Chemistry: Definitions AS

Atomic number

The number of protons in the nucleus of of an atom

Bohr model

Describes an atom as a small dense nucleus with electrons orbiting around the nucleus

This model explains different periodic properties of atoms

Electron

A negatively charged subatomic particle which orbits the nucleus at various energy levels

The relative mass of an electron is 1/1836

Ion

A charged atom or molecule

Isotopes

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

Isotopes of an element have different masses

Mass number

The total number of protons and neutrons in the nucleus of an atom

Mass spectrometry

An instrument which gives accurate information about relative isotopic mass and the relative abundance of isotopes

Neutron

A neutral subatomic particle found in the nucleus of an atom

The relative mass of a neutron is 1

Proton

A positively charged subatomic particle found in the nucleus of the atom

The relative mass of a proton

Relative abundance

The amount of one substance compared with another

Relative atomic mass

The weighted mean mass of an atom compared with 1/12th mass of an atom of carbon-12

Relative isotopic mass

The mass of an atom of an isotope compared with 1/12th mass of an atom of carbon-12

Relative formula mass

The mass of the formula unit of a compound with a giant structure

E.g: NaCl has a relative formula mass of 58.44g

Relative molecular mass (Mr)

The mass of a simple molecule

Ammonium ion

Ion with the formula NH4^+

Carbonate

Ion with the formula CO3^ 2+

Hydroxide

An ion with the formula OH^-

Ionic compound

A compound made up of oppositely charged ions hel

Nitrate

An ion with the formula NO3^-

Silver ion

Has the formula Ag^+

Sulfate

Ion with the formula SO4^ 2-

Zinc ion

Has the formula Zn^ 2+

Amount of substance

The quantity that has moles as its units, used as a way of counting atoms

The amount of substance can be calculated using mass (n=m/M), gas volumes (n=pV/(RT)) or solution volume and concentration (n=CV)

Anhydrous

Crystalline compound containing no water

Atom economy

Measure of the amount of starting materials that end up as useful products

A high atom economy means a process is more sustainable as there is less waste produced

Atom economy equation

Percentage atom economy = (molecular mass of desired product/sum of molecular masses of all reactants) x100

Avogadro constant (NA)

The number of particles per mole of substance (6.02×10²³ mol^-1)

Composition by mass

The relative mass of each element in a compound

Empirical formula

The simplest whole number ratio of atoms of each element present in a compound

Hydrated

Crystalline compound containing water

Ideal gas

A gas which has molecules that occupy negligible space with no interactions between them

Ideal gas equation

pV=nRT

Molar gas volume

The volume of 1 mole of gas (dm³ mol^-1)

Mole (mol)

The amount of any substance containing as many particles as there are carbon atoms in exactly 12g of carbon-12 isotope

Molecular formula

The number and type of atoms of each element in a molecule

Percentage yield

The percentage ratio of the actual yield of product from a reaction compared with the theoretical yield

Percentage yield equation

Percentage yield = (actual yield/theoretical yield) x100

Relative molecular mass

The average mass of one molecule of an element or compound compared to 1/12th the mass of an atom of carbon-12

Stoichiometry

The relative quantities of substances in a reaction

Water of crystallisation

Water molecules that form part of the crystalline structure of a compound

Acid

Compounds that release H^+ ions in aqueous solution

Alkali

Water soluble bases

Alkalis release OH^- ions into aqueous solution

Base

Substance that can accept H^+ ions from another substance

Neutralisation

A reaction between H^+ and OH^- forming water

This could be a reaction between an acid and a base to form a salt (types of bases include carbonates, metal oxides and alkalis)

Strong acid

An acid that completely dissociates in solution

Weak acid

An acid that only partially dissociates in solution

Titration

Technique used to determine the amount of one solution of a known concentration required to completely react with a known volume of another solution of unknown concentration

Oxidation

Loss of electrons/increase in oxidation number

Oxidation number

Number that represents the number of electrons lost or gained by an atom of an element

Positive oxidation number indicates the loss of electrons

Roman numerals are used to indicate the oxidation number of elements that may have different oxidation states (e.g: iron (II) and iron (III))

Redox reaction

Reaction in which one element is oxidised and another is reduced

Reduction

The gain of electrons/decrease on oxidation number

Atomic orbital

Region of space around the nucleus holding up to 2 electrons with opposite spins

Orbitals are filled in order of increasing energy, with orbitals of the same energy occupied singularly before pairing

(1 orbital in the s sub shell, 3 orbitals in the p sub shell etc…)

Electron configuration

The arrangement of electrons into orbitals and energy levels around the nucleus of an atom/ion

Energy level

The shell that an electron is in

Shell

The orbit that an orbital is in around the nucleus of an atom

The shell closest to the nucleus is the first shell and the outermost shell occupied by electrons is the valence shell

Sub shell

Subdivision of the electronic shells into different orbitals

Types of sub shell are s, p, d and f etc…

Average bond enthalpy

Average energy required to break a bond used as a measurement of the strength of a covalent bond

Average bond enthalpy is measured using a variety of molecules that contain a specific bond

