CHEMISTRY Notes
Chapter 1: The Nature of Matter
1.1 The states of matter
● Matter is anything that occupies space and has mass. It exists in three states:
solid, liquid, and gas.
● Solids have a fixed volume and shape, liquids have a fixed volume but take
the shape of their container, and gases have neither a fixed volume nor shape.
● The state of matter can be changed by altering temperature and pressure.
● Changes of state include melting (solid to liquid), freezing (liquid to solid),
sublimation (solid to gas), evaporation (liquid to gas), condensation (gas to liquid),
and boiling (liquid to gas at a specific temperature).
● Pure substances have distinct melting and boiling points, which can be used to
test purity. Impurities lower the melting point and raise the boiling point.
● Heating curves illustrate the relationship between temperature and time during
a state change.
1.2 Separating and purifying substances
● Mixtures can be separated based on differences in physical properties.
● Methods for separating mixtures include filtration, centrifugation, decantation,
and distillation.
● Solutions can be separated using evaporation, crystallization, and fractional
distillation techniques.
● Chromatography, such as paper chromatography, is used to separate and
identify components in a mixture.
● The purity of substances can be assessed using melting point, boiling point, and
chromatography.
1.3 Atoms and molecules
● Elements are pure substances that cannot be broken down into simpler
substances by chemical means.
● Compounds are formed when two or more elements chemically combine.
● The kinetic theory explains the behavior of matter in terms of the movement of
particles.
● Diffusion is the movement of particles from an area of high concentration to
an area of low concentration.
● Brownian motion is the random movement of particles in a fluid due to
collisions with other particles.
1.4 The structure of the atom
● Atoms are composed of protons, neutrons, and electrons.
● Protons have a positive charge, neutrons have no charge, and electrons have a
negative charge.
● The atomic number is the number of protons in an atom's nucleus.
● The mass number is the sum of protons and neutrons in the nucleus.
● Isotopes are atoms of the same element with different numbers of neutrons.
● Relative atomic mass is the average mass of an atom of an element compared
to 1/12th the mass of a carbon-12 atom.
● Radioactivity is the spontaneous emission of radiation from the nucleus of an
unstable atom.
1.5 Electron arrangements in atoms
● Electrons are arranged in energy levels or shells around the nucleus.
● The arrangement of electrons determines an atom's chemical properties.
Summary:
Chapter 1 explores the nature of matter, its states, and the changes it undergoes. It
details methods for separating mixtures and purifying substances, emphasizing the
importance of understanding physical properties. The chapter delves into the atomic
structure, introducing the concept of elements, compounds, and the kinetic theory. It
explains how the arrangement of electrons in atoms influences their chemical
behavior. Finally, it touches on radioactivity and its applications.
Chapter 2: Elements and Compounds
2.1 The Periodic Table - classifying the elements
● The periodic table organizes elements based on their properties, with groups
(vertical columns) and periods (horizontal rows).
● Elements in the same group have similar chemical properties, while those in the
same period have varying properties.
2.2 Trends in Groups
● Within a group, elements share similar properties due to having the same
number of electrons in their outer shell.
● Reactivity increases down a group for metals (lose electrons more easily) and
decreases for non-metals (gain electrons less readily).
2.3 Trends across a period
● Across a period, elements have different properties due to varying numbers of
electrons in their outer shell.
● Atomic size decreases from left to right (due to increased nuclear charge), while
ionization energy (energy to remove an electron) generally increases.
2.4 Chemical bonding in elements and compounds
● Elements can form compounds by bonding through the sharing or transfer of
electrons.
● Ionic bonding involves the transfer of electrons, resulting in positive and
negative ions that attract each other.
● Covalent bonding involves the sharing of electrons between atoms.
2.5 The chemical formulae of elements and
compounds
● Chemical formulae represent the composition of substances using element
symbols and subscripts.
● They show the types and ratios of atoms in a compound (e.g., H₂O for
water).
2.6 Metals, alloys and crystals
● The structure of substances at the atomic level determines their properties.
