Full set of Chemistry CGP summer set

Neutron

A subatomic particle found in the nucleus of an atom, with no charge.

Ion

A charged particle formed when an atom loses or gains electrons.

Proton

A subatomic particle found in the nucleus of an atom, with a positive charge.

Electron

A subatomic particle with a negative charge that orbits the nucleus of an atom.

Atomic Number

The number of protons in the nucleus of an atom, denoted by Z.

Mass Number

The total number of protons and neutrons in the nucleus of an atom, denoted by A.

Isotope

Atoms of the same element with the same number of protons but different numbers of neutrons.

Relative Atomic Mass

The average mass of an element considering the abundance of its isotopes, denoted by Ar.

Relative Formula Mass

The sum of the relative atomic masses of all atoms in a chemical formula, denoted by Mr.

Energy Shells

The regions around the nucleus of an atom where electrons are arranged in different levels or energy states.

Subshells

Subdivisions of energy levels within an atom, denoted by the letters s, p, d, f, etc., each with a specific maximum number of electrons it can hold.

Electron Configuration

The distribution of electrons in the energy levels and subshells of an atom, represented by numbers and letters indicating the energy level, subshell type, and number of electrons in that subshell.

Electron Arrangement

The specific organization of electrons in the energy levels and subshells of an atom, often depicted using diagrams or simple notation to show the number of electrons in each level or subshell.

Periodic Table

A tabular arrangement of the chemical elements, organized by atomic number, electron configuration, and recurring chemical properties.

S-block

The section of the periodic table consisting of groups 1 and 2, where elements have their outermost electrons in an s subshell.

P-block

The section of the periodic table consisting of groups 3 to 0, where elements have their outermost electrons in a p subshell.

Atomic Number

The number of protons in the nucleus of an atom, determining the element's identity in the periodic table.

Group

Vertical columns in the periodic table where elements share similar chemical properties due to having the same number of electrons in their outer shell.

Period

Horizontal rows in the periodic table where properties of elements change gradually across the row.

Electron Configuration

The distribution of electrons of an atom or molecule in atomic or molecular orbitals.

Full Outer Shell

A stable electron configuration achieved by some elements when they gain or lose electrons to have a complete outer shell.

Ionisation Energy

The energy required to remove an electron from an atom or ion in the gaseous state.

Sodium Ion

An ion of sodium with a 1+ charge formed by losing one electron.

Three Factors Affecting Ionisation Energy

Nuclear charge, distance from the nucleus, and shielding by inner electrons affect the ease of removing an electron.

Periodic Table Trends in Ionisation Energies

Ionisation energy decreases down a group and generally increases across a period due to changes in nuclear charge, distance from the nucleus, and shielding.

Simple Ions

Elements in the s-block and p-block form ions with full outer electron shells, leading to predictable charges based on group number.

Transition Metals

Metals between Groups 2 and 3 that form ions with varying charges, known as oxidation numbers, unlike s-block metals.

Oxidation Number

The oxidation number of an atom tells you how many electrons the atom has donated or accepted when it has reacted, also known as the charge on an atom.

Roman Numerals

Roman numerals are used to show the oxidation number of certain elements, where (I) = +1, (II) = +2, (III) = +3, and so on.

Intermolecular Bonds

Weak forces of attraction that form between molecules, distinct from the strong bonds within a molecule.

Electronegativity

The ability of an atom to attract electrons in a covalent bond, with fluorine being the most electronegative element.

Ionic Bonding

The type of bond that forms when electrons are transferred from one atom to another, resulting in the attraction between oppositely charged ions.

Covalent Bonding

The bond that forms when atoms share pairs of electrons to achieve a full outer shell, resulting in the formation of a molecule.

Ions

Charged particles that form when atoms lose or gain electrons.

Electron Transfer

The process where electrons are moved from one atom to another to form ions.

Dot-and-Cross Diagram

A diagram that represents the transfer of electrons between atoms using dots for electrons from one atom and crosses for electrons from another.

Ionic Compound

A compound formed by the attraction between positively and negatively charged ions.

Ionic Formula

The representation of an ionic compound showing the ratio of positive to negative ions.

Melting Point

The temperature at which a solid substance changes into a liquid.

Electrical Conductivity

The ability of a substance to conduct electricity.

Solubility

The ability of a substance to dissolve in a solvent, usually water.

Giant Ionic Structures

Large, closely packed regular arrays of ions formed by ionic compounds.

Covalent Molecule

A molecule formed by atoms sharing electrons in covalent bonds.

Multiple Covalent Bond

A bond formed when two atoms share more than one pair of electrons.

Dative Covalent Bond

A covalent bond where both shared electrons come from the same atom.

Hydrogen Fluoride (HF)

A small covalent molecule where hydrogen is bonded to fluorine. It has a dot-and-cross diagram showing the bonding electrons and lone pairs of electrons.

Nitrogen Gas

Despite nitrogen atoms being strongly bonded to each other in each molecule, nitrogen is a gas at room temperature due to the weak intermolecular bonds between the molecules.

Lone Pairs

Electron pairs in a molecule that are not shared between atoms, affecting the physical properties of covalent molecules.

