Topic C1 The periodic table 2
Topic C2: Transition elements 7
Topic H: Acids and Bases 10
Topic 9: Redox Processes 19
Topic D: Models of Bonding and Structure (AHL) 30
Topic J Organic Chemistry 35
, Topic C1 The periodic table
- Horizontal rows (periods) corresponds to principal quantum number, n, of the
highest occupied energy level in the elements of the period
- Across the period: reactive metals ⇒ metalloids ⇒ nonmetals (except
period 1)
- AKA proton number, atomic number
- Vertical columns (groups) correspond to number of valence (outer-shell) electrons
- Elements in same group have same number of valence electrons
Atomic radius
- Atomic radius is half the distance between the nuclei of neighbouring atoms in
pure element (if noble gas, AKA non-bonding atomic radius, van der Waals’ radius)
- Non-metals: covalent radius
- Metals: metallic radius
- Anionic radius always larger than cationic radius in same period - have one more
occupied quantum shell of electrons
Ionisation energy
- First Ionisation energy is defined as the minimum energy required in removing
one mole of valence electrons from one mole of gaseous atoms to form 1 mole of
singly positively charged gaseous ions
- Always positive (endothermic) - energy always required to overcome
electrostatic attraction between electron and protons in nucleus
Electron affinity
- First electron affinity is the enthalpy change when one mole of gaseous atoms
acquires one mole of electrons to form one mole of singly negatively charged
gaseous ions
- Energy released when nucleus attracts electron added to outer shell
- Mostly negative (exothermic) - energy released
- Second electron affinity is the enthalpy change when one mole of singly negatively
charged gaseous ions acquires one mole of electrons to form one mole of doubly
negatively charged gaseous atoms
- Always positive - electron being added to anion; energy supplied to
overcome repulsion between two negatively charged species
- Trend varies depending on group
- e.g. Fluorine is less exothermic due to electron-electron repulsion too small)
- Exception: nitrogen has positive 1st E.A. - 3 unpaired electrons in valence shell;
extra repulsion when new electron “pairs up” - energy needs to be supplied to
overcome interelectronic repulsion between paired electrons
- Larger nuclear charges of fluorine and oxygen cancel out repulsion; EA are
still exothermic
, Periodic trends across period and down group
Across the period Down the group Exceptions/additional
Nuclear charge increases due to increase in Nuclear charge increases due to increase in Nuclear charge - total charge of protons in nucleus - more
atomic number atomic number protons = more electrostatic attraction on electrons
Shielding effect remains relatively constant - Shielding effect increase due to more inner Inner shell electrons repel valence electrons
inner quantum shells of electrons remain quantum shells Given by number of electrons in inner shells
the same Greater shielding effect = lower attraction
Effective nuclear charge increases - Effective nuclear charge decreases Difference between nuclear charge and shielding effect –
electrostatic attraction between nucleus and measure of how tightly electrons attracted to nucleus
valence electrons increases
Atomic radius decreases - valence electrons Atomic radius increases - increased Cationic radius smaller than parents - fewer electrons in
attracted closer to nucleus distance reduces electrostatic attraction cation; shielding effect decreases
Ionic radius decreases - increase in nuclear between protons in nucleus and valence Anionic radius larger than parents - increase in
charge electrons repulsion between electrons
Ionic radii increase - more electron energy
levels
Ionisation energy increases - more energy ionisation energy decreases - size of atoms - Less energy required to remove 2p electron than 2s
required to remove valence electrons increases; reduced electrostatic attraction; electron - 2p electron further from nucleus
less energy required to remove - Less energy required to remove paired 2p electron than
unpaired 2p due to repulsion
Electronegativity increases - atoms get Electronegativity decreases - decreased
smaller, electrostatic attraction between electrostatic attraction between bonding
bonding electrons and nuclei increases electrons and nuclei of atoms