● Metals form giant structures with a 'sea' of delocalized electrons, making them
good conductors.
● Alloys are mixtures of metals with improved properties compared to pure
metals.
● Crystals have regular, repeating arrangements of atoms, ions, or molecules.
Summary:
Chapter 2 explores the organization of elements in the periodic table and the trends
in their properties. It explains how elements combine through chemical bonding to
form compounds, represented by chemical formulae. The arrangement of atoms, ions,
or molecules in substances, whether in metals, alloys, or crystals, is crucial in
determining their characteristics and behavior.
Chapter 3: Chemical Reactions
3.1 Chemical reactions and equations
● Chemical reactions involve the rearrangement of atoms to form new substances.
● Chemical equations represent these reactions, with reactants on the left and
products on the right, separated by an arrow.
3.2 Equations for chemical reactions
● Equations include state symbols: (s) for solid, (l) for liquid, (g) for gas, and
(aq) for aqueous solution.
● Balanced equations have the same number of each type of atom on both sides.
● Balancing involves adjusting coefficients (numbers in front of formulas) but not
changing the formulas themselves.
3.3 Types of chemical reaction
● Combination reactions: Two or more substances combine to form a single
product.
● Decomposition reactions: A single compound breaks down into two or more
simpler substances.
● Displacement reactions: One element displaces another from a compound.
● Redox reactions: Reactions involving oxidation (loss of electrons) and
reduction (gain of electrons).
3.4 A closer look at reactions, particularly redox
reactions
● Redox reactions can be identified by changes in oxidation states.
● Oxidation involves an increase in the oxidation state, while reduction involves a
decrease.
3.5 Electrolysis
● Electrolysis is the decomposition of a compound using electricity.
● It takes place in an electrolytic cell, which includes electrodes (anode and
cathode) and an electrolyte.
● The electrolyte is the compound being decomposed, often in molten or aqueous
solution form.
● Positive ions move to the cathode, and negative ions move to the anode.
3.6 A closer look at electrode reactions
● Electrode reactions involve the transfer of electrons.
● At the cathode, positive ions gain electrons and are reduced.
● At the anode, negative ions lose electrons and are oxidized.
● The reactivity of elements determines the products formed during electrolysis.
Summary:
● Chemical reactions involve the rearrangement of atoms.
● Chemical equations represent these reactions.
● Reactions can be classified into combination, decomposition, displacement, and
redox.
● Redox reactions involve changes in oxidation states.
● Electrolysis is the decomposition of a compound using electricity.
● Electrode reactions involve the transfer of electrons.
Chapter 4: Acids, Bases, and Salts
4.1 What is an acid?
● Acids are substances that produce hydrogen ions (H+) when dissolved in water.
● Examples of acids include hydrochloric acid (HCl), sulfuric acid (H2SO4),
nitric acid (HNO3), and ethanoic acid (CH3COOH).
● Acids can be classified as strong or weak based on their ability to ionize in
water. Strong acids ionize completely, while weak acids ionize partially.
4.2 Acid and alkali solutions
● The pH scale is used to measure the acidity or alkalinity of a solution. It
ranges from 0 to 14, with 7 being neutral. Solutions with a pH less than 7 are acidic,
while solutions with a pH greater than 7 are alkaline.
● A universal indicator is a mixture of dyes that changes color depending on the
pH of the solution. It can be used to estimate the pH of a solution.
4.3 Metal oxides and non-metal oxides
● Metal oxides are bases, meaning they react with acids to form salts and water.
● Non-metal oxides are acids, meaning they react with bases to form salts and
water.
4.4 Acid reactions in everyday life
● Acids have various applications in everyday life, such as:
○ In the manufacture of fertilizers
○ In car batteries
○ In the production of food and drinks
○ In the digestion process
4.5 Alkalis and bases
● Alkalis are bases that dissolve in water.
● Examples of alkalis include sodium hydroxide (NaOH), potassium hydroxide
(KOH), and calcium hydroxide (Ca(OH)2).
