Exam Notes

Alkanes and Alkenes

• 9.10C Alkanes Formulae and Structures: Recall the formulae of molecules of the alkanes, methane, ethane, propane, and butane, and be able to draw their structures, showing all covalent bonds.

  • Methane: $CH_4$

  • Ethane: C2H6

  • Propane: C3H8

  • Butane: C4H10

• 9.11C Saturated Hydrocarbons: Explain why alkanes are saturated hydrocarbons.

  • Alkanes are saturated because they contain only single carbon-carbon bonds and the maximum number of hydrogen atoms.

• 9.16C Complete Combustion: Describe how the complete combustion of alkanes and alkenes involves the oxidation of the hydrocarbons to produce carbon dioxide and water.

  • General equation: Hydrocarbon + O2 —> CO2 + H_2O$</p></li></ul></li></ul><h3 id="65a6b1ea-0cf2-4016-ace6-2c341d9f76a2" data-toc-id="65a6b1ea-0cf2-4016-ace6-2c341d9f76a2" collapsed="false" seolevelmigrated="true">Core Practical: Reaction Rates</h3><ul><li><p><strong>7.1 Core practical investigation:</strong> Investigate the effects of changing conditions of a reaction on the rates of chemical reactions by:</p><ul><li><p>Measuring the production of gas (reaction between hydrochloric acid and marble chips).</p><ul><li><p>CaCO<em>3(s) + 2HCl(aq) —> CaCl</em>2(aq) + H<em>2O(l) + CO</em>2(g)$

• Observing a color change (reaction between sodium thiosulfate and hydrochloric acid).

  • Na2S2O3(aq) + 2HCl(aq) —> 2NaCl(aq) + S(s) + SO2(g) + H2O(l)

Gases

• 8.23 Oxygen Test: Describe the chemical test for oxygen.

  • Oxygen relights a glowing splint.

• 8.24 Greenhouse Effect: Describe how various gases in the atmosphere, including carbon dioxide, methane, and water vapor, absorb heat radiated from the Earth, subsequently releasing energy which keeps the Earth warm; this is known as the greenhouse effect.

  • Greenhouse gases trap infrared radiation.

• 8.25 Climate Change Evidence: Evaluate the evidence for human activity causing climate change, considering:

  • The correlation between the change in atmospheric carbon dioxide concentration, the consumption of fossil fuels, and temperature change.

  • The uncertainties caused by the location where these measurements are taken and historical accuracy.

• 8.26 Atmospheric Composition and Climate Change: Describe:

  • The composition of today’s atmosphere.

    • Approximately 78% nitrogen, 21% oxygen, and small amounts of other gases, including carbon dioxide, argon, and water vapor.

  • The potential effects on the climate of increased levels of carbon dioxide and methane generated by human activity, including burning fossil fuels and livestock farming.

  • Mitigation of these effects, considering scale, risk, and environmental implications.

Homologous Series

• 8.6 Homologous Series: Explain an homologous series as a series of compounds which:

  • Have the same general formula.

  • Differ by $CH_2$ in molecular formulae from neighboring compounds.

  • Show a gradual variation in physical properties, as exemplified by their boiling points.

  • Have similar chemical properties.

Combustion

• 8.7 Complete Combustion: Describe the complete combustion of hydrocarbon fuels as a reaction in which:

  • Carbon dioxide and water are produced.

  • Energy is given out (exothermic reaction).

  • $Hydrocarbon + Oxygen \rightarrow Carbon Dioxide + Water$

• 8.8 Incomplete Combustion: Explain why the incomplete combustion of hydrocarbons can produce carbon and carbon monoxide.

  • Occurs when there is a limited supply of oxygen.

  • $Hydrocarbon + Limited Oxygen \rightarrow Carbon Monoxide + Carbon + Water$

• 8.9 Carbon Monoxide as a Toxic Gas: Explain how carbon monoxide behaves as a toxic gas.

