Chemistry Paper 1

? Atoms, elements and compounds

  • All substances are made of atoms.
  • An atom is the smallest part of an element that can exist.
  • Atoms of each element are represented by a chemical symbol, eg O represents an atom of oxygen, Na represents an atom of sodium.
  • There are about 100 different elements.
  • Elements are shown in the periodic table.
  • Compounds are formed from elements by chemical reactions.
  • Chemical reactions always involve the formation of one or more new substances, and often involve a detectable energy change.
  • Compounds contain two or more elements chemically combined in fixed proportions and can be represented by formulae using the symbols of the atoms from which they were formed.
  • Compounds can only be separated into elements by chemical reactions.
  • Chemical reactions can be represented by word equations or equations using symbols and formulae.
1 of 67 Mixtures

  • A mixture consists of two or more elements or compounds not chemically combined together.
  • The chemical properties of each substance in the mixture are unchanged.
  • Mixtures can be separated by physical processes such as filtration, crystallisation, simple distillation, fractional distillation and chromatography.
  • These physical processes do not involve chemical reactions and no new substances are made.
2 of 67 The development of the model of the atom

  • New experimental evidence may lead to a scientific model being changed or replaced.
  • Before the discovery of the electron, atoms were thought to be tiny spheres that could not be divided.
  • The discovery of the electron led to the plum pudding model of the atom.The plum pudding model suggested that the atom is a ball of positive charge with negative electrons embedded in it.
  • The results from the alpha particle scattering experiment led to the conclusion that the mass of an atom was concentrated at the centre (nucleus) and that the nucleus was charged. This nuclear model replaced the plum pudding model.
  • Niels Bohr adapted the nuclear model by suggesting that electrons orbit the nucleus at specific distances. The theoretical calculations of Bohr agreed with experimental observations.
  • Later experiments led to the idea that the positive charge of any nucleus could be subdivided into a whole number of smaller particles, each particle having the same amount of positive charge. The name proton was given to these particles.
  • The experimental work of James Chadwick provided the evidence to show the existence of neutrons within the nucleus. This was about 20 years after the nucleus became an accepted scientific idea.
3 of 67 Relative electrical charges of subatomic p

  • The relative electrical charges of the particles in atoms are:
  • In an atom, the number of electrons is equal to the number of protons in the nucleus.
  • Atoms have no overall electrical charge.
  • The number of protons in an atom of an element is its atomic number.
  • All atoms of a particular element have the same number of protons.
  • Atoms of different elements have different numbers of protons.
4 of 67 Size and mass of atoms

  • Atoms are very small, having a radius of about 0.1 nm (1 x 10-10 m).
  • The radius of a nucleus is less than 1/10 000 of that of the atom (about 1 x 10-14 m). Almost all of the mass of an atom is in the nucleus.
  • The relative masses of protons, neutrons and electrons are:
  • The sum of the protons and neutrons in an atom is its mass number.
  • Atoms of the same element can have different numbers of neutrons; these atoms are called isotopes of that element.
  • Atoms can be represented as shown in this example:
5 of 67 Relative atomic mass

  • The relative atomic mass of an element is an average value that takes account of the abundance of the isotopes of the element.  
6 of 67 Electronic structure

  • The electrons in an atom occupy the lowest available energy levels (innermost available shells). The electronic structure of an atom can be represented by numbers or by a diagram. For example, the electronic structure of sodium is 2,8,1 or

showing two electrons in the lowest energy level, eight in the second energy level and one in the third energy level.

7 of 67 The periodic table

  • The elements in the periodic table are arranged in order of atomic (proton) number and so that elements with similar properties are in columns, known as groups.
  • The table is called a periodic table because similar properties occur at regular intervals.
  • Elements in the same group in the periodic table have the same number of electrons in their outer shell (outer electrons) and this gives them similar chemical properties.
8 of 67 Development of the periodic table

  • Before the discovery of protons, neutrons and electrons, scientists attempted to classify the elements by arranging them in order of their atomic weights.
  • The early periodic tables were incomplete and some elements were placed in inappropriate groups if the strict order of atomic weights was followed.
  • Mendeleev overcame some of the problems by leaving gaps for elements that he thought had not been discovered and in some places changed the order based on atomic weights.
  • Elements with properties predicted by Mendeleev were discovered and filled the gaps. Knowledge of isotopes made it possible to explain why the order based on atomic weights was not always correct.
9 of 67 Metals and non-metals

