Chemistry Definitions

Dalton's Atomic Theory

1) All matter is made up of very small particles called atoms. 2) All atoms are indivisible. They cannot be broken down into simpler particles

Cathode rays

Streams of negatively charged particles called electrons. They travel in straight lines from the cathode to the anode, are deflected by electric and magnetic fields, and have sufficient energy to move a small object such as a paddle wheel.

Energy Level

The fixed energy value that an electron in an atom may have.

Ground state

Electrons occupy the lowest available energy level in an atom.

Excited state

Electrons occupy higher energy levels than those available in the ground state in an atom.

Equation for light of definite frequency

E2 - E1 = hf

Heisenberg's Uncertainty Principle

It is impossible to measure at the same time both the velocity and the position of an electron.

Orbital

A region in space within which there is a high probability of finding an electron.

Sublevel

A subdivision of a main energy level and consists of one or more orbitals of the same energy.

Element

A substance that cannot be split into simpler substances by chemical means.

Triad

A group of three elements with similar chemical properties in which the atomic weight (relative atomic mass) of the middle element is approximately equal to the average of the other two.

Newlands' Octaves

Arrangements of elements in which the first and the eighth element, counting from a particular element, have similar properties.

Mendeleev's Periodic Law

When elements are arranged in order of increasing atomic weight (relative atomic mass) the properties of the elements recur periodically, i.e. the properties displayed by an element are repeated at regular intervals in other elements.

Atomic Number

The number of protons in the nucleus of an atom.

Modern Periodic Table

Arrangement of elements in order of increasing atomic number.

Modern Periodic Law

When elements are arranged in order of increasing atomic number, the properties of the elements recur periodically, i.e. the properties displayed by an element are repeated at regular intervals in other elements.

Mass number

The sum of the number of protons and neutrons in the nucleus of an atom of that element.

Isotopes

Atoms of the same element (i.e. they have the same atomic number) which have different mass numbers due to the different number of neutrons in the nucleus.

Relative atomic mass

The average of the mass of the isotopes of that element as they occur naturally. Taking their abundances into account and expressed on a scale in which the atoms of the carbon-12 isotope have a mass of exactly 12 units.

Principle of Mass Spectrometry

Charged particles moving in a magnetic field are deflected to different extents according to their masses and are thus separated according to these masses.

Electron configuration

The arrangement of electrons in an atom of an element.

Aufbau's Principle

When building up the electron configuration of an atom in its ground state the electrons occupy the lowest available energy levels.

Hund's Rule of Maximum Multiplicity

When two or more orbitals of equal energy are available the electrons occupy them singly before filling them in pairs.

Pauli Exclusion Principle

No more than two electrons may occupy an orbital and they must have opposite spin.

Compound

A substance that is made up of two or more different elements combined together chemically.

Octet Rule

When bonding occurs, atoms tend to reach an electron arrangement with eight electrons in the outermost energy level.

Ion

A charged atom or group of atoms

Ionic bond

The force of attraction between oppositely charged ions in a compound. Ionic bonds are always formed by the complete transfer of electrons from one atom to another.

Transition metal

One that forms at least one ion with a partially filled sublevel.

Molecule

A group of atoms joined together. It is the smallest particle of an element or compound that can exist independently.

Valency

The number of atoms of hydrogen or any other monovalent element with which each atom of the element combines.

Sigma bond

The head-on overlap of two orbitals.

Pi Bond

The sideways overlap of p orbitals.

Electronegativity

The relative attraction that an atom in a molecule has for the shared pair of electrons in a covalent bond.

Polar covalent bond

A bond in which there is unequal sharing of the pair (or pairs) of electrons. This causes one end of the bond to be slightly positive and the other end to be slightly negative.

Electronegativity difference for ionic bonding

Greater than 1.7

Electronegativity difference for covalent bonding

Less than or equal to 1.7

Electronegativity difference for polar covalent bonding

Greater than 0.4, Less than or equal to 1.7

Electronegativity difference for non polar bonding

Less than or equal to 0.4

Intramolecular bonding

Bonding that takes place within a molecule, i.e. it holds the atoms together. Covalent bonding and polar covalent bonding are examples of intramolecular bonding.

Intermolecular forces

Forces of attraction that exist between molecules. Van der Waals forces, dipole-dipole forces and hydrogen bonding are examples of intermolecular forces.

Van der Waals forces

Weak attractive forces between molecules resulting from the formation of temporary dipoles. They are the only forces of attraction between non-polar molecules.

Dipole-dipole forces

Forces of attraction between the negative pole of one polar molecule and the positive pole of another polar molecule.

Hydrogen Bonding

Particular types of dipole-dipole attractions between molecules in which hydrogen atoms are bonded to nitrogen,, oxygen or fluorine. The hydrogen atom carries a partial positive charge and is attracted to the electronegative atom in another molecule. Thus, the hydrogen bond acts as a bridge between two electronegative atoms in separate molecules.

Law of Conservation of Mass

The total mass of the products of a chemical reaction is the same as the total mass of the reactants.