  • Carbon monoxide binds to hemoglobin in red blood cells, preventing oxygen transport.

• 8.10 Problems with Incomplete Combustion: Describe the problems caused by incomplete combustion producing carbon monoxide and soot in appliances that use carbon compounds as fuels.

• 8.11 Sulfur Dioxide Production: Explain how impurities in some hydrocarbon fuels result in the production of sulfur dioxide.

• 8.12 Acid Rain: Explain some problems associated with acid rain caused when sulfur dioxide dissolves in rainwater.

  • Acid rain can damage buildings, harm aquatic life, and affect plant growth.

Reaction Rates

• 7.2 Determining Reaction Rate: Suggest practical methods for determining the rate of a given reaction.

  • Measuring the volume of gas produced over time.

  • Measuring the change in mass over time.

  • Measuring the color change over time using colorimetry.

• 7.3 Collision Theory: Explain how reactions occur when particles collide and that rates of reaction are increased when the frequency and/or energy of collisions is increased.

• 7.4 Factors Affecting Reaction Rates: Explain the effects on rates of reaction of changes in temperature, concentration, surface area to volume ratio of a solid, and pressure (on reactions involving gases) in terms of frequency and/or energy of collisions between particles.

  • Temperature: Higher temperature increases the kinetic energy of particles, leading to more frequent and energetic collisions.

  • Concentration: Higher concentration means more particles in a given volume, increasing collision frequency.

  • Surface Area: Smaller particle size (larger surface area to volume ratio) increases the area available for collisions.

  • Pressure: Higher pressure increases the concentration of gases, leading to more frequent collisions.

• 7.5 Interpreting Graphs: Interpret graphs of mass, volume, or concentration of reactant or product against time.

Noble Gases

• 6.14 Chemical Inertness: Explain why the noble gases are chemically inert, compared with the other elements, in terms of their electronic configurations.

  • Noble gases have a full outer electron shell, making them stable and unreactive.

• 6.15 Uses of Noble Gases: Explain how the uses of noble gases depend on their inertness, low density, and/or non-flammability.

  • Helium is used in balloons due to its low density and inertness.

  • Argon is used as a shielding gas in welding due to its inertness.

  • Neon is used in advertising signs because it emits light when an electric current is passed through it.

Halogens

• 6.6 Color and State of Halogens: Recall the colors and physical states of chlorine, bromine, and iodine at room temperature.

  • Chlorine: Green gas

  • Bromine: Red-brown liquid

  • Iodine: Grey-black solid

• 6.7 Trends in Physical Properties: Describe the pattern in the physical properties of the halogens, chlorine, bromine, and iodine, and use this pattern to predict the physical properties of other halogens.

  • Melting and boiling points increase down the group.

  • Color becomes darker down the group.

• 6.8 Test for Chlorine: Describe the chemical test for chlorine.

  • Chlorine bleaches damp litmus paper, turning it white.

• 6.9 Reactions with Metals: Describe the reactions of the halogens, chlorine, bromine, and iodine, with metals to form metal halides, and use this pattern to predict the reactions of other halogens.

  • $Metal + Halogen \rightarrow Metal Halide$

• 6.10 Formation of Hydrogen Halides: Recall that the halogens, chlorine, bromine, and iodine, form hydrogen halides which dissolve in water to form acidic solutions, and use this pattern to predict the reactions of other halogens.

  • $Hydrogen + Halogen \rightarrow Hydrogen Halide$

• 6.11 Relative Reactivity: Describe the relative reactivity of the halogens chlorine, bromine, and iodine, as shown by their displacement reactions with halide ions in aqueous solution, and use this pattern to predict the reactions of astatine.

  • Reactivity decreases down the group: Chlorine > Bromine > Iodine

• 6.12 Redox Reactions: Explain why these displacement reactions are redox reactions in terms of gain and loss of electrons, identifying which of the substances are oxidized and which are reduced.

  • Oxidation is the loss of electrons.

