Chemistry

atom — The smallest part of an element that can take part in a chemical change.

Atoms are the fundamental building blocks of all matter. They consist of a central nucleus and electrons orbiting around it. Imagine an atom as a tiny solar system, where the nucleus is the sun and the electrons are planets orbiting it. Most of the 'solar system' is empty space.

nucleus — A tiny region in the centre of the atom where nearly all of its mass is concentrated.

The nucleus contains protons and neutrons and carries a positive charge due to the protons. Its small size and high density were deduced from Rutherford's alpha particle scattering experiment. If an atom were the size of a football stadium, its nucleus would be like a pea in the center, yet containing almost all the mass.

nucleons — Particles that make up the nucleus, specifically protons and neutrons.

Protons and neutrons are the subatomic particles found within the atomic nucleus. The total number of nucleons in an atom is its mass number. Think of nucleons as the 'bricks' that build the central part of the atom, the nucleus.

protons — Positively charged particles found in the nucleus of an atom.

Protons have a relative charge of +1 and a relative mass of 1. The number of protons defines the atomic number of an element and determines its identity. Protons are like the 'identity cards' of an atom; their number tells you exactly which element it is.

neutrons — Neutral (uncharged) particles found in the nucleus of an atom.

Neutrons have a relative charge of 0 and a relative mass of 1. They contribute to the mass of the atom but not its charge, and their varying numbers lead to isotopes. Neutrons are like 'spacers' or 'ballast' in the nucleus; they add mass but don't affect the charge or chemical identity.

electrons — Negatively charged particles that move around the nucleus in regions of space called orbitals or shells.

Electrons have a relative charge of -1 and a very small relative mass (approximately 1/2000 of a proton). They are responsible for chemical bonding and the overall charge of an ion. Electrons are like the 'workers' of the atom, involved in all the chemical interactions and determining how atoms bond together.

orbitals — Regions of space outside the nucleus where there is a probability of finding a particular electron.

The orbital model is a more accurate representation of electron location, used for discussing bonding. It contrasts with the simpler shell model which places electrons at fixed distances. Think of an orbital not as a fixed path, but as a 'cloud' or 'fuzzy region' where an electron is most likely to be found.

electron shells — A simpler model of the atom where electrons move around the nucleus at certain distances, each at its own particular energy level.

This model is convenient for understanding what happens to electrons during chemical reactions and when discussing ionisation energies. Each shell can hold a specific maximum number of electrons. Imagine electron shells as concentric layers of an onion around the nucleus, with electrons occupying these layers.

atomic number (proton number) — The number of protons in the nucleus of an atom, denoted by Z.

The atomic number uniquely identifies an element; every atom of a given element has the same atomic number. It also equals the number of electrons in a neutral atom. The atomic number is like an element's unique ID number; no two elements have the same one.

mass number (nucleon number) — The number of protons plus neutrons in the nucleus of an atom, denoted by A.

The mass number represents the total number of nucleons in an atom. It is used to distinguish between isotopes of the same element. The mass number is like the 'total count' of the heavy particles in the atom's core.

isotopes — Atoms of the same element that have different numbers of neutrons.

Isotopes have the same atomic number (same number of protons) but different mass numbers. They exhibit identical chemical properties due to the same electron configuration but differ in physical properties like mass and density. Think of isotopes as 'siblings' of the same element; they share the same family name (element) and core characteristics (protons/electrons) but have slightly different 'weights' (neutrons).

ions — Electrically charged particles formed from atoms or molecules by the gain or loss of electrons.

Positive ions (cations) are formed when an atom loses electrons, resulting in more protons than electrons. Negative ions (anions) are formed when an atom gains electrons, resulting in more electrons than protons. Ions are like atoms that have either 'donated' or 'received' electrons, giving them an overall electrical charge.

principal quantum shell — These principal energy levels or principal quantum shells (symbol n) are numbered according to how far they are from the nucleus.

These shells are numbered (n=1, 2, 3, etc.) with n=1 being closest to the nucleus and having the lowest energy. Electrons in shells further from the nucleus have more energy and are held less tightly. The higher the floor number in a building, the further it is from the ground (nucleus) and the more energy it takes to get there.

sub-shell — The principal quantum shells, apart from the first, are split into sub-shells (sub-levels).

These sub-shells are designated by letters s, p, d, and f, each having a specific maximum number of electrons (s=2, p=6, d=10, f=14). Within a principal shell, the energy of sub-shells increases in the order s < p < d. If a principal quantum shell is a 'floor' in an apartment building, then sub-shells are the different 'types of apartments' on that floor, each with a different capacity.

atomic orbital — An atomic orbital is a region of space around the nucleus of an atom that can be occupied by one or two electrons.

Each sub-shell contains a specific number of orbitals (s=1, p=3, d=5), and each orbital can hold a maximum of two electrons with opposite spins. Orbitals have distinct three-dimensional shapes. If a sub-shell is an 'apartment type', then an atomic orbital is a specific 'room' within that apartment, which can accommodate up to two electrons.

s orbital — An s orbital has a spherical shape.

There is one s orbital in each principal quantum shell. Electrons in s orbitals are found within a spherical region around the nucleus. Higher energy s orbitals (e.g., 2s, 3s) are larger but retain their spherical shape. Imagine an s orbital as a perfectly round balloon centered on the nucleus, where the electron is most likely to be found within the balloon's volume.

p orbital — The shape of a p orbital is like an hourglass with two 'lobes'.

There are three p orbitals in each principal quantum shell from n=2 onwards, arranged at right angles to each other along the x, y, and z axes (px, py, pz). All three p orbitals within a sub-shell have the same energy. Think of a p orbital as a dumbbell or a figure-eight shape, with the nucleus located at the narrow 'waist' between the two lobes.

electronic configuration — The arrangement of electrons in an atom is called its electronic structure or electronic configuration.

This describes how electrons are distributed among the atom's shells, sub-shells, and orbitals. It dictates an atom's chemical properties and reactivity. Think of electronic configuration as an atom's 'address book' for its electrons, specifying which 'street' (shell), 'house' (sub-shell), and 'room' (orbital) each electron resides in.

spin-pair repulsion — Electrons in the same atomic orbital in a sub-shell repel each other more than electrons in different atomic orbitals.

This increased repulsion makes it easier to remove an electron from a paired orbital, thus decreasing the first ionisation energy. Electrons prefer to occupy separate orbitals within a sub-shell to minimize this repulsion. Imagine two people sharing a small room versus having their own rooms; the two people in the same room experience more 'repulsion'.

free radical — A free radical is a species with one or more unpaired electrons.

Free radicals are highly reactive due to their unpaired electron(s), which seek to form a pair by reacting with other species. They are often formed by homolytic fission of covalent bonds. Think of a free radical as a person with one hand free in a crowded room, constantly looking for another hand to hold to feel 'complete' or stable.

first ionisation energy — The first ionisation energy of an element is the energy needed to remove one electron from each atom in one mole of atoms of the element in the gaseous state to form one mole of gaseous 1+ ions.

This is an endothermic process, meaning energy must be supplied. It is a key indicator of how strongly the outermost electron is held by the nucleus and is influenced by nuclear charge, distance, shielding, and spin-pair repulsion. Imagine trying to pull a single magnet off a large metal surface; this is the effort needed to pull off the very first magnet.

successive ionisation energies — We can continue to remove electrons from an atom until only the nucleus is left, and this sequence of ionisation energies is called successive ionisation energies.

Each successive ionisation energy is higher than the last because electrons are removed from increasingly positive ions, requiring more energy to overcome the stronger electrostatic attraction. Large jumps indicate removal from a new, inner principal quantum shell. Continuing the magnet analogy, successive ionisation energies are the increasing effort needed to pull off the second, third, and subsequent magnets.

s-block — Elements in Groups 1 and 2 have outer electrons in an s sub-shell and are therefore together called the s-block.

These elements are highly reactive metals, readily losing their one or two s-electrons to form positive ions. Their chemical properties are largely determined by these outer s-electrons. In the Periodic Table 'apartment building', the s-block is like the 'ground floor' or 'entrance wing' where the first electrons always go into the s-type rooms.

p-block — Elements in Groups 13 to 18 (apart from He) have outer electrons in a p sub-shell and are therefore together called the p-block.

This block includes a wide range of elements from metals to non-metals and metalloids, with varying chemical properties depending on how many p-electrons they have and their position in the group. The p-block is like the 'upper floors' or 'main residential wing' of the Periodic Table apartment building, where electrons start filling the p-type rooms after the s-type rooms are full.

d-block elements — Elements that add electrons to the d sub-shells are called the d-block elements.

Most d-block elements are transition elements, characterized by their ability to form ions with variable oxidation states and colored compounds, due to the involvement of their d-electrons in bonding. The d-block is like a 'middle section' or 'special amenities wing' in the Periodic Table apartment building, where electrons fill the d-type rooms after the s-type rooms of the next floor are already occupied.

unified atomic mass unit — One twelfth of the mass of an unbound neutral atom of the carbon-12 isotope in its ground state.

This unit (symbol u or Da) provides a standard reference for comparing the masses of atoms and molecules, as individual atomic masses are too small to measure directly. It is based on carbon-12 because it is a stable and abundant isotope, much like comparing the weight of all fruits to a standard slice of a specific apple.

relative isotopic mass — The mass of a particular isotope of an element which has the Avogadro number of atoms (6.02 × 10^23).

This term refers to the mass of a specific isotope, such as carbon-13, rather than the average mass of all isotopes of an element. It is used in mass spectrometry to identify individual isotopes, similar to knowing the exact height of one specific student in a class.

relative atomic mass — The ratio of the average mass of the atoms of an element to the unified atomic mass unit.