● Bases react with acids to form salts and water, a process called neutralization.
4.6 Characteristic reactions of acids
● Acids react with metals to produce hydrogen gas and salt.
● Acids react with metal carbonates to produce carbon dioxide gas, salt, and
water.
● Acids react with metal hydroxides (bases) to produce salt and water
(neutralization).
4.7 Acids and alkalis in chemical analysis
● Acids and alkalis can be used to identify unknown substances based on their
characteristic reactions.
● For example, the presence of carbon dioxide gas (effervescence) when an
unknown substance reacts with an acid indicates that the substance is a metal
carbonate.
4.8 Salts
● Salts are formed when acids react with bases, metal oxides, or metal
carbonates.
● The type of salt formed depends on the acid and base used in the reaction.
● Salts are composed of positively charged metal ions and negatively charged
non-metal ions.
4.9 Preparing soluble salts
● Soluble salts can be prepared by reacting an acid with a metal, metal oxide, or
metal carbonate, and then evaporating the water to leave behind the solid salt.
● The method used to prepare a soluble salt depends on the reactivity of the
metal and the desired salt.
4.10 Preparing insoluble salts
● Insoluble salts can be prepared by precipitation reactions, where two soluble
salts react to form an insoluble salt and a soluble salt.
● The insoluble salt precipitates out of the solution and can be collected by
filtration.
4.11 Strong and weak acids and alkalis
● Strong acids and alkalis ionize completely in water, while weak acids and
alkalis ionize partially.
● The strength of an acid or alkali affects its reactivity and its pH.
Summary
● Acids are substances that produce hydrogen ions in water, while bases are
substances that neutralize acids.
● The pH scale measures the acidity or alkalinity of a solution.
● Acids and bases react to form salts and water.
● Salts can be soluble or insoluble in water.
● The strength of an acid or alkali depends on its degree of ionization in water.
● Acids and alkalis have various applications in everyday life and chemical
analysis.
Chapter 5: Quantitative Chemistry
5.1 Chemical analysis and formulae
● Chemical analysis is used to identify the elements and compounds present in a
sample.
● Empirical formula: simplest whole-number ratio of the atoms of each element
in a compound
● Molecular formula: actual number of atoms of each element in a compound
● To find the empirical formula, you need to know the masses of each element in
a sample of the compound.
● To find the molecular formula, you need to know the empirical formula and
the relative formula mass of the compound.
5.2 The mole and chemical formulae
● The mole is a unit of measurement used to express amounts of a chemical
substance.
● One mole of a substance contains 6 x 10^23 particles (Avogadro’s constant).
● Molar mass: the mass of one mole of a substance, expressed in grams.
● To find the number of moles in a sample of a substance, divide the mass of the
sample by the molar mass of the substance.
● To find the mass of a sample of a substance, multiply the number of moles in
the sample by the molar mass of the substance.
5.3 The mole and chemical equations
● Chemical equations can be used to represent chemical reactions.
● The coefficients in a balanced chemical equation show the relative numbers of
moles of each substance involved in the reaction.
● To find the number of moles of a product formed in a reaction, multiply the
number of moles of the limiting reactant by the ratio of the coefficient of the product
to the coefficient of the limiting reactant.
● To find the mass of a product formed in a reaction, multiply the number of
moles of the product by the molar mass of the product.
5.4 Calculations involving gases
● The volume of one mole of any gas at room temperature and pressure (RTP)
is 24 dm^3 (molar gas volume).
● To find the volume of a gas at RTP, multiply the number of moles of the gas
by the molar gas volume.
● To find the number of moles of a gas at RTP, divide the volume of the gas by
the molar gas volume.
5.5 Moles and solution chemistry
● The concentration of a solution is the amount of solute dissolved in a given
volume of solution.
● Concentration is usually expressed in moles per dm^3 (mol/dm^3) or grams
per dm^3 (g/dm^3).
● To find the concentration of a solution, divide the number of moles of solute by
the volume of the solution.