  • Reduction is the gain of electrons.

Alkali Metals

• 6.1 Classification: Explain why some elements can be classified as alkali metals (group 1), halogens (group 7), or noble gases (group 0), based on their position in the periodic table.

• 6.2 Properties of Alkali Metals: Recall that alkali metals:

  • Are soft.

  • Have relatively low melting points.

• 6.3 Reactions with Water: Describe the reactions of lithium, sodium, and potassium with water.

  • $2Li(s) + 2H<em>2O(l) \rightarrow 2LiOH(aq) + H</em>2(g)$

  • $2Na(s) + 2H<em>2O(l) \rightarrow 2NaOH(aq) + H</em>2(g)$

  • $2K(s) + 2H<em>2O(l) \rightarrow 2KOH(aq) + H</em>2(g)$

• 6.4 Pattern in Reactivity: Describe the pattern in reactivity of the alkali metals, lithium, sodium, and potassium, with water; and use this pattern to predict the reactivity of other alkali metals.

  • Reactivity increases down the group.

• 6.5 Electronic Configuration and Reactivity: Explain this pattern in reactivity in terms of electronic configurations.

  • Alkali metals lose one electron to form a +1 ion. Down the group, the outermost electron is further from the nucleus, making it easier to lose.

Percentage Yield and Atom Economy

• 5.11C Percentage Yield: Calculate the percentage yield of a reaction from the actual yield and the theoretical yield.

  • $Percentage Yield = (Actual Yield / Theoretical Yield) * 100$

• 5.12C Actual Yield vs. Theoretical Yield: Describe that the actual yield of a reaction is usually less than the theoretical yield and that the causes of this include:

  • Incomplete reactions.

  • Practical losses during the experiment.

  • Competing, unwanted reactions (side reactions).

• 5.13C Atom Economy: Recall the atom economy of a reaction forming a desired product.

• 5.14C Atom Economy Calculation: Calculate the atom economy of a reaction forming a desired product.

  • $Atom Economy = (Mass of Desired Product / Total Mass of Reactants) * 100$

• 5.15C Choosing Reaction Pathways: Explain why a particular reaction pathway is chosen to produce a specified product, given appropriate data such as atom economy, yield, rate, equilibrium position, and usefulness of by-products.

• 5.16C Molar Volume: Describe the molar volume of any gas at room temperature and pressure, as the volume occupied by one mole of molecules of any gas at room temperature and pressure.

  • Molar volume is 24 dm³ or 24000 cm³.

• 5.17C Molar Volume Calculations: Use the molar volume and balanced equations in calculations involving the masses of solids and volumes of gases.

• 5.18C Avogadro's Law: Use Avogadro's law to calculate volumes of gases involved in a gaseous reaction, given the relevant equation.

Purity and Separation Techniques

• 2.5 Purity in Chemistry: Explain the difference between the use of 'pure' in chemistry compared with its everyday use and the differences in chemistry between a pure substance and a mixture.

• 2.6 Melting Point Data: Interpret melting point data to distinguish between pure substances which have a sharp melting point and mixtures which melt over a range of temperatures.

• 2.7 Separation Techniques: Explain the types of mixtures that can be separated by using the following experimental techniques:

  • Simple distillation

  • Fractional distillation

  • Filtration

  • Crystallization

  • Paper chromatography

• 2.8 Choosing Separation Techniques: Describe an appropriate experimental technique to separate a mixture, knowing the properties of the components of the mixture.

• 2.9 Paper Chromatography: Describe paper chromatography as the separation of mixtures of soluble substances by running a solvent (mobile phase) through the mixture on the paper (the paper contains the stationary phase), which causes the substances to move at different rates over the paper.

• 2.10 Interpreting Chromatograms: Interpret a paper chromatogram:

  • To distinguish between pure and impure substances.

  • To identify substances by comparison with known substances.

  • To identify substances by calculation and use of Rf values.