Hydrogen Bonds

The strongest type of intermolecular bond formed between covalent molecules with lone pairs on nitrogen, fluorine, or oxygen atoms bonded to hydrogen(s), leading to high boiling and melting points and increased solubility.

Period 3 oxides

Compounds formed by elements in Period 3 of the periodic table, showing a transition in bonding types from ionic to covalent across the period.

Ionic bonding

A type of chemical bonding that involves the transfer of electrons from one atom to another, resulting in the formation of ions held together by electrostatic forces.

Covalent bonding

A type of chemical bonding where atoms share pairs of electrons to achieve a stable electron configuration, forming molecules.

Melting point

The temperature at which a solid substance changes into a liquid state at standard pressure.

Chemical equation

A symbolic representation of a chemical reaction showing the reactants on the left side and the products on the right side.

Balancing equations

The process of ensuring that the number of atoms of each element is the same on both sides of a chemical equation by adjusting coefficients.

State symbols

Symbols used in chemical equations to indicate the physical state of a substance (solid, liquid, gas, or aqueous).

Symbol equation

A type of chemical equation that includes the chemical formulae of the reactants and products, showing the atoms involved in the reaction.

Ionic equation

A type of chemical equation that represents the species that are actually involved in a reaction in solution, focusing on ions.

Charge balance

Ensuring that the total charge on the reactant side of an ionic equation is equal to the total charge on the product side.

Metallic Bonds

Bonds that hold metals together in a lattice structure, formed by the attraction between positive ions and free electrons.

Electron Shielding

The phenomenon where inner electrons shield outer electrons from the attraction of the positive nucleus, leading to weaker metallic bonds.

Group 2 Metals

Metals in Group 2 of the periodic table that lose two electrons during reactions, becoming more reactive as you go down the group.

Reactivity

The tendency of an element to undergo chemical reactions, increasing as you go down Group 2 due to easier donation of outer electrons.

Ionization Energy

The energy required to remove an electron from an atom, affected by factors like nuclear charge, electron shielding, and distance from the nucleus.

Potassium hydroxide

KOH

Sodium hydroxide

NaOH

Acids

Substances with a pH less than 7.

Bases

Substances with a pH greater than 7.

pH Scale

Measures how acidic or basic a substance is, ranging from 0 to 14.

Neutralization reaction

Reaction between an acid and a base to form a salt and water.

Proton donors

Acids release hydrogen ions (H+) when mixed with water.

Proton acceptors

Bases accept hydrogen ions (H+).

Skeletal formula

Representation of a molecule showing the carbon skeleton and functional groups.

Displayed formula

Shows all atoms and bonds in a molecule.

Molecular formula

Indicates the number of atoms of each element in a molecule.

Homologous series

Groups of organic compounds with similar properties and functional groups.

Alkanes

Hydrocarbons with only single bonds between carbon atoms.

Saturated molecules

Molecules like alkanes where all available bonds are formed.

Complete combustion

Reaction of a fuel with oxygen to produce carbon dioxide and water.

Methane

CH4

Alkenes

Hydrocarbons containing a carbon-carbon double covalent bond (C=C) somewhere in the carbon chain, making them unsaturated molecules.

Isomers

Molecules with the same molecular formula but different structures.

Polymers

Long, chain-like molecules built up from lots of repeating units (monomers), formed by linking together small alkenes under the right conditions.

Hydrogen Bonds

Strong intermolecular bonds formed between the slightly positive hydrogen atom in the -OH group of alcohols and lone pairs of electrons on oxygen atoms in other alcohol molecules.

Disproportionation

A rare type of chemical reaction where an element in a reactant is both oxidized and reduced simultaneously.

Displacement

A reaction where one element displaces another, less reactive element from a compound.

Dehydration

The removal of water from a compound by heating, often resulting in the formation of a C=C bond in organic molecules.

Cracking

The thermal decomposition of long-chain hydrocarbon molecules into shorter-chain alkanes and alkenes, requiring high temperatures, pressures, and a catalyst.

Condensation

A reaction where atoms are added to an unsaturated bond to make it saturated, often forming a simple molecule like water.

Combustion

The chemical reaction between a fuel and oxygen, producing carbon dioxide and water in complete combustion.

Precipitation

The formation of a solid (precipitate) in a solution due to a chemical reaction or change in temperature affecting solubility.

Oxidation

The loss of electrons or gain of oxygen in a chemical reaction, opposite of reduction.

Hydrogenation

The addition of hydrogen across a C=C bond, resulting in the saturation of the bond.

Neutralisation

The reaction between a basic compound and an acid, producing a salt, water, and other products.

Exothermic

Any chemical reaction that releases heat energy, with products having less energy than the reactants.

Elimination

The removal of a small molecule from a larger molecule, often involving the removal of H2O or H2.

Electrolysis

A process using electricity to break down a compound, requiring the reactants to be in a liquid state.

Endothermic

Any chemical reaction that absorbs heat energy, with products having more energy than the reactants.

Substitution

A reaction where an atom or group of atoms in a molecule is replaced by a different atom or group of atoms.

Reduction

The gain of electrons or loss of oxygen in a chemical reaction, always occurring together with oxidation.