Chemistry Key Definitions

Relative Atomic Mass

The weighted average mass of an atom of an element relative to one-twelfth of the mass of an atom of carbon-12.

Relative Isotopic Mass

The mass of an atom of a particular isotope relative to one-twelfth of the mass of an atom of carbon-12

Relative Molecular Mass

The average mass of a molecule relative to one-twelfth of the mass of an atom of carbon-12.

Relative Formula Mass

The average mass of a formula unit relative to one-twelfth of the mass of an atom of carbon-12.

Avogadro Constant

One mole is defined as the same number of particles there are atoms of carbon in exactly 12 grams of carbon 12

Empirical Formula

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

Molecular Formula

The actual number of atoms in each element of a molecule

Molar volume and equation

Volume per mole of gas
n=v÷24.5

PV=nRT (units)

Pressure - Pa
Volume - m³
Temperature - Kelvin
R - 8.31

Concentration

A measurement of how much solute exists within a certain volume of solvent

Standard Solution

A solution of known concentration

Reversible reaction

A reaction that can go either way depending on the conditions

Dynamic Equilibrium

An equilibrium that exists in a closed system when the rate of the forward reaction is equal to the rate of the reverse reaction

Chatelier's Principle

When a system at equilibrium is subject to a change, the position of the equilibrium will shift to minimise the change

Position of Equilibrium

The proportion of products to reactants in an equilibrium mixture.

An acid

Proton donor

A base

Proton acceptor

Lone pair

Non-bonding pair of electrons

Ionic bond

A electrostatic force of atraction between oppositely charged ions, one metallic and one non-metallic

Covalent Bond

A bond involving the sharing of a pair of electrons between 2 non-metals

Dative Covalent/Coordiate Bond

A bond which both electrons have been donated from the same atom

Polar covalent bond

A bond that shows a difference in charge

Intermolecular Force

a force that occurs between molecules

Permanent Dipole

A difference in charge across a covalent bond that arise due to the difference in electonegativity bewteen two atoms

van de Waal's forces

weak attractions between nonpolar molecules

Hydrogen Bond

weak attraction between a hydrogen atom and another atom

Molar first ionisation energy

The energy required to remove one mole of electrons from one mole of its gaseous ions

Electronegativity

The ability of an atom in a covalent bond to attract the bonding pair of electrons towards itself

Compound

Two or more different atoms chemically joined together

Molecule

Two or more atoms chemically bonded together

Sodium Chloride

Example of ionic compound
Face centred cubic
6:6 coordination

Caesium chloride

Body centred cubic
8:8 coordination
NB: Caesium ions are bigger than sodium ions so more ions and surround it

Physical Properties of ionic compounds

Melting Point -
Strength -
Conductability -
Solubility -

High - large energy to overcome strong electrostatic attractions

Very brittle - ions arragned by similarity to avoid repulsion splitting the crystal

Does not conduct while solid - ions held strongly
Conduct when molten or in aqueous solution

Insoluble in polar solvents - soluble in water

Repulsion in covalent bonds

Attraction between oppositely charged nuclei overcomes repulsion between two positively charged nuclei.

Physical properties of covalent compounds
Conductability -
Solubility -
Boiling Point -

Does not conduct electricity as no mobile ions or electrons
Organic solvents>water
Low - vdw weak

Electronegativity in Bonds

Cl-CL - purely covalent
H-CL - Cl more electronegative so bonding electrons closer to it
Difference in charge (permenant dipole)
Hδ+ - Clδ

Two types of intermolecular forces and how it forms

Hydrogen bonding
Van der waal's forces
Caused by weak attractive forces between dipoles in different molecules

van der Waals

How they are caused -
How they are increased
Strength

Caused by movement of electrons in shells.
This movement causes an instantaneous dipole.
Dipoles attract each other by VDW.

Increases with number of electrons

Weaker than all other bonds in this topic

Hydrogen Bonding

HON - formation
Causes permanent dipole-dipole

NOF

Ice relative to water

Less dense due to hydrogen bonds holding water molecules further apart.

Properties of hydrogen bonding

Allows pondskaters to walk on water

Properties of water

Hydrogen bonds are an extra force - hgher than vdw

Iodine

Grey at room temp, purple when sublimed (gas)
Composed of diatomic molecules
Each molecule only attracted by vdw
Little energy needed to separate molecules

Properties of Diamond

Very high melting point due to strong numerous covalent bonds
Strong as carbon bonded to 4 others
Cannot conduct electricity due to no free electrons
Used in drills

Graphite

Very high melting point due to strong numerous covalent bonds.
Soft - each carbon bonded to 3 others
Layers held by vdw - can slide over each other
Can conduct due to free electron that can carry a charge.
Used in lubricants

Silica

Very high melting point due to numerous strong covalent bonds
Cannot conduct as no mobile electrons
Strong

VSEPR Principle

The shape of a molecule is determined by the number of electron pairs.

Electrons repel each other due to negative charges.

Electrons spread out to minimise repulsion.

Lone pairs repel more than bonding pairs

BF3

3 bonding pairs
120°
Trigonal Planar

CH4

4 bonding pairs
0 lone pairs
109.5°
Tetrahedral

SF6

6 bonding pairs
90°
Octahedral

NH3

3 bonding pairs
1 lone pair
107°
Pyramidal

H20

2 bonding pairs
2 lone pairs
104.5°
Non linear

Element arrangement

In order of increasing atomic number

S Block

Group 1 and Group 2.

Basic oxides

Cations.

