3.1 Formulae
1. Define the molecular formula of a compound.
The number and type of different atoms present in one molecule of the compound.
2. Define the empirical formula of a compound.
The simplest whole-number ratio of the different atoms or ions present in a compound.
3. How is the charge of an ion determined for Group 1, 2, and 3 elements?
The charge is based on valency: Group 1 loses 1 electron (+), Group 2 loses 2 electrons (2+), and Group 3 loses 3 electrons (3+).
4. How is the charge of an ion determined for Group 5, 6, and 7 elements?
Group 5 gains 3 electrons (3–), Group 6 gains 2 electrons (2–), and Group 7 gains 1 electron (–).
5. What is the formula and charge for the ammonium ion?
NH4+.
6. What is the formula and charge for the hydroxide ion?
OH–.
7. What is the formula and charge for the nitrate ion?
NO3–.
8. What is the formula and charge for the sulfate ion?
SO42–.
9. What is the formula and charge for the carbonate ion?
CO32–.
10. What are the charges for the transition metals iron(II) and iron(III)?
Fe2+ and Fe3+ respectively.
11. What is the charge for a zinc ion and a silver ion?
Zinc is Zn2+ and Silver is Ag+.
12. What is the formula for aluminium oxide?
Determine the charges (Al3+, O2–), cross the charges to become subscripts, resulting in Al2O3.
13. What is the formula for iron(II) nitrate, when should brackets be used?
Use brackets for a polyatomic ion if the ratio is more than one (e.g., Fe(NO3)2).
14. State the formula for lithium fluoride.
LiF.
15. State the formula for magnesium sulfide.
MgS.
16. State the formula for calcium nitride.
Ca3N2.
17. State the formula for barium sulfate.
BaSO4.
18. State the formula for silver chloride.
AgCl.
19. In naming covalent compounds, what prefixes are used for one, two, three, and four atoms?
mono- (one), di- (two), tri- (three), and tetra- (four).
20. State the formula for carbon monoxide and carbon dioxide.
CO and CO2.
21. State the formula for sulfur trioxide.
SO3.
22. Define valency.
The number of electrons lost, gained, or shared to achieve a full outer shell.
23. What is the valency of elements in Group 4?
4.
24. Construct a word equation and symbol equation for the reaction of carbon and oxygen.
carbon + oxygen → carbon dioxide; C + O2 → CO2.
25. What do the state symbols (s), (l), (g), and (aq) represent?
(s) solid, (l) liquid, (g) gas, and (aq) aqueous solution.
26. State the principle of conservation of mass in a chemical reaction.
The mass of any one element at the beginning of a reaction equals the mass of that element at the end of the reaction.
27. In what form are metals and non-metals typically written in a chemical equation?
Metals are written in atom form (e.g., Fe, Na), while non-metals are often in molecule form (e.g., O2, Cl2, H2).
28. Balance the equation for the formation of water: H2 + O2 → H2O.
2H2 + O2 → 2H2O.
29. Balance the equation for the formation of ammonia: N2 + H2 → NH3.
N2 + 3H2 → 2NH3.
30. Balance the equation for the reaction between sodium hydroxide and sulfuric acid.
2NaOH + H2SO4 → Na2SO4 + 2H2O.
3.2 Relative masses of atoms and molecules
31. Define relative atomic mass (Ar).
The average mass of the isotopes of an element compared to 1/12th of the mass of an atom of 12C.
32. Define relative molecular mass (Mr) and when is it specifically used?.
It is the mass of one molecule when compared with 1/12th of the mass of a Carbon-12 atom; it is used specifically for molecules.
33. What is the definition of relative formula mass, and when is it specifically used?
It is the mass of one formula unit when compared with 1/12th of the mass of a Carbon-12 atom; it is used specifically for ionic compounds.
34. What term is used for the Mr of ionic compounds?
Relative formula mass.
35. What is the relative mass of an electron compared to a proton or neutron?
An electron has a relative mass of 1/1840.
36. What is the unit for relative atomic mass (Ar) and relative molecular mass (Mr)?
These values have no units because they are relative comparisons.
3.3 The mole and the Avogadro constant
37. State the average mass of a chlorine atom given its isotopes.
35.5.
38. Describe how to calculate the percentage by mass of an element in a compound.
Divide the mass of the specific element in the formula by the total relative formula mass (Mr) of the compound and multiply by 100%.
39. Define the Avogadro constant.
It is the number of units (6.02 × 1023) in one mole of a substance.
40. What three types of "units" can a mole represent?
Atoms (e.g., Ne), molecules (e.g., O2), or formula units (e.g., NaCl).
41. Calculate the number of atoms in 2.5 moles of zinc (Zn).
2.5 × 6.02 × 1023 = 1.505 × 1024 atoms.
42. How many atoms are present in 0.01 mole of ammonium sulfate, (NH4)2SO4?
0.01 × 15 × 6.02 × 1023 = 9.03 × 1022 atoms (since there are 15 atoms in one formula unit).