  • Elements that react to form positive ions are metals.
  • Elements that do not form positive ions are non-metals.
  • The majority of elements are metals.
  • Metals are found to the left and towards the bottom of the periodic table.
  • Non-metals are found towards the right and top of the periodic table.
10 of 67 Group 0

  • The elements in Group 0 of the periodic table are called the noble gases.
  • They are unreactive and do not easily form molecules because their atoms have stable arrangements of electrons.
  • The noble gases have eight electrons in their outer shell, except for helium, which has only two electrons.
  • The boiling points of the noble gases increase with increasing relative atomic mass (going down the group).
11 of 67 Group 1

  • The elements in Group 1 of the periodic table are known as the alkali metals and have characteristic properties because of the single electron in their outer shell.
  • In Group 1, the reactivity of the elements increases going down the group.
12 of 67 Group 7

  • The elements in Group 7 of the periodic table are known as the halogens and have similar reactions because they all have seven electrons in their outer shell.
  • The halogens are non-metals and consist of molecules made of pairs of atoms. 
  • In Group 7, the further down the group an element is the higher its relative molecular mass, melting point and boiling point.
  • In Group 7, the reactivity of the elements decreases going down the group.
  • A more reactive halogen can displace a less reactive halogen from an aqueous solution of its salt.
13 of 67 Comparison with Group 1 elements

  • The transition elements are metals with similar properties which are different from those of the elements in Group 1.
14 of 67 Typical properties

  • Many transition elements have ions with different charges, form coloured compounds and are useful as catalysts.
15 of 67 Chemical bonds

  • There are three types of strong chemical bonds: ionic, covalent and metallic.
  • For ionic bonding the particles are oppositely charged ions.
  • For covalent bonding the particles are atoms which share pairs of electrons.
  • For metallic bonding the particles are atoms which share delocalised electrons.
  • Ionic bonding occurs in compounds formed from metals combined with non-metals.
  • Covalent bonding occurs in most non-metallic elements and in compounds of non-metals. Metallic bonding occurs in metallic elements and alloys.
16 of 67 Ionic bonding

  • When a metal atom reacts with a non-metal atom electrons in the outer shell of the metal atom are transferred. Metal atoms lose electrons to become positively charged ions.
  • Non-metal atoms gain electrons to become negatively charged ions.
  • The ions produced by metals in Groups 1 and 2 and by non-metals in Groups 6 and 7 have the electronic structure of a noble gas (Group 0).
  • The electron transfer during the formation of an ionic compound can be represented by a dot and cross diagram, eg for sodium chloride.
  • The charge on the ions produced by metals in Groups 1 and 2 and by non-metals in Groups 6 and 7 relates to the group number of the element in the periodic table.
17 of 67 Ionic compounds

  • An ionic compound is a giant structure of ions.
  • Ionic compounds are held together by strong electrostatic forces of attraction between oppositely charged ions.
  • These forces act in all directions in the lattice and this is called ionic bonding.
  • The structure of sodium chloride can be represented in the following forms:
18 of 67 Covalent bonding

  • When atoms share pairs of electrons, they form covalent bonds.
  • These bonds between atoms are strong.
  • Covalently bonded substances may consist of small molecules.
  • Some covalently bonded substances have very large molecules, such as polymers.
  • Some covalently bonded substances have giant covalent structures, such as diamond and silicon dioxide.
  • The covalent bonds in molecules and giant structures can be represented in the following forms:
19 of 67 Metallic bonding

  • Metals consist of giant structures of atoms arranged in a regular pattern.
  • The electrons in the outer shell of metal atoms are delocalised and so are free to move through the whole structure.
  • The sharing of delocalised electrons gives rise to strong metallic bonds.
  • The bonding in metals may be represented in the following form: 
20 of 67 The three states of matter