Law of Conservation of Matter

Any chemical reaction, matter is neither created nor destroyed but merely changes from one form into another.

Atomic radius (covalent radius)

Half the distance between the nuclei of two atoms of the same element that are joined together by a single covalent bond

Atomic radius increases down the group for two reasons

New energy level, screening effect of inner electrons

Atomic radius decreases across a period for two reasons

Increase in effective nuclear charge, no increase in screening effect

First ionisation energy

The minimum energy required to completely remove the most loosely bound electron from a neutral gaseous atom in its ground state

1st ionisation energy decreases down a group for two reasons

Increasing atomic radius, screening effect of inner electrons

1st ionsation increasies across a period for two reasons

Increase in effective nuclear charge, decreasing atomic radius

Exception to the general trend across a period

Any atom that has a fully filled or half-filled sublevel has extra stability (e.g. Beryllium and Nitrogen)

2nd Ionisation energy

The energy required to remove an electron from an ion with one positive charge in the gaseous state

Electronegativity decreases down a group for two reasons

Increasing atomic radius, screening effect of inner electrons

Electronegativity increases across a period for two reasons

Increase in effective nuclear charge, decreasing atomic radius

Radioactivity

The spontaneous breaking up of unstable nuclei with the emission of oner or more types of radiation

Nuclear reaction

A process that alters the composition, structure or energy of an atomic nucleus

Half-life

The time taken for half of the nuclei in any given sample to decay

Radioisotope

Radioactive isotope

Radiocarbon dating (carbon dating)

Technique used to determine the age of an object containing carbon. It is based on the ratio of carbon-14 to carbon-12 in an object

One Mole of a substance

The amount of that substance that containes 6x10^23 particle of that substance

Mass of one mole of an element...

Is equal to relative atomic mass in grams

Relative atomic mass of a compound

The average mass of one molecule of that compound compared with one twelfth of the mass of one atom of the carbon-12 isotope

Mass of one mole of one mole in a compound...

Is equal to relative molecular mass in grams

Mass (in g) is equal to...

Mr x n

n of an element is equal to...

Mass of element/Relative Atomic Mass

n of a compound is equal to...

Mass of compound/Relative Molecular Mass

One mole of a substance contains...

6x10^23 particles of that substance

Gas

A substance that has no well-defined boundaries but diffuses rapidly to fill any container in which it is placed

Conversion of Celsius to Kelvin

K = C + 273

Normal atmospheric pressure

1x10^5 Pa or 100kPa

Measurement of volume

1L or 1000cm^3 or 1dm^3

Standard Temperature

s.t. is 273K

Standard Pressure

s.p. is 1x10^5Pa or 100kPa

Boyle's Law

At constant temperature, the volume of a fixed mass of gas is inversely proportional to its pressure

Equation for Boyle's Law

pV = k

Charles' Law

At constant pressure, the volume of a fixed mass of gas is directly proportional to its temperature measured on the Kelvin scale

Equation for Charles' Law

V/T = k

Equation for the Combined (General) Gas Law

(p1 x V1)/T1 = (p2 x V2)/T2

Gay-Lussac's Law of Combining Volumes

In a reaction between gases, the volumes of the reacting gases and the volumes of any gaseous products are in the ratio of small whole numbers provided the volumes are measured at the same temperature and pressure

Avogadro's Law

Equal volumes of gases contain equal numbers of molecules under the same conditions of temperature and pressure

S.t.p. of one mole of gas occupies a volume of...

22.4L

Ideal gas

A gas that perfectly obeys all the assumptions of the kinetic theory of gases under all conditions of temperature and pressure

Real gases

Gases that differ from ideal gases because (i) forces of attraction and repulsion do exist between the molecules and (ii) the volume of the molecules is not negligible

Equation for an Ideal Gas

pV = nRT

Molecular formula

A formula which shows the number and type of each atom present in a molecule of that compound

Empirical Formula

The formula showing the simplest while number ratio of the numbers of different atoms present in the molecule

To find the molecular formula

Molecular formula = Empirical formula x Whole Number

Arrhenius definition of an acid

A substance that dissociates in water to produce H+ ions

Arrhenius definition of a strong acid

A substance that almost completely dissociates in water to give hydrogen ions

Arrhenius definition of a weak acid

A substance that only slightly dissociates in water to give hydrogen ions

Arrhenius definition of a base

A substance that dissociates in water to produce OH- ions

Arrhenius definition of a strong base

A substance that almost completely dissociates in water to give hydroxide (hydroxyl) ions

Arrhenius definition of a weak base

A substance that only slightly dissociates in water to give hydroxide (hydroxyl) ions

Restrictions of the Arrhenius Theory

Hydronium ions exist rather than bare H+ ions, restricted his definitions to aqueous solutions i.e. reactions occurring in water

Bronsted-Lowry definition of an acid

A proton donor

Bronsted-Lowry definition of a base

A proton acceptor

Bronsted-Lowry definition of a strong acid

A good proton donor

Bronsted-Lowry definition of a weak acid

A poor proton donor