IB 1 CHEMISTRYTeacher: Annika Nyberg , annika.nyberg@mattliden.fi

Course book:

Pearson Baccalaureate:Chemistry for the IB Diploma, HL or SL2nd edition

1. Stoichiometric Relationships

1.1 Introduction

NOS: The atomic theory

All matter is composed of atoms. These atoms cannot be created or destroyed during chemical reactions,

they can only be rearranged. Physical and chemical properties of matter depend on the bonding and

arrangement of these atoms.

The atom

Atoms are composed of subatomic particles: protons, neutrons and electrons.

In a neutral atom, the number of electrons equals the number of protons!

Matter:

States of matter

Temperature

The way the particles of matter move depends on the temperature.

1

Changes of state

Heating curve for water

Element:

Compound:

Pure substance:

Mixture:

Homogenous mixture:

Heterogenous mixture:

2

The language of chemistry

The International Union of Pure and Applied Chemistry (IUPAC) is an organization that develops a system of standardized nomenclature for chemical compounds.

See the IUPAC gold book for chemical terminology: http://goldbook.iupac.org/index.html

Ion:

Anion:

Cation:

Polyatomic ion:

Common polyatomic ions:

Name Formula

ammonium ion

carbonate ion

hydrogencarbonate ion

hydroxide ion

nitrate(V) ion

nitrate(III) ion

phosphate(V) ion

phosphonate ion

sulphate(VI) ion

sulphate(IV) ion

ethanedioate ion

peroxide ion

Common acids:

Name Formulahydrochloric acid

nitric(V) acid

phosphoric(V) acid

sulphuric(VI) acid

ethanoic acid

Ex.1 Write the formula for the hydrogen atom, the hydrogen molecule and the hydrogen ion.

3

Writing equations

Write the correct formulas for all the reacting species, reactants on the left-hand side and the products on the right-hand side.

Write the correct coefficients in front of each species. The reaction is then said to be stoichiometrically balanced.

It is good practice to include the state symbols: (s), (l), (g), (aq)

Ions that remain unreacted in the reaction (= spectator ions) can be left out from the equation.

Ex 2. Write an equation for the reaction between an aqueous solution of silver nitrate and an aqueous solution of potassium chloride.

• Law of conservation of mass: Atoms cannot be created or destroyed during a chemical reaction.

The atom economy

In an ideal chemical process the amount of products produced equals the amount of reactants and no atoms are “wasted”.

Some types of reaction

Combination or synthesis reactions: two or more substances are combined to produce a single product.

Decomposition reactions: a single reactant is broken down into two or more products.

Single replacement reactions: one element replaces another in a compound.

Double replacement: occur between ions in solution to form insoluble substances and weak or non-electrolytes.

4

Ex 3. Write equations for the following chemical reactions:

a) Zinc metal reacts with hydrochloric acid to form the salt zinc chloride. Hydrogen gas is evolved.

b) Hydrogen gas and oxygen gas react to form water.

c) Magnesium burns in oxygen to form a white powder known as magnesium oxide.

d) Iron reacts with chlorine gas to produce iron(III)chloride.

Ex 4. Hydrogen sulphide, H2S, reacts with oxygen to form sulphur dioxide and water as shown below.

2 H2S + _ O2 → _ SO2 + _ H2O

What is the whole number coefficient for oxygen when this equation is balanced?

A 1 C 3

B 2 D 6

Ex 5.

Ex 6.

Ex 7.

Ex 8.

5

Ionic compound:

Molecule:

1.2 The mole concept

The international system of measurement (SI)

The SI (Systeme International d´Unités) system has seven base units. All the other units are derived from them.

Property Unit Symbol

Useful prefixes:

Prefix Abbreviation Scale

Amount of a substance, n

Ex.9 Calculate the mass of one hydrogen-1 atom.

A single atom of an element has an extremely small mass, far too small to weigh.

Chemists measure amount of substances in moles, by counting particles (atoms, molecules or ions).

12,00 grams of carbon-12 contains 6,02 x 1023 atoms. This number is called the Avogadro's constant (NA or L).