  • The three states of matter are solid, liquid and gas. Melting and freezing take place at the melting point, boiling and condensing take place at the boiling point.
  • The three states of matter can be represented by a simple model. In this model, particles are represented by small solid spheres. Particle theory can help to explain melting, boiling, freezing and condensing.
  • The amount of energy needed to change state from solid to liquid and from liquid to gas depends on the strength of the forces between the particles of the substance. The nature of the particles involved depends on the type of bonding and the structure of the substance. The stronger the forces between the particles the higher the melting point and boiling point of the substance.
  • Limitations of the simple model above include that in the model there are no forces, that all particles are represented as spheres and that the spheres are solid.
21 of 67 State symbols

  • In chemical equations, the three states of matter are shown as (s), (l) and (g), with (aq) for aqueous solutions.
22 of 67 Properties of ionic compounds

  • Ionic compounds have regular structures (giant ionic lattices) in which there are strong electrostatic forces of attraction in all directions between oppositely charged ions.
  • These compounds have high melting points and high boiling points because of the large amounts of energy needed to break the many strong bonds.
  • When melted or dissolved in water, ionic compounds conduct electricity because the ions are free to move and so charge can flow
23 of 67 Properties of small molecules

  • Substances that consist of small molecules are usually gases or liquids that have relatively low melting points and boiling points.
  • These substances have only weak forces between the molecules (intermolecular forces). It is these intermolecular forces that are overcome, not the covalent bonds, when the substance melts or boils.
  • The intermolecular forces increase with the size of the molecules, so larger molecules have higher melting and boiling points.
  • These substances do not conduct electricity because the molecules do not have an overall electric charge.
24 of 67 Polymers

  • Polymers have very large molecules.
  • The atoms in the polymer molecules are linked to other atoms by strong covalent bonds.
  • The intermolecular forces between polymer molecules are relatively strong and so these substances are solids at room temperature.
25 of 67 Giant covalent structures

  • Substances that consist of giant covalent structures are solids with very high melting points.
  • All of the atoms in these structures are linked to other atoms by strong covalent bonds.
  • These bonds must be overcome to melt or boil these substances.
  • Diamond and graphite (forms of carbon) and silicon dioxide (silica) are examples of giant covalent structures.
26 of 67 Properties of metals and alloys

  • Metals have giant structures of atoms with strong metallic bonding.
  • This means that most metals have high melting and boiling points.
  • In pure metals, atoms are arranged in layers, which allows metals to be bent and shaped.
  • Pure metals are too soft for many uses and so are mixed with other metals to make alloys which are harder. 
27 of 67 Metals as conductors

  • Metals are good conductors of electricity because the delocalised electrons in the metal carry electrical charge through the metal.
  • Metals are good conductors of thermal energy because energy is transferred by the delocalised electrons.
28 of 67 Diamond

  • In diamond, each carbon atom forms four covalent bonds with other carbon atoms in a giant covalent structure, so diamond is very hard, has a very high melting point and does not conduct electricity.
29 of 67 Graphite

  • In graphite, each carbon atom forms three covalent bonds with three other carbon atoms, forming layers of hexagonal rings which have no covalent bonds between the layers.
  • In graphite, one electron from each carbon atom is delocalised.
30 of 67 Graphene and fullerenes

  • Graphene is a single layer of graphite and has properties that make it useful in electronics and composites.
  • Fullerenes are molecules of carbon atoms with hollow shapes.
  • The structure of fullerenes is based on hexagonal rings of carbon atoms but they may also contain rings with five or seven carbon atoms.
  • The first fullerene to be discovered was Buckminsterfullerene (C60) which has a spherical shape.
  • Carbon nanotubes are cylindrical fullerenes with very high length to diameter ratios.
  • Their properties make them useful for nanotechnology, electronics and materials.
31 of 67 Sizes of particles and their properties

  • Nanoscience refers to structures that are 1–100 nm in size, of the order of a few hundred atoms.
  • Nanoparticles, are smaller than fine particles (PM2.5), which have diameters between 100 and 2500 nm (1 x 10-7 m and 2.5 x 10-6 m).
  • Coarse particles (PM10) have diameters between 1 x 10-5 m and 2.5 x 10-6 m. Coarse particles are often referred to as dust.
  • As the side of cube decreases by a factor of 10 the surface area to volume ratio increases by a factor of 10.
  • Nanoparticles may have properties different from those for the same materials in bulk because of their high surface area to volume ratio.
  • It may also mean that smaller quantities are needed to be effective than for materials with normal particle sizes.
32 of 67 Uses of nanoparticles