This value (Ar) is a weighted average that accounts for the natural abundance of an element's isotopes. It has no units because it is a ratio. Values are typically found in the Periodic Table, much like the average weight of marbles in a bag compared to a standard marble.

relative molecular mass — The ratio of the weighted average mass of a molecule of a molecular compound to the unified atomic mass unit.

This value (Mr) is calculated by summing the relative atomic masses of all atoms in a molecule. It applies to molecular compounds and has no units, similar to the total weight of all LEGO bricks in a specific model.

relative formula mass — The sum of the relative atomic masses of all atoms present in one formula unit of an ionic compound.

This term (also symbol Mr) is used for compounds containing ions or giant covalent structures, calculated in the same way as relative molecular mass by adding up the Ar values of all atoms in the simplest formula unit. It has no units, like the total weight for a single serving of a dish that doesn't exist as discrete molecules.

water of crystallisation — Water that is part of the structure of some crystalline compounds.

Compounds containing this water are called hydrated compounds, and the water molecules are chemically bonded within the crystal lattice. Heating can remove this water, converting the compound to its anhydrous form, much like a sponge that has absorbed water.

hydrated compound — A compound containing water of crystallisation as part of its structure.

These compounds have a specific number of water molecules associated with each formula unit, indicated by a dot in their chemical formula (e.g., CuSO4·5H2O). They often have distinct physical properties from their anhydrous forms, similar to a house with a built-in swimming pool.

anhydrous compound — A compound which does not contain water of crystallisation.

This is the form of a compound after all water of crystallisation has been removed, typically by heating. Anhydrous compounds can often absorb water to become hydrated again, much like a dried-out sponge.

mass spectrometer — An instrument used to measure the mass of each isotope present in an element and their relative abundances.

It works by converting atoms into ions, deflecting them with a magnetic field, and detecting them based on their mass-to-charge ratio. The output is a mass spectrum, much like a strong wind blowing different sized balls, where heavier balls are deflected less.

relative abundance — The proportion of each isotope present in a sample of an element.

This is measured by a mass spectrometer and is displayed on the vertical axis of a mass spectrum. It is crucial for calculating the weighted average relative atomic mass of an element, similar to the percentage of boys and girls in a class.

mass spectrum — A display from a mass spectrometer showing relative abundance on the vertical axis and mass to ion charge ratio (m/e) on the horizontal axis.

Each peak in the spectrum corresponds to an ion with a specific m/e ratio, allowing for the identification of isotopes, molecular ions, and fragments. For singly charged ions, m/e equals the mass number, much like a bar chart where each bar represents a different 'weight' of atom or molecule.

molecular ion — An ion, M+, formed when one electron is removed from a molecule to form an ion with a single positive charge.

This ion corresponds to the intact molecule and typically gives the peak at the highest mass-to-charge ratio in a mass spectrum, which represents the relative molecular mass of the compound, like a picture of a whole car before it's broken down.

fragmentation — The process where molecules are broken apart by electron bombardment in a mass spectrometer, producing smaller ions.

These smaller ions, or fragments, appear as peaks at lower m/e values in the mass spectrum. Analysing these patterns can help deduce the structure of an organic molecule, similar to how the pieces of a shattered bottle can tell you about the original bottle.

[M + 1] peak — A very small peak just beyond the molecular ion peak at a mass of [M + 1], caused by molecules in which one of the carbon atoms is the 13C isotope.

Since natural carbon contains 1.10% carbon-13, this peak is always present. Its relative abundance compared to the M+ peak can be used to deduce the number of carbon atoms in a molecule, like finding a rare, slightly heavier version of a common item.

[M + 2] peak — A peak two units beyond the molecular ion peak, often indicating the presence of chlorine or bromine atoms in a compound due to their heavier isotopes.

Chlorine has 35Cl and 37Cl (ratio 3:1), and bromine has 79Br and 81Br (ratio 1:1). The relative height of the [M+2] peak compared to the M peak helps identify which halogen is present and how many, like having two types of identical twins, one slightly heavier.

Avogadro constant — The number of particles equivalent to the relative atomic mass or relative molecular mass of a substance in grams, with a numerical value of 6.02 × 10^23.

This constant (L or NA) defines the number of specified particles (atoms, molecules, ions, electrons) in one mole of any substance. It links the microscopic world of atoms to the macroscopic world of measurable masses, acting like a 'chemist's dozen'.

mole — The amount of substance which contains 6.02 × 10^23 specified particles.

The mole (mol) is the SI unit for amount of substance. One mole of any substance has a mass in grams equal to its relative atomic mass or relative molecular mass. It allows chemists to work with measurable quantities of substances, like a 'standard package' of particles.

molar mass — The mass in grams of 1 mole of a compound, whether ionic, simple molecules or giant covalent structures.

The units of molar mass are g mol^-1. It is numerically equal to the relative molecular mass or relative formula mass of a substance. It is used to convert between mass and moles, representing the weight of a 'standard package' of particles.

balanced equation — A chemical equation where the number of each type of atom is the same on both the reactants side and the products side.

This adheres to the law of conservation of mass, stating that atoms cannot be created or destroyed in a chemical reaction. Stoichiometric numbers are placed in front of formulae to achieve balance, like ensuring all LEGO pieces are used when building a new model.

state symbols — Symbols used to specify the physical states of reactants and products in a chemical reaction.

The symbols are (s) for solid, (l) for liquid, (g) for gas, and (aq) for aqueous solution. They provide important information, especially for understanding chemical equilibria and reaction rates, like adding labels to ingredients in a recipe.

limiting reagent — The reactant which is not in excess and is completely consumed in a chemical reaction.

This reagent determines the maximum amount of product that can be formed. Once it is used up, the reaction stops, even if other reactants are still present, like the cheese in a sandwich recipe if you run out of it first.

excess reagent — The reactant which has the number of moles in excess in a chemical reaction.

This reagent is not completely consumed during the reaction; some of it will remain after the limiting reagent has been used up. Its quantity does not determine the maximum amount of product formed, like having extra bread when making sandwiches.

empirical formula — The simplest whole number ratio of the elements present in one molecule or formula unit of the compound.

It is determined from the mass or percentage composition of elements in a compound. For ionic compounds, the empirical formula is always the formula unit, similar to simplifying a recipe ratio like 6 eggs to 12 cups of flour to 1 egg to 2 cups of flour.

molecular formula — The actual number of each of the different atoms present in a molecule.

It is always a multiple of the empirical formula. To deduce it, both the empirical formula and the relative molecular mass of the compound are needed, like knowing the actual recipe (6 eggs and 12 cups of flour) from a simplified ratio.

oxidation numbers — Roman numerals after ions that show the charge on the ion, especially for transition elements.

They are used in naming compounds to indicate the specific charge of a metal ion, particularly when an element can form ions with different charges (e.g., iron(II) vs. iron(III)), similar to a suffix on a name like 'King Henry VIII'.

compound ions — Ions that contain more than one type of atom.

These ions are polyatomic and carry an overall positive or negative charge (e.g., sulfate ion SO4^2-, ammonium ion NH4^+). Their formulae and charges must be recalled to write correct ionic compound formulae, like a small team of different players carrying a single 'charge'.

spectator ions — Ions that play no part in a chemical reaction and remain unchanged in solution.

These ions are present in the reaction mixture but do not participate in the actual chemical change. They are omitted when writing ionic equations, much like people watching a football game who are present but not actively playing.

ionic equation — A simplified chemical equation that shows only the ions or other particles taking part in a reaction, with spectator ions omitted.

Ionic equations are particularly useful for reactions in aqueous solution, precipitation reactions, and reactions involving changes in oxidation state. Both atoms and charges must be balanced, similar to a highlight reel of a sports game showing only key plays.

concentration — The amount of solute dissolved in a solvent to make 1 dm^3 (one cubic decimetre) of solution.

It is typically expressed in mol dm^-3. A high concentration means a large amount of solute per unit volume, while a low concentration means a small amount, like how much cordial syrup you put into a litre of water.

titration — A procedure used to determine the amount of substance present in a solution of unknown concentration.

It typically involves reacting a solution of known concentration (titrant) with a solution of unknown concentration (analyte) until the reaction is complete, usually indicated by a colour change from an indicator, like precisely measuring paint to get a desired shade.

Electronegativity — Electronegativity is the power of a particular atom that is covalently bonded to another atom to attract the bonding pair of electrons towards itself.

It is a measure of an atom's ability to attract electrons in a covalent bond. Electronegativity increases across a period and up a group in the Periodic Table, with fluorine being the most electronegative element. Imagine a tug-of-war over a rope (the shared electrons) between two people (atoms); the stronger person (more electronegative atom) pulls the rope closer to themselves.

Ionic bonding — Ionic bonding is the electrostatic attraction between positive ions (cations) and negative ions (anions) in an ionic crystal lattice.

This type of bonding typically occurs between metals and non-metals, where electrons are transferred from the metal atom to the non-metal atom. The resulting oppositely charged ions are held together by strong electrostatic forces, forming a stable crystal lattice. Imagine two magnets, one positively charged and one negatively charged, strongly attracting each other and holding together in a fixed arrangement, like bricks in a wall.

Covalent bonds — Covalent bonds are formed when the outer electrons of two atoms are shared.

This type of bonding typically occurs between non-metal atoms, allowing each atom to achieve a stable noble gas electronic configuration by sharing electron pairs. The shared electrons are attracted to the nuclei of both atoms, holding them together. Think of two friends sharing a toy; both benefit from having access to the toy, even though neither fully owns it.

Single covalent bond — A shared pair of electrons is called a single covalent bond, or a bond pair.

This is the simplest form of covalent bonding where two atoms contribute one electron each to form a shared pair. It is represented by a single line between the atoms in a displayed formula. Two people each contributing one item to a shared potluck dish.

Double covalent bond — A double covalent bond is formed when some atoms bond together by sharing two pairs of electrons.

