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.