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Chapter 1.1: Chemical Changes 1

Chapter 1.1: Chemical Changes

In this chapter, you will learn how we may monitor the rate of a chemical reaction. You will also consolidate your prior learning concerning identifying and modifying the rate of a chemical reaction.

Activity 1.1: Prior Learning

Based on the work you have completed in S1 to S3, complete Prior Learning 1.1.

Rate of Chemical Reactions

• The rate of a chemical reaction is an indicator as to the speed of a chemical reaction.

• The rate of a chemical reaction is defined as unit change per unit time. • In practice, we cannot calculate the absolute rate of reaction at any given

point in time. Instead we must calculate the average rate of reaction between two time points.

Activity 1.2: Mass of Reactant

Mass is measured in the units of grams (g) or kilograms (kg). We can use changes in mass to track the progress of a chemical reaction.

As calcium carbonate reacts with hydrochloric acid, calcium chloride, water and carbon dioxide gas are produced. This can be represented as a word equation. (a) Write a word equation for this reaction.

Calcium chloride remains in the reaction vessel along with any water produced. Carbon dioxide gas is released into the atmosphere. As the carbon dioxide gas is released, atoms are being lost from the original matter present in the reaction vessel. As the reaction proceeds we see a decrease in the mass of the reaction vessel. The mass lost over time is equivalent to the mass of carbon dioxide gas released.

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• During the experiment we were in control of time – we decided when we were going to make an observation. In this case time is the independent variable. A variable is a factor we can change or observe during an experiment.

• We recorded the mass (in grams) of the apparatus at regular time points throughout the experiment. In this case the mass is the dependent variable.

• When we wish to display the information recorded during an experiment, it is often best to use a graph. The type of graph drawn is dictated by the nature of the independent variable.

• Time is a continuous variable – with values stretching from 0 to infinity. For all continuous independent variables we must draw a line graph.

• When drawing a line graph – the independent variable always forms the x-axis; the dependent variable always forms the y-axis.

(c) Draw a line graph of your results.

(d) Based on the shape of your graph, describe what you discovered during the experiment.

(e) Explain why the graph formed was this particular shape? Why was mass lost from the apparatus?

Mass (g)

Time (seconds)



• The rate of a chemical reaction can be described as the rate of change during the experiment.

• As the reaction between the calcium carbonate and acid proceeds, mass is lost from the apparatus. The loss of mass is due to the release of carbon dioxide gas from the reaction mixture.

• During the reaction we are observing the decrease in mass per unit time. We can calculate the average rate of the chemical reaction using, rate = mass lost (g) ÷ time (s)

• The units of rate are grams per second (g/s)

(f) Calculate the rate of the chemical reaction you observed by copying and completing the following table.

Time Interval (s)

Initial Mass (g)

Final Mass (g)

Change in Mass (g)

Length of time (s)

Rate of change (g/s)

0 to 60

120 to 240

360 to 540

Activity 1.3: Volume of Product

Volume is measured in cm3 and litres. Note 1 cm3 = 1 ml. We can use volume of product to track the rate of a chemical reaction.

As you will have witnessed in the last experiment, calcium carbonate reacts with acid to release a gas. This time instead of measuring the mass of gas lost from the apparatus, we will measure the volume of gas produced per unit time.

Materials:

Test tube rack stopwatch Boiling tube 3 marble chips Stopper with delivery tube 2 mol l-1 hydrochloric acid plastic container 50 cm3 plastic measuring cylinder

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(b) Draw a line graph of your results, ensure you use a suitable scale for each axis. Your line graph must use at least half the graph paper provided.

(c) Copy and complete the following table. Calculate the average rates of reaction for each time interval in cm3 per second.

Time Interval (s)

Initial volume (cm3)

Final volume (cm3)

Change in volume (cm3)

Length of time (s)

Rate of change (cm3/s)

0 to 60

120 to 240

200 to 240

(d) When was the rate of reaction greatest?

(e) Why is the rate of reaction greatest at this time?

• As a chemical reaction proceeds, the rate of the chemical reaction decreases. The initial rate of a chemical reaction is the greatest.

• A line graph shows illustrates how the rate of a chemical reaction changes as a reaction proceeds. The steeper the line on the graph, the greater the rate of the chemical reaction at that point in time.

Activity 1.4: Colour Change

For some chemical reactions it is only possible to record the time taken for the reaction to reach its conclusion or end point. In these reactions we can calculate the average relative rate of reaction for the whole reaction.

If this is the case, the average rate of reaction is represented as a function of the time taken.

Rate is calculate as 1 ÷ time taken to decolourise (units of relative rate are s-1).



• The shorter the time taken, the greater the value of the relative rate. • The longer the time taken, the lower the value of the relative rate.

An acidified solution of potassium permanganate is caused to change colour on contact with glucose solution. Potassium permanganate solution changes colour from purple to colourless as the reaction proceeds. If all other variables are kept constant, the time taken for the colour change reflects the concentration of glucose molecules present.

Materials:

test tube rack water 5 boiling tubes 0.0025 mol l-1 potassium permanganate solution stopper 1.0 mol l-1 hydrochloric acid 10 cm3 syringe 0.5 mol l-1 glucose solution 5 cm3 syringe 2 cm3 syringe

Method:

1. Label your boiling tubes A to E. 2. To each test tube add the following:

Tube Volume of Water (cm3) Volume of 0.5 mol l-1 Glucose

solution (cm3)

Volume of 1.0 mol l-1 hydrochloric

acid (cm3)

A 8 2 5 B 6 4 5 C 4 6 5 D 2 8 5 E 0 10 5

Note: first add the water using the 10 cm3 syringe, followed by the glucose using the same syringe. Use the 5 cm3 syringe for the hydrochloric acid.

3. To tube A, add 2cm3 of potassium permanganate. Immediately start the stop watch. Time how long in seconds it takes for the solution to become colourless.

4. Repeat with each of the other tubes B to E one tube at a time.



Results:

(a) Copy and complete the following table.

Tube Concentration of Glucose (mol l-1)

Time taken to decolourise (s)

Average Relative Rate of reaction (s-1)

A 0.1 B 0.2 C 0.3 D 0.4 E 0.5

(b) Draw a line graph of concentration of glucose (mol l-1) against average relative rate of reaction (s-1).

(c) What conclusion can you draw based on the results from your experiment?

(d) Why do your results present this pattern? What is causing the relative rate of reaction to be greatest?

• The rate of a chemical reaction is determined by the relative concentrations of reactant particles.