This type of bond is stronger and shorter than a single covalent bond due to the greater quantity of negative charge between the two atomic nuclei. Examples include oxygen (O2) and carbon dioxide (CO2). Two friends sharing two toys, doubling their connection and making it harder to separate them.

Triple covalent bond — A triple covalent bond is formed when atoms bond together by sharing three pairs of electrons.

This is the strongest and shortest type of covalent bond, involving the sharing of six electrons between two atoms. Nitrogen (N2) is a classic example, where each nitrogen atom achieves an octet by forming a triple bond. Three friends sharing three toys, creating an even stronger and more compact connection.

Co-ordinate bond — A co-ordinate bond (or dative covalent bond) is formed when one atom provides both the electrons needed for a covalent bond.

This occurs when one atom has a lone pair of electrons and a second atom has an unfilled orbital to accept that lone pair. The bond is still a shared pair of electrons, but both electrons originate from the same atom. One person bringing both ingredients for a shared meal, rather than each person bringing one.

Lone pairs — The pairs of outer-shell electrons not used in bonding are called lone pairs.

These electron pairs are located on an atom but are not involved in forming covalent bonds with other atoms. They contribute to the overall electron density around an atom and significantly influence molecular shape due to their greater repulsive forces compared to bonding pairs. Imagine a person having extra hands that aren't holding anyone else's hands; these extra hands still take up space and push others away.

Octet of electrons — When ions or atoms have 8 electrons in their outer shell like this it is called an octet of electrons.

This refers to the stable electronic configuration of having eight electrons in the outermost principal energy level, typically achieved by noble gases. Atoms often gain, lose, or share electrons to attain this stable configuration. Like having a full set of eight pieces in a board game, making your position very stable and complete.

Electron deficient — An electron-deficient compound is one where an atom has an unfilled orbital to accept a lone pair.

This term describes molecules where the central atom has fewer than eight electrons in its outer shell, making it capable of accepting a lone pair of electrons from another atom to form a co-ordinate bond. Boron trifluoride (BF3) is a common example. Imagine a person with an empty seat at their table, ready for someone to join them.

Expanded octet — An expanded octet occurs when an atom has more than eight electrons around its central atom.

This phenomenon is observed in elements from Period 3 and beyond (e.g., phosphorus, sulfur, chlorine) which can use their unfilled d-orbitals to accommodate more than eight electrons in their outer shell when forming compounds like SF6 or PCl5. Like a person having more than two hands to hold things, allowing them to bond with more partners than usual.

Bond length — The internuclear distance between two covalently bonded atoms is called the bond length.

It is the distance between the nuclei of two atoms joined by a covalent bond. Bond length is inversely related to bond strength; generally, shorter bonds are stronger bonds due to greater electron density between the nuclei. The distance between two people holding hands; the closer they are, the tighter their grip.

Bond energy — Bond energy is the energy needed to break one mole of a given covalent bond in the gaseous state.

It is a measure of the strength of a covalent bond, expressed in kilojoules per mole (kJ mol−1). Higher bond energy indicates a stronger bond, requiring more energy to break, which can influence a compound's reactivity. The amount of effort required to pull apart two LEGO bricks that are strongly connected.

Sigma bonds — Sigma bonds (σ bonds) are formed from end-on overlap of atomic orbitals.

These are the strongest type of covalent bond, characterized by electron density that is symmetrical about the internuclear axis. They can be formed by the overlap of s-s, s-p, or p-p (end-on) atomic orbitals, or hybrid orbitals. Imagine two people shaking hands directly in front of them, creating a strong, direct connection.

Pi bonds — Pi bonds (π bonds) are formed from sideways overlap of atomic orbitals.

These bonds are typically formed by the sideways overlap of p orbitals, resulting in electron density above and below the internuclear axis. A single π bond consists of two electron clouds and is generally weaker than a σ bond. Imagine two people holding hands above and below a table, creating a connection that is less direct than a handshake across the table.

Hybridisation — The mixing of atomic orbitals is called hybridisation.

This process involves the combination of atomic orbitals (e.g., s and p orbitals) to form new, degenerate hybrid orbitals (e.g., sp, sp2, sp3) that are more suitable for forming covalent bonds. Hybridisation helps explain observed molecular geometries and bond angles. Like blending different colors of paint to create new, uniform colors that are better suited for a specific painting.

Metallic bonding — Metallic bonding is the electrostatic attraction between positive ions and delocalised electrons.

In metals, atoms lose their outer shell electrons, forming positive ions arranged in a lattice. These electrons become delocalised, meaning they are free to move throughout the entire metal structure, creating a 'sea' of electrons that holds the positive ions together. Imagine a crowd of people (positive ions) holding hands in a swimming pool (delocalised electrons); the water allows them to move but keeps them connected.

Delocalised electrons — Delocalised electrons (mobile electrons) are electrons that are not associated with any one particular atom or bond.

In metallic bonding, these are the outer shell electrons that are free to move throughout the entire metal lattice, rather than being confined to a specific atom or covalent bond. Their mobility is responsible for the electrical and thermal conductivity of metals. Imagine a group of children playing freely in a park, not tied to any specific parent, able to move anywhere within the park boundaries.

Polar — A bond is polar (or that it has a dipole) when a covalent bond is formed between two atoms having different electronegativity values, causing the more electronegative atom to attract the pair of electrons in the bond towards it.

This results in an asymmetric electron distribution, where one atom gains a partial negative charge (δ–) and the other a partial positive charge (δ+). The degree of polarity is measured as a dipole moment. Like a slightly uneven tug-of-war, where one side pulls the rope a bit closer, creating a slight imbalance.

Van der Waals’ forces — Van der Waals’ forces is a general term used to describe all intermolecular forces.

These are weak attractive forces that exist between molecules, distinct from the strong intramolecular forces (covalent, ionic, metallic bonds). They include instantaneous dipole–induced dipole forces, permanent dipole–permanent dipole forces, and hydrogen bonding. Think of them as very weak sticky notes holding molecules together, easily pulled apart compared to superglue (covalent bonds) holding atoms within a molecule.

Instantaneous dipole–induced dipole (id–id) forces — Instantaneous dipole–induced dipole (id–id) forces, or London dispersion forces, are very weak forces of attraction that exist between all atoms or molecules due to temporary dipoles.

These forces arise from the constant movement of electron clouds, which can momentarily create an uneven distribution of charge, forming a temporary dipole. This temporary dipole can then induce a dipole in a neighbouring molecule, leading to a weak, fleeting attraction. Imagine a group of people randomly shifting their weight; for a brief moment, more weight might be on one side, causing a slight tilt, which then influences the person next to them to tilt slightly in response.

Permanent dipole–permanent dipole (pd–pd) forces — Permanent dipole–permanent dipole (pd–pd) forces are attractive forces between two molecules having permanent dipoles.

These forces occur in polar molecules where there is a permanent separation of charge due to differences in electronegativity between bonded atoms. The partial positive end of one molecule is attracted to the partial negative end of a neighbouring molecule. Like tiny, fixed magnets, where the north pole of one molecule is always attracted to the south pole of another, creating a constant, albeit weak, pull.

Hydrogen bonding — Hydrogen bonding is the strongest form of intermolecular bonding, a type of permanent dipole–permanent dipole bonding.

It occurs when a hydrogen atom covalently bonded to a highly electronegative atom (F, O, or N) is attracted to a lone pair of electrons on another highly electronegative atom (F, O, or N) in a neighbouring molecule. This creates a particularly strong intermolecular attraction. Imagine a very strong, specific handshake between molecules, where one hand (hydrogen) is strongly attracted to a specific glove (lone pair) on another molecule.

Ideal gas — A theoretical gas that fits the description of the kinetic theory of gases, having zero particle volume and no intermolecular forces of attraction.

Ideal gases are a theoretical model used to simplify gas behaviour, assuming negligible molecular volume and no interactions between molecules. Real gases approximate ideal behaviour under specific conditions, typically high temperature and low pressure, where intermolecular forces and molecular volume become less significant. Imagine a room full of perfectly bouncy, invisible ping-pong balls that never stick to each other and take up no space themselves, only the space they move in. That's an ideal gas.

Melting point — The temperature at which a solid changes to a liquid at 1 atmosphere pressure.

At the melting point, particles in a solid gain enough kinetic energy to overcome the forces holding them in a fixed lattice, allowing them to slide past each other. For ionic compounds, this requires high temperatures due to strong ionic bonds, while molecular solids have lower melting points due to weaker intermolecular forces. Think of ice cubes in a glass. As the temperature rises to 0 °C, the ice starts to turn into water, but both solid and liquid can coexist at that specific temperature.

Vapour pressure — The pressure exerted by a vapour in equilibrium with its liquid.

In a closed container, liquid molecules evaporate to form vapour, and vapour molecules condense back into liquid. When the rates of evaporation and condensation become equal, an equilibrium is established, and the pressure exerted by the vapour at this point is the vapour pressure. It increases with temperature as more molecules have enough energy to escape the liquid phase. Imagine a sealed soda bottle. Even if it's not fizzing, there's always some CO2 gas above the liquid, creating pressure. That's the vapour pressure of CO2 in equilibrium with the dissolved CO2.

Boiling point — The temperature at which a liquid changes to a gas at 1 atmosphere (101 325 Pa) pressure, or the temperature at which the vapour pressure is equal to the atmospheric pressure.

At the boiling point, the vapour pressure of the liquid becomes equal to the external atmospheric pressure, allowing bubbles of vapour to form throughout the liquid, not just at the surface. This requires sufficient energy to overcome all intermolecular forces holding the liquid together. When you boil water, you see bubbles forming throughout the liquid, not just steam rising from the surface. This is because the water's internal pressure (vapour pressure) is finally strong enough to push against the air pressure above it.

