GRADE 9 HOMEWORK BOOKLET – CHEMISTRY

Week 1 Unit – Acid/ Bases and Salts Convert the following equations into chemical equations, complete and balance them.

1. zinc + hydrochloric acid ___________________ + hydrogen

2. ?_______________________ + sulphuric acid copper(II) sulphate + water + carbon dioxide

3. magnesium oxide + ?___________________________ magnesium nitrate + water

4. zinc + sulphuric acid zinc sulphate + ?___________________________

5. magnesium hydroxide + hydrochloric acid = ?___________________________ + water

6. ?__________________ + nitric acid copper(II) nitrate + water + ?__________________

7. zinc carbonate + sulphuric acid ?______________ + ?____________ + carbon dioxide

8. iron + ?___________________________ iron(II) chloride + hydrogen

9. magnesium oxide + hydrochloric acid ?____________________ + ?________________

10. ?_________________ + hydrochloric acid calcium chloride + water + ?__________________

11. ?___________________________ + sulphuric acid magnesium sulphate + hydrogen

12. magnesium + nitric acid ?___________________________ + hydrogen

13. zinc hydroxide + ?___________________________ zinc chloride + water

14. ?____________________ + hydrochloric acid magnesium chloride + water + carbon dioxide

15. aluminium + hydrochloric acid ?___________________________ + hydrogen

16. sodium hydroxide + hydrochloric acid ?_______________________ + ?________________

17. sodium carbonate + ?_________________ sodium sulphate + water + ?________________

18. ?___________________________ + nitric acid calcium nitrate + water + carbon dioxide

19. ?___________________________ + sulphuric acid iron(II) sulphate + hydrogen

20. zinc oxide + ?___________________________ zinc chloride + water

21. copper(II) carbonate + sulphuric acid?_____________________ + water + carbon dioxide

22. aluminium + ?___________________________ aluminium sulphate + hydrogen

23. calcium oxide + hydrochloric acid ?______________________ + ?_________________

24. magnesium hydroxide + ?________________ magnesium sulphate + ?________________

25. ammonia + sulphuric acid ?___________________________

26. ammonia + ?___________________________ ammonium chloride

27. zinc hydroxide + sulphuric acid ?_____________________ + ?______________

28. copper(II) oxide + hydrochloric acid ?___________________________ + water

29. sodium hydroxide + sulphuric acid ___________________________ + water

30. sodium hydrogencarbonate + hydrochloric acid

?________________ + ?_________________ + ?________________

31. ammonia + ?___________________________ ammonium nitrate

32. calcium oxide + ?_____________________ calcium chloride + ?_____________________

33. aluminium hydroxide + nitric acid ?___________________________ + water

34. calcium carbonate + ?________________ calcium ethanoate + ?_______ + ?_________

35. ethanoic acid + magnesium ?____________________ + hydrogen

36. _________________ + magnesium oxide magnesium ethanoate + ?____________

37. sodium hydroxide + ?_______________ ?_____ ethanoate + ?_____________

Week 2 Unit – Acid/ Bases and Salts

1) Copper sulphate crystals can be prepared in the lab by adding excess copper(II) oxide which is insoluble in water ,to dilute sulphuric acid. The mixture is heated.

a) Why is it necessary to use excess copper(II) oxide? b) How is the excess copper(II) oxide removed? c) Describe how copper sulphate crystals are formed from this solution.

2) What is a universal indicator? How is the use of universal indicator advantageous compared to any other indicator.

3) Explain the differences between strong acids and weak acids. Give two examples each. 4) Give one example each for a strong base and a weak base. 5) What are oxides? Mention the four types of oxides. Give an example each. 6) Give the ionic equation for the following reactions a) Zn + H2SO4 → ZnSO4 + H2 b)Na2CO3 + H2SO4 → Na2SO4 + CO2 + H2O

7) Change the ionic equation into a symbol equation. CuO + 2H+ → Cu

2+ + H2O

Qn 8)

[4] Qn 9)

Qn 10)

Week 3 Unit – Acid/ Bases and Salts

1) What happens when Al3+

, Zn2+

, Ca2+

, Fe2+

, Fe3+

, Cu2+

salt solutions are tested with the following a) NaOH is added (i) in drops (ii) in excess b) NH4OH is added (i) in drops (ii) in excess

2)Lead carbonate and lead iodide are insoluble in. which two soluble salts could you use in the

preparation of each substance? Write

a. A word equation

b. A symbol equation

c. An ionic equation to represent the reactions taking place.