• The higher the concentration of reactant particles the greater opportunity there is for collisions between particles to occur.

• The more collisions there are, the more product(s) will be made in a shorter period of time.

• The more product(s) made in a shorter period of time, the greater the rate of a chemical reaction.

• Initially in any chemical reaction we will see the greatest rate. Rate decreases as time progresses as the concentration of reactant particles decreases.

Activity 1.5: Summary

Based on the work you have completed in this chapter, complete Summary 1.1


Chapter 1.2: Chemical Equations 9

Chapter 1.2: Chemical Equations

In this chapter you will learn how to write balanced chemical equations as a method of summarising what is occurring in a chemical reaction. As part of this, you will also consolidate your prior learning on writing chemical formulae, valency, word equations and formulae equations.



Chemical formulae are used as a shorthand method of representing different chemicals. Each substance has its own specific chemical formula, which shows the types and number of each type of atom present in the substance.

Activity 1.7: Chemical Formulae

1. Using valency rules, calculate the chemical formulae of the following simple substances.

a. sodium chloride b. magnesium oxide c. carbon chloride d. hydrogen oxide e. calcium phosphide f. aluminium chloride

2. Using valency rules, calculate the chemical formulae of the following substances.

You may find it useful to consult page 4 of the data booklet to help you.

a. lithium hydroxide b. magnesium sulphate c. aluminium hydroxide d. sodium hydrogencarbonate e. calcium phosphate f. potassium permanganate



3. Using valency rules, calculate the chemical formulae of the following substances, each of which contains a transition metal. As transition metals do not belong to a particular group, you may find it useful to consult page 7 of the data booklet to help you.

a. iron (III) oxide b. copper (II) chloride c. chromium (IV) sulphate d. mercury (I) dichromate e. zinc sulphate f. gold phosphide

4. Some substances do not always follow valency rules when it comes to writing

their chemical formulae. In these cases the names reflect the formulae of the substance. The names contain prefixes. The prefixes tell us the number of each type of atom present. Write the chemical formulae of the following substances.

a. carbon monoxide b. nitrogen trioxide c. dihydrogen pentoxide d. disodium tetraphosphate e. calcium pentphosphide f. carbon trichloride

Word equations are used as a summary for a chemical reaction. A word equation shows what you start out with (the reactants) and what you end up with (the products). An arrow is used to separate the reactants from the products, this tells us there has been a chemical change. We never use an equals sign between the reactants and products.

Activity 1.8: Word Equations

Write word equations for the following chemical reactions. If you need further assistance see “Thinking About Equations” pages 3 to 6.

1. Sodium metal reacts with oxygen from the air to form sodium oxide. 2. Sodium metal reacts with water to release hydrogen and sodium hydroxide. 3. Sulphur dioxide gas is released when the sulphur found in coal is burnt in the

presence of an unlimited supply of oxygen from the air.



4. Ammonia is manufactured by the Haber process, during which hydrogen and nitrogen combine over an iron catalyst at a high temperature.

5. Hydrogen gas and a salt are produced on reaction of a metal with an acid. 6. Sodium sulphate and water are produced when sodium hydroxide solution

neutralises sulphuric acid. The reaction is exothermic.

Formulae equations are used to display more information regarding a chemical reaction. They use chemical formula in place of words. They are the next step in working towards writing a balanced chemical equation.

When writing formulae equations it is important to remember some elements are diatomic in their natural state – those elements whose names end in –gen and –ine are all diatomic molecules in their natural state – H2, N2, O2, F2, Cl2, Br2, I2.

Activity 1.9: Formulae Equations

Write formulae equations for each of the following chemical reactions. You will first need to write a word equation and use valency or prefix rules to calculate the formula of each substance BEFORE writing the formula equation.

Your teacher may demonstrate some of these reactions to you.

If you need further assistance see “Thinking About Equations” pages 7 to 10.

1. Carbon powder burnt in an unlimited supply of oxygen produces carbon dioxide. 2. Sulphuric acid reacts with magnesium metal to produce hydrogen gas and

magnesium sulphate solution. 3. Aluminium powder is able to react with copper chloride solution to produce

aluminium chloride solution and solid copper metal. 4. Phosphoric acid reacts with zinc metal to produce hydrogen gas and zinc

phosphate solution. 5. Calcium oxide powder reacts with water to produce calcium hydroxide solution. 6. Aluminium and iodine react together to form aluminium iodide.

Balanced chemical equations form the basis of many different calculations. The ability to balance a chemical equation is an essential skill to carryout calculations. Before attempting to balance equations it is important you are able to write chemical formulae and form formula equations appropriately.



If you require further practice and consolidation for these please see Chapters 16 to 18 of “Questions for Standard Grade Chemistry” by Norman Conquest.

A balanced chemical equation is a detailed statement of what is occurring during a chemical reaction. It notes the chemicals and their relative quantities. From a balanced chemical equation reaction quantities can be calculated.

Process for writing a balanced chemical equation:

a) Write a word equation for the chemical reaction b) Substitute the names of the chemicals with their appropriate chemical formulae

derived from valency rules or prefixes. c) Check for balancing (some equations are already balanced and need no further

attention). d) Make a single alteration at a time – in the order metals, non-metals (except

oxygen and hydrogen), oxygen then hydrogen. Check for balancing after each change. Repeat the process until a balanced chemical equation is achieved.

Worked Examples:

(1) When nitric acid is neutralised by ammonium hydroxide solution, ammonium nitrate and water are produced.

a) Word Equation: nitric acid + ammonium → ammonium + water hydroxide nitrate

b) Formulae Equation: HNO3 + NH4OH → NH4NO3 + H20

c) Check for balancing: Reactants: Products: H x 6 H x 6 N x 2 N x 2 O x 4 O x 4

Equal numbers of each type of atom are found in the reactants and products, therefore the equation is already balanced.



(2) When magnesium is added to phosphoric acid, magnesium phosphate and hydrogen gas are produced.

a) Word Equation: magnesium + phosphoric → magnesium + hydrogen acid phosphate

b) Formulae Equation: Mg + H3PO4 → Mg3(PO4)2 + H2

c) Check for balancing: Reactants: Products: Mg x 1 Mgx 3 H x 3 H x 2 P x 1 P x 2 O x 4 O x 8

d) Change the number of Mg so that the metal atoms are balanced. 3Mg + H3PO4 → Mg3(PO4)2 + H2

Reactants: Products: Mg x 3 Mgx 3 H x 3 H x 2 P x 1 P x 2 O x 4 O x 8

e) Change the number of P so that the non-metal atoms (except O and H) are

balanced. 3Mg + 2H3PO4 → Mg3(PO4)2 + H2


The numbers of oxygen atoms are balanced, leaving only the number of hydrogen atoms to be balanced.



f) Change the number of H so that all atoms are balanced.