  • Nanoparticles have many applications in medicine, in electronics, in cosmetics and sun creams, as deodorants, and as catalysts.
  • New applications for nanoparticulate materials are an important area of research.
33 of 67 Conservation of mass and balanced chemical

  • The law of conservation of mass states that no atoms are lost or made during a chemical reaction so the mass of the products equals the mass of the reactants.
  • This means that chemical reactions can be represented by symbol equations which are balanced in terms of the numbers of atoms of each element involved on both sides of the equation.
34 of 67 Relative formula mass

  • The relative formula mass (Mr) of a compound is the sum of the relative atomic masses of the atoms in the numbers shown in the formula.
  • In a balanced chemical equation, the sum of the relative formula masses of the reactants in the quantities shown equals the sum of the relative formula masses of the products in the quantities shown.
35 of 67 Mass changes when a reactant or product is

  • Some reactions may appear to involve a change in mass but this can usually be explained because a reactant or product is a gas and its mass has not been taken into account.
  • For example: when a metal reacts with oxygen the mass of the oxide produced is greater than the mass of the metal or in thermal decompositions of metal carbonates carbon dioxide is produced and escapes into the atmosphere leaving the metal oxide as the only solid product.
36 of 67 Chemical measurements

  • Whenever a measurement is made there is always some uncertainty about the result obtained.
37 of 67 Moles (HT only)

  • Chemical amounts are measured in moles.
  • The symbol for the unit mole is mol.
  • The mass of one mole of a substance in grams is numerically equal to its relative formula mass.
  • One mole of a substance contains the same number of the stated particles, atoms, molecules or ions as one mole of any other substance.
  • The number of atoms, molecules or ions in a mole of a given substance is the Avogadro constant.
  • The value of the Avogadro constant is 6.02 x 1023 per mole.
38 of 67 Amounts of substances in equations

  • The masses of reactants and products can be calculated from balanced symbol equations.
  • Chemical equations can be interpreted in terms of moles.
  • For example: 

Mg + 2HCI          MgCI2 + H2

shows that one mole of magnesium reacts with two moles of hydrochloric acid to produce one mole of magnesium chloride and one mole of hydrogen gas.

39 of 67 Using moles to balance equations

  • The balancing numbers in a symbol equation can be calculated from the masses of reactants and products by converting the masses in grams to amounts in moles and converting the numbers of moles to simple whole number ratios.
40 of 67 Limiting reactants

  • In a chemical reaction involving two reactants, it is common to use an excess of one of the reactants to ensure that all of the other reactant is used.
  • The reactant that is completely used up is called the limiting reactant because it limits the amount of products.
41 of 67 Concentration of solutions

  • Many chemical reactions take place in solutions.
  • The concentration of a solution can be measured in mass per given volume of solution, eg grams per dm3 (g/dm3 ).
42 of 67 Percentage yield

  • Even though no atoms are gained or lost in a chemical reaction, it is not always possible to obtain the calculated amount of a product because:
  • the reaction may not go to completion because it is reversible
  • some of the product may be lost when it is separated from the reaction mixture
  • some of the reactants may react in ways different to the expected reaction.
  • The amount of a product obtained is known as the yield.
  • When compared with the maximum theoretical amount as a percentage, it is called the percentage yield. 
43 of 67 Atom economy

  • The atom economy (atom utilisation) is a measure of the amount of starting materials that end up as useful products.
  • It is important for sustainable development and for economic reasons to use reactions with high atom economy.
  • The percentage atom economy of a reaction is calculated using the balanced equation for the reaction as follows:
44 of 67

4.3.4 Using concentrations of solutions in mol/dm3

  • The concentration of a solution can be measured in mol/dm3 .
  • The amount in moles of solute or the mass in grams of solute in a given volume of solution can be calculated from its concentration in mol/dm3 .
  • If the volumes of two solutions that react completely are known and the concentration of one solution is known, the concentration of the other solution can be calculated.
45 of 67

4.3.5 Use of amount of substance in relation to vo

  • Equal amounts in moles of gases occupy the same volume under the same conditions of temperature and pressure.
  • The volume of one mole of any gas at room temperature and pressure (20oC and 1 atmosphere pressure) is 24 dm3 .
  • The volumes of gaseous reactants and products can be calculated from the balanced equation for the reaction.
46 of 67

