The Acidic Environment – 9.3.4 Review Questions

Part 1: Acid Theory

1. Outline the historical development of ideas about acids proposed by:

Antoine Lavoisier (1780s)

Humphry Davy (1815)

Svante Arrhenius (1884)

2. Outline three limitations of Arrhenius’ definition of acids and bases.

3. Outline the Bronsted-Lowry theory of acids and bases (1923). In your answer include definitions of acids and bases according to this theory and examples of acids and bases.

4. State three advantages of the Bronsted-Lowry theory of acids and bases over Arrhenuis’ theory.

5. Describe the relationship between an acid and its conjugate base, and a base and its conjugate acid. Show at least two examples of reactions that demonstrate conjugate acid/base pairs.

6. Complete the following table:

Acid Conjugate Base Base Conjugate AcidHCl Br-

CH3COOH CO32-

H2SO4 NH3

HNO3 H2OHNO2 OH-

H2SO3 HSO4-

H2CO3 HCO3-

H3PO4 H2PO4-

HSO4- HPO4

2-

HCO3- PO4

3-

H3O+ CH3COO-

H2OOH-

NH4+

NH3

7. Write equations to show reactions of the following acids in water. Only show the transfer of one (1) proton to water. Remember to ensure you correctly use single or reversible arrows.

HCl

CH3COOH

H2SO4

HNO3

H2SO3

H2CO3

NH4+

8. Write equations to show the reaction of the following bases in water. Only show the transfer of one (1) proton to water. They are all reversible reactions so the correct arrows must be used.

CO32-

NH3

PO43-

CH3COO-

9. Identify the species from the lists above which can act as both acids and bases according to the Bronsted-Lowry theory.

10. Define the term amphiprotic species. Include equations to show that water is amphiprotic.

11. Distinguish between the term amphiprotic and amphoteric.

12. Use equations to show that the hydrogen sulfate ion (HSO4-) is amphiprotic.

13. All bases must contain a pair of electrons to accept the H+ from an acid. Draw electron dot structures for the following bases, and thus demonstrate how they accept protons from acids.

NH3

CH3COO-

H2O

OH-

14. Because salts contain ions which can act as acids or bases, salts do not necessarily form neutral solutions

Strong acids form very weak bases while strong bases form very weak acids As a result, chloride ion from hydrochloric acid are very weak bases and do not

react with water Metal ions such as Na+, K+, Ca2+ and Ba2+ do not react with water. A weak acid, like acetic acid, forms a strong base, the acetate ion CH3COO-.

Use the above rules to identify and explain a salt that forms

A neutral solution

An acidic solution

A basic solution

15. A salt formed from a strong acid and strong base is neutral or almost neutrala. Give examples of 2 neutral salts

b. Write equations for the neutralisation reactions which form these salts

16. A salt formed from a weak acid and a strong base is alkaline (as the ions in the salt react with water to produce hydroxide ions).

a. Give examples of 2 alkaline salts

b. Write equations for the neutralisation reactions which form these salts

c. Write equations to show the reactions of ions that cause the solutions to be alkaline.

17. A salt formed from a strong acid and weak base is acidic (as the ions in the salt react with water to produce hydronium ions).

a. Give 2 examples of acidic salts

b. Write equations for the neutralisation reactions that form these salts.

c. Write equations to show the reactions of the ion that causes these solutions to be acidic.

18. A salt formed from a strong acid and a strong base is close to neutral (as one of the ions reacts to produce hydroxide ions and the other reacts to produce hydronium ions).

a. Give an example of a neutral salt formed from a strong and weak acid.

b. Write an equation for the neutralisation reaction that forms this salt.

c. Write equations to show the reactions of the ions that cause this solution to be close to neutral.

19. Identify neutralisation as a proton transfer reaction that is exothermic. Write an ionic equation to represent all neutralisation reaction and show the enthalpy term.