3Mg + 2H3PO4 → Mg3(PO4)2 + 3H2


Equal numbers of each type of atom are found in the reactants and products, therefore the equation is now balanced. As you become more confident in your abilities, you do not need to write out each step. Instead, only the balanced chemical equation is required as a final answer.

Activity 1.10: Balanced Chemical Equations Write balanced chemical equations for each of the following chemical reactions.

1. Copper metal reacts with chlorine gas to produce copper (II) chloride. 2. On heating silver (I) oxide decomposes to form silver metal and oxygen gas. 3. Magnesium metal reacts with dilute hydrochloric acid to produce magnesium

chloride and hydrogen gas. 4. When calcium chloride solution and sodium phosphate solution are mixed, solid

calcium phosphate and sodium chloride solution are formed. 5. Calcium carbonate reacts with hydrochloric acid to form carbon dioxide gas and

calcium chloride solution. 6. Ammonia solution reacts with water to form ammonium hydroxide solution.

Further practice of balancing chemical equations can be achieved by completing the following exercises from “Assessment Tests for Standard Grade Chemistry” by Chemcord.


Based on the work you have completed for this chapter, complete Summary 1.2.


Chapter 1.3: Bonding 15

Chapter 1.3: Bonding

In this chapter you will learn about the structure of a variety of different covalent and ionic substances. You will be able to classify chemicals dependent on their physical properties.



Covalent substances are made up of two or more non-metal elements joined together by strong covalent bonds. Ionic substances are made up of a metal and at least one non-metal element held together by electrostatic attractions.

By studying the physical properties of different substances we can classify them into three distinct groups – covalent molecular, covalent network and ionic.

Activity 1.13: Physical Properties

Copy and complete the following table and answer the questions which follow. You may find page 6 of the data booklet will help you.

Name of Compound Covalent or Ionic? Melting Point (°C) Boiling Point (°C) Ammonia Water

Silicon Dioxide Sodium Chloride Barium Chloride Calcium Oxide

1. Which of the substances shown above would conduct electricity in the solid state?

2. Which of the substances shown above would conduct electricity when molten or dissolved in water?

3. Draw a sectional diagram of the apparatus you would use to test the electrical conductivity of a molten compound.

4. What general statement can you make regarding the melting and boiling points of ionic compounds?

5. What general statement(s) can you make regarding the melting and boiling points of covalent compounds.

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Molecules of ammonia are held together by strong covalent bonds between the atoms which make up the molecule. Each covalent bond is a shared pair of electrons.

Covalent network solids, such as silicon dioxide, are held together by strong covalent bonds between the atoms which make up the network.

Ionic solids, such as sodium chloride, are held together by the electrostatic attractions between oppositely charged ions. Each sodium ion is attracted to six different chloride ions, and each chloride ion is attracted to six different sodium ions.

To melt or boil a covalent molecular solid requires us to give the molecules sufficient energy to move further apart. The molecules remain intact, and the strong covalent bonds within the molecule are not affected. Weak forces of attraction between the molecules are broken as the molecules move further apart. The weak forces of attraction between molecules are known as van der Waals forces.

To melt or boil a covalent network substance we must break the strong covalent bonds between the atoms which make up the substance, giving the atoms sufficient energy to move further apart. To break a strong covalent bond requires a great deal of energy. Heating a substance will transfer energy to its atoms, making them vibrate faster until eventually an atom has sufficient energy to break away.

To melt of boil an ionic substance requires a great deal of energy. This is due to fact that multiple electrostatic attractions between adjacent ions must be overcome before an individual ion has sufficient energy to move away from its neighbouring ions.

Activity 1.15: Shapes of Molecules

a) Collect a molymod kit. Make models of Carbon Dioxide (CO2), water (H2O), Ammonia (NH3) and Methane (CH4).

b) Draw a picture of each molecule, beside each draw the corresponding outer electron diagram and structural formula for each molecule.

• If the atoms are arranged in a straight line, we say the molecule is a linear molecule.

• If the atoms are arranged in such a way that when viewed in different planes (lines of vision) it appears differently, we say the molecule is planar.

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Activity 1.16: Formulae

1. Write the formula of each of the following substances. a. Silicon Dioxide b. Calcium Oxide c. Nitrogen Dioxide d. Potassium Iodide e. Ethane

2. Based on your knowledge of chemical structure, and the information from page 6

of the data booklet, answer the following questions. a. Which of the substances in question 1 exist as covalent molecules? b. Which of the substances in question 1 exist as covalent networks? c. Which of the substances in question 1 exist as ionic solids? d. How are the atoms arranged in covalent networks? e. How are the ions arranged in ionic solids? f. Is it always appropriate to use the simplest (empirical) formulae for

every substance?

For covalent molecular substances, the chemical formula gives a statement of the number of each type of atom found in the molecule. It is absolute.

For covalent network substances and ionic substances, the chemical formula gives the relative ratio of each type of particle expressed as a whole number. It is only a ratio and usually there are many thousands if not millions of each type of particle present in a sample of the substance.

When writing the formula of an ionic substance, it is sometimes required to include the charges of each ion. For an ionic compound, each ionic formula should be electrically neutral overall – the number of positive charges should equal the number of negative charges.

Potassium Chloride Potassium Sulphate

Symbols K Cl K SO4

Valency 1 1 1 2

Cross-over

Formula KCl K2SO4

Ionic Formula K+Cl- (K+)2SO42-

Brackets are used to “lock” the charge to the ion and avoid confusion.



Activity 1.17: Ionic Formulae

1. Using valency rules, calculate the simplest formula for each of the following ionic compounds.

a. Silver Chloride b. Magnesium Oxide c. Calcium Sulphate d. Lithium Hydroxide e. Sodium Phosphate f. Zinc Chloride

2. Re-write the formulae as ionic formulae – use brackets to “lock” charges to ions if multiples of the same ion are present.


Based on the work you have completed this chapter, complete Summary 1.3.

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Chapter 1.4: Moles and other Quantities 22

e.g. Formula mass of water H2O

• Step 1: Formula Water H2O

• Step 2: Elements hydrogen x 2 oxygen x 1

• Step 3: Relative Atomic Masses H = 1 x 2 = 2 O = 16 x 1 = + 16

• Step 4: Total 18

• Step 5: Answer Water has a formula mass of 18 a.m.u.