Crystal lattice — A regularly repeating arrangement of ions, atoms or molecules within a crystal.

Many solids, including ionic, metallic, and covalent compounds, form crystals due to the regular, ordered packing of their constituent particles. This repeating pattern extends in three dimensions, defining the macroscopic shape and properties of the crystal. Think of a perfectly stacked pyramid of oranges in a grocery store; each orange is in a specific, repeating position relative to its neighbours, forming a larger, ordered structure.

Giant ionic structures — Compounds with ionic lattices, which have three-dimensional arrangements of alternating positive and negative ions.

These structures are characterised by strong electrostatic forces of attraction between large numbers of oppositely charged ions, acting in all directions. This results in high melting/boiling points, hardness, brittleness, and electrical conductivity only when molten or in solution. Imagine a massive, intricate 3D checkerboard where every black square is a positive ion and every white square is a negative ion, all strongly attracted to each other in a vast, repeating pattern.

Allotropes — Different crystalline or molecular forms of the same element.

Allotropes are different structural modifications of an element, where the atoms are bonded together in a different manner. This leads to different physical and chemical properties, even though they are composed of the same element. Examples include diamond, graphite, and fullerenes for carbon. Think of different ways you can arrange LEGO bricks of the same colour. You can build a tall tower or a flat wall; both are made of the same bricks but have very different shapes and properties.

Nanoparticles — Individual particles in fullerenes that may have one of their dimensions between 0.1 and 100 nanometres (1 nanometre = 10⁻⁹ m).

Nanoparticles are extremely small particles, typically in the nanometre range, which often exhibit unique properties compared to their bulk material due to their high surface area to volume ratio. Fullerenes are a class of carbon allotropes that can exist as nanoparticles. Imagine a grain of sand (bulk material) versus a single atom (molecular level). Nanoparticles are somewhere in between, like tiny dust motes that are still much smaller than what you can see with the naked eye.

Buckminsterfullerene — The first fullerene discovered, C60, which is a simple molecular structure with the shape of a football.

C60 molecules consist of 60 carbon atoms arranged at the corners of 20 hexagons and 12 pentagons, forming a hollow sphere. It is a simple molecular substance with weak intermolecular forces, leading to a relatively low sublimation point, softness, and poor electrical conductivity compared to graphite. It's literally shaped like a soccer ball, with the carbon atoms forming the vertices where the panels meet.

Nanotubes — A class of fullerenes described as hexagonally arranged carbon atoms like a single layer of graphite bent into the form of a cylinder.

Nanotubes are cylindrical fullerenes with high electrical conductivity along their long axis due to delocalised electrons, and very high tensile strength. They have strong covalent bonding throughout their structure, resulting in high melting points. Imagine taking a sheet of chicken wire (representing graphite) and rolling it up into a seamless tube. That's essentially a nanotube.

Graphene — A single isolated layer of graphite, consisting of a hexagonally arranged sheet of carbon atoms.

Graphene is the most chemically reactive form of carbon, extremely strong for its mass, and an excellent conductor of electricity and heat. Its properties are exaggerated compared to graphite due to its two-dimensional nature. Imagine peeling off a single, atom-thin layer from a piece of graphite. That incredibly thin, strong, and conductive sheet is graphene.

enthalpy — Enthalpy is the total energy associated with the materials that react.

It is a state function, meaning its value depends only on the current state of the system, not on the path taken to reach that state. We cannot measure absolute enthalpy, only changes in enthalpy (ΔH). Think of the total potential energy stored in a stretched spring. You can't measure the 'absolute' potential energy, but you can measure the change in potential energy when it's stretched or released.

enthalpy change — The energy exchange between a chemical reaction and its surroundings at constant pressure is called the enthalpy change.

It is represented by ΔH and measured in kilojoules per mole (kJ mol⁻¹). A negative ΔH indicates an exothermic reaction (energy released), while a positive ΔH indicates an endothermic reaction (energy absorbed). Think of a bank account balance. If you deposit money (endothermic), your balance (enthalpy) goes up. If you withdraw money (exothermic), your balance goes down.

exothermic — Chemical reactions that transfer heat energy to the surroundings are described as exothermic.

In an exothermic reaction, the temperature of the surroundings increases, and the enthalpy change (ΔH) is negative because the enthalpy of the products is less than the enthalpy of the reactants. Burning a candle is exothermic; it releases heat and light, making the surroundings warmer.

endothermic — Chemical reactions that absorb heat energy from the surroundings are described as endothermic.

In an endothermic reaction, the temperature of the surroundings decreases, and the enthalpy change (ΔH) is positive because the enthalpy of the products is greater than the enthalpy of the reactants. An instant cold pack is endothermic; it absorbs heat from its surroundings, making the pack feel cold.

activation energy — The activation energy, EA, is the minimum energy that colliding particles must possess for a reaction to happen.

It is always a positive value because energy must be absorbed to increase the kinetic energy of reactant molecules, allowing them to collide with enough force to break existing bonds and initiate the reaction. Imagine pushing a ball up a hill. The activation energy is the initial push needed to get the ball over the top of the hill before it can roll down the other side (forming products).

standard conditions — Standard conditions are a pressure of 101 kPa, a temperature of 298 K (25 °C), and each substance involved in the reaction being in its normal physical state at these conditions.

These conditions are used to make comparisons of enthalpy changes fair and consistent. The symbol ⊖ is used to indicate that an enthalpy change refers to a reaction carried out under standard conditions. Like setting a baseline for comparing athletes' performance – everyone runs on the same track, at the same temperature, with the same equipment.

standard enthalpy change of reaction — The standard enthalpy change of reaction, ΔHᵣ⊖, is the enthalpy change when the amounts of reactants shown in the stoichiometric equation react to give products under standard conditions.

The reactants and products must be in their standard states. This value can be exothermic or endothermic and is specific to the balanced chemical equation given. It's like the total energy cost or gain for a specific recipe, where the ingredients and final dish are all at a standard 'room temperature' and 'atmospheric pressure'.

standard enthalpy change of formation — The standard enthalpy change of formation, ΔHᵠf, is the enthalpy change when one mole of a compound is formed from its elements under standard conditions.

The reactants (elements) and products (compound) must be in their standard states. By definition, the standard enthalpy change of formation of any element in its standard state is zero. Imagine building a house from scratch (elements) to a finished house (compound). The ΔHᵠf is the energy change for building exactly one house from its basic components.

standard enthalpy change of combustion — The standard enthalpy change of combustion, ΔHᵠc, is the enthalpy change when one mole of a substance is burnt in excess oxygen under standard conditions.

The reactants and products must be in their standard states. Enthalpy changes of combustion are always exothermic (ΔH is negative) because burning releases energy. It's the energy released when you completely burn one standard-sized log (one mole of substance) in a bonfire (excess oxygen).

standard enthalpy change of neutralisation — The standard enthalpy change of neutralisation, ΔHᵠneut, is the enthalpy change when one mole of water is formed by the reaction of an acid with an alkali under standard conditions.

This is essentially the enthalpy change for the reaction H⁺(aq) + OH⁻(aq) → H₂O(l). It is typically an exothermic process, with a value around −57.1 kJ mol⁻¹ for strong acid-strong alkali reactions. It's the specific energy released when one 'unit' of acid and one 'unit' of alkali combine to form one 'unit' of water, like two puzzle pieces fitting together perfectly.

standard enthalpy change of atomisation — The standard enthalpy change of atomisation, ΔHᵠat, is the enthalpy change when one mole of gaseous atoms is formed from its element under standard conditions.

This process always requires energy input, so ΔHᵠat values are always positive (endothermic). It represents the energy needed to break all bonds in one mole of an element to form individual gaseous atoms. Imagine taking a solid block of LEGO bricks (an element) and breaking it down into individual, separated LEGO bricks (gaseous atoms). The energy needed for this is the enthalpy of atomisation.

bond dissociation energy — The amount of energy needed to break a specific covalent bond is called the bond dissociation energy.

This is also known as exact bond energy or bond enthalpy. Bond dissociation energies are always positive (endothermic) because energy is required to break bonds. It's like the exact amount of force needed to snap a specific type of stick. Each stick (bond) might require a slightly different force depending on its exact composition.

bond energy — The amount of energy needed to break a specific covalent bond is called the bond dissociation energy, sometimes called bond energy or bond enthalpy.

Bond energies are always positive values, representing the energy absorbed to break a bond. When new bonds are formed, the same amount of energy is released (negative value). Think of it as the 'strength' of a connection between two atoms. A high bond energy means a strong connection that requires a lot of effort to break.

average bond energy — Average bond energies are values taken from a number of bonds of the same type but in different environments.

These are used because the exact bond energy of a particular bond (e.g., O–H) can vary slightly depending on the other atoms in the molecule. Using average values allows for estimations of enthalpy changes. It's like calculating the average height of all students in a school rather than the exact height of one specific student. It gives a good general idea, but not precise individual data.

specific heat capacity — The energy required to raise the temperature of 1 g of a substance by 1 °C (1 K) is called the specific heat capacity, c, of the liquid.

For water, its specific heat capacity is 4.18 J g⁻¹ °C⁻¹. This value is crucial for calculating heat transferred (q) in calorimetry experiments. Think of it as how 'stubborn' a substance is to heat up. Water has a high specific heat capacity, meaning it takes a lot of energy to change its temperature, like a large, heavy object is hard to get moving.

Hess’s law — Hess’s law states that 'the total enthalpy change in a chemical reaction is independent of the route by which the chemical reaction takes place as long as the initial and final conditions are the same'.

This law is a direct consequence of the law of conservation of energy and allows for the calculation of enthalpy changes that cannot be measured directly by experiment, using enthalpy cycles. Imagine climbing a mountain. The total change in your altitude (enthalpy change) from base camp to the summit is the same, whether you take a direct path or a winding, multi-stage route.

