co(H2O)5(OH)^+2Co(H2O)6^+3

8. Determine the conjugates for each of the acids.

HF F^-

HCN CN^-

Al(H2O)6^+3Al(H2O)5(OH)^+2

NH4^+NH3

HPO4^-2PO4^-3

9. Describe a strong and weak acid as well as a strong and weak base in terms of each of the following:

Conductivity ( strong acids or bases are good conductors strong while weak acids or bases are poor conductors )and, Size of Ka or Kb (strong acids or bases have large Kas or Kbs while weak acids or bases have small Kas or Kbs) , Degree of Ionization (strong acids or bases completely ionize while weak acids or bases do not completely ionize) and pH (Assuming 1.0M, strong acids pH=0, strong bases pH=14, weak acids pH>0, weak bases pH<14 ) .

10. Why is the strongest acid in water H₃O⁺? Explain! HCl + H₂O <--------> H₃O⁺ + Cl ^-

11. Why is the strongest base in water OH^-? Explain! NaOH ----> Na⁺ + OH^-

12. Which has the higher pH H₂S0₃ or H₃BO₃? Explain!

H₃BO₃ as it is the weakest acid.

13. Which has the higher pH NaCN or NaF? Explain!

NaCN as it is the strongest base.

14. A buffer has a pH of 9.00. 2 drops of a dilute strong acid are added. Estimate how the pH changes?

It will remain the same or drop slightly.

15. a) Complete the chart by indicating the pairs required to make buffer solutions. For example HCN (weak acid) and NaCN (salt containing the conjugate of the weak acid) will make a buffer solution. b) Write an equation the describes the equilibrium for each buffer. c) Circle the formulas that have high concentrations.

Weak Acid or Base	Salt
HF	NaF
CH3COOH	NaCH3COO
NH3	NH4Cl
HCN	NaCN
H2CO3	NaHCO3
H3PO4	KH2PO4
HCH3COO	NaCH3COO

HF + H₂O <--------> H₃O⁺ + F ^-

CH₃COOH + H₂O <--------> H₃O⁺ + CH₃COO ^-

NH₃ + H₂O <----> NH₄ ⁺ + OH^-

HCN + H₂O <--------> H₃O⁺ + CN ^-

H₂CO₃ + H₂O <--------> H₃O⁺ + HCO₃ ^-

H₃PO₄ + H₂O <--------> H₃O⁺ + H₂PO₄ ^-

HCH₃COO + H₂O <--------> H₃O⁺ + CH₃COO ^-

16. Match each equation with its type:

Acid/base formula equation	HCl + NaOH ----->NaCl + HOH(l )
Acid/base net ionic equation	H^+ + OH^- -----> HOH(l)
Solubility product ionization equation	AgCl(s)< -----> Ag^+ + Cl^-
Hydrolysis of a weak acid	NH4^+ + H20 <-------> NH3 + H3O^+
Hydrolysis of a weak base	F^- + HOH(l) <-----> HF + OH^-
Ionization of water	H20(l) <-----> H^++ OH^-

17. Write the equilibrium expression for each of the above equations.

Kb=[HF][OH^-]/[F^- ]

Ksp=[Ag⁺][Cl^-]

Kw=[ H⁺ ][OH^-]

Ka=[NH₃ ][H₃O⁺/[NH₄⁺ ]

18. A student tested the electrical conductivity of two acid solutions. One solution was a strong acid and the other a weak acid. Both solutions had the same conductivity. Explain how this could be possible.

The strong acid was dilute and the weak acid was concentrated.

19. Describe in terms of hydrolysis how NaCH₃COO can be added to potato chips in order to produce the vinegar flavour. CH₃COO^- + H₂O <--------> CH₃COOH + OH^- the acetate ion hydrolyzes with the water in your mouth to produce vinegar (CH₃COOH ).

20. Describe what happens to the H⁺ and the OH^- when the pH increases by 2 units.

[H⁺] decreases by a factor of 100. [OH^- ] increases by a factor of 100.

21. Describe two gases responsible for acid rain. SO₂ or SO₃ and NO₂ Write aquations to show how they react with water. SO₂ + H₂O -----> H₂SO₃ or SO₃ + H₂O -----> H₂SO₄

2NO₂ + H₂O------> HNO₃ + HNO₂What gas naturally lowers the pH of normal rain? CO₂

22. Complete the following reaction using a formula equation, complete ionic equation and net ionic equation.

H₂C₂O₄ + 2NaOH ----->Na₂C₂O₄ + 2H₂O

H₂C₂O₄ + 2Na⁺ + 2OH^- -------> 2Na⁺ + C₂O₄^-2 + 2H₂O

H₂C₂O₄ + 2OH^- -------> C₂O₄^-2 + 2H₂O

23. Write the equilibrium expression for phosphoric acid in water. H₃PO₄ <----> H⁺ + H₂PO₄^-

24.

