1. Stoichiometry I
Chapter 11

2. Learning Objectives
By the end of this chapter you should be able to… 1. Calculate the % of each element in a compound 2. Calculate the empirical formula of a compound given the % composition. 3. Calculate the empirical formula of a compound given the masses of reactants and products. 4. Calculate the masses of reactants or products from balanced chemical equations. 5. Calculate volumes of gases from balanced chemical equations.

3. Stoichiometry
is the the branch of chemistry that tells us the relationship between the amounts of reactants and products in a chemical reaction.
stoicheoin = element
metron = measure

4. 1. To Calculate the % of each element in a compound
Steps:
Calculate the total mass of the element in the compound
Calculate the relative molecular mass of the compound
Divide:
Mass of element
x 100
Mass of compound

5. Example 1(a)
Calcium oxide, CaO is commonly called lime or quicklime. Calculate the percentage composition by mass of each element in calcium oxide, CaO.

6. Example 1(b)
Ammonium nitrate, NH4NO3, is often used in fertilisers to supply nitrogen to the soil. What is the percentage of nitrogen in ammonium nitrate?

7. Example 1(c)
Calculate the percentage of H2O in copper(II) sulfate crystals, CuSO4.5H2O.

8. DIY
Page 149: Q11.1 and Q11.2

9. Calculations involving the Empirical Formula of a Compound
The Molecular Formula of a compound is a formula which shows the number and type of each atom present in a molecule of that compound. E.g: H2O.
The Empirical Formula of a compound is the formula showing the simplest whole number ratio of the numbers of different atoms present in the molecule.
Rarely used in chemistry but we have to be able to calculate it!

10. e.g. the molecular formula for ethene is C2H4 the empirical formula for ethene is CH2 e.g. the molecular formula for hydrogen peroxide is H2O2 the empirical formula for hydrogen peroxide is HO

11. Molecular formula = empirical formula x n
2. How to Calculate Empirical Formula when given percentage composition:
Molecular formula = empirical formula x n
(where n is a whole number)
Molecular mass = empirical mass x n

12. Example 2(a)
Glucose is the name of the sugar found in blood. The empirical formula of glucose is CH2O. When the relative molecular mass of glucose is measured by a mass spectrometer, it is found to be 180. What is the molecular formula of glucose?

13. Example 2(b) – Variation of the Same Q
Glucose contains 40% carbon, 6.67% hydrogen and 53.33% oxygen by mass. Find it’s empirical formula. Lay out the question using a table:
Element %
Percentage
Simplest ratio
Carbon
Hydrogen
Oxygen
Atomic Mass
Take smallest number from this column and divide everything by it
You get marks for this table in the exam!

14. Example 2(b) – Variation of the Same Q
Glucose contains 40% carbon, 6.67% hydrogen and 53.33% oxygen by mass. Find it’s impirical formula. Lay out the question using a table:
Element %
Percentage
Simplest ratio
Carbon
Hydrogen
Oxygen
Atomic Mass
Empirical Formula =

15. Example 2(c)
A compound containing only carbon, hydrogen and oxygen was analysed and found to contain 65.11% carbon, 8.83% hydrogen, 26.06% oxygen. What is the empirical formula of the compound?

16. Example 2(d)
A compound containing only carbon, hydrogen and oxygen was found on analysis to contain 52.17% carbon, 13.04% hydrogen, and 34.79% oxygen. What is the empirical formula of the compound? Using a mass spectrometer, it was found that the relative molecular mass of the compound was 92. What is the molecular formula?

17. DIY
Page 149: Q11.3, 11.5, 11.6

18. Analysis based on the masses of compounds that react
3. Calculate the Empirical formula of a Compound given the masses of reactants and products
Gravimetric Analysis
Analysis based on the masses of compounds that react
Step 1: Write a word equation
Step 2: Work out the mass of the missing reactant
Step 3: Work out number of moles of each reactant to see the ratio they react in

19. Example 3(a)
A student heated 1.02g of copper powder in air and noted that the brown copper powder turned black. It was found that the mass of black powder formed was 1.28g. Assuming that the black powder is an oxide of copper, calculate the empirical formula of this oxide. (Cu = 63.5, O = 16).

20. Take Note:
If the ratio had turned out as 2:1 the empirical formula of copper oxide would have been Cu2O (see example 11.8, pg 144)

21. Example 3(b)
A student heated 1.61g of hydrated sodium sulfate crystals, Na2SO4.xH2O over a bunsen burner until no further loss in mass occurred. It was found that 0.71g of the anhydrous salt was obtained. Calculate the formula of the hydrated salt.

22. Example 3(c)
When 1.2g of metal M were reacted with oxygen, it was found that 2.0g of the oxide MO were formed. Calculate the relative atomic mass of M.

23. DIY
Pages 149 – 150 Questions 11.7, 11.8, 11.9, 11.10

24. 4. Calculating Masses of Reactants or Products from Balanced Chemical Equations
Example 4(a)
Magnesium reacts with oxygen to produce magnesium oxide according to the equation
2Mg + O2  2MgO
If a student burns 9g of magnesium in excess oxygen (i.e. there is enough oxygen present to ensure that all of the magnesium reacts), what mass of magnesium oxide will be formed?

25. Example 4(a) cont
Step 1: write equation and identify what you know and what you want to know.
2Mg + O2  2MgO
Step 2: Make shortened versions of equation

26. Example 4(b)
A student is told to make 100g of magnesium oxide by burning magnesium in excess oxygen. How much magnesium should the student burn?
Magnesium reacts with oxygen according to the equation
2Mg + O  2MgO

27. Example 4(c)
Lime (calcium oxide) is manufactured in a lime kiln by heating limestone (calcium carbonate) to bring about the following reaction: CaCO3  CaO + CO2 Calculate the amount of lime that could theoretically be obtained by heating 5kg of limestone.

28. DIY
Page 150, Q11.11, 11.12, 11.13

29. 5.Calculating Volumes of Gases from Balanced Chemical Equations
1 mole of any gas = 22.4L at s.t.p.
We can use this to calculate the volumes of gases used up or formed in chemical reactions.

30. Example 5(a)
Carbon dioxide is formed when carbon is burned in air according to the equation: C + O2  CO2 What volume of carbon dioxide (at s.t.p.) is formed when 30g of carbon are burned in excess oxygen?

31. Example 5(b)
Hydrogen gas may be prepared by reacting zinc metal with dilute sulfuric acid. The following reaction occurs:
Zn + H2SO4  ZnSO4 + H2
What volume of hydrogen gas at s.t.p. is evolved when 7g of Zn are added to excess dilute sulfuric acid?

32. Example 5(c)
Methane (natural gas) burns in air to form carbon dioxide and water. The equation for the reaction is: CH4 + 2O2  CO2 + 2H2O What volume of oxygen, measured at s.t.p., would be required for the complete combustion of 25g of methane?

33. DIY
Page 150 Q11.14, 11.15, 11.16