Oxidation — Oxidation is an increase in oxidation number or loss of electrons.

This definition extends beyond simple oxygen gain or hydrogen removal to encompass electron transfer and changes in oxidation states, making it universally applicable to all redox reactions. It always occurs simultaneously with reduction. Think of oxidation as 'losing money' in a transaction; the atom gives away electrons, becoming more positive or less negative.

Reduction — Reduction is a decrease in oxidation number or gain of electrons.

This definition, like oxidation, is based on electron transfer or changes in oxidation states, providing a comprehensive understanding of the process. It always occurs simultaneously with oxidation. Think of reduction as 'gaining money' in a transaction; the atom accepts electrons, becoming more negative or less positive.

Redox reactions — Redox reactions are reactions in which both oxidation and reduction take place together.

These reactions are fundamental to many chemical and biological processes, such as photosynthesis and respiration. They involve the transfer of electrons from one species to another, resulting in changes in their oxidation numbers. Imagine a trade where one person gives an item (electrons) and another receives it; you can't have one without the other. Redox is a simultaneous 'give and take' of electrons.

Oxidation numbers — An oxidation number is a number given to each atom or ion in a compound that shows its degree of oxidation.

Oxidation numbers can be positive, negative, or zero, and are crucial for identifying oxidation and reduction in covalent compounds and balancing complex redox equations. A higher positive number indicates a more oxidised state. Consider oxidation numbers as a 'score' for how many electrons an atom has effectively 'lost' or 'gained' compared to its elemental state, helping to track electron shifts in reactions.

Half-equations — Half-equations are separate equations showing either oxidation or reduction in a redox reaction.

They represent the electron transfer for a single species, allowing for easier balancing of complex redox reactions. When combined, with electrons balanced, they form the overall ionic equation. Think of half-equations as two sides of a financial ledger; one side shows money leaving (oxidation) and the other shows money entering (reduction), and both must balance to show the complete transaction.

Oxidising agent — An oxidising agent (oxidant) is a substance which brings about oxidation by removing electrons from another atom or ion.

In doing so, the oxidising agent itself gets reduced, meaning its own oxidation number decreases. Common oxidising agents include oxygen and potassium manganate(VII). An oxidising agent is like a 'thief' that takes electrons from another substance, causing that substance to be oxidised, while the thief itself gains electrons and is reduced.

Reducing agent — A reducing agent (reductant) is a substance that brings about reduction by donating (giving) electrons to another atom or ion.

In doing so, the reducing agent itself gets oxidised, meaning its own oxidation number increases. Common reducing agents include hydrogen and reactive metals. A reducing agent is like a 'donor' that gives electrons to another substance, causing that substance to be reduced, while the donor itself loses electrons and is oxidised.

Disproportionation — Disproportionation is a type of redox reaction where some atoms of an element are reduced and some are oxidised in the same reaction.

This 'self-reduction/oxidation' reaction occurs when an element in an intermediate oxidation state reacts to form products with both higher and lower oxidation states. Chlorine reacting with alkali is a classic example. Imagine a person who both buys and sells the same stock; they are simultaneously acting as a buyer (reduction) and a seller (oxidation) of that stock.

sacrificial protection — Sacrificial protection is a method where a more reactive metal is oxidised in preference to another metal, preventing the latter from rusting.

The more reactive metal acts as a reducing agent, losing electrons and increasing its oxidation number, thereby protecting the less reactive metal. This is commonly used to protect iron structures like ship hulls. Think of a bodyguard taking a hit for someone else; the bodyguard (sacrificial metal) gets 'damaged' (oxidised) to protect the principal (iron).

reversible reaction — A reaction in which the products can react to re-form the original reactants.

In a reversible reaction, the forward and backward reactions occur simultaneously, leading to an equilibrium state. This continuous movement in both directions is indicated by the ⇌ symbol in chemical equations, meaning the reaction never stops completely.

dynamic equilibrium — A state where the rate of the forward reaction equals the rate of the backward reaction, and the concentrations of reactants and products remain constant.

At dynamic equilibrium, although concentrations are constant, molecules are continuously reacting, making the system dynamic. This state is analogous to people moving up and down an escalator at the same rate, where the number of people on each floor remains constant, but individuals are continuously in motion.

closed system — A system in which none of the reactants or products escapes from the reaction mixture.

For dynamic equilibrium to be established, it is essential that matter cannot be lost to the surroundings. This ensures that all species involved in the equilibrium remain within the system, allowing their concentrations to become constant. A sealed bottle of fizzy drink, where carbon dioxide gas is in equilibrium with dissolved carbon dioxide, serves as an example of a closed system.

Le Chatelier’s principle — If a change is made to a system at dynamic equilibrium, the position of equilibrium moves to minimise this change.

This principle is a powerful tool for predicting how an equilibrium system will respond to disturbances such as changes in temperature, concentration, or pressure. The system will shift to counteract the applied stress, much like a seesaw tilting to rebalance itself when weight is added to one side.

position of equilibrium — The relative amounts of products and reactants present in an equilibrium mixture.

The position of equilibrium indicates whether there is a greater concentration of products (shifted to the right) or reactants (shifted to the left) in the mixture. Changes in external conditions can cause this position to shift, altering the relative quantities of substances at equilibrium.

equilibrium constant, Kc — A constant value that relates the equilibrium concentrations of products and reactants in a specific way for a given reaction at a particular temperature.

Kc is calculated by dividing the product of the concentrations of products (each raised to its stoichiometric power) by the product of the concentrations of reactants (each raised to its stoichiometric power). Its value is unique for a given reaction at a specific temperature and is not affected by changes in concentration or pressure. Solids and liquids are excluded from the Kc expression because their concentrations are considered constant.

partial pressure — The pressure exerted by any one gas in a mixture of gases.

In a mixture of gases, each gas contributes to the total pressure in proportion to its mole fraction. The sum of partial pressures equals the total pressure, much like individual friends contributing their 'push' to the total effort of moving a heavy box.

mole fraction — The ratio of the number of moles of a specific gas to the total number of moles of all gases in a mixture.

This dimensionless quantity is crucial for determining the partial pressure of a gas in a mixture, as partial pressure is directly proportional to mole fraction. For example, if you have a bag of marbles, the mole fraction of red marbles is the number of red marbles divided by the total number of marbles.

equilibrium constant, Kp — A constant value that relates the equilibrium partial pressures of products and reactants for a given gas-phase reaction at a particular temperature.

Kp is calculated similarly to Kc, but uses partial pressures instead of concentrations. Like Kc, its value is only affected by temperature and remains constant even if the total pressure changes, provided the temperature is constant. Only gaseous species are included in the Kp expression.

acid — A substance that releases hydrogen ions (H+) when it dissolves in water, or a proton donor according to Brønsted–Lowry theory.

Acids typically have pH values below 7. They can be classified as strong (fully dissociated) or weak (partially dissociated) depending on their extent of ionisation in water. An acid acts like a generous giver, always ready to donate its proton (H+) to another molecule.

base — A substance that neutralises an acid, or a proton acceptor according to Brønsted–Lowry theory.

Bases typically contain oxide or hydroxide ions and react with acids to form a salt and water. A base is like a welcoming receiver, ready to accept a proton (H+) from an acid. Alkalis are a specific type of base that are soluble in water.

alkali — A base that is soluble in water and forms hydroxide ions (OH−) in solution.

Alkalis have pH values above 7 and can be strong or weak. All alkalis are bases, but not all bases are soluble in water. They are a subset of bases, much like 'waterproof' shoes are a subset of all shoes.

Brønsted–Lowry acid — A proton donor.

This definition is more general than the Arrhenius definition as it does not require the reaction to take place in aqueous solution. A proton is simply a hydrogen ion, H+. It's like a person giving away a small, important item (the proton) to someone else.

Brønsted–Lowry base — A proton acceptor.

This definition is also more general than the Arrhenius definition. A proton is a hydrogen ion, H+. Any species that can accept a proton is a Brønsted–Lowry base, not just those containing OH-. It's like a person receiving a small, important item (the proton) from someone else.

hydroxonium ion — The H3O+ ion, formed when a hydrogen ion (H+) combines with a water molecule.

In aqueous solution, hydrogen ions are highly reactive and associate with water molecules, forming H3O+. While often simplified to H+(aq), using H3O+(aq) demonstrates a more accurate understanding of the hydrated proton, like a lonely H+ ion finding a 'buddy' in a water molecule.

amphoteric — A substance that can act as either an acid or a base.

Water is a common example of an amphoteric substance, capable of acting as a base by accepting a proton from HCl or as an acid by donating a proton to NH3. It's like a 'chameleon' molecule that changes its role depending on the reaction partner.

strong acid — An acid that dissociates almost completely in aqueous solution.

This means nearly all acid molecules donate their protons to water, leading to a high concentration of H3O+ ions and a very low pH for a given concentration. Examples include HCl, H2SO4, and HNO3. A strong acid is like a very effective 'proton dispenser'.

weak acid — An acid that is only partially dissociated in aqueous solution.

Only a small fraction of weak acid molecules donate their protons to water, resulting in a low concentration of H3O+ ions and a higher pH compared to a strong acid of the same concentration. Most organic acids are weak acids. A weak acid is like a 'reluctant proton dispenser'.

strong base — A base that dissociates almost completely in aqueous solution.

This means nearly all base molecules or ions accept protons or release hydroxide ions, leading to a high concentration of OH− ions and a high pH for a given concentration. Group 1 metal hydroxides are strong bases. A strong base is like a very effective 'hydroxide dispenser'.

weak base — A base that dissociates to only a small extent in aqueous solution.