• 2.11 Core Practical: Composition of Inks: Core Practical: Investigate the composition of inks using simple distillation and paper chromatography.

• 2.12 Potable Water: Describe how:

  • Waste and groundwater can be made potable, including the need for sedimentation, filtration, and chlorination.

  • Seawater can be made potable by using distillation.

  • Water used in analysis must not contain any dissolved salts.

Empirical Formula and Conservation of Mass

• 1.43 Calculating Relative Formula Mass: Calculate:

  • Relative formula mass given relative atomic masses.

  • Percentage by mass of an element in a compound given relative atomic masses.

  • $Percentage \space by \space mass = (Ar * Number \space of \space atoms \space of \space the \space element) / Mr \space of \space the \space compound *100$

• 1.44 Empirical Formula Calculation: Calculate the formulae of simple compounds from reacting masses or percentage composition and understand that these are empirical formulae.

• 1.45 Deducing Empirical and Molecular Formula: Deduce:

  • The empirical formula of a compound from the formula of its molecule.

  • The molecular formula of a compound from its empirical formula and its relative molecular mass.

• 1.46 Experiment for Empirical Formula: Describe an experiment to determine the empirical formula of a simple compound such as magnesium oxide.

• 1.47 Law of Conservation of Mass: Explain the law of conservation of mass applied to:

  • A closed system including a precipitation reaction in a closed flask.

  • A non-enclosed system including a reaction in an open flask that takes in or gives out a gas.

• 1.48 Mass Calculations: Calculate masses of reactants and products from balanced equations, given the mass of one substance.

Types of Bonding and Structure

• 1.32 Classifying Substances: Explain why elements and compounds can be classified as:

  • Ionic

  • Simple Molecular (Covalent)

  • Giant Covalent

  • Metallic
    and how the structure and bonding of these types of substances results in different physical properties, including relative melting point and boiling point, relative solubility in water, and ability to conduct electricity (as solids and in solution).

• 1.33 Properties of Ionic Compounds: Explain the properties of ionic compounds limited to:

  • High melting points and boiling points, in terms of forces between ions.

  • Whether or not they conduct electricity as solids, when molten, and in aqueous solution.

• 1.34 Properties of Simple Molecular Compounds: Explain the properties of typical covalent, simple molecular compounds limited to:

  • Low melting points and boiling points, in terms of forces between molecules (intermolecular forces).

  • Poor conduction of electricity.

• 1.35 Graphite and Diamond: Recall that graphite and diamond are different forms of carbon and that they are examples of giant covalent substances.

• 1.36 Structures of Graphite and Diamond: Describe the structures of graphite and diamond.

• 1.37 Properties and Uses of Graphite and Diamond: Explain, in terms of structure and bonding, why graphite is used to make electrodes and as a lubricant, whereas diamond is used in cutting tools.

• 1.38 Properties of Fullerenes and Graphene: Explain the properties of fullerenes including $C_{60}$ and graphene in terms of their structures and bonding.

Covalent Bonding

• 1.28 Formation of Covalent Bonds: Explain how a covalent bond is formed when a pair of electrons is shared between two atoms.

• 1.29 Formation of Molecules: Recall that covalent bonding results in the formation of molecules.

• 1.30 Size of Atoms and Molecules: Recall the typical size (order of magnitude) of atoms and small molecules.

• 1.31 Dot and Cross Diagrams: Explain the formation of simple molecular, covalent substances, using dot and cross diagrams, including:

  • Hydrogen ($H_2$)

  • Hydrogen chloride ($HCl$)

  • Water ($H_2O$)

  • Methane ($CH_4$)

  • Oxygen ($O_2$)

  • Carbon dioxide ($CO_2$)

Ions and Ionic Compounds

• 1.22 Definition of an Ion: Recall that an ion is an atom or group of atoms with a positive or negative charge.

• 1.23 Calculating Protons, Neutrons, and Electrons in Ions: Calculate the numbers of protons, neutrons, and electrons in simple ions given the atomic number and mass number.

