Chemistry
Formulae and equations: Unit 1.1
Useful definitions
Element A substance which cannot be splitchemicallyintoanything simpler
The atom The smallest part of an element that can take part in a chemical reactions
Molecule T
he smallest part of a compound
The smallest part of atoms that don't exist in the free state i.e H2
An ion An electrically charged atom or molecule, cations are +, anions are -
A symbol Represents an atom or mole of an element
Formula Represents one molecule of a compound or the simplest ratio of ions present
Valency numerical measure of the combining power of an atom/ion
A
Historically the number of hydrogen atoms which will combine with one or a group of
atoms (because hydrogen has only 1 electron in its outer shell)
Also the number of positive or negative charges in an ion
Oxidation and reduction
● Oxidation ● Reduction
− + 2+ − +
𝐿𝑖 − 1𝑒 → 𝐿𝑖
○ ○ 𝐶𝑢 + 1𝑒 → 𝐶𝑢
○ Addition of oxygen ○ Removal of oxygen
○ Loss of hydrogen ○ Gain of hydrogen
○ Loss of electrons ○ Gain of electrons
Oxidation numbers and oxidation states
● The number of electrons needed to be gained or lost for an element to become
neutral
○ Mg2+ needs to gain 2 electrons, the O.N is +2
● O.N are given to various elements by applying the following rules
1. The O.N of an element is 0
2. The sum of oxidation number in a neutral compound is 0
3. In a compound containing hydrogen, hydrogen has an O.N of +1
○ Except in metal hydrides such as NaH and KH, hydrogen has an O.N
of -1 and the metal becomes +1
4. In a compound containing oxygen, oxygen has an O.N of -2
○ Except with fluoride or peroxides1, where the O.N is multiplied by the
number of oxygen atoms
5. In a compound containing a halogen, halogens have an O.N of -1
6. Group 1 metals have oxidation numbers of +1 in compounds
7. Group 2 metals have oxidation numbers of +2 in compounds
8. In species with atoms of different elements, the most electronegative element
is given the negative O.N
1
Compounds with more than one oxygen
1
, Oxidation agent Causes oxidation by gaining electrons in a reaction
Reducing agent Causes reduction by losing electrons in a reaction
Ionic equations
● Provide a shorter equation which focuses on the changes taking place
● Spectator ions- have not taken part in any chemicalchange and can be removed
from the equation
2
,Basic ideas about atoms:Unit 1.2
Orbitals
● Region of space where there is a 95% chance of finding an electron𝑥 ∈ 𝑅
Orbit/shell/energy level The path of an electron around the nucleus
Charge cloud he orbital is regarded as a spread of charge a the electron is likely to be
T
anywhere in the orbital
Greatest electron density at the centre
● Electrons can occupy 4 types orbitals within an orbit
○ s,p,d and f orbitals
1. s-orbitals are spherical
2. p-orbitals are dum-bell shaped
Energy levels of sub-shells
● Shells are represented by n and are numbered
○ n=1, n=2 etc
● Electrons exist in pairs
○ Spin in opposite directions↑↓
● Orbitals can hold 2 electrons each
○ s subshell - 1 s-orbital
○ p subshell - 3 p-orbitals
○ d subshell - 5 d-orbitals
● The 4s subshell is exceptional as it has a lower energy level than the 3d subshell
Electronic configuration
● Describes how electrons are arranged in their shells, subshells and orbitals
● Excited state configuration -Atoms with one or moreexcited electrons
● 3 rules apply to atoms in theground state
The Aufbau principle lectrons in ground state occupy orbitals in order of orbital levels with
E
the lowest being filled first
- Intuitive except 4s has a lower energy level than 3d and is filled
first
The Pauli exclusion principle Each orbital can hold 2 electrons which must have opposite spins
- 2 parallel spins is not allowed
Hund’s rule ach orbital of the same energy level must be occupied singly with
E
parallel spin before pairing can take place
● Chromium and copper are exceptional
○ More stable to have a ½ full or full 3d subshell than to have a full 4s subshell
and a 3d subshell with an empty orbital