One mole of ANY substance contains 6,02 x 1023 particles.

6

Isotopes

The atoms of an element are made up of an mixture of isotopes. Isotopes have:

Relative atomic mass, Ar (no unit)

The relative atomic mass of an element is the weighted mean mass of all the naturally occuring isotopes of that element relative to the mass of carbon-12.

For example, hydrogen has 1/12 of the mass of carbon-12 → 1 mol of hydrogen atoms has a mass of 1.01 g.

proton electron neutron

Mass (kg)

Relative mass

Ex 10. Rubidium exists as a mixture of two isotopes 85Rb and 87Rb. The percentage abundances are 72.1% and 27.9%, respectively. Calculate the relative atomic mass of rubidium.

Relative molecular mass, Mr

The relative molecular mass, Mr, of a molecule is the sum of the relative atomic masses of the atoms in the molecule (found in the periodic table).

Mr (C2H5OH) =

Relative formula mass, Mr

The relative formula mass, Mr, is similar to the relative molecular mass, but can be used with non-molecular substances such as ionic compounds.

Mr (AgNO3) =

Molar mass, M (g mol-1)

The mass of one mole of ANY substance is known as the molar mass.

For example, 1 mol of iron contains 6.02 x 1023 iron atoms and has a mass of 55.85 g mol-1.

Ex 11.

Calculate the mass of one mole of C2H5OH.

7

Ex 12. Calculate the molar mass of the following substances and ions:

a) Mg(NO3)2

b) Na2CO3

c) Fe2(SO4)3

d) I2

h) MgSO4 ▪ 7H2O

Ex 13. Calculate the amount (in mol) of carbon dioxide in a sample of 1.50 x 1023 molecules.

Ex 14.

Ex 15. a) How many atoms are there in 8 g of oxygen?

b) What is the mass of 1 mol of methane, CH4, molecules?

c) What is the mass of 0,0100 mol of caffeine, C8H9O2N4?

Formulas of compounds

Empirical formula:

Molecular formula:

Structural formula:

8

Percentage composition by mass

Ex 16. Calculate the percentage of mass of magnesium in magnesiumoxide.

Ex.18 A compound contains 12.79% carbon, 2.15% hydrogen and 85.06% bromine by mass. Its relative molecular mass is 187,9. Determine a) the empirical formula and b) the molecular formula of the compound.

Ex.19 Upon analysis, a sample of an acid with a molar mass of 194.13 g mol-

1 was found to contain 0.25 g of hydrogen, 8.0 g of sulfur, and 16.0 g of oxygen. Determine the empirical formula and the molecular formula.

Ex 20-23

9

1.3 Reacting masses and volumes

Chemical equations show reactants combining in fixed ratios (moles) to form products, e.g. methane burns in air:

CH4 (g) + 2 O2(g) → CO2 (g) + 2H2O (g)

Ex 24. Calculate the mass of carbon dioxide produced from the complete combustion of 1,00 g of methane.

The limiting reagent (often the most expensive) will be completely consumed during the reaction and determines the amount of products formed.

The remaining reactants are said to be in excess.

The balanced chemical equation can be used to predict how much product can theoretically be produced from given masses of starting material.

Theoretical yield:

Example:

1) Write the balanced equation

2) Convert the given mass of reactants to moles

3) Compare the given mole ratios with the ratio of coefficients in the equation

4) Identify the limiting reactant from the above ratios

5) Calculate the moles of product from the given moles of limiting reactant.

10

Ex 25. Calculate the mass of magnesium that can be obtained from the reaction between 4,8 g of magnesium and 4,8 g of sulfur. Identify the limiting reactant and calculate the mass of the unreacted element present in excess.

Ex 26. a) Write the balanced chemical equation for the reaction where ethene and steam react to produce ethanol (C2H5OH).

b) What is the maximum amount of ethanol that can be produced when 1,0 kg of ethene and 0,010 kg of steam are placed into the reaction vessel?