Only a small fraction of weak base molecules or ions accept protons or release hydroxide ions, resulting in a low concentration of OH− ions and a lower pH compared to a strong base of the same concentration. Ammonia is a common weak base. A weak base is like a 'reluctant hydroxide dispenser'.

neutralisation reaction — A reaction in which an acid reacts with a base (or alkali) to form a salt and water.

The simplest ionic equation for neutralisation is H+(aq) + OH−(aq) → H2O(l), representing the combination of hydrogen ions from the acid and hydroxide ions from the base. It's like two opposing forces cancelling each other out to form water and a salt.

salt — A compound formed when an acid reacts with a base (or alkali).

Salts are typically ionic compounds formed from the cation of a base and the anion of an acid. Neutralisation reactions always produce a salt and water. For example, mixing baking soda (base) and vinegar (acid) produces sodium acetate, which is a salt.

acid–base indicator — A dye or mixture of dyes that changes colour over a specific pH range.

Indicators are typically weak acids where the acid form and its conjugate base form have different colours. The observed colour depends on the relative concentrations of these forms, which in turn depends on the pH of the solution. It's like a 'pH traffic light' signalling a threshold crossing.

end-point — The point in a titration where the indicator changes colour, signalling the completion of the reaction.

For an indicator to be suitable, its end-point (the pH at which it changes colour) should coincide with the equivalence point of the titration, which is the theoretical stoichiometric point where the acid and base have exactly reacted. It's like the 'finish line' in a race.

rate of reaction — The rate of a reaction is defined as the change in concentration of reactant or product per unit time.

It is typically measured in mol dm⁻³ s⁻¹. The rate usually changes as the reaction proceeds because reactant concentrations decrease, and it can be determined experimentally by monitoring changes in concentration or physical properties over time. Imagine a queue at a checkout; the 'rate of reaction' is how quickly people are processed through it.

frequency of collisions — The frequency of collisions is the number of collisions per unit time, e.g. number of collisions per second.

According to collision theory, an increase in the frequency of collisions between reactant particles can lead to an increased rate of reaction, provided these collisions are effective. This is influenced by factors like concentration, pressure, and temperature. Think of bumper cars at an amusement park; the 'frequency of collisions' is how often the cars hit each other.

ineffective collision — An ineffective collision takes place if the colliding particles do not have enough energy to react.

In an ineffective collision, reactant particles simply bounce off each other without transforming into products. This occurs when the kinetic energy of the colliding particles is less than the activation energy required for the reaction. If two billiard balls just tap each other gently, they bounce off without significant change, much like an ineffective collision.

effective collision — If the reactant particles do have enough energy to react, they may change into product particles when they collide.

Also known as a successful collision, this type of collision leads to the formation of products. For a collision to be effective, particles must collide with energy equal to or greater than the activation energy and with the correct orientation. Consider two puzzle pieces; for them to 'react' and join, they must not only touch but also fit together perfectly and be pushed with enough force.

activation energy — The minimum energy that colliding particles must possess for a collision to be effective is called the activation energy, EA, for that particular reaction.

It represents an energy barrier that must be overcome for reactants to transform into products. A higher activation energy means fewer particles have enough energy to react, leading to a slower reaction rate. Think of pushing a ball up a hill; the 'activation energy' is the minimum energy needed to get it over the top.

Boltzmann distribution — The distribution of energies at a given temperature can be shown on a graph called the Boltzmann distribution.

This curve illustrates that in a sample of substance, particles possess a range of energies, with most having an intermediate amount. It is crucial for explaining the effect of temperature and catalysts on reaction rates by showing the proportion of particles with energy greater than the activation energy. Imagine a class of students taking a test; the 'Boltzmann distribution' is like a graph showing how many students got each possible score.

catalyst — A catalyst is a substance that increases the rate of a reaction but remains chemically unchanged itself at the end of the reaction.

Catalysts work by providing an alternative reaction pathway (mechanism) with a lower activation energy. This increases the proportion of reactant molecules that have sufficient energy to react, thereby increasing the frequency of effective collisions. Think of a shortcut on a journey; the 'catalyst' is like building a tunnel through a mountain, providing an easier path.

catalysis — Catalysis is the process by which a catalyst increases the rate of a reaction.

This process involves the catalyst interacting with reactants to facilitate a new reaction pathway with a lower activation energy, without being consumed itself. Catalysis can be homogeneous or heterogeneous, depending on the phases of the catalyst and reactants. If a coach helps a team improve their performance by teaching them a more efficient strategy, the 'catalysis' is the act of coaching.

homogeneous catalyst — When a catalyst and the reactants in a catalysed reaction are in the same phase, we describe the catalyst as a homogeneous catalyst.

For example, if both the catalyst and reactants are dissolved in water as an aqueous solution, it is a homogeneous catalyst. The reaction occurs uniformly throughout the mixture. Imagine adding a special ingredient to a soup that dissolves completely and helps other ingredients blend faster.

heterogeneous catalyst — If the catalyst is in a different phase to the reactants, we describe the catalyst as a heterogeneous catalyst.

Many heterogeneous catalysts are solids that catalyse gaseous or liquid reactants, with reactions typically occurring on the surface of the solid catalyst. Examples include iron in the Haber process or nickel in hydrogenation. Think of a non-stick pan (solid catalyst) used for cooking (liquid/solid reactants); they are in different phases.

Periodicity — Periodicity is the recurrence of the same pattern in properties across periods in the Periodic Table.

This term describes the repeating trends in physical and chemical properties of elements as their atomic number increases, particularly when moving across a period. These patterns are fundamental to understanding the organization and behavior of elements. Think of the days of the week repeating every seven days; Monday always follows Sunday, just as certain chemical properties repeat in a cycle across the Periodic Table.

atomic radius — The atomic radius is a measure of the size of an atom, often determined from its single covalent radius.

It is typically half the distance between the nuclei of two identical atoms joined by a single covalent bond. Across a period, atomic radius decreases due to increasing nuclear charge pulling outer electrons closer, while down a group it increases due to additional electron shells. Imagine a series of concentric circles, where the atomic radius is the distance from the center to the outermost circle. As you add more protons to the center (across a period), the pull on the outer circle gets stronger, making it shrink.

ionic radius — The ionic radius is a measure of the size of an ion.

Cations are smaller than their parent atoms because they lose outer electrons and have less shielding, while anions are larger than their parent atoms due to increased electron-electron repulsion in the same shell. Across a period, ionic radius generally decreases for both cations and anions due to increasing nuclear charge. Think of a balloon: if you let some air out (lose electrons to form a cation), it shrinks. If you blow more air in (gain electrons to form an anion), it expands.

electronegativity — Electronegativity is the power of an atom to attract the electrons in a covalent bond towards itself.

It increases across a period due to increasing nuclear charge and decreasing atomic radius, leading to a stronger attraction for bonding electrons. It decreases down a group as the outer electrons are further from the nucleus and experience more shielding. Imagine two people pulling on a rope (a covalent bond). The person with higher electronegativity is stronger and pulls the rope (bonding electrons) closer to themselves.

amphoteric — Compounds that can act as both acids and bases are called amphoteric.

Aluminium oxide and aluminium hydroxide are examples of amphoteric substances, reacting with acids to form salts and water (acting as a base) and reacting with hot, concentrated alkalis to form salts (acting as an acid). This dual behaviour provides evidence of intermediate bonding character. Imagine a person who can speak two different languages fluently, switching between them depending on who they are talking to. An amphoteric substance 'speaks' both acid and base 'languages' depending on the reactant.

metallic radius — The metallic radius is half the distance between the nuclei in a giant metallic lattice.

This measurement describes the size of atoms in metallic elements. As you move down Group 2, the metallic radius increases because additional electron shells are added, positioning the outermost electrons further from the nucleus, much like adding more nested Russian dolls increases the overall size.

lime — Lime is another name for calcium oxide, CaO.

Calcium oxide is formed by the thermal decomposition of calcium carbonate (limestone) in a lime kiln. It is a basic oxide used in cement manufacture and reacts vigorously with water, similar to how a cooked potato is different from a raw one, though both come from the same source.

quicklime — Quicklime is another name for calcium oxide, CaO.

It is produced by heating limestone (calcium carbonate) to high temperatures. Quicklime is a basic oxide that reacts exothermically with water to form calcium hydroxide, releasing a lot of heat and expanding rapidly, similar to 'quick-setting' cement.

slaked lime — Slaked lime is calcium hydroxide, Ca(OH)2.

It is formed when calcium oxide (quicklime) reacts with water. Slaked lime is a basic compound used by farmers to neutralise acidic soil, raising its pH, much like a thirst quenched by water makes quicklime a more stable, less reactive compound.

limewater — Limewater is a weakly alkaline solution of slightly soluble calcium hydroxide, Ca(OH)2.

It is formed when calcium reacts with water or when calcium oxide reacts with excess water. Limewater is commonly used to test for carbon dioxide, which turns it cloudy, much like dissolving a tiny bit of chalk in water would appear cloudy if too much is present.

halogens — The elements in Group 17 of the Periodic Table.

Halogens are non-metals with seven electrons in their outer principal quantum shell, existing as diatomic molecules (F2, Cl2, Br2, I2) at room temperature. They are highly reactive and tend to gain one electron to achieve a stable noble gas configuration, much like a group of seven friends looking for an eighth to complete their team.

volatility — The ease with which a substance evaporates.

For halogens, volatility decreases down the group as the number of electrons increases, leading to stronger instantaneous dipole–induced dipole forces between molecules. This requires more energy to overcome, making evaporation harder. This is similar to how quickly water evaporates compared to honey; water is more volatile because its molecules are less attracted to each other.

oxidising agents — Substances that accept electrons from another substance, causing the other substance to be oxidised.

Halogens are oxidising agents because their atoms readily gain one electron to achieve a stable outer shell. In doing so, their oxidation number is reduced from 0 to -1. An oxidising agent is like a 'borrower' of electrons; it takes electrons from another substance, making that substance 'lose' electrons (oxidise).

reducing agents — Substances that donate electrons to another substance, causing the other substance to be reduced.