Reducing agents

D block

transition metals

P block

Group 6 and 7

Acidic oxides

Anions

Oxidising agents

Oxidation States

Elements = 0
Hydrogen = +1
Oxygen = -2
Group 1 = +1
Group 2 = +2
Group 7 = -1

S block elements

Reacting with oxygen
Reacting with Acids

All burn
All form a salt and water

Group 1 reacting with water

Equation
Forms

2X(s) + 2H2O(l) → 2LiOH(aq) + H2(g)
Colourless alkaline X hydroxide solution and H gas

Equation of Group 2 metals reacting with water/steam

Reactivity of Group 1 and 2 metals

Reactivity increases down the group
Group 1 more reactive than 2

Solubility of Group 1 Salts in water

All are soluble

Solubility of Group 2 Salts in water

Hydroxides more soluble down the group
Sulphates more soluble down the group
All carbonates insoluble
All nitrates and chlorides are soluble

How to perform a flame test

Dip nichrome wire into HCL
Place in roaring flame.
Repeat till no colour produced - clean
Dip in acid and then sample

Flame Test colours

Lithium
Sodium
Potassium
Magnesium
Calcium
Strontium
Barium

Red
Orange
Lilac
Colourless
Brick Red
Crimson Red
Apple green

Colours of Halogens (important)

Fluorine - yellow gas
Chlorine - yellow to green gas
Bromine - Orange/brown liquid
Iodine - Grey Solid

Displacement reactions of Halogens with Metals

Reactivity decreases down the group.
E.g 2NaBr + Cl2 - 2NaCl + Br2
Feflects the decrease in oxidising power with chlorine oxidising bromide ion to bromine and reducing to a chloride

Melting Points of Halogens

Increases down the group.
Due to increasing molecular forces as more electrons in induced dipole-dipole.

Test with Ammonia

Chlorine dissolves in dilute ammonia
Bromine dissolves in conc ammonia
Iodine does not dissolve in either.

Use of Chlorine and Fluorine in water

Kills bacteria and viruses.
Reduces tooth decay
Unethical but outweighs risks

E=hf

Energy in KJ
H is Plank's constant
In hertz

λ=c/f

Wavelength in m
Speed of light
F of light (Hz)

Absorbtion Spectrum

Arises when the electron is excited to higher energy levels

Emission Spectrum

Arises when the electron emits energy as it drops to lower energy levels.

Balmer Series

Occurs in the visible part of the Spectrum.
Arises from electronic transitions from n=2 and above
Energy levels grow closer as we approach infinity.
Shows the quantisation of energy - how electrons fall to lower discrete levels.

Lyman Series

Between n=1 and above
If electron falls back to n=1, series of lines are shown in ultraviolet region.
Lines converge as energy levels become less

EM Spectrum

Radio, micro , IR, UV, X, Gamma
Gamma has the highest energy and frequency.
Radiowaves has the highest wavelength

Relationship of Lyman Series and Ionisation energy of Hydrogen

Frequency of convergence limit
Use E=hf
Measure the convergence frequency

Calculation Ionisation Energy

E=Lhc/λ

Behavior of radiation in radioactive waves

Alpha has a small deflection to the negative in an electric field
Beta has a large deflection to the postive in an electric field.
Gamma is unaffected by the electric field.

Penetrating power of radioactive waves

Alpha is the most ionising. Travels a few cm in air and stopped by a sheet of paper.

Beta travels a few metres in air. Stopped by a sheet of aluminium.

Gamma is the least ionising. Travels a few km in air, stopped by a thick sheet of lead.

Hazards of radiation

Ionises cells which can cause cancer.
Caues mutations.
Reduces cell growth rate.

Uses of radiation

In medicine, radium used to treat cancers.
Used to destroy tumours in the body.

In metal fatigue, gamma waves used from colbalt 60

Electron Capture

Electron enters the nucleus.
Neutrino emitted from the nucleus
Atomic number goes down by 1

81 Kr 0 e ---- 81 Br
36 -1 35

Positron decay

Emits a postiron and a neutrino.
Atomic number decreases by 1

23 M ---- 23 Na + 0 e
12 11 +1

Loss of Alpha Particle

226 Ra ---- 222 Rn + 4 He
88 86 2

Mass number decreases by 4.
Atomic number decreases by 2

Loss of Beta particles

14 C ---- 14 N + e-
6 7

A neutron is converted to a proton and an electron. The proton remains in the nucleus and the electron is ejected as a beta particle

Electronic arrangements

Orbital -
How orbitals can hold electrons

A region in space where one can find an electron - orbtal can hold two electrons in opposite spin.

Factors affecting ionisation energy

Atomic radius
Nuclear charge
Electron shielding

Atomic Radius

The further away the electron is from the nucleus, the weaker the nuclear attraction

Nuclear Charge

The total charge of all the protons in the nucleus.
The higher the NC, the greater the attractive force on the electrons.

Electron Shielding

Caused by the inner shells of electrons repelling the outer electrons.

First ionisation Energy

The energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous ions

Describing ionisation trends
Start from Hydrogen to Magnesium

Hydrogen - relatively high as in first shell with no shielding.

Helium - Higher than H due proton increasing nuclear charge. No increase in shielding.

Lithium has a significant drop. Increase in shielding outweighs the increased nuclear charge.

Berylium has a higher IE due to increased nuclear charge, no increase in shielding.

Drop in Boron due to increase shielding as entered P orbital. Proves existence of subshells.

Increase in carbon due to increase in nuclear charge. Does not pair with other electrons to avoid repulsion, goes into another P orbital.

Increase for Nitrogen due to increase nuclear charge.

Drop for oxygen as the electron is now forced to pair up. Repulsion means less energy needed to remove.

Increase in fluorine due to greater nuclear charge.

Increase in Neon due to greater nuclear charge.

Cycle repeats.