Activity 1.20: Formula Mass Calculation

Calculate the formula mass of each of the following examples.

1. chlorine gas 9. copper (II) sulphate

2. copper metal 10. mercury (II) nitrate

3. sodium chloride 11. ammonium phosphate

4. carbon dioxide 12. silver (I) nitrate

5. sodium phosphate 13. sodium hydroxide

6. ammonia 14. propane

7. sodium carbonate 15. butene

8. calcium nitrate 16. dibromoethane

Gram Formula Mass (GFM) of a substance is the same numerical value as the formula mass of the substance. Instead of the units a.m.u. GFM is expressed with the unit of grams (g).

The Gram Formula Mass of a substance is the same value (in grams) as the mass of one mole of the substance.

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e.g. 1. How many moles of sodium chloride are present in 46.8 g of the substance?

• Step 1: Formula NaCl

• Step 2: GFM 23 + 35.5 = 58.5 g

• Step 3: No. Moles = Mass ÷ GFM = 46.8 ÷ 58.5 = 0.8 moles

2. What is the mass of 2.4 moles of sodium hydroxide?

• Step 1: Formula NaOH

• Step 2: GFM 23 + 16 + 1 = 40 g

• Step 3: Mass = no. moles x GFM = 2.4 x 40 = 96 g

Activity 1.22: Mole Calculations

Calculate the following quantities, using the triangle shown above to help you.

1. Mass of 0.6 moles of hydrochloric acid.

2. No. moles of ammonia is present in 1.7 g of the gas.

3. Mass of 0.048 moles of iron (III) oxide

4. No. moles of ammonium sulphate in 52.8 g of the solid.

5. Mass of 0.6 moles of sodium hydroxide.

6. No. moles of sulphuric acid present in 49 g of the substance.

7. Mass of 3.8 moles of propane.

8. No. moles of mercury (II) nitrate present in 77.88g of the solid.

Concentration is a way of describing how much of a substance is dissolved in a solution. Concentration is measured using the units of moles per litre or mol/l or mol l-1.

Concentration = no. moles ÷ volume

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We may use Gram Formula Mass (mass of 1 mole of a substance) to calculate percentage composition and empirical (simplest) formulae. Example: Calculate the percentage composition of calcium bromide.

1. Find the formula for calcium bromide

calcium bromine

Ca Br

2 1

CaBr2

2. Calculate the Gram Formulae Mass (GFM) of calcium bromide

Ca x 1 = 40 x 1 = 40 Br x 2 = 80 x 2 = 160

TOTAL = 200 g

3. Calculate the % by mass of each of the elements, calcium and bromine, contributing towards the mass of 1 mole. Note it will always be 100% in total.

Calcium: 40 g out of 200g is made up of calcium atoms

Therefore % contribution Ca = (40 ÷ 200) x 100 = 20 %

Bromine: 160 g out of 200 g is made up of bromine atoms

Therefore % contribution of Br = (160 ÷ 200) x 100 = 80 %

Activity 1.25: Percentage Composition

1. What is the percentage composition by mass of the elements in ammonia?

2. What is the percentage composition by mass of the elements in aluminium sulphate?

3. What is the percentage composition by mass of ammonium phosphate?



An empirical formula is the simplest formula (ratio of atoms) which may be calculated for a chemical, based on experimental results.

Example:

Analysis of a compound in a fertiliser was found to be 26.1% nitrogen, 7.5% hydrogen and 66.4% chlorine by mass. What is the empirical formula of the compound?

The calculation is carried out using the following steps. It is often helpful to use a table to present the information carryout the calculation stepwise.

a. List the elements present.

b. Calculate the mass of each element present in 100 g of compound or the % composition of 1 mole of the compound.

c. Calculate the number of moles of each substance present in a 100 g sample of the compound, by dividing the mass present in 100 g of compound by the RAM of the element.

d. Calculate the mole ratio, by dividing the number of moles of each element by the smallest number of moles calculated.

e. Simplify and round up or down each number in the ratio to the nearest whole number. State empirical formula.

a. Elements Nitrogen Hydrogen Chlorine b. % composition

for each element

26.1 7.5 66.4

c. Number of moles of element present in 1 mole

= 26.1/14 = 1.86

= 7.5 / 1 = 7.5

= 66.4 / 35.5 = 1.87

d. Mole ratio = 1.86 / 1.86 = 1

= 7.5 / 1.86 = 4.03

= 1.87 / 1.86 = 1.01

e. Formula ratio 1 4 1 Empirical formula = NH4Cl



Activity 1.26: Empirical Formulae

1. A sample of an oxide of nitrogen was found to contain 1.88g of nitrogen and 3.16g of oxygen. Calculate the empirical formula for this compound.

2. Urea is a compound which is used as a fertiliser. Its composition by mass is 46.7% nitrogen, 6.7% hydrogen, 19.9% carbon and 26.7% oxygen. Calculate its empirical formula.

Further practice in calculating % composition of compounds and empirical formulae based on experimental results are available in “Questions for Standard Grade Chemistry” by Normal Conquest, page 83 onwards.


Based on the work you have completed in this chapter of work, complete Summary 1.4.


Chapter 1.5: Acids and Bases 31

Chapter 1.5: Acids and Bases

In this chapter you will consolidate prior learning regarding pH, acids and alkalis. You will learn to define a base, and describe the chemical reactions which they may be involved in.



pH is a numerical scale used to determine how acidic an aqueous solution is.

When a molecule of acid dissolves in water, it breaks apart to release at least one hydrogen ion. This increases the concentration of hydrogen ions in solution. When we are measuring the pH of a solution we are measuring the concentration of hydrogen ions in the solution.

We may use pH paper, indicators or electronic methods to determine the pH of a solution.

The pH number of a solution is a simple way of stating the concentration of hydrogen ions (H+ (aq)) found in a solution.

If we are using indicators to determine pH number, the indicator molecules change colour depending on the concentration of hydrogen ions in a solution.

Activity 1.29: Serial Dilution of an Acid

When we carry out a dilution, we take a small quantity of the concentrated solution and dilute it with a large quantity of water. This results in a more dilute solution. If we carefully monitor the volumes used, we can calculate the concentration of the resultant solution.

pH = -log10[H+(aq)] Note: we use [ ] to say concentration without having to write out the words

If we change the concentration of hydrogen ions in a solution by a factor of ten, we change the solutions pH by 1 pH unit.