Halide ions (Cl-, Br-, I-) act as reducing agents because they can donate their extra electron. Their power as reducing agents increases down Group 17 as the ionic radius increases and the outer electron is less strongly attracted to the nucleus. A reducing agent is like a 'giver' of electrons; it donates electrons to another substance, making that substance 'gain' electrons (reduce).

electronegativity — The ability of a covalently bonded atom to attract the bonding pair of electrons towards itself.

Electronegativity decreases down Group 17 because atomic radius increases, meaning the outer shell is further from the nucleus and experiences more shielding, thus having a weaker attraction for bonding electrons. This is like a stronger person in a tug-of-war pulling harder on a rope (the bonding electrons).

cyclohexane — A colourless liquid that is immiscible in water and can dissolve halogens.

Cyclohexane is used in displacement reactions to extract dissolved halogens from the aqueous layer, making their characteristic colours (pale green for Cl2, orange for Br2, purple for I2) more clearly visible. It acts like a 'colour enhancer' for halogens, pulling them out of water so their true colours can be seen without interference.

silver chloride — A white precipitate formed when silver ions react with chloride ions.

AgCl is insoluble in water but dissolves in dilute aqueous ammonia to form a soluble complex ion, [Ag(NH3)2]+. This is like a white cloud that disappears completely when a light mist (dilute ammonia) touches it, distinguishing it from other silver halides.

silver bromide — A cream precipitate formed when silver ions react with bromide ions.

AgBr is insoluble in water and dilute aqueous ammonia, but dissolves in concentrated aqueous ammonia to form a soluble complex ion, [Ag(NH3)2]+. This is like a cream-coloured cloud that only disappears when a strong downpour (concentrated ammonia) hits it, distinguishing it from AgCl and AgI.

silver iodide — A pale yellow precipitate formed when silver ions react with iodide ions.

AgI is insoluble in water, dilute aqueous ammonia, and concentrated aqueous ammonia, making it distinct from silver chloride and silver bromide. This is like a pale yellow cloud that remains visible no matter how much mist or rain (ammonia) falls on it.

hydrogen chloride — A gas with the formula HCl, formed from the reaction of hydrogen with chlorine or sodium chloride with concentrated sulfuric acid.

HCl is a thermally stable hydrogen halide, not decomposed by concentrated sulfuric acid. It is visible as white fumes and is denser than air. It's like the 'tough' hydrogen halide that resists breaking down easily, unlike its heavier counterparts.

disproportionation — A type of redox reaction where some atoms of an element are reduced and some are oxidised in the same reaction.

Chlorine undergoes disproportionation when it reacts with alkali or water, where its oxidation number changes from 0 to both -1 (reduction) and +1 or +5 (oxidation). The products depend on temperature. This is like a person who both gives and receives gifts at the same party; the element acts as both an oxidising and reducing agent simultaneously.

chloric(I) acid — The compound HClO, formed when chlorine undergoes disproportionation in water.

Chloric(I) acid is a sterilising agent used in water purification, killing bacteria. It can also dissociate to produce ClO-(aq) ions, which also act as sterilising agents. It's like a 'clean-up crew' that chlorine forms in water, actively destroying harmful microorganisms.

primary pollutants — Pollutants that are given off directly into the air from the source of pollution.

These are substances emitted directly from human activities or natural processes, such as nitrogen oxides from vehicle exhaust fumes or sulfur dioxide from power plants. They are the initial substances released into the atmosphere, like raw ingredients added directly to a recipe.

secondary pollutants — Pollutants that are formed by reactions between primary pollutants or other atmospheric components, rather than being given off directly from human activity.

These pollutants are not directly emitted but are created in the atmosphere through chemical reactions, often initiated by sunlight. Peroxyacetyl nitrate (PAN) in photochemical smog is a key example, formed from nitrogen oxides and volatile organic compounds. They are like a cooked dish, formed by combining and transforming raw ingredients through a chemical 'cooking' process.

photochemical smog — A type of air pollution formed when volatile organic compounds react in sunlight with oxides of nitrogen to make harmful secondary pollutants like peroxyacetyl nitrate (PAN).

This smog is characterized by a hazy appearance and is common in urban areas with high vehicle emissions. Sunlight provides the energy for the photochemical reactions that convert primary pollutants into more harmful secondary pollutants, affecting respiratory health and plant life. It's like a chemical 'soup' in the air where sunlight acts as the heat, stirring up reactions between car exhaust fumes to create a new, harmful mixture.

peroxyacetyl nitrate (PAN) — A harmful secondary pollutant with the chemical formula CH3CO3NO2, found in photochemical smog.

PAN is formed from the reaction of volatile organic compounds (VOCs) and nitrogen oxides (NOx) in the presence of sunlight. It is a strong eye and lung irritant and can also damage plant life, contributing significantly to the adverse effects of photochemical smog. PAN is like a toxic byproduct created when two different types of waste (VOCs and NOx) mix and react under specific conditions (sunlight).

eutrophication — The process where nitrates leached from fertilisers into rivers and lakes promote the growth of algae on the surface, leading to a reduction in dissolved oxygen and harm to aquatic ecosystems.

When excess nitrates enter water bodies, they act as nutrients for algae, causing rapid growth (algal bloom). When these algae die, bacteria decompose them, consuming large amounts of dissolved oxygen, which can kill fish and other aquatic life. Imagine overfeeding a fish tank; the excess food causes a rapid growth of bacteria and algae, which then consume all the oxygen, harming the fish.

hydrocarbons — Hydrocarbons are compounds of carbon and hydrogen only.

These compounds serve as the basic building blocks of organic molecules. Alkanes, for instance, are simple hydrocarbons that lack any functional group, forming the foundation for understanding more complex organic structures.

functional group — A functional group is a particular atom, or grouping of atoms, within a molecule that determines the characteristic chemical properties of the compounds that contain it.

Just as an engine dictates a car's behavior, a functional group dictates a molecule's reactivity. Different classes of organic compounds are defined by their specific functional groups, such as the C=C double bond in alkenes or the -COOH group in carboxylic acids, which also influence physical properties like boiling point and solubility.

homologous series — Homologous series are classes of related organic compounds where all compounds consist of molecules with a particular functional group.

Members of a homologous series share a common functional group, leading to similar chemical properties. They also possess the same general formula and exhibit a gradual change in physical properties as the carbon chain length increases, much like family members sharing traits but varying in age.

alkyl group — An alkyl group is a hydrocarbon side-chain named by adding '-yl' to the normal alkane stem.

These groups, such as methyl (-CH3) or ethyl (-C2H5), act as branches on the main carbon chain. They are often shown in brackets in structural formulae and play a role in naming branched organic compounds.

aryl groups — Aryl compounds contain at least one benzene ring.

A benzene ring is a stable, hexagonal structure of six carbon atoms, each typically bonded to one hydrogen. These groups are fundamental to aromatic chemistry, possessing unique stability due to delocalised pi electrons.

empirical formula — The empirical formula gives us the least detail, telling us the simplest ratio of the different types of atoms present in the molecule.

This formula is derived from experimental data, converting mass percentages of elements into a simplified whole-number ratio of moles. It's like a simplified recipe, showing the ratio of ingredients rather than the exact quantities.

molecular formula — The molecular formula shows us the actual numbers of each type of atom in a molecule.

To determine the molecular formula, the empirical formula and the relative molecular mass are required. It represents the true composition of one molecule, indicating the actual count of each atom.

structural formula — The structural formula tells us about the atoms bonded to each carbon atom in the molecule.

This representation offers more detail than the molecular formula by indicating how atoms are connected, particularly highlighting carbon-carbon double bonds. It's a simplified map of connectivity.

displayed formula — The displayed formula shows all the bonds within a molecule.

This 2D representation makes all covalent bonds visible, akin to a detailed blueprint. Every atom and every bond, whether single, double, or triple, is explicitly drawn.

skeletal formula — A simplified version of the displayed formula where all the symbols for carbon and hydrogen atoms are removed, as well as the carbon to hydrogen bonds, but carbon to carbon bonds are left in place.

This quick representation of complex structures includes all other non-carbon or hydrogen atoms and their bonds. It's like drawing only the main supports of a building, with other details implied.

sigma (σ) bonds — Sigma (σ) bonds are single covalent bonds where the pair of electrons is found in a region of space between the nuclei of the two atoms sharing the electrons.

These bonds are formed by the direct, head-on overlap of atomic orbitals. In many organic compounds, sp3 hybridised carbon atoms form four σ bonds, leading to a tetrahedral arrangement with bond angles of 109.5°.

pi (π) bonds — Pi (π) bonds are formed when two p orbitals overlap sideways, with the two lobes of the π bond lying above and below the plane of the atoms.

A C=C double bond consists of one σ bond and one π bond, while a C≡C triple bond comprises one σ bond and two π bonds. The presence of a π bond restricts rotation around the bond, which is crucial for geometrical isomerism.

structural isomers — Structural isomers are compounds with the same molecular formula but different structural formulae.

These isomers are like having the same set of LEGO bricks but building completely different structures. They can differ in the carbon skeleton (chain isomerism), the position of a functional group (position isomerism), or even the type of functional group present (functional group isomerism).

chain isomerism — Chain isomers differ in the structure of their carbon 'skeleton'.

This means the arrangement of carbon atoms in the main chain or branches is different, even though the molecular formula is the same. It's like having the same number of train cars but arranging them in a straight line versus having some cars on a side track.

position isomerism — In position isomerism, the location of the functional group varies in each isomer.

The carbon skeleton remains consistent, but the position of a substituent or functional group changes, leading to distinct compounds with different properties. This is akin to moving a bathroom from one floor to another in the same house layout.

functional group isomerism — In functional group isomerism, there are different functional groups present for the same molecular formula.