• 1.24 Formation of Ions: Explain the formation of ions in ionic compounds from their atoms, limited to compounds of elements in groups 1, 2, 6, and 7.

• 1.25 Naming Compounds: Explain the use of the endings -ide and -ate in the names of compounds.

• 1.26 Formulae of Ionic Compounds: Deduce the formulae of ionic compounds (including oxides, hydroxides, halides, nitrates, carbonates, and sulfates) given the formulae of the constituent ions.

• 1.27 Structure of Ionic Compounds: Explain the structure of an ionic compound as a lattice structure

  • consisting of a regular arrangement of ions

  • held together by strong electrostatic forces (ionic bonds) between oppositely-charged ions.

Periodic Table

• 1.16 Atomic Number: Explain the meaning of the atomic number of an element in terms of position in the periodic table and number of protons in the nucleus.

• 1.17 Arrangement of Periodic Table: Describe that in the periodic table

  • elements are arranged in order of increasing atomic number, in rows called periods

  • elements with similar properties are placed in the same vertical columns called groups

• 1.18 Metals and Non-metals: Identify elements as metals or non-metals according to their position in the periodic table, explaining this division in terms of the atomic structures of the elements.

• 1.19 Electronic Configurations: Predict the electronic configurations of the first 20 elements in the periodic table as diagrams and in the form, for example 2.8.1.

• 1.20 Position and Electronic Configuration: Explain how the electronic configuration of an element is related to its position in the periodic table.

Atomic Structure

• 1.2 Structure of an Atom: Describe the structure of an atom as a nucleus containing protons and neutrons, surrounded by electrons in shells.

• 1.3 Relative Charge and Mass: Recall the relative charge and relative mass of:

  • a proton

    • Charge: +1

    • Mass: 1

  • a neutron

    • Charge: 0

    • Mass: 1

  • an electron

    • Charge: -1

    • Mass: negligible

• 1.4 Equal Numbers of Protons and Electrons: Explain why atoms contain equal numbers of protons and electrons.

• 1.5 Size of Nucleus: Describe the nucleus of an atom as very small compared to the overall size of the atom.

• 1.6 Mass in Nucleus: Recall that most of the mass of an atom is concentrated in the nucleus.

• 1.7 Mass Number: Recall the meaning of the term mass number of an atom, which is the total number of protons and neutrons in the atom's nucleus.

• 1.8 Atomic Number and Identity: Describe atoms of a given element as having the same number of protons in the nucleus and that this number is unique to that element.

• 1.9 Isotopes: Describe isotopes as different atoms of the same element containing the same number of protons but different numbers of neutrons in their nuclei.

• 1.10 Calculating Subatomic Particles: Calculate the numbers of protons, neutrons, and electrons in atoms given the atomic number and mass number.

• 1.11 Relative Atomic Mass: Explain how the existence of isotopes results in relative atomic masses of some elements not being whole numbers.

• 1.12 Calculating Relative Atomic Mass: Calculate the relative atomic mass of an element from the relative masses and abundances of its isotopes.

  • $Ar = \frac{\Sigma (isotope \space mass * isotope \space abundance)}{\Sigma isotope \space abundance}$

Chemical Equations and Safety

• 0.1 Formulae of Elements and Compounds: Recall the formulae of elements, simple compounds, and ions.

• 0.2 Word Equations: Write word equations.

• 0.3 Balanced Chemical Equations: Write balanced chemical equations, including the use of the state symbols (s), (l), (g), and (aq)

• 0.4 Balanced Ionic Equations: Write balanced ionic equations.

• 0.5 Hazard Symbols: Describe the use of hazard symbols on containers

  • to indicate the dangers associated with the contents

  • to inform people about safe-working precautions with these substances in the laboratory.

• 0.6 Risk Evaluation and Precautions: Evaluate the risks in a practical procedure and suggest suitable precautions for a range of practicals including those mentioned in the specification.