Materials:

100 cm3 measuring cylinder 1.0 mol l-1 hydrochloric acid 10 cm3 measuring cylinder universal indicator 250 cm3 beaker 100 cm3 beaker stirring rod 8 test tubes test tube rack

Method:

1. Wash all equipment in cold tap water. 2. Half fill the 250 cm3 beaker with distilled water. This is your stock beaker. 3. Label your test tubes 0 to 7. 4. ¾ fill test tube 0 with 1.0 mol l-1 hydrochloric acid. 5. Using the 10 cm3 measuring cylinder, measure out 10 cm3 of 1.0 mol l-1

hydrochloric acid. Pour this into your 100 cm3 beaker. 6. Using the 100 cm3 measuring cylinder, collect 90 cm3 distilled water from your

stock beaker. Add this to your 100 cm3 beaker. 7. Stir well with the stirring rod. 8. ¾ fill test tube 1 with this solution. 9. Collect 10 cm3 of the contents of the 100 cm3 beaker, using the 10 cm3

measuring cylinder. 10. Pour the rest of the solution in the 100 cm3 beaker down the sink. Rinse the

beaker with cold tap water. 11. Pour the 10 cm3 of reserved solution into the 100 cm3 beaker. Add 90 cm3 of

distilled water and mix. 12. ¾ fill test tube 2 with this solution. 13. Repeat steps 9 to 12 until all test tubes contain a diluted solution. 14. Add 2 drops of universal indicator to each of the test tubes. Invert with your

thumb over the mouth of the test tube to mix. 15. Copy and complete the table shown overleaf. Use a colour chart to complete the

estimated pH number. Compare the colour achieved when the solution is mixed with universal indicator and read of the appropriate number from the colour chart.



tube number colour with universal indicator

estimated pH number

[H+(aq)] mol l-1

log10[H+(aq)] pH number

(-log10[H+(aq)]) 0 1 2 3 4 5 6 7

Water conducts electricity very slightly when present in its pure form. This is due to water molecules breaking apart to form hydrogen and hydroxide ions.

H2O (l) H+(aq) + OH-(aq)

The special double arrow tells us that there is a balanced situation or an equilibrium.

Pure water is neutral and has a pH of 7. This tells us that [H+(aq)] = 1 x 10-7 mol l-1.

One water molecule splits up to release a hydrogen ion and a hydroxide ion. Therefore in pure water, the concentration of hydrogen ions and hydroxide ions are equal.

[H+(aq)] = [OH-(aq)] = 1 x 10-7 mol l-1

For pure water we can calculate a constant – the ionic product of water.

[H+(aq)] x [OH-(aq)] = 1 x 10-7 mol l-1 x 1 x 10-7 mol l-1 = 1 x 10-14 mol2 l-2

Note: when multiplying using standard notation we ADD the powers when dividing using standard notation we SUBTRACT the powers

The ionic product of water remains constant in any aqueous solution. We can therefore use it to help us calculate the pH of a solution, given the concentration of hydroxide ions.

[H+(aq)] = 1 x 10-14 ÷ [OH-(aq)]

pH = -log10[H+(aq)]

[H+(aq)] = 1 x 10-pH

[H+(aq)] (mol l-1)

[OH-(aq)] (mol l-1)

1 x 10-14

(mol2 l-2)



Activity 1.30: Calculations

1. Calculate the pH of the following solutions. a. 0.01 mol l-1 H+(aq) b. 0.0001 mol l-1 H+(aq) c. 0.1 mol l-1 HNO3 solution d. 0.01 mol l-1 H2SO4 solution

2. Calculate the concentration of H+ (aq) ions in each of the following solutions.

a. pH 0 b. pH 3 c. pH 5 d. pH 14

3. Calculate the concentration of OH-(aq) ions in each of the following solutions.

a. pH 7 b. pH 9 c. 0.1 mol l-1 HCl solution d. 0.00001 mol l-1 HNO3 solution

Activity 1.31: Serial Dilution of an Alkali

Materials:

100 cm3 measuring cylinder 1.0 mol l-1 sodium hydroxide solution 10 cm3 measuring cylinder universal indicator 250 cm3 beaker 100 cm3 beaker stirring rod 8 test tubes test tube rack

Method:

1. Wash all equipment in cold tap water. 2. Half fill the 250 cm3 beaker with distilled water. This is your stock beaker. 3. Label your test tubes 14 to 7. 4. ¾ fill test tube 14 with 1.0 mol l-1 sodium hydroxide solution.



5. Using the 10 cm3 measuring cylinder, measure out 10 cm3 of 1.0 mol l-1 sodium hydroxide solution. Pour this into your 100 cm3 beaker.

6. Using the 100 cm3 measuring cylinder, collect 90 cm3 distilled water from your stock beaker. Add this to your 100 cm3 beaker.

7. Stir well with the stirring rod. 8. ¾ fill test tube 13 with this solution. 9. Collect 10 cm3 of the contents of the 100 cm3 beaker, using the 10 cm3

measuring cylinder. 10. Pour the rest of the solution in the 100 cm3 beaker down the sink. Rinse the

beaker with cold tap water. 11. Pour the 10 cm3 of reserved solution into the 100 cm3 beaker. Add 90 cm3 of

distilled water and mix. 12. ¾ fill test tube 12 with this solution. 13. Repeat steps 9 to 12 until all test tubes contain a diluted solution. 14. Add 2 drops of universal indicator to each of the test tubes. Invert with your

thumb over the mouth of the test tube to mix. 15. Copy and complete the table below. Use a colour chart to estimate the pH

number. Compare the colour achieved when the solution is mixed with universal indicator and read of the appropriate number from the colour chart. Does your estimated pH number agree with the calculated pH number?

tube number colour with universal indicator

[OH-(aq)] mol l-1

[H+(aq)] mol l-1

log10[H+(aq)] pH number

(-log10[H+(aq)]) 14 13 12 11 10 9 8 7

Activity 1.32: Conductivity and Concentration

As we dilute a solution of an acid we change the mass of acid present in a given volume of the solution. We are therefore changing the number of ions found in the solution.

What effect does this have on the ability of a solution to conduct electricity?

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• As we dilute an acid or alkali, we are decreasing the concentration of ions in the solution.

• As the concentration of ions decreases, fewer ions are available to carry an electrical charge through the solution so electrical conductivity decreases.

• We can use the electrical conductivity of an acid or alkali solution to estimate its concentration.

• Electronic methods of measuring pH use conductivity to determine the pH of a solution.