20. Explain how acid spills on a highway could be cleaned up minimum damage to the environment.

Part 2: Buffers

1. Define what is meant by the term buffer, and identify two examples.

2. Qualitatively describe the effect if buffers with reference to a specific example in a natural system.

3. Blood must remain in the pH range 7.35 – 7.45. The blood contains a carbonic acid/hydrogen carbonate buffer system.

a. Write equations to show how the carbonic acid is formed and how the hydrogen carbonate ion is formed.

b. A high concentration of carbon dioxide in the blood leads to an abnormal medical condition. Use the concepts of buffers and the above equations to explain how a healthy person prevents this condition.

c. Explain what would happen if the blood of a healthy person was exposed to high concentrations of an alkali.

Part 3: Titration

1. Explain the difference between qualitative and quantitative analysis.

2. Explain the difference between volumetric and gravimetric analysis. Are these examples of qualitative and quantitative analysis?

3. Define the term titration.

4. What is the equivalence point of a titration?

5. What is the endpoint of a titration?

6. How is an indicator chosen for a specific titration?

7. What is a standard solution?

8. Primary solutions are important in titrations.a. What is a primary solution?

b. Give two examples of primary standards (one should be a suitable acid and one should be a suitable base).

c. What should a primary standard be made from?

9. The following equipment is used in all tradition. Name each piece and state how each piece of equipment is rinsed just prior to its use.

1)

2)

3)

4)

10. Describe the techniques used in preparing a standard solution of a base and using this solution in a titration. Ensure that you include details of all steps involved in accurate determination of the unknown concentration of an acidic solution.

11. 15.43g of pure barium hydroxide was dissolved in water and made up to exactly 500mL in a volumetric flask. Calculate the concentration of the solution.

12. What mass of pure sulfuric acid must be dissolved in 100mL to make a 0.550 mol/L solution?

13. How many moles of HCl are there in 45.3mL of 0.148 mol/L hydrochloric acid solution?

14. 25.0mL of a solution of sodium hydroxide was pipetted into a flask, a few drops of a suitable indicator was added and the solution was titrated with 0.123mol/L sulfuric acid solution from a burette. 27.4mL of sulfuric acid was required to reach equivalence point. Calculate the concentration of the sodium hydroxide solution in mol/L and g/L.

15. 5.24g of anhydrous carbonate was dissolved in water in a volumetric flask and the volume made up to 250mL. 10.0mL was pipetted into a conical flask and titrated with HCl. 21.6mL was required to reach the end point. Calculate the concentration of the HCl solution.

The above HCl solution was then used to determine the concentration of an unknown barium hydroxide solution. 25.0mL of the barium hydroxide solution required 28.4mL HCl for exact neutralisation. Calculate the concentration of the barium hydroxide solution.

16. In order to determine the concentration of an HCl solution, a student carefully diluted 10.0mL to 250mL, then titrated 10.0mL of the diluted solution with 0.147mol/L of sodium hydroxide solution. 21.5mL was needed to reach the end point. Calculate the concentration of HCl.

17. Vinegar is a solution of acetic acid in water. The concentration of the acetic acid in the vinegar was determined by titration. The vinegar was systematically diluted by a factor of 5.

a. State the equipment needed for the systematic dilution.

b. 25.0mL of the diluted vinegar was placed in a conical flask. What piece of glassware would be used?

c. The dilute acetic acid was titrated with 0.100 mol/L NaOH. An endpoint was reached after a titre of 37.1mL.

i. What indicator should be used for this titration?

ii. Calculate the concentration of acetic acid in both the diluted and original vinegar solution.

iii. Calculate the concentration of acetic acid in the original vinegar in grams per 100mL.

d. Suggests a sequence of steps that would have been done prior to this titration to standardise the sodium hydroxide solution. Include in your explanation reasons why the sodium hydroxide could not have been used as a primary standard.

Part 4: Titration pH Graphs

For each graph below, assume the following: Acid in the burette is added to the base in the receiving flask All solutions have a concentration of 0.1mol/L

1. Describe and explain the shape of this curve. Suggest suitable indicator(s) for use in this titration.

2. Describe and explain the shape of this curve. Suggest suitable indicator(s) for use in this titration.

3. Describe and explain the shape of this curve. Suggest suitable indicator(s) for use in this titration.

4. Describe and explain the shape of this curve. Suggest suitable indicator(s) for use in this titration.

Consider the following titration curves.