43. Define molar mass and state its unit.
Molar mass is the mass of one mole of a substance in grams; its unit is g/mol.
44. State the calculation triangle for Mass, Mole, and Molar mass.
Mass = mole × molar mass; Mole = mass / molar mass.
45. Calculate the mass of 0.5 mole of hexane (C6H14) molecules.
43 g.
46. Calculate the number of moles in 4.61 g of lead iodide (PbI2).
0.01 mol.
47. What is the simplest ratio for the molecular formula of ethene (C2H4)?
CH2.
48. Define molecular formula.
A formula that shows the actual number of atoms of each element present in one molecule of the compound.
49. State the mathematical relationship between molecular formula and empirical formula.
Molecular Formula = (Empirical Formula)n.
50. In the experimental determination of the empirical formula for magnesium oxide, what is the purpose of the crucible lid?
The lid prevents the product from escaping as white smoke while being lifted occasionally to allow air to enter for the reaction.
51. In the magnesium oxide experiment, why must the crucible be cooled before weighing and then reheated?
To ensure all the magnesium has reacted and a constant mass is achieved.
52. If a sample of antifreeze is 38.7% carbon, 9.7% hydrogen, and 51.6% oxygen, what is its empirical formula?
CH3O.
53. Calculate the molecular formula of the antifreeze in the previous question if its Mr is 62.
C2H6O2.
54. Define molar volume of a gas.
It is the volume occupied by one mole of any gas.
55. What are the standard conditions for r.t.p. ?
25 °C and 1 atmosphere of pressure.
56. State the volume of 1 mole of any gas at r.t.p.
24 dm3.
57. How many cubic centimeters (cm3) are in one cubic decimeter (dm3)?
1000 cm3 = 1 dm3.
58. Find the volume of 0.01 mole of ammonia gas (NH3) at r.t.p.
0.24 dm3.
59. Calculate the number of moles in 1920 cm3 of hydrogen chloride (HCl) gas at r.t.p.
0.08 mol.
60. Describe the principle of conservation of mass in chemical reactions.
Matter cannot be created or destroyed; the mass of elements at the beginning of a reaction equals the mass at the end.
61. In stoichiometry, what is the limiting reactant?
The reactant that is completely consumed in a reaction and determines the maximum amount of product formed.
62. In stoichiometry, what is the excess reactant?
The reactant that remains after the reaction has stopped because there is more than enough to react with the limiting reactant.
63. Define concentration and its two common units.
It is the amount of solute per unit volume of solvent, measured in g/dm³ or mol/dm³.
64. State the formula for molar concentration.
Concentration (mol/dm3) = amount of solute (mol) / volume of solvent (dm3).
65. Calculate the concentration of a solution containing 10 g of NaOH in 500 cm3 of water.
0.5 mol/dm3.
66. How many moles of solute are in 100 cm3 of 16 g/dm3 copper(II) sulfate solution?
0.01 mol.
67. In a titration calculation, if 500 cm3 of HCl is neutralised by 250 cm3 of 1.0 mol/dm3 NaOH, what is the acid's concentration?
0.5 mol/dm3.
68. State the formula for percentage purity.
Percentage purity = (mass of pure substance / mass of impure substance) × 100%.
69. If a 64 g sample of impure calcium carbonate contains 51.2 g of the pure compound, what is its purity?
80%.
70. A 24 g sample of impure CaCO3 decomposes to produce 4800 cm3 of CO2. Calculate the percentage purity.
83.3%.
71. Define percentage yield.
Percentage yield = (actual yield / theoretical yield) × 100%.
72. Explain the difference between theoretical yield and actual yield.
Theoretical yield is the maximum amount of product calculated from the balanced equation; actual yield is the amount of product physically obtained from the experiment.
73. Calculate the percentage yield if 100 g of CaCO3 theoretically produces 56 g of CaO but only 50.4 g is obtained.
90%.
74. What is the percentage yield of magnesium oxide if 9 g of magnesium produces 14 g of MgO?
93.3%.
75. Define a hydrated substance and an anhydrous substance.
A hydrated substance is chemically combined with water; an anhydrous substance contains no water.
76. What is the specific term for the water molecules present in hydrated crystals?
Water of crystallisation.
77. How do you determine the value of 'x' in a hydrated formula like CuSO4 · xH2O?
Calculate the moles of the anhydrous salt and the moles of water removed during heating, then find their simplest whole-number ratio.
78. If 20.5 g of hydrated copper(II) sulfate is heated to leave 13.12 g of anhydrous salt, what is the value of 'x'?
5 (resulting in CuSO4 · 5H2O).
79. Define relative formula mass specifically for ionic compounds.
It is the mass of one formula unit compared with 1/12th the mass of a Carbon-12 atom.
80. Why is the term "relative formula mass" used for NaCl instead of "relative molecular mass"?
Because NaCl is an ionic compound and does not exist as discrete molecules.
81. What is the molar volume of any gas at room temperature and pressure (r.t.p)?
24 dm3.
82. State the units of molar mass.
g/mol.