• The lower the concentration of a solution, the slower the rate of a chemical reaction it can participate in.

• 4 mol l-1 hydrochloric acid reacts with calcium carbonate powder much faster than 0.1 mol l-1 hydrochloric acid.

Neutralisation reactions all involve increasing the pH of an acid towards 7. The reagent used to neutralise an acid is known as a base. Bases include metal oxides powders, metal carbonate powders and alkalis. Alkalis are solutions of soluble bases.

Activity 1.33: Reaction of an Acid and Metal Oxide

Metal oxides react with an acid to produce a salt and water.

(a) Write a word equation for this general reaction.

The name of the salt produced is dependent on the metal oxide and acid used in the reaction:

• hydrochloric acid produces chloride salts • sulphuric acid produces sulphate salts • nitric acid produces nitrate salts • phosphoric acid produces phosphate salts



Materials:

two boiling tubes copper (II) oxide powder test tube rack 0.1 mol l-1 sulphuric acid spatula filter funnel stirring rod filter paper 10 cm3 measuring cylinder 250 cm3 beaker

Method:

1. Using the measuring cylinder, measure 10 cm3 of 0.1 mol l-1 sulphuric acid into one of the boiling tubes.

2. Half fill the beaker with hot water from a freshly boiled kettle. 3. Place your boiling tube of acid in the beaker, and leave to heat for 5 minutes. 4. Add ¼ spatula measure of copper (II) oxide powder to your hot acid, and stir

gently using your glass rod. Observe for signs of a chemical reaction. 5. Continue adding ¼ spatula measures of copper (II) oxide until no further

chemical changes occur, and black powder collects at the bottom of the boiling tube.

6. Filter your reaction mixture. Collecting your filtrate in your second boiling tube. 7. Leave your filtrate in a crystallising dish on the window ledge for 24 hours.

(b) Write a balanced chemical equation for the reaction of copper (II) oxide with

sulphuric acid, producing copper (II) sulphate and water.

(c) What signs were there that a chemical reaction had occurred?

(d) How did you know when the reaction was complete?

(e) Draw a sectional diagram of the filtration apparatus you used.

(f) Why is it possible to recover the unreacted copper(II) oxide from the reaction mixture?

(g) Draw a diagram of your crystals. Write a description of your crystals.



Activity 1.34: Reaction of an Acid with a Metal Carbonate

Metal carbonates react with an acid to produce a salt, carbon dioxide gas and water.

(a) Write a word equation for this general reaction

Materials:

two boiling tubes zinc (II) carbonate powder test tube rack 0.1 mol l-1 sulphuric acid spatula filter funnel stirring rod filter paper 10 cm3 measuring cylinder

Method:

1. Using the measuring cylinder, measure 10 cm3 of 0.1 mol l-1 sulphuric acid into one of the boiling tubes.

2. Add ¼ spatula measure of zinc (II) carbonate powder to your acid, and stir gently using your glass rod. Observe for signs of a chemical reaction.

3. Continue adding ¼ spatula measures of zinc (II) carbonate until no further chemical changes occur, and white powder collects at the bottom of the boiling tube.

4. Filter your reaction mixture. Collecting your filtrate in your second boiling tube. 5. Leave your filtrate in a crystallising dish on the window ledge for 24 hours.

(b) Write a balanced chemical equation for the reaction of zinc (II) carbonate with

sulphuric acid, producing zinc (II) sulphate, carbon dioxide and water.


(d) How did you know when the reaction was complete?

(e) Draw a diagram of your crystals. Write a description of your crystals.



Activity 1.35: Reaction of an Acid and Alkali Alkalis are solutions produced by dissolving a soluble base in water. When an acid reacts with an alkali, a salt and water are produced. The reaction is exothermic.

(a) Write a word equation for this general reaction.

Materials:

Dimple tile pH paper polystyrene cup 1.0 mol l-1 hydrochloric acid two 50 cm3 measuring cylinders 1.0 mol l-1 sodium hydroxide thermometer pH colour chart

Method:

1. Place a drop of hydrochloric acid in a dimple of the dimple tile and add a small piece of pH paper.

2. Compare the colour achieved to the colour chart and read off the pH number. 3. Repeat steps 1 and 2 using the sodium hydroxide solution. 4. Using separate measuring cylinders, measure out 25 cm3 volumes of each

solution. 5. Find the temperature of each solution – washing and drying the thermometer

between solutions. 6. The average of the two temperatures gives the starting temperature. 7. Pour both solutions into the polystyrene cup and stir gently with the

thermometer. Record the maximum temperature reached during the reaction between the acid and alkali. Calculate the change in temperature during the reaction.

(b) Write a balanced chemical equation for the reaction of sodium hydroxide with hydrochloric acid, producing sodium chloride and water.




(d) In the example, the salt produced was sodium chloride. Which chemicals could be reacted together to produce the following salts:

i. potassium nitrate ii. lithium chloride iii. barium sulphate iv. ammonium sulphate

(e) Write balanced chemical equations for each of the examples above.


Based on the work you have completed this chapter, complete Summary 1.5.


Chapter 1.6: Quantities from Experiments 42

Chapter 1.6: Quantities from Experiments

In this chapter you will learn how to use a balanced chemical equation to interpret the results of an experiment and calculate quantities chemicals involved. You will be required to know how to balance a chemical equation before commencing this topic, although in the worked examples a balanced chemical equation is provided. The methods set out in this chapter may be applied to any chemical reaction as we use the universal unit of chemistry – the mole.

When asked to write ionic formulae for an ionic compound, we simply write the chemical formulae and add in the charges of the ions. Use brackets to lock charges to ions when writing the formula of a compound containing multiples of the same ion.

e.g. sodium phosphide

Chemical Formula: Na3P

Ionic Formula: (Na+)3P3-

If we dissolve an ionic compound in water, the ions are released from the lattice and are free to move.

Ionic formulae are empirical or simplest ratios. Solid sodium chloride (NaCl) contains one sodium ion for every chloride ion. If we dissolve 1 mole of solid sodium chloride in water, this will release 1 mole each of sodium ions and chloride ions into the solution.

We can represent this as an ionic equation:

Na+Cl-(s) + H20(l) → Na+(aq) + Cl-(aq) + H2O(l)

H2O(l) appears on both sides of the arrow, suggesting it is not chemically involved in the break-up of the ionic lattice. We can simplify the equation by omitting H2O(l).