This type of isomerism results in compounds with very different chemical properties because they possess distinct reactive groups. For example, an alcohol and an ether can have the same molecular formula but vastly different chemistries.

stereoisomers — Stereoisomers are compounds whose molecules have the same atoms bonded to each other, but with different arrangements of the atoms in space.

These isomers share the same structural formula but differ in their 3D orientation. They are categorised into geometrical (cis/trans) isomerism and optical isomerism, much like a left-hand drive and right-hand drive version of the same car model.

geometrical (cis/trans) isomerism — Geometrical (cis/trans) isomerism arises in unsaturated organic compounds due to restricted rotation about a C=C double bond, or in substituted cyclic compounds due to limited rotation about C-C single bonds in the ring.

In cis isomers, identical groups are on the same side of the double bond or ring, while in trans isomers, they are on opposite sides. These spatial differences lead to distinct physical and sometimes chemical properties.

optical isomerism — Optical isomerism occurs if a molecule contains a carbon atom that is bonded to four different atoms or groups of atoms, forming two non-superimposable mirror images called enantiomers.

These non-superimposable mirror images, like your left and right hands, rotate the plane of polarised light by equal amounts but in opposite directions. The central carbon atom responsible for this is known as a chiral centre.

chiral centre — A chiral centre is a carbon atom bonded to four different atoms or groups of atoms.

The presence of a chiral centre is the prerequisite for optical isomerism. Identifying these centres involves systematically checking each carbon atom to ensure it has four unique substituents attached.

homolytic fission — Homolytic fission is a type of bond breaking where both atoms at each end of the bond leave with one electron from the pair that formed the covalent bond.

This process, often initiated by energy like UV light, produces highly reactive species known as free radicals, each possessing an unpaired electron. It's like two friends sharing a pizza and each taking exactly half when they split.

free radicals — Free radicals are species produced when a bond breaks homolytically, each having an unpaired electron and being very reactive.

These 'lone wolves' are highly unstable and participate in free-radical reactions, which typically involve initiation, propagation, and termination steps. It is crucial to represent the unpaired electron with a dot.

initiation step — The initiation step is the formation of free radicals to start a reaction off, requiring an input of energy to break a covalent bond.

This is the crucial first stage of a free-radical mechanism, where stable molecules are converted into the reactive free radicals that will drive the subsequent reactions. It's the spark that starts the fire.

propagation steps — Propagation steps are reactions in a mechanism that regenerate more free radicals, continuing a chain reaction.

In these steps, a free radical reacts with a stable molecule to form a new molecule and another free radical, sustaining the reaction in a domino effect. A radical is consumed and another is produced.

termination step — The termination step is the final step in a mechanism, when two free radicals meet and form a product molecule, with no free radicals generated.

This step removes free radicals from the reaction mixture, effectively bringing the chain reaction to an end. It involves two free radicals combining to form a stable molecule.

heterolytic fission — Heterolytic fission is the 'uneven' breaking of a covalent bond, where the more electronegative atom takes both electrons in the covalent bond.

This process results in the formation of ions: a positively charged species, such as a carbocation, and a negatively charged species, like an anion. It's like one friend taking the entire pizza, leaving the other with nothing.

carbocation — A carbocation is a positively charged alkyl ion where the carbon atom with the positive charge only has three covalent bonds, making it electron deficient.

These electron-deficient intermediates are crucial in many organic reaction mechanisms. Their stability increases with the number of electron-donating alkyl groups attached to the positively charged carbon (tertiary > secondary > primary).

electrophile — An electrophile is an electron-deficient species that accepts a pair of electrons, resulting in the formation of a new covalent bond.

Electrophiles are 'electron seekers', typically positively charged ions or molecules with a partial positive charge. They are attacked by electron-rich species, such as nucleophiles.

nucleophile — A nucleophile is an electron-rich species that donates a pair of electrons, leading to the formation of a new covalent bond with an electron-deficient atom.

Nucleophiles are 'lovers of nuclei', typically carrying a negative or partial negative charge and possessing lone pairs of electrons available for donation. They attack electron-deficient centres.

addition reactions — Addition reactions involve the formation of a single product from two or more reactant molecules.

These reactions typically occur across double or triple bonds, where the multiple bond breaks, and new single bonds are formed with the adding species. It's like two puzzle pieces fitting together to form one larger piece.

elimination reactions — Elimination reactions result in the removal of a small molecule from a larger reactant molecule.

This process typically leads to the formation of a double or triple bond. For example, the dehydration of an alcohol to an alkene involves the elimination of a water molecule.

substitution reactions — Substitution reactions involve the replacement of one atom, or a group of atoms, by another.

In these reactions, a part of the molecule is exchanged for a different atom or group. A common example is the free-radical substitution of alkanes with halogens.

hydrolysis — Hydrolysis is the breakdown of a molecule by water.

This reaction involves water acting as a reactant to cleave a bond in another molecule, often resulting in the formation of an alcohol and an acid or a salt. It can be accelerated by alkali.

condensation reactions — Condensation reactions involve a first step where addition takes place, followed by a second step where elimination occurs to form the final product.

These reactions typically result in the formation of a larger molecule with the simultaneous elimination of a small molecule, such as water or HCl. They are crucial in forming polymers.

oxidation — Oxidation is defined as the loss of electrons from a species, but in organic reactions, it is often simpler to think of it in terms of the number of oxygen and/or hydrogen atoms before and after a reaction.

In organic chemistry, oxidation often involves gaining oxygen atoms or losing hydrogen atoms. For example, an alcohol can be oxidised to an aldehyde or a carboxylic acid.

reduction — Reduction is defined as the gain of electrons by a species, but in organic reactions, it is often simpler to think of it in terms of the number of oxygen and/or hydrogen atoms before and after a reaction.

Conversely, reduction in organic chemistry often involves gaining hydrogen atoms or losing oxygen atoms. For instance, a ketone can be reduced to a secondary alcohol.

Hydrocarbons — Hydrocarbons are compounds containing carbon and hydrogen only.

Hydrocarbons are the fundamental building blocks of many organic compounds, forming the main components of crude oil. They serve as fuels like petrol and diesel, and as starting materials for plastics. Their structure can be either saturated, like alkanes, or unsaturated, like alkenes.

Saturated hydrocarbons — Saturated hydrocarbons are hydrocarbons that have the maximum number of hydrogen atoms in their molecules, containing only single covalent bonds.

These hydrocarbons, such as alkanes and cycloalkanes, have only carbon-carbon single bonds, meaning all carbon atoms exhibit sp3 hybridisation. This 'saturation' implies they cannot hold any more hydrogen atoms, similar to a sponge saturated with water.

Cycloalkanes — Cycloalkanes are saturated hydrocarbons in which there is a 'ring' consisting of three or more carbon atoms.

Unlike straight-chain alkanes, cycloalkanes form a ring structure, which means two hydrogen atoms are lost compared to their open-chain counterparts. Despite having the general formula CnH2n, similar to alkenes, they are saturated because all carbon-carbon bonds are single bonds.

Homolytic fission — Homolytic fission of a covalent bond is when each atom takes one electron from the pair of electrons in the bond as it breaks.

This process is like two friends sharing a toy and each taking exactly one half when they split up. It typically requires energy, such as ultraviolet light, and results in the formation of two highly reactive species called free radicals, each possessing an unpaired electron.

Free radicals — Free radicals are atoms or groups of atoms, each with an unpaired electron.

These species are highly reactive, much like a person with one hand free in a crowded room eager to grab onto something. They are formed during homolytic fission, for example, when a Cl-Cl bond breaks under ultraviolet light, and are crucial intermediates in chain reactions like free-radical substitution.

Free-radical substitution — Free-radical substitution is an overall reaction between alkanes and halogens, involving initiation, propagation and termination steps.

This reaction occurs in the presence of ultraviolet light, where a hydrogen atom in an alkane is replaced by a halogen atom. It proceeds via a chain reaction involving highly reactive free radicals, akin to a game of 'tag' where the 'tagger' keeps creating new 'taggers'.

Cracking — Cracking is the process where large, less useful hydrocarbon molecules are broken down into smaller, more useful molecules.

This process is like chopping a long log into smaller pieces of firewood and kindling. It is carried out at high temperatures over an aluminium oxide catalyst in the absence of oxygen, producing smaller alkanes (useful fuels) and alkenes (valuable feedstock for the chemical industry).

Unsaturated hydrocarbons — Unsaturated hydrocarbons are hydrocarbons that contain one or more carbon-carbon double or triple bonds.

Unlike saturated hydrocarbons, these molecules have 'room' for more hydrogen atoms because their double or triple bonds can be broken to form single bonds. Alkenes, with their C=C double bond, are a key example, and their reactivity stems from this area of high electron density.

Electrophile — An electrophile is an acceptor of a pair of electrons.

Electrophiles are 'electron-love-seekers' that are attracted to areas of high electron density, such as the π bond in alkenes. They are typically positively charged ions or molecules with a partial positive charge, or molecules that can be induced to have one, seeking to form a new covalent bond by accepting an electron pair.

Heterolytic fission — Heterolytic fission is when a covalent bond breaks and one atom takes both electrons from the shared pair.

This process is like one friend selfishly taking the entire toy when two friends split up. It results in the formation of a positively charged ion (carbocation) and a negatively charged ion, and is characteristic of electrophilic addition reactions in alkenes.

Hydrogenation — Hydrogenation is the addition reaction of alkenes with hydrogen, producing an alkane.

This reaction is like adding extra seats to a car, where hydrogen atoms are added across the double bond of an alkene, turning it into a saturated alkane. It requires heating the alkene and hydrogen gas over a finely divided platinum or nickel catalyst, and is used in the manufacture of margarine.