■ Only 1 electron is given to the 4s subshell
3
, Noble gas cores
● Noble gases can be used to represent full shells when writing electronic configuration
● Only works using noble gases due to their full outer shells
● O2 - can be written as [Ne]
Ionisation energies
● Energy needed to remove electrons from an atom
● Molar first ionisation energy (1st I.E)
○ Energy required to remove one mole of electrons from 1 mole of gaseous
atoms to form one mole of gaseous positively charged ions
● The electrons lost to form positive ions are always taken from the 4s subshell and
then the 3d subshell
Nuclear charge otal positive charge of all the protons in a nucleus of an atom
T
Greater nuclear charge = greater attractive force
Effective nuclear charge The nuclear charge and outer electron experiences
● Energy required depends on:
Distance from nucleus Weaker attraction force therefore less effective nuclear charge
Shielding illed inner shells weaken the attraction force, more shielding results in less
F
effective nuclear charge
- Filled subshells slightly increase shielding
Spin pair repulsion Paired e- is slightly easier to remove because itsrepelled by its pair
Nuclear charge
○ Total positive charge of all the protons in a nucleus of an atom
■ Greater nuclear charge = greater attractive force
● Effective nuclear charge
○ The nuclear charge an outer shell electron experiences
● 1st I.E generally increases across a row due to more protons for the same shielding,
therefore increase in nuclear charge
○ Some decreases in 1st I.E needed due to shielding from subshells and spin
pair repulsion
● 1st I.E generally decreases down a group due to less effective nuclear charge from
an increase in shielding
● Noble gases have the highest 1st I.E for their period
○ Greatest number of protons in the nucleus and therefore greatest nuclear
charge for the same shielding
● Alkali metals have the lowest 1st I.E for their period
○ Lowest number of protons in the nucleus therefore lowest nuclear charge for
the same shielding
4
Formulae and equations: Unit 1.1
Useful definitions
Element A substance which cannot be splitchemicallyintoanything simpler
The atom The smallest part of an element that can take part in a chemical reactions
Molecule T
he smallest part of a compound
The smallest part of atoms that don't exist in the free state i.e H2
An ion An electrically charged atom or molecule, cations are +, anions are -
A symbol Represents an atom or mole of an element
Formula Represents one molecule of a compound or the simplest ratio of ions present
Valency numerical measure of the combining power of an atom/ion
A
Historically the number of hydrogen atoms which will combine with one or a group of
atoms (because hydrogen has only 1 electron in its outer shell)
Also the number of positive or negative charges in an ion
Oxidation and reduction
● Oxidation ● Reduction
− + 2+ − +
𝐿𝑖 − 1𝑒 → 𝐿𝑖
○ ○ 𝐶𝑢 + 1𝑒 → 𝐶𝑢
○ Addition of oxygen ○ Removal of oxygen
○ Loss of hydrogen ○ Gain of hydrogen
○ Loss of electrons ○ Gain of electrons
Oxidation numbers and oxidation states
● The number of electrons needed to be gained or lost for an element to become
neutral
○ Mg2+ needs to gain 2 electrons, the O.N is +2
● O.N are given to various elements by applying the following rules
1. The O.N of an element is 0
2. The sum of oxidation number in a neutral compound is 0
3. In a compound containing hydrogen, hydrogen has an O.N of +1
○ Except in metal hydrides such as NaH and KH, hydrogen has an O.N
of -1 and the metal becomes +1
4. In a compound containing oxygen, oxygen has an O.N of -2
○ Except with fluoride or peroxides1, where the O.N is multiplied by the
number of oxygen atoms
5. In a compound containing a halogen, halogens have an O.N of -1
6. Group 1 metals have oxidation numbers of +1 in compounds
7. Group 2 metals have oxidation numbers of +2 in compounds
8. In species with atoms of different elements, the most electronegative element
is given the negative O.N
1
Compounds with more than one oxygen
1