A titration was conducted by adding NaOH from a Teflon-coated burette to HCL in a conical flask. The pH in the flask was recorded during the titration and Curve A was produced.

The student was instructed to use either litmus of bromothymol blue as a suitable indicator to determine the endpoint.

a) Explain which of the two above indicators is the most appropriate indicators for this titration.

A second titration was conducted by adding NaOH to a different acid. The pH in the flask was recorded during the titration and Curve B was produced.

b) Nominate a suitable indicator for this second titration.

c) Explain why the results of the first titration are not inaccurate if an indicator other than litmus or bromothymol blue is used.

d) State a general principle you must use in choosing a suitable indicator for any given titration.

2015 Paper

2014 Paper

2013 Paper

2012 Paper

CLOZE PASSAGE No 3 - Acids in our body and our foodComplete the following sentences using appropriate words or short phrases

a) Acids are proton …………………….

b) When HCl is added to water the ionisation produces what 2 ions

c) The IUPAC name for acetic acid

d) 2-hydroxypropane-1,2,3-trcarboxylic acid is commonly called

e) The pH scale begins at ………….. and ends at ………….

f) Complete the following equation pH =

g) How much more acidic is an acid pH 2 compared to pH 5

h) An acid that completely dissociates is strong or weak?

i) Equal concentrations of citric, acetic and Hydrochloric. Which will have the lowest pH

j) Cloudy ammonia is acid, alkaline or neutral

k) Before using a pH probe it must first be c……………………..

l) Phenolphthalein in an alkali would be what colour

m) A 0.1M acid has a pH of 2.5 Is the acid weak or strong

n) Choose 2 words to describe 10M Acetic acid. Weak,Strong,Dilute,Concentrated

o) When H2SO4 is added to water equilibrium lies to the L or R

p) The [H+] of 0.2M sulfuric acid is

q) Seawater has a pH of

r) An acidic oxide commonly used as a food additive

s) An acid found in the stings of ants and bees

t) Lactic acid is commonly found in ………

u) Vitamin C is a source of what acid

v) What mass of citric acid is required to make up 200ml of a 0.1M solution

w) Calculate the pH of 0.0035M HCl solution

x) Calculate the pH of a 0.0035M H2SO4 solution

y) Between readings, a pH probe should be washed with …………………………..

z) What is the new concentration when 20ml of 10M HCl is made up to 500ml(use eqn M1V1 = M2V2)

CLOZE PASSAGE No 4 - The History of AcidsComplete the following sentences using appropriate words or short phrases

a) Chemist who proposed that acids were oxygen containing substances

b) Humphry Davy proposed that ……………….. was the element present in all acids

c) Arrhenius theory states “An acid is a substance that ………….. in water to produce ………….. ions as the only …………………..ions in solution”

d) A water soluble base should contain ……………………… ions

e) Name the 2 chemists that poposed an acid is a proton donor

f) Using the LB model complete the equation HCl + H2O

g) State the conjugate acid for CO32-

h) State the conjugate base for NH4+

i) A species that can be either an acid or base is said to be ……………..

j) Identify if the following salts when added to water are Acid, Neutral or BasesSodium Chloride, Ammonium Chloride, Silver Acetate

k) The reaction of any species with water is called …………………..

l) The colour of universal indicator in Sodium Carbonate

m) The name of the domestic acid substance you used to determined its concentration

n) 2 pieces of calibrated glassware is in titrations

o) Used to rinse the conical flask in a titration

p) The chemical used for our primary standard

q) The point at which the indicator changes colour

r) Give 2 examples of amphiprotic species

s) A suitable indicator for determining the conc of acetic acid in vinegar

t) A buffer is a mixture of a w……………. acid and its c……………………… base

u) An example of a naturally buffered system

v) 2 chemicals that could be added to dilute and neutralise an acid spill

w) A computer based technology used in titration experiments

x) The name given to the volume of liquid delivered from a burette

y) The name of the flask used to prepare a primary standard

z) In an ideal titration the ……………………….. point and end point coincide