a) An indicator HInd is red in acid and blue in base. Write the equation for the indicator and explain the colours.	HInd(red) + H2O <-----> H3O^++ Ind^-(blue) When [HInd] =[Ind^-] the color is orange
b) What is true at the transition point?	[HInd]=[Ind^-] and Ka = [H3O^+]
c)What is the color of this indicator in a solution of AlCl3?	AlCl3 is an acid salt- indicator is red.
d) In the above solution, what is larger [HInd] or [Ind^-] ?	[HInd] > [Ind^-]
e) Calculate the Ka for Phenolphthalein.	8 X 10^-10
f) What indicator has a Ka of approximately 1.0 X 10^-10?	Thymolpthalein

25. Give an example of a monoprotic HCl → H⁺ + Cl^-, diprotic H₂SO₄ → H⁺ + HSO₄^-1 , and triprotic acid H₃PO₄ ⇌ H⁺ + H₂PO₄^-1. Write an equation for each to show how they ionize in water.

26. Give the approximate pH of the equivalence for each titration. Choose an appropriate indicator.

Acid	Base	pH of Equvalence Point	Indicator
HCl	NaOH	7	Bromo thymol blue
H2SO4	NH3	5	Methyl Red
HF	KOH	9	Phenolpthalein

27. Which of the following will have the lowest pH?

HCLO

HClO2

HClO3

HClO4 Strongest acid more O's &lowest pH

28. Pick a formula that is amphriprotic. Formulas in red can both gain or lose a proton.

H2SO4	H20	F^-	HCO3^-
CO3^-2	KOH	H2PO4^-	HPO4^-2

CALCULATIONS

1. Calculate the quantities required to complete the table. In the first row write the general equations for each quantity. Watch your significant figures.

[H^+] = 10^-pH	[OH^-] = 10^-pOH	pH = -Log[H^+]	pOH = -Log[OH^-]
[H^+]	[OH^-]	pH	pOH	Acid/base/neutral
5.0 x 10^-3M	2.0 x 10^-12M	2.30	11.70	acid
7.7 x 10^-10M	1.3 x 10^-5M	9.11	4.89	basic
8 x 10^-4M	1 x 10^-11M	3.1	10.9	acid
3.22 x 10^-12M	3.10 x 10^-3M	11.492	2.508	basic
1.0 x 10^-7M	1.0 x 10^-7M	7.00	7.00	neutral (2sig figs)

2. What volume of .20M H₂SO₄ is needed to neutralize 50.00 ml .30M NaOH? 0.038L

3. What mass of NaF is required to prepare 100ml of .300M solution? 1.26g

4. 35.5 ml of .300M NaOH is required to neutralize 10.0 ml of H₂SO₄. What is the acid concentration? .533M

5. 100 ml of .200 M HCl is mixed with 120ml of .200M NaOH. Calculate the pH of the resulting solution. pH=12.26

6. Calculate the Ka for phenolpthalein. 8x10^-10

7. The Ksp of AgOH is 6.8x10^-12. Calculate the pH. pH=8.42

8. The OH^- concentration in .10M NaCN is 2.7x10^-6M. Calculate the Kb from this information only. Kb=7.3x10^-11

9. Calculate the pH for .20M HCl. pH= 0.70

10. Calculate the pH for .10M Ba(OH)₂. pH=13.30

11. Calculate the pH for .40M HCN. pH=4.85

12. Calculate the pH for .40M Na₂CO₃. pH=11.93

13. What is the pH for .30M NaCl? pH=7.00

14. Calculate the pH of .20M NH₃. pH=11.28

15. Calculate the pH of .20M NH₄Cl. pH=4.98

16. Show by calculation if H₂PO₄^- is an acid or base (compare the Kb and Ka).

Ka (H₂PO₄^- ) =6.2x10^-8> Kb (H₂PO₄^- ) =1.3x10^-12 therefore an acid

17. Calculate the pH of a saturated solution of Mg(OH)₂ if the Ksp is 1.2 X 10^-11. pH = 10.46

18. A .50M NH₃ solution is found to have a OH^- concentration of 1.86 x 10^-3 M. Using this data only calculate the Kb. 6.9 x 10^-6

19. A .18M acid HX has a pH of 2.40. What is the Ka? 9.0 x 10^-5

20. The following data were recorded when 25.00ml of H₂SO₄ were titrated with .1030M NaOH. The volumes of NaOH used in three runs were: 46.06ml, 44.52ml, 44.54ml. Calculate the acid concentation. 9.173 x 10-2 M

21. The following data were recorded when 10.00ml of NaOH were titrated with .1030M H₂SO₄. The volumes of H₂SO₄ used in three runs were: 12.55ml, 12.55ml, 12.10ml. Calculate the base concentration. 0.2585 M

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