Na+Cl-(s) → Na+(aq) + Cl-(aq)

In a reaction mixture, there will often be ions which remain unchanged during the chemical reaction. Such as ions are present but are not directly involved in the chemical reaction. These ions are said to be spectator ions.

Consider the reaction between 0.1 mol l-1 sodium hydroxide solution and 0.1 mol l-1 hydrochloric acid solution.



a. Word equation:

sodium hydroxide + hydrochloric → sodium chloride + water solution solution solution

b. Balanced chemical equation:

NaOH + HCl → NaCl + H2O

c. Ionic equation (including state symbols):

Na+(aq) + OH-(aq) + H+(aq) + Cl-(aq) → Na+(aq) + Cl-(aq) + H2O(l)

d. Identify and remove any spectator ions:

Na+(aq) + OH-(aq) + H+(aq) + Cl-(aq) → Na+(aq) + Cl-(aq) + H2O(l)

e. Write the final overall ionic equation:

OH-(aq) + H+(aq) → H2O(l)

Activity 1.37: Spectator Ion Equations

For each of the following chemical reactions write an ionic equation, remembering to remove any spectator ions. Follow the steps laid out above.

a. word equation b. balanced chemical equation c. ionic equation (including state symbols) d. identify and remove an spectator ions e. write the final overall ionic equation

1. potassium hydroxide solution reacting with nitric acid solution. 2. sodium hydroxide solution reacting with sulphuric acid solution. 3. ammonium hydroxide solution reacting with hydrochloric acid solution.



4. calcium hydroxide solution reacting with hydrochloric acid solution.

In S1 to S3 you performed a number of qualitative chemistry experiments. Qualitative experiments focus on if a noticeable change has occurred e.g. has a gas been released or not. Some of the experiments examined earlier in S4 were quantitative. Quantitative experiments require you to measure and record a specific quantity e.g. volume, mass, concentration or optical density.

Scientists use the quantitative technique of titration to determine the concentration of a solution. With all titrations, a standard solution of known concentration is required to be used to determine the unknown concentration using a known reaction mechanism.

Activity 1.38: Titration Technique

Aim: To determine the concentration of sodium hydroxide of unknown concentration, by titration with 0.10 mol l-1 hydrochloric acid solution.

Materials:

50 cm3 burette burette clamp and stand 10 cm3 pipette 0.10 mol l-1 Hydrochloric Acid solution pipette filler Sodium Hydroxide solution three 100 cm3 beakers phenolphthalein solution filter funnel white tile

Method:

1. Assemble the apparatus as demonstrated by your teacher. 2. Place the filter funnel in the burette, and ensure the tap is closed. 3. Lift the burette stand onto the floor – to ensure the filter funnel is below eye

level. 4. Fill the burette with acid – until just below the 0 cm3 mark at the top of the

burette. 5. Label one beaker “acid” and another “alkali”.



6. Collect half a beaker of hydrochloric acid and half a beaker of sodium hydroxide, into the appropriate beakers.

7. Copy the table below. 8. Using the pipette transfer 10 cm3 of sodium hydroxide to the conical flask from

your “alkali” beaker. 9. Add 3 drops of phenolphthalein solution, swirl to mix. Place the conical flask on

the white tile. 10. Note the burette reading and record this in your table in the “starting” value

for the rough titration. Remember the scale starts with 0 cm3 at the top of the burette and runs to 50 cm3 at the bottom of the burette.

11. Add approximately 1 cm3 of acid to the conical flask by gently opening the tap at the bottom of the burette. Close the tap and gently swirl the contents of the flask. Repeat this until the indicator changes from purple to colourless.

12. Note the burette reading and record this in your table in the “final” value for the rough titration.

13. Calculate the titre (volume of acid) in cm3 required to neutralise 10 cm3 of sodium hydroxide solution.

14. Empty the contents of your flask down the sink and rinse your flask with cold tap water.

15. Repeat steps 8 and 9. 16. Note the burette reading and record this in your table in the “starting” value

for the 1st accurate titration. 17. Open the tap of the burette and run acid into the alkali, gently swirling, until

within 2 cm3 of the rough titre value. e.g. rough titre of 9.6 cm3, 1st accurate starting reading was 12.4 cm3, therefore run burette until reading reaches 20.0 cm3.

18. Begin to add acid drop by drop, swirling with after each addition, until the indicator changes colour from purple to colourless.

19. Note the burette reading and record this in your table in the “final” value for the 1st accurate titration.

20. Calculate the titre (volume of acid) in cm3 required to neutralise 10 cm3 of sodium hydroxide solution.

21. Repeat steps 14 to 20 until you have two accurate titre values to within 0.2 cm3.



Results:

Rough Titration 1st Accurate Titration

2nd Accurate Titration

Starting (cm3) Final (cm3) Titre (cm3)

Calculate Average Accurate Titre (cm3) = (1st accurate titre) + (2nd accurate titre) 2

We may use this value in further calculations – to determine the concentration of the sodium hydroxide solution.

In a demonstration titration, it was found that an average accurate titre value was equal to 10.5 cm3. As we know the concentration of the acid, we are able to calculate the concentration of the sodium hydroxide solution.

a. Word Equation

sodium hydroxide + hydrochloric → sodium chloride + water solution solution solution

b. Balanced Chemical Equation

NaOH + HCl → NaCl + H2O

c. Mole ratio

1 mole + 1 mole → 1 mole + 1 mole Moles are the universal unit in Chemistry. We use them to calculate reaction quantities. The numbers we use to balance an equation refer to the numbers of each type of particle. We are able to assume 1 particle of a formula unit can be multiplied up to represent 1 mole of a substance.



If a number two is in front of a particle in the formula equation, when writing the mole ratio we say there are two moles of a substance present. If no numbers are written in front of a particle to balance the chemical equation, then we assume there is only one of that particle required. Hence, 1 mole of the substance is required to balance the equation. d. Calculate the “known” number of moles

Before attempting to calculate numbers of moles, it is often useful to evaluate what information we already have relating to the reactants and products.

sodium hydroxide: volume = 10 cm3 concentration = ?

hydrochloric acid: volume = 10.5 cm3 concentration = 0.10 mol l-1

sodium chloride: volume = ? concentration = ?

water: volume = ? concentration = ?

Using the information from the experiment, it is possible to calculate the number of moles of hydrochloric acid taking part in the reaction.

No. moles = concentration x volume (volume is always in litres)

= 0.10 x (10.5/1000)

= 0.00105 moles HCl

e. Calculate the “unknown” number of moles

Using the mole ratio above, it is possible to calculate the number of moles of sodium hydroxide which would react with the hydrochloric acid.