4.3.5 Use of amount of substance in relation to vo

  • Equal amounts in moles of gases occupy the same volume under the same conditions of temperature and pressure.
  • The volume of one mole of any gas at room temperature and pressure (20oC and 1 atmosphere pressure) is 24 dm3 .
  • The volumes of gaseous reactants and products can be calculated from the balanced equation for the reaction.
47 of 67 Metal oxides

  • Metals react with oxygen to produce metal oxides.
  • The reactions are oxidation reactions because the metals gain oxygen.
48 of 67 The reactivity series

  • When metals react with other substances the metal atoms form positive ions.
  • The reactivity of a metal is related to its tendency to form positive ions.
  • Metals can be arranged in order of their reactivity in a reactivity series.
  • The metals potassium, sodium, lithium, calcium, magnesium, zinc, iron and copper can be put in order of their reactivity from their reactions with water and dilute acids.
  • The non-metals hydrogen and carbon are often included in the reactivity series.
  • A more reactive metal can displace a less reactive metal from a compound.
49 of 67 Extraction of metals and reduction

  • Unreactive metals such as gold are found in the Earth as the metal itself but most metals are found as compounds that require chemical reactions to extract the metal.
  • Metals less reactive than carbon can be extracted from their oxides by reduction with carbon.
  • Reduction involves the loss of oxygen.
50 of 67 Oxidation and reduction in terms of electr

  • Oxidation is the loss of electrons and reduction is the gain of electrons.
51 of 67 Reactions of acids with metals

  • Acids react with some metals to produce salts and hydrogen.
52 of 67 Neutralisation of acids and salt productio

  • Acids are neutralised by alkalis (eg soluble metal hydroxides) and bases (eg insoluble metal hydroxides and metal oxides) to produce salts and water, and by metal carbonates to produce salts, water and carbon dioxide.
  • The particular salt produced in any reaction between an acid and a base or alkali depends on:
  • the acid used (hydrochloric acid produces chlorides, nitric acid produces nitrates, sulfuric acid produces sulfates) 
  • the positive ions in the base, alkali or carbonate.
53 of 67 Soluble salts

  • Soluble salts can be made from acids by reacting them with solid insoluble substances, such as metals, metal oxides, hydroxides or carbonates.
  • The solid is added to the acid until no more reacts and the excess solid is filtered off to produce a solution of the salt.
  • Salt solutions can be crystallised to produce solid salts.
54 of 67 The pH scale and neutralisation

  • Acids produce hydrogen ions (H+) in aqueous solutions.
  • Aqueous solutions of alkalis contain hydroxide ions (OH– ).
  • The pH scale, from 0 to 14, is a measure of the acidity or alkalinity of a solution, and can be measured using universal indicator or a pH probe.
  • A solution with pH 7 is neutral.
  • Aqueous solutions of acids have pH values of less than 7 and aqueous solutions of alkalis have pH values greater than 7.
  • In neutralisation reactions between an acid and an alkali, hydrogen ions react with hydroxide ions to produce water.
  • This reaction can be represented by the equation:
55 of 67 Titrations

  • The volumes of acid and alkali solutions that react with each other can be measured by titration using a suitable indicator.
56 of 67 Strong and weak acids

  • A strong acid is completely ionised in aqueous solution.
  • Examples of strong acids are hydrochloric, nitric and sulfuric acids.
  • A weak acid is only partially ionised in aqueous solution.
  • Examples of weak acids are ethanoic, citric and carbonic acids.
  • For a given concentration of aqueous solutions, the stronger an acid, the lower the pH.
  • As the pH decreases by one unit, the hydrogen ion concentration of the solution increases by a factor of 10.
57 of 67 The process of electrolysis

  • When an ionic compound is melted or dissolved in water, the ions are free to move about within the liquid or solution.
  • These liquids and solutions are able to conduct electricity and are called electrolytes. Passing an electric current through electrolytes causes the ions to move to the electrodes. Positively charged ions move to the negative electrode (the cathode), and negatively charged ions move to the positive electrode (the anode).
  • Ions are discharged at the electrodes producing elements.
  • This process is called electrolysis.
58 of 67 Electrolysis of molten ionic compounds