● To find the number of moles of solute in a solution, multiply the concentration
of the solution by the volume of the solution.
Summary
Chapter 5 covers the calculations involved in chemical reactions and chemical
analysis. It introduces the concept of the mole, a unit of measurement used to express
amounts of a chemical substance. The mole is a fundamental concept in quantitative
chemistry and is used in many calculations, including those involving chemical
formulae, chemical equations, gases, and solutions. The chapter also covers topics
such as empirical and molecular formulae, molar mass, and concentration.
Chapter 6: How far? How fast?
6.1 Energy changes in chemical reactions
● Exothermic reactions: release heat energy to the surroundings.
● Endothermic reactions: absorb heat energy from the surroundings.
● Examples: Combustion of fuel (exothermic), Thermal decomposition of
limestone (endothermic).
● Energy level diagrams: used to illustrate energy changes in reactions.
● Activation energy: the minimum energy required for a reaction to occur.
6.2 Rates of reaction
● Factors affecting reaction rates: temperature, concentration, particle size,
catalysts, and surface area.
● Collision theory: explains how reactions occur when particles collide with
sufficient energy and correct orientation.
● Measuring reaction rates: change in mass, the volume of gas produced, change
in color, or turbidity.
● Rate equations: express the relationship between the rate of a reaction and the
concentrations of reactants.
6.3 Catalysts
● Catalysts: substances that speed up reactions without being used up.
● How catalysts work: provide an alternative pathway with lower activation
energy.
● Types of catalysts: homogeneous (same phase as reactants) and heterogeneous
(different phase).
● Examples: Catalytic converters in cars, and enzymes in biological systems.
6.4 Photochemical reactions
● Photochemical reactions: reactions that occur when light energy is absorbed.
● Examples: Photosynthesis, the formation of ozone in the upper atmosphere.
6.5 Reversible reactions and chemical equilibria
● Reversible reactions: reactions that can proceed in both forward and reverse
directions.
● Chemical equilibrium: a state where the rates of the forward and reverse
reactions are equal.
● Equilibrium position: the relative amounts of reactants and products at
equilibrium.
● Factors affecting equilibrium position: temperature, concentration, and pressure.
● Le Chatelier's principle: predicts the effect of changes in conditions on the
equilibrium position.
Summary:
● Chemical reactions involve energy changes, either releasing or absorbing heat.
● Reaction rates are influenced by factors like temperature, concentration, and
catalysts.
● Catalysts speed up reactions by providing an alternative pathway.
● Photochemical reactions are initiated by light energy.
● Reversible reactions can reach a state of equilibrium where forward and
reverse reactions occur at equal rates.
● Le Chatelier's principle helps predict how changes in conditions affect the
equilibrium position of a reversible reaction.
Chapter 7: Patterns and Properties of
Metals
7.1 The Alkali Metals
● Alkali metals are soft and have relatively low melting points.
● They react vigorously with water, releasing hydrogen gas and forming
hydroxides.
● Reactivity increases down the group.
● Found in Group 1 of the periodic table.
● Includes lithium, sodium, potassium, rubidium, cesium, and francium.
7.2 Aluminum
● Aluminum is a light, strong, and corrosion-resistant metal.
● Extracted from its ore, bauxite, using electrolysis.
● Used in a wide range of applications, including construction, transportation,
and packaging.
7.3 The Transition Elements
● Transition elements are found in the central block of the periodic table (Groups
3-12).
● They are hard, strong, and dense metals with high melting points.
● They often form colored compounds and ions.
● They have variable oxidation states and can act as catalysts.
● Includes elements like iron, copper, gold, and silver.
7.4 The Reactivity of Metals
● The reactivity of metals can be determined by their reactions with water, acids,
and other metal ions.
● The reactivity series is a list of metals arranged in order of their reactivity.
● More reactive metals can displace less reactive metals from their compounds.
7.5 Electrical Cells and Energy
● Electrical cells convert chemical energy into electrical energy.