3)An analytical chemist working for an environmental health organization has been given a sample

of water which is thought to have been contaminated by a sulphate, a carbonate and a chloride.

a. Describe how she could confirm the presence of these three types of salt by simple chemical

tests.

b. Write ionic equations to help you explain what is happening during the testing process.

4)copy out and complete the table, which covers the different methods of preparing salts.

Method of preparation

Name of salt prepared Two substances used in the preparation

Acid + alkali Potassium sulphate …………………and………….......

Acid + metal ………………………………. …………………and dilute HCl

Acid + insoluble base

Magnesium sulphate …………………and………….......

Acid + carbonate Copper……………………… …………………and………….......

Precipitation Lead iodide …………………and………….......

5)Write word and balanced chemical equation for each reaction shown in your table. Also write ionic

equations where appropriate.

1. Study the following scheme.

Sodium hydroxide calcium hydroxide ↓dil.H2SO4 ↓dil.HNO3 Solution A Solution B + + Water --------------------------------------------------------------- ---- Liquid C ↓ ↓ Mixed ↓ White precipitate D + Solution E

a. Give the names and formulae of substances A to E.

b. Describe a test which could be used to identify the presence of water.

c. Which indicator is suitable for the initial reaction between the hydroxides and the dilute acids

shown?

d. Write balanced chemical equations for the reactions taking place in the scheme.

e. Write an ionic equation for the production of the white precipitate D.

6) Kerry has four test tubes containing oxygen, hydrogen, carbon dioxide and nitrogen. The tubes were unlabeled. Suggest how she would identify each gas by using chemical tests. 7)

Qn 8

Week 4 : Unit – Acid/ Bases and Salts Prescribed text book pgs 130 & 131- Nos 2,3,8,9,10. Week 5 Unit –Chemical reactions. 1. Dilute nitric acid was added to a large amount of magnesium carbonate in a conical flask as shown.

The flask was placed on a balance and the mass of the flask and contents recorded every minute. The results are shown in the table.

(a) Make a similar grid with the same axes and on it Plot the results and draw a smooth line graph.[3]

(a) Which result appears to be inaccurate? Why have you selected this result? [2] (b) Why does the mass of the flask and contents decrease? [1] (c) Suggest the purpose of the cotton wool. [1] (d) At what time did the reaction finish? [1] (e) On the grid, sketch the graph you would expect if the experiment were repeated using nitric acid at a higher temperature. [2] (f)Explain how the rate of this reaction may be increased. [2] (g)Explain how the rate of this reaction may be decreased. [2]

(h)Explain how the graph changes when (i)temp (ii)conc of nitric acid (iii)particle size of magnesium carbonate is changed. [6] 2) In an experiment, a 2g lump of iron and 2g of powdered iron are added separately to equal volumes of dilute sulphuric acid. The solid line on the graph shows the volume of gas given off when the 2g lump is used. Which dotted line is obtained when the iron is powdered?

Prescribed text Pg 158,159- Qns 4,5,6,7,10. Week 6 Unit – Chemical reactions Prescribed text Pg 170,171- Qn nos , 4,5,6,9,10. Week 7 Unit – Chemical reactions Worksheet on redox reactions.

Week 8 Unit – Metals

1) Mention five general properties of metals. Not all metals share the typical properties. Name 2 metals each which are NOT a) hard and strong b) malleable at room temperature c) having a high density .

2)Describe the general characteristics of transition metals , and compare with those alkali metals with respect to the following melting point (b) boiling point (c) colour (d) colour of their compounds (e) use as catalyst (f) oxidation states.