  • When a simple ionic compound (eg lead bromide) is electrolysed in the molten state using inert electrodes, the metal (lead) is produced at the cathode and the non-metal (bromine) is produced at the anode.
59 of 67 Using electrolysis to extract metals

  • Metals can be extracted from molten compounds using electrolysis.
  • Electrolysis is used if the metal is too reactive to be extracted by reduction with carbon or if the metal reacts with carbon.
  • Large amounts of energy are used in the extraction process to melt the compounds and to produce the electrical current.
  • Aluminium is manufactured by the electrolysis of a molten mixture of aluminium oxide and cryolite using carbon as the positive electrode (anode).
60 of 67 Electrolysis of aqueous solutions

  • The ions discharged when an aqueous solution is electrolysed using inert electrodes depend on the relative reactivity of the elements involved.
  • At the negative electrode (cathode), hydrogen is produced if the metal is more reactive than hydrogen.
  • At the positive electrode (anode), oxygen is produced unless the solution contains halide ions when the halogen is produced.
  • This happens because in the aqueous solution water molecules break down producing hydrogen ions and hydroxide ions that are discharged.
61 of 67 Representation of reactions at electrodes

  • During electrolysis, at the cathode (negative electrode), positively charged ions gain electrons and so the reactions are reductions.
  • At the anode (positive electrode), negatively charged ions lose electrons and so the reactions are oxidations.
  • Reactions at electrodes can be represented by half equations, for example:
62 of 67 Energy transfer during exothermic and endo

  • Energy is conserved in chemical reactions.
  • The amount of energy in the universe at the end of a chemical reaction is the same as before the reaction takes place.
  • If a reaction transfers energy to the surroundings the product molecules must have less energy than the reactants, by the amount transferred.
  • An exothermic reaction is one that transfers energy to the surroundings so the temperature of the surroundings increases.
  • Exothermic reactions include combustion, many oxidation reactions and neutralisation.
  • Everyday uses of exothermic reactions include self-heating cans and hand warmers. An endothermic reaction is one that takes in energy from the surroundings so the temperature of the surroundings decreases.
  • Endothermic reactions include thermal decompositions and the reaction of citric acid and sodium hydrogencarbonate.
  • Some sports injury packs are based on endothermic reactions.
63 of 67 Reaction profiles

  • Chemical reactions can occur only when reacting particles collide with each other and with sufficient energy.
  • The minimum amount of energy that particles must have to react is called the activation energy.
  • Reaction profiles can be used to show the relative energies of reactants and products, the activation energy and the overall energy change of a reaction.
64 of 67 The energy change of reactions

  • During a chemical reaction:
  • energy must be supplied to break bonds in the reactants
  • energy is released when bonds in the products are formed.
  • The energy needed to break bonds and the energy released when bonds are formed can be calculated from bond energies.
  • The difference between the sum of the energy needed to break bonds in the reactants and the sum of the energy released when bonds in the products are formed is the overall energy change of the reaction.
  • In an exothermic reaction, the energy released from forming new bonds is greater than the energy needed to break existing bonds.
  • In an endothermic reaction, the energy needed to break existing bonds is greater than the energy released from forming new bonds.
65 of 67 Cells and batteries

  • Cells contain chemicals which react to produce electricity.
  • The voltage produced by a cell is dependent upon a number of factors including the type of electrode and electrolyte.
  • A simple cell can be made by connecting two different metals in contact with an electrolyte.
  • Batteries consist of two or more cells connected together in series to provide a greater voltage.
  • In non-rechargeable cells and batteries the chemical reactions stop when one of the reactants has been used up.
  • Alkaline batteries are non-rechargeable.
  • Rechargeable cells and batteries can be recharged because the chemical reactions are reversed when an external electrical current is supplied.
66 of 67 Fuel cells

  • Fuel cells are supplied by an external source of fuel (eg hydrogen) and oxygen or air.
  • The fuel is oxidised electrochemically within the fuel cell to produce a potential difference.
  • The overall reaction in a hydrogen fuel cell involves the oxidation of hydrogen to produce water.
  • Hydrogen fuel cells offer a potential alternative to rechargeable cells and batteries.
67 of 67




HI, thank you for this these look amazing. I was just wondering what grade you got?



evies3 she probably got a grade 9 or A* 

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