● They consist of two electrodes made of different metals and an electrolyte.
● Oxidation occurs at the anode, and reduction occurs at the cathode.
● The flow of electrons between the electrodes generates an electric current.
● Examples include simple cells (e.g., zinc-copper cells) and rechargeable cells
(e.g., lithium-ion batteries).
Summary:
● Chapter 7 focuses on the properties and patterns of metals in the periodic table.
● It covers the alkali metals, aluminum, transition elements, and the reactivity
series of metals.
● The chapter also explains how electrical cells work and their use in energy
conversion.
Chapter 8: Industrial Inorganic Chemistry
8.1 The extraction of metals by carbon reduction
● Metal ores are finite resources found in the Earth’s crust.
● Ores are often oxides, and metals can be extracted from them through
reduction by heating with carbon.
● Carbon is more reactive than metals and forms strong bonds with oxygen.
● The method of extraction depends on the reactivity of the metal.
● Highly reactive metals like aluminum are extracted by electrolysis.
● Moderately reactive metals like iron are extracted in a blast furnace.
● Less reactive metals like copper can be extracted by heating the ore alone.
8.2 The extraction of metals by electrolysis
● Electrolysis is used when carbon reduction is not feasible, like for very reactive
metals.
● The electrolyte is often the molten ore or the ore dissolved in another molten
salt.
● The positive electrode attracts negatively charged non-metal ions, and the
negative electrode attracts positively charged metal ions.
8.3 Ammonia and fertilizers
● Ammonia is manufactured by the Haber process, combining nitrogen from the
air with hydrogen derived from natural gas.
● The reaction is reversible, so conditions are optimized for maximum yield.
● Ammonia is used to produce fertilizers, nitric acid, and nylon.
8.4 Sulfur and sulfuric acid
● Sulfur is found in volcanic regions and as an impurity in fossil fuels.
● Sulfuric acid is produced by the Contact process, involving the catalytic
oxidation of sulfur dioxide and then reaction with water.
● Sulfuric acid is used in car batteries, making fertilizers and detergents, and
cleaning metal surfaces.
8.5 The chloralkali industry
● The chlor-alkali industry uses electrolysis of brine to produce chlorine,
hydrogen, and sodium hydroxide.
● These products are used in various applications, including water treatment,
plastics, and paper manufacturing.
8.6 Limestone
● Limestone is a versatile raw material used in construction, agriculture, and
industry.
● It is heated to produce quicklime, which reacts with water to form slaked lime.
● These compounds have numerous applications, such as neutralizing acidic soil
and making cement.
8.7 The economics of the chemical industry
● The chemical industry is crucial to the economy, producing a wide range of
essential products.
● Factors influencing the location of chemical plants include availability of raw
materials, energy, and transportation.
● Environmental considerations and safety regulations are also important in the
chemical industry.
Summary:
Chapter 8 covers the industrial processes for extracting metals and producing
important inorganic chemicals. It emphasizes the economic and environmental
factors that influence these processes and the diverse applications of the products.
Chapter 9: Organic Chemistry
9.1 The unique properties of carbon
● Carbon atoms can form four covalent bonds, allowing a diversity of organic
compounds.
● Carbon-carbon bonds can be single, double, or triple.
● Organic compounds have chains or rings of carbon atoms.
9.2 Alkanes
● Alkanes are hydrocarbons with only single C-C bonds.
● General formula: C(n)H(2n+2)
● Examples: methane (CH4), ethane (C2H6), propane (C3H8)
9.3 Alkenes
● Alkenes are hydrocarbons with at least one double C=C bond.
● General formula: C(n)H(2n)
● Examples: ethene (C2H4), propene (C3H6)
9.4 Hydrocarbon structure and isomerism
● Isomers are compounds with the same molecular formula but different
structural arrangements.
● Structural isomers have different arrangements of atoms.
● Stereoisomers have the same arrangement of atoms but different spatial
arrangements.
9.5 Chemical reactions of the alkanes
● Alkanes are generally unreactive but undergo combustion and substitution
reactions.