3) Explain why the reaction of iron with hydrochloric acid is a redox reaction. Prescribed text – Pg 182 – Nos 1 & 2 Week 9 Unit – Metals

1) An investigation was carried out on the reactions of four different metals. Equal masses of copper, magnesium, iron and zinc were used.

Experiment 1 A 15 cm

3 sample of dilute sulphuric acid was added to each of four boiling tubes. The initial

temperature of the acid was measured. Zinc was added to the first tube, iron to the second tube, magnesium to the third tube and copper to the fourth tube. The maximum temperature reached in each tube was measured and any observations were recorded in the table.

(a) Use the thermometer diagrams to complete the results table.

Use your results and observations to answer the following questions. (i) Which metal is most reactive with sulphuric acid? (ii) Give two reasons why you chose this metal. (iii) Name the gas given off. The reaction between magnesium and aqueous copper(II) sulphate was then investigated. Experiment 2 A 5 cm

3 sample of aqueous copper(II) sulphate was measured into a test-tube. The initial temperature of

the solution was measured.Magnesium powder was added to the test-tube and the maximum temperature reached was measured. Use the thermometer diagrams to complete the results table.

(b) How do your observations show that the reaction of magnesium with aqueous copper(II) sulphate is exothermic? (c) What type of exothermic reaction occurs when magnesium is added to aqueous copper(II) sulphate? (d) Use your results from Experiments 1 and 2 to put the four metals in order of reactivity.

2. What happens when the product of burning sodium is added to water containing universal indicator and named the product formed in water.

3. (a) Give three observations that could be made when magnesium reacts with sulphuric acid.

(b) Give the word and symbol equation for zinc reacting with sulphuric acid.

(c) Give the test, and result, for the gas formed in the reaction between a metal and the acid.

4. (a) What substances must be in contact with iron before it will rust?

(b) What is the chemical structure of rust? (name and formula) and explain why it is an example of oxidation.

(c) Car bodies can be doubly protected from rusting with paint and galvanising. Describe and explain the two methods. Will the car body automatically rust if it is deeply scratched? explain your answer.

(d) Blocks of magnesium are bolted to a ships hull. Explain how this reduces steel corrosion. Why must there be no paint between the magnesium and the steel hull?

5. (a) Give three observations that could be made after adding zinc to copper(II) sulphate solution.

(b) give the symbol equation for the reaction between zinc and copper(II) sulphate solution.

(c) explain why this is an example of a displacement reaction.

(d) give the ionic equation for the reaction between zinc and copper(II) sulphate solution and explain the reaction in terms of oxidation, reduction and electron transfer.

6. Copper(II) carbonate is heated with carbon powder. (i) The green carbonate powder changes to a black powder at first, and (ii) after further strong heating, reddish-brown bits of a metal are formed.

(a) Describe with the aid of symbol equations what is happening in stage (i) and stage (ii).

(b) Why is stage (i) called a thermal decomposition? and stage (ii) called a reduction reaction?

7. A mixture of aluminium powder and iron(III) oxide is made and a magnesium strip placed in it.

(a) describe what happens after the magnesium fuse is ignited and what do the products look like.

(b) Give the symbol equation for the reaction and explain it in terms of oxidation and reduction.

(c) Mg + CuO ==> ??? Will these two react on heating Y/N? explain your answer.

(d) Given part of the metal reactivity series: (high) sodium magnesium zinc iron copper gold (low)

(i) where should hydrogen be placed? and explain why?

(ii) give a full equation to show iron dissolving in dilute sulphuric acid and name the products.

(iii) give the ionic equation to show iron dissolving in any acid.

(iv) why is gold found as the pure element?