Bonding pair

Pair of outer shell electrons involved in bonding

Covalent bond

A strong bond formed between 2 atoms due to the electrostatic attraction between a shared pair of electrons and the atomic nuclei

Dative covalent (coordinate) bond

A type of covalent bond in which both of the electrons in the shared pair come from one atom

Electronegativity

The ability of an atom to attract bonding electrons in a covalent bond

Often quantified using Pauling’s electronegativity values

Electronegativity increases towards F in the periodic table

Electron pair repulsion theory

Pairs of electrons around a nucleus repel each other so the shape that a molecule adopts has these pairs of electrons positioned as far apart as possible

Lone pairs offer more repulsion than bonding pairs as they are closer to the nucleus of the central atom

Hydrogen bonding

Type of intermolecular bonding that occurs between molecules containing N, O or F and a H atom (-NH, -OH, HF)

A lone pair on the electronegative atom (N, O or F) allows the formation of a hydrogen bond

Intermolecular forces

Interactions between different molecules

Types include permanent dipole-dipole interactions and induced dipole-dipole interactions as well as hydrogen bonds

Ionic bond

Electrostatic attraction between positive and negative ions

Ionic compounds

Compounds made of oppositely charged ions

High melting and boiling points

Soluble and can conduct electricity in the liquid or aqueous state

Ionic lattice

Giant structure in which oppositely charged ions are strongly attracted in all directions

Linear

Shape of a molecule in which the central atom has 2 bonding pairs

London (dispersion) forces

Induced dipole-dipole interactions caused when the random movement of electrons creates a temporary dipole in one molecule which then induces a dipole in a neighbouring molecule

Lone pair

Pair of outer shell electrons not involved in bonding

Macroscopic properties

Properties of a bulk material rather than the individual atoms/molecules that make up the material

Non-linear

Shape of a molecule in which the central atom has 2 bonding pairs and 2 lone pairs

Octahedral

Shape of a molecule in which the central atom has 6 bonding pairs

Permanent dipole

Permanent uneven distribution of charge

Polar bond

Covalent bond that has a permanent dipole due to the different electronegativities of the atoms that make up the bond

Polar molecule

Molecule that contains polar bonds with dipoles that don’t cancel out due to their direction (must be unsymmetrical)

Pyramidal

Shape of a molecule in which the central atom has 3 bonding pairs and 1 lone pair

Simple molecular lattice

Solid structure made up of covalently bonded molecules attracted by intermolecular force

Relatively low melting and boiling points

Insoluble in water but soluble in inorganic solvents

Don’t conduct electricity

Tetrahedral

Shape of a molecule in which the central atom has 4 bonding pairs

Trigonal bipyramidial

Shape of a molecule in which the central atom has 5 bonding pairs

Trigonal planar

Shape of a molecule in which the central atom has 3 bonding pairs

Atomic (proton) number

The number of protons in the nucleus of an atom

Bohr model

Describes an atom as a small dense nucleus with electrons orbiting around the nucleus

This explains the different periodic properties of atoms

Cations

Positively charged ions

D-block

The part of the periodic table in which the elements have their highest energy electron in and-orbital

Electron configuration

The arrangement of electrons into orbitals and energy levels around the nucleus of an atom/ion

First ionisation energy

The removal of one mole of electrons from one mole of gaseous atoms

Factors affecting it:

• Nuclear attraction

• Nuclear charge

• Atomic radius

Small decrease in first ionisation energy due to s- and p- sub shell energies (between Be and B) and p-orbital repulsion (between N and O)

Giant covalent lattice

Network of atoms bonded by strong covalent bonds (e.g: carbon (diamond, graphite and graphene) and silicon)

Insoluble with high melting and boiling points due to the presence of strong covalent bonds

Poor electrical charge conductors because of the lack of mobile charge carriers (apart from graphene and graphite)

Giant metallic lattice structure

The structure of all metals made up of cations and delocalised electrons

Insoluble with high melting and boiling points due to the strong electrostatic attraction between the cations and electrons

Good electrical conductors because of the presence of delocalised electrons which can carry charge

Group

Column in the periodic table

Melting point

The temperature at which a solid melts to become a liquid

Increases from giant metallic to giant covalent structures then decreases to simple molecular structures

Metallic bonding

Strong electrostatic attraction between cations and delocalised electrons

P-block

Part of the periodic table in which the elements have their highest energy electron in a p-orbital

Period

Row in the periodic table

Periodicity

Repeating trend in physical and chemical properties across the periods of the periodic table

S-block

Part of the periodic table in which the elements have their highest energy electron in the s-orbital

Successive ionisation energy

Energy required to remove each electron one-by-one from one mole of gaseous atoms/ions

Base (group 2)

Substance that can accept H^+ ions from another substance

Group 2 compounds can be used as bases

Ca(OH)2 is used to neutralise acidic soils in agriculture

Mg(OH)2 and CaCO3 are used as antacids to treat indigestion

Electron configuration (group 2)

The arrangement of electrons into orbitals and energy levels around the nucleus of an atom/ion

Group 2 elements have an s² outer shell electron configuration

First ionisation energy

Removal of one mole of electrons from one mole of gaseous atoms/ions

Factors affecting:

• Nuclear attraction

• Nuclear charge

• Atomic radius