Mole ratio 1 mole hydrochloric acid : 1 mole sodium hydroxide

Therefore 0.00105 moles HCl : 0.00105 moles NaOH



f. Calculate the “unknown” quantity required

We were asked to find the concentration of the sodium hydroxide solution.

We used 10 cm3 for each titration.

Concentration = no. moles ÷ volume (volume is in litres)

= 0.00105 ÷ (10/1000)

= 0.105 mol l-1

Following the stages outlined above, calculate the concentration of the sodium hydroxide you used in your own titration. Check your answer with your teacher once you have attempted the calculation.

Activity 1.39: Titration Calculations

Use the same approach to answer the following questions. The steps are also listed below to remind you.

a. word equation b. balanced chemical equation c. mole ratio d. calculate the “known” number of moles e. calculate the “unknown” number of moles f. calculate the “unknown” quantity required

1. What volume of 0.5 mol/l sodium hydroxide solution would exactly neutralise

40 cm3 of 0.2 mol/l sulphuric acid?

2. 1.96 g of pure phosphoric acid is dissolved in water and the solution is exactly neutralised by a quantity of 0.2 mol/l sodium hydroxide solution. What volume of sodium hydroxide solution will be required for complete neutralisation of the acid?

3. A 20 cm3 sample of vinegar is titrated with standard 0.1 mol/l sodium hydroxide solution. Exactly 36.4 cm3 of the alkali is found to neutralise the acid sample. What is the concentration in mol/l, of the acid in the vinegar solution?



Moles are a universal and transferable unit in Chemistry. We can use moles to help us answer a variety of different quantitative calculations in Chemistry. If we have a balanced chemical equation we can calculate any quantity relating to a chemical reaction.

Activity 1.40: Conservation of Mass

Materials:

100 cm3 measuring cylinder sodium chloride 250 cm3 beaker Stirring rod Spatula Balance

Method:

1. Using the balance, find the mass of each beaker and the measuring cylinder. 2. Place the beaker on the balance pan, do not “zero” the balance. Place

approximately 5 grams of sodium chloride into the beaker. 3. Measure a volume of 80 cm3 of cold tap water into the measuring cylinder. Find

the mass of the measuring cylinder and water together. 4. Pour the water into the beaker and stir until the sodium chloride crystals are

dissolved. 5. Find the mass of the beaker and sodium chloride solution together. 6. Record your observations in a copy of the table below.

Results:

Mass (grams) Beaker

Beaker + sodium chloride crystals Mass of sodium chloride crystals

Cylinder Cylinder + water

Mass of 80 cm3 water Mass of solution

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c) Copy and complete the following table.

Mass (grams) a. flask b. hydrochloric acid (acid) c. calcium carbonate (marble) d. flask + acid + marble (at start) e. flask + reaction mixture (at end) f. “lost” mass g. gas syringe h. gas syringe + carbon dioxide i. carbon dioxide

d) Calculate the number of moles calcium carbonate present at the start of the experiment.

e) Calculate the number of moles of carbon dioxide produced during the experiment.

The process of collecting and measuring the mass of product produced during an experiment can be difficult and is open to error. Gas can easily be lost from the equipment. Errors may be made when finding the mass.

If we know the mass of calcium carbonate we start with, and the balanced chemical equation for the reaction, we may calculate the expected mass of carbon dioxide which could be produced.

e.g. If 2.03 g of calcium carbonate is reacted with excess hydrochloric acid, what mass of carbon dioxide gas would be produced?

a. Word Equation

calcium + hydrochloric → calcium chloride + carbon dioxide + water carbonate acid solution gas

b. Balanced Chemical Equation

CaCO3 + 2HCl → CaCl2 + CO2 + H2O



c. Mole ratio

1 mole + 2 moles → 1 mole + 1 mole + 1 mole d. Calculate the “known” number of moles

Before attempting to calculate numbers of moles, it is often useful to evaluate what information we already have relating to the reactants and products.

calcium carbonate: mass = 2.03 g GFM = can be calculated

hydrochloric acid*: volume = 80 cm3 concentration = 1.0 mol l-1

calcium chloride: volume = ? concentration = ?

water: volume = ? concentration = ?

carbon dioxide: mass = ? GFM = can be calculated

Using the information from the experiment, it is possible to calculate the number of moles of calcium carbonate taking part in the reaction.

*The hydrochloric acid is present in excess meaning more acid is present that is required for the reaction to occur.

No. moles = mass ÷ GFM (GFM = 40 + 12 + 48 = 100 g )

= 2.03 ÷ 100

= 0.0203 moles CaCO3

e. Calculate the “unknown” number of moles

Using the mole ratio above, it is possible to calculate the number of moles of carbon dioxide which would be produced with the number of moles of calcium carbonate present.

Mole ratio 1 mole calcium carbonate : 1 mole carbon dioxide

Therefore 0.0203 moles CaCO3 : 0.0203 moles CO2



f. Calculate the “unknown” quantity required

We were asked to find the mass of carbon dioxide gas produced.

mass = no. moles x GFM (GFM = 12 + 32 = 44 g)

= 0.0203 x 44

= 0.89 g of CO2 are produced

f) Following the stages outlined above, calculate the mass of carbon dioxide which should have been produced during the demonstration experiment. Check your answer with your teacher once you have attempted the calculation.

g) Was the calculated value close to the recorded mass of carbon dioxide?

h) What could have caused the difference in mass of carbon dioxide gas obtained?

Activity 1.43: Mass Problems

Use the processes outlined above to answer the following questions.

1. 2SO2 + O2 → 2SO3

7.8g of sulphur dioxide is completely converted to sulphur trioxide in the Contact Process. What mass of sulphur trioxide is obtained?

2. 3NO2 + H2O → 2HNO3 + NO The final stage in the industrial production of nitric acid (HNO3) involves the above reaction. What mass of nitrogen dioxide must react to produce 2.52g of nitric acid.

3. Iron (III) oxide can be reduced to iron by carbon monoxide. 20g of iron (III) oxide is treated with carbon monoxide, but only 80% by mass reacts. What mass of carbon monoxide will have been used in the process?

4. Ethene reacts with hydrogen iodide to form iodoethane, C2H5I. Under certain conditions, 3.5 g of ethene enters the reaction chamber every 24 hours, with 7.8 g of iodoethane being produced in that time. Calculate the percentage yield of product, by mass, for this process.




Based on the work you have completed in this chapter, complete Summary 1.6.

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