Week 10 Metals 1)A) List and describe the four principal raw materials (including air) which are used in the blast furnace. Carefully explain the function of each raw material used. (B) Draw a diagram to show the outline of the structure of a blast furnace and mark on it the following labels: hot air, raw materials (ore, coke, limestone), hot waste gases , reaction zone, melting zone, slag, molten iron tapped off. (C) Describe, with the help of word/symbol equations, the four chemical reactions for the various stages of the process. (D) (ii) Describe some of the potential pollution problems from a blast furnace and how they might be reduced. (ii) What can be done with waste heat from the furnace? (E) the limestone reaction to remove silica impurities occurs in two stages (some texts combine these equations to give CaCO3 + SiO2 ==> CaSiO3 + CO2

(i) CaCO3 ==> CaO + CO2 then (ii) CaO + SiO2 ==> CaSiO3 What type of reaction is (i)? (iii) If calcium oxide is a basic oxide, what sort of oxide is silicon dioxide? (iv)Why do they react together and what sort of compound is calcium silicate? (F) Carefully explain why reaction for stage 3: iron(III) oxide + carbon monoxide ==> …. is an example of a REDOX reaction i.e. it involves an oxidation and a reduction. (G) (i) State what is oxidised in the reaction for stage 1. (ii) State what is oxidised, and to what, and what is reduced, and to what, in the reaction for stage 2 . (H) Other mineral oxides of iron can be used in a blast furnace e.g. iron(II) oxide FeO and tri-iron tetroxide Fe3O4 Give balanced symbol equations for the following word equations …

(i) Iron(II) oxide + carbon ==> iron + carbon dioxide (ii) Tri-iron tetroxide + carbon ==> iron + carbon dioxide (iii) Tri-iron tetroxide + carbon monoxide ==> iron + carbon dioxide

(2) Give two good reasons for recycling aluminium cans.

(3) Why should aluminium not be used for water pipes? [aluminium saucepans are out of favour too!]

(4) Give three good reasons for using aluminium alloys for aircraft wings.

(5) Aluminium is used for window frames but theoretically it is reactive enough to corrode away fairly rapidly. Explain why this does not happen.

(6) Explain why aluminium can displace iron from iron(III) oxide in the Thermit reaction. Why is this NOT an economic method for making iron?

(7) Aluminium (atomic number 13) is in Group III of the Periodic Table. What is the electron configuration of aluminium and its ion? What is the formula of aluminium chloride? Give the word/symbol equation for Al + HCl ==>

(8) Neglecting impurities, calculate the % of aluminium in bauxite ore.

(9) Chromium can be extracted from its compound by displacement with aluminium. Write an equation to show this reaction and calculate how many tonnes of aluminium are needed to make 108 tonnes of chromium. [Atomic masses: Al = 27; Cr = 52; formula of chromium oxide Cr2O3]

For Extra Questions refer to text book pg 182, 183 , 200, 201.

Week 11 Electrolysis Prescribed text Pg 142 – nos 1,2,3,4. Week 12 Electrolysis 1) The diagram shows the apparatus used to pass an electric current through concentrated hydrochloric acid.

(a) Label the electrodes.

(b) Give two observations when the current is switched on. (c) Give a test for the product at the negative electrode (cathode).

Prescribed text pg 142 – nos 5,6,7. Prescribed text pg 182 – nos 10,11,12 Week 13 Electrolysis 1)(A) What is the main ore of aluminium called? What is the name and formula of the principal aluminium compound in it? (B) Aluminium is extracted from its ore by electrolysis. What does this mean? (C) Sketch in outline the electrolysis cell, add labels to show: bauxite input, carbon anodes (+), electricity supply, aluminium output, carbon cathode (-), waste gases. (D) (i) Why is cryolite added to the bauxite ore in the electrolysis cell? (ii) Why is it expensive to produce aluminium by this method? (iii) Give the names and symbols of the two ions free to move in the molten ore. Where and why do theses ions move to? (iv) Give the two electrode equations to show the formation of aluminium at the cathode(-) and oxygen at the anode(+). Explain the changes in terms of oxidation and reduction . (v) Why do the carbon anodes have to be replaced regularly? Is there any danger of a toxic gas being produced? (vi) Is the electrolysis of molten aluminium oxide an exothermic or endothermic chemical change and explain your answer. Prescribed text pg 143 – Nos 8,9,10,11.