Percentage yield

Ex 27. 100,0 g of I2O5 reacts with 33,6 g of CO. Calculate the theoretical yield of carbon dioxide and given an experimental yield, in mol, of 0,900 mol CO2, calculate the percentage yield.

I2O5 (s) + 5CO (g) → I2 (g) + 5 CO2 (g)

Ex 28. Nitrogen gas (N2) can be prepared from this reaction:

2 NH3 (g) + 3 CuO(s) → N2(g) + 3Cu(s) + 3H2O(g)

a) If 18,1 g NH3 are reacted with 90,40 g CuO, determine the mass of N2 that can be formed.

b) An experiment produced 8,35 g N2. Determine the percentage yield.

Ex 29. In an experiment to produce a sample of hex-1-ene, 20.4 grams of hexan-1-ol was heated with an excess of phosphoric (V) acid. The phosphoric (V) acid acted as a dehydrating agent, removing water from the alcohol to form hex-1-ene.

CH3CH2CH2CH2CH2CH2OH → CH3CH2CH2CH2CH=CH2 + H2O

After purification of the hex-1-ene, 10.08 grams was produced. Calculate the percentage yield. (A: 60%)

11

Solutions

Some liquids are pure substances, but more commonly liquids are solutions containing two or more components.

Solution:

- solid/solid:

- solid/liquid:

- liquid/liquid:

- gas/liquid:

Solvent:

Solute:

Concentration, c

Standard solution:

Ex 30. Calculate the number of moles of chloride ions in 0.250 dm3 of 0.0200 mol dm-3 zinc chloride, ZnCl2, solution.

Ex 31. 250 cm3 of a solution contains 5.85 g of sodium chloride. Calculate the concentration of sodium chloride in mol dm-3.

Dilution of solutions

Stock solution:

12

Ex. 32. If 25,0 cm-3 of 0,56 mol dm-3 H2SO4 is diluted to a volume of 125 cm3, what is the concentration of the resulting solution?

Ex 33. Describe how you would prepare 2,00 dm3 of the following solutions:

a) 0,250 mol dm-3 of NaOH from solid NaOH

b) 0,250 mol dm-3 NaOH from of a stock solution of 1.00 mol dm-3.

c)1 mol dm-3 H2SO4 from concentrated (18 mol dm-3) sulfuric acid.

Density

density = __mass___ volume

Ex 34. A solution of ethanol (C2H5OH) in water is prepared by dissolving 75,0 cm3 of ethanol (density = 0.79 g/cm3) in enough water to make 250,0 cm3 of solution. What is the concentration (mol/l) of ethanol in this solution?

Ex 35. Calculate the number of moles of HCl in 1 dm3 of a concentrated solution of HCl (37 mass-% HCl) if the density of the solution is 1,18 kg dm-3.

Acid-base titration

Ex 35. Sodium hydroxide reacts with hydrochloric acid according to the following equation:

NaOH (aq) + HCl (aq) → NaCl (aq) + H2O (l)

Calculate the volume of 0,0500 mol dm-3 sodium hydroxide solution to react exactly with 25 cm3 of 0,20 mol dm-3 hydrochloric acid.

13

Avogadro´s law and the molar volume of a gas

Properties of gases:

Pressure:

Avogadro´s law:

14

Molar volume of a gas:

Standard temperature and pressure, STP:

Room temperature, SATP:

Molar volume of an ideal gas at STP:

Boyle´s law

Describes the relationship between volume and pressure At constant temperature, increasing the pressure on a fixed mass of gas

decreases its volume.

15

Charles' law

Describes the relationship between volume and temperature. If the pressure is held constant, increasing the temperature of a fixed

mass of gas increases its volume.

Gay-Lussac's law

Describes the relationship between pressure and temperature. If the volume is held constant, increasing the temperature of a fixed mass

of gas proportionately raises its pressure

= an increase in temperature increases the kinetic energy of the particles, which means that they will move faster and collide with the walls with more energy and more frequency

The combined gas law

16

The ideal gas equation

The ideal gas law relates pressure, volume, temperature and amount of substance:

R is the gas constant = 8,31 J K-1 mol-1

17