● Combustion: complete combustion produces carbon dioxide and water;
incomplete combustion produces carbon monoxide and/or carbon.
● Substitution: hydrogen atoms are replaced by other atoms or groups (e.g.,
halogenation).
9.6 Chemical reactions of the alkenes
● Alkenes are more reactive than alkanes due to the presence of the double bond.
● Addition reactions: atoms or groups are added to the carbon atoms across the
double bond (e.g., hydrogenation, hydration, halogenation).
9.7 Alcohols
● Alcohols have the -OH functional group.
● General formula: C(n)H(2n+1)OH
● Examples: methanol (CH3OH), ethanol (C2H5OH), propanol (C3H7OH)
9.8 The reactions of ethanol
● Ethanol undergoes combustion, oxidation, and esterification.
● Combustion: produces carbon dioxide and water.
● Oxidation: can be oxidized to ethanal and then ethanoic acid.
● Esterification: reacts with organic acids to form esters.
9.9 Organic acids and esters
● Organic acids have the -COOH functional group.
● General formula: C(n)H(2n+1)COOH
● Examples: methanoic acid (HCOOH), ethanoic acid (CH3COOH)
● Esters are formed by the reaction of an alcohol and an organic acid.
Summary:
● Organic chemistry is the study of carbon compounds, which can form a wide
variety of structures due to carbon's ability to form four covalent bonds.
● Alkanes and alkenes are hydrocarbons with different properties based on their
bonding.
● Isomerism occurs when compounds have the same molecular formula but
different structures.
● Alkanes undergo combustion and substitution reactions, while alkenes undergo
addition reactions due to their reactive double bond.
● Alcohols, organic acids, and esters are important organic compounds with
distinct properties and reactions.
Chapter 11: Petrochemicals and Polymers
11.1 Petroleum
● Petroleum is a fossil fuel formed over millions of years from the remains of
ancient marine organisms.
● It is a complex mixture of hydrocarbons, mainly alkanes.
● Fractional distillation is used to separate petroleum into fractions with different
boiling points and uses.
● Some important fractions include gasoline, naphtha, kerosene, diesel oil,
lubricating oil, and bitumen.
11.2 Alternative Fuels and Energy Sources
● The need for alternative fuels: Due to the finite nature of fossil fuels and their
environmental impact, there is a growing need for alternative, renewable energy
sources.
● Biofuels: Produced from recently living organisms, offering a renewable
alternative to fossil fuels. Ethanol (from fermentation) and biodiesel (from vegetable
oils) are examples.
● Hydrogen: Can be burned as a fuel or used in fuel cells. Produces only water
when burned, but storage and transportation are challenging due to its low density.
11.3 Addition Polymerisation
● Monomers and Polymers: Polymers are large molecules made up of repeating
subunits called monomers.
● Addition polymers: Formed when unsaturated monomers (containing double
bonds) react together.
● Examples: Polyethene (polythene) from ethene, polychloroethene (PVC) from
chloroethene, polypropene from propene, and polytetrafluoroethene (PTFE) from
tetrafluoroethene.
● Properties and uses of addition polymers: Properties vary depending on the
type of polymer. Uses include plastics, packaging materials, clothing, and electrical
insulation.
11.4 Condensation Polymerisation
● Condensation polymers: Formed when monomers with two functional groups
react together, eliminating a small molecule (often water) in the process.
● Examples: Polyesters (like Terylene) from dicarboxylic acids and diols, and
polyamides (like nylon) from dicarboxylic acids and diamines.
● Properties and uses of condensation polymers: Often have good heat resistance
and are used in textiles, carpets, and engineering plastics.
Summary
● Petroleum is a valuable source of hydrocarbons, separated into useful fractions
by fractional distillation.
● The need for alternative fuels has led to the development of biofuels and
hydrogen as potential replacements for fossil fuels.
● Addition polymerization and condensation polymerization are two important
ways of creating polymers with a wide range of properties and applications.