, Oxidation agent Causes oxidation by gaining electrons in a reaction
Reducing agent Causes reduction by losing electrons in a reaction
Ionic equations
● Provide a shorter equation which focuses on the changes taking place
● Spectator ions- have not taken part in any chemicalchange and can be removed
from the equation
2
,Basic ideas about atoms:Unit 1.2
Orbitals
● Region of space where there is a 95% chance of finding an electron𝑥 ∈ 𝑅
Orbit/shell/energy level The path of an electron around the nucleus
Charge cloud he orbital is regarded as a spread of charge a the electron is likely to be
T
anywhere in the orbital
Greatest electron density at the centre
● Electrons can occupy 4 types orbitals within an orbit
○ s,p,d and f orbitals
1. s-orbitals are spherical
2. p-orbitals are dum-bell shaped
Energy levels of sub-shells
● Shells are represented by n and are numbered
○ n=1, n=2 etc
● Electrons exist in pairs
○ Spin in opposite directions↑↓
● Orbitals can hold 2 electrons each
○ s subshell - 1 s-orbital
○ p subshell - 3 p-orbitals
○ d subshell - 5 d-orbitals
● The 4s subshell is exceptional as it has a lower energy level than the 3d subshell
Electronic configuration
● Describes how electrons are arranged in their shells, subshells and orbitals
● Excited state configuration -Atoms with one or moreexcited electrons
● 3 rules apply to atoms in theground state
The Aufbau principle lectrons in ground state occupy orbitals in order of orbital levels with
E
the lowest being filled first
- Intuitive except 4s has a lower energy level than 3d and is filled
first
The Pauli exclusion principle Each orbital can hold 2 electrons which must have opposite spins
- 2 parallel spins is not allowed
Hund’s rule ach orbital of the same energy level must be occupied singly with
E
parallel spin before pairing can take place
● Chromium and copper are exceptional
○ More stable to have a ½ full or full 3d subshell than to have a full 4s subshell
and a 3d subshell with an empty orbital
■ Only 1 electron is given to the 4s subshell
3
, Noble gas cores
● Noble gases can be used to represent full shells when writing electronic configuration
● Only works using noble gases due to their full outer shells
● O2 - can be written as [Ne]
Ionisation energies
● Energy needed to remove electrons from an atom
● Molar first ionisation energy (1st I.E)
○ Energy required to remove one mole of electrons from 1 mole of gaseous
atoms to form one mole of gaseous positively charged ions
● The electrons lost to form positive ions are always taken from the 4s subshell and
then the 3d subshell
Nuclear charge otal positive charge of all the protons in a nucleus of an atom
T
Greater nuclear charge = greater attractive force
Effective nuclear charge The nuclear charge and outer electron experiences
● Energy required depends on:
Distance from nucleus Weaker attraction force therefore less effective nuclear charge
Shielding illed inner shells weaken the attraction force, more shielding results in less
F
effective nuclear charge
- Filled subshells slightly increase shielding
Spin pair repulsion Paired e- is slightly easier to remove because itsrepelled by its pair
Nuclear charge
○ Total positive charge of all the protons in a nucleus of an atom
■ Greater nuclear charge = greater attractive force
● Effective nuclear charge
○ The nuclear charge an outer shell electron experiences
● 1st I.E generally increases across a row due to more protons for the same shielding,
therefore increase in nuclear charge
○ Some decreases in 1st I.E needed due to shielding from subshells and spin
pair repulsion
● 1st I.E generally decreases down a group due to less effective nuclear charge from
an increase in shielding
● Noble gases have the highest 1st I.E for their period
○ Greatest number of protons in the nucleus and therefore greatest nuclear
charge for the same shielding
● Alkali metals have the lowest 1st I.E for their period
○ Lowest number of protons in the nucleus therefore lowest nuclear charge for
the same shielding
4