Electron affinity becomes more exothermic Electron affinity becomes less exothermic - Fluorine less exothermic than chlorine (electron-electron
- added electron strongly attracted to decrease in attraction between nucleus and repulsion)
nucleus added electron Nitrogen less exothermic (positive first EA)
, Melting points
- Normal melting point is the temperature at which a pure solid is in equilibrium
with its pure liquid at one atmospheric pressure
- Overcome intermolecular force of attraction
- Depends on structure and electrostatic force of attraction between
particles
Trend across period 3
Na to Al Giant metallic Increase in metallic bond strength:
structure - Decrease in metallic radius
- increase in number of electrons donated to sea
of electrons
Increased charge density of cation
Si Giant molecular Covalent bonds throughout giant covalent structure
P to Cl Simple molecular Instantaneous dipole-induced dipole (london
(dispersion)) forces
- Increases when Mr increases [S8 > P4 > Cl2]
Ar Monoatomic Instantaneous dipole-induced dipole (london
(dispersion)) forces
Trends in group 1 (alkali metals)
Physical properties Chemical properties
- Good conductors of electricity and heat - Very reactive
- Low densities - Form ionic compounds with
- Shiny grey surface when cut nonmetals
- Low melting point - Good reducing agents
Topic C2: Transition elements 7
Topic H: Acids and Bases 10
Topic 9: Redox Processes 19
Topic D: Models of Bonding and Structure (AHL) 30
Topic J Organic Chemistry 35
, Topic C1 The periodic table
- Horizontal rows (periods) corresponds to principal quantum number, n, of the
highest occupied energy level in the elements of the period
- Across the period: reactive metals ⇒ metalloids ⇒ nonmetals (except
period 1)
- AKA proton number, atomic number
- Vertical columns (groups) correspond to number of valence (outer-shell) electrons
- Elements in same group have same number of valence electrons
Atomic radius
- Atomic radius is half the distance between the nuclei of neighbouring atoms in
pure element (if noble gas, AKA non-bonding atomic radius, van der Waals’ radius)
- Non-metals: covalent radius
- Metals: metallic radius
- Anionic radius always larger than cationic radius in same period - have one more
occupied quantum shell of electrons
Ionisation energy
- First Ionisation energy is defined as the minimum energy required in removing
one mole of valence electrons from one mole of gaseous atoms to form 1 mole of
singly positively charged gaseous ions
- Always positive (endothermic) - energy always required to overcome
electrostatic attraction between electron and protons in nucleus
Electron affinity
- First electron affinity is the enthalpy change when one mole of gaseous atoms
acquires one mole of electrons to form one mole of singly negatively charged
gaseous ions
- Energy released when nucleus attracts electron added to outer shell
- Mostly negative (exothermic) - energy released
- Second electron affinity is the enthalpy change when one mole of singly negatively
charged gaseous ions acquires one mole of electrons to form one mole of doubly
negatively charged gaseous atoms
- Always positive - electron being added to anion; energy supplied to
overcome repulsion between two negatively charged species
- Trend varies depending on group
- e.g. Fluorine is less exothermic due to electron-electron repulsion too small)
- Exception: nitrogen has positive 1st E.A. - 3 unpaired electrons in valence shell;
extra repulsion when new electron “pairs up” - energy needs to be supplied to
overcome interelectronic repulsion between paired electrons
- Larger nuclear charges of fluorine and oxygen cancel out repulsion; EA are
still exothermic
, Periodic trends across period and down group
Across the period Down the group Exceptions/additional
Nuclear charge increases due to increase in Nuclear charge increases due to increase in Nuclear charge - total charge of protons in nucleus - more
atomic number atomic number protons = more electrostatic attraction on electrons
Shielding effect remains relatively constant - Shielding effect increase due to more inner Inner shell electrons repel valence electrons
inner quantum shells of electrons remain quantum shells Given by number of electrons in inner shells
the same Greater shielding effect = lower attraction
Effective nuclear charge increases - Effective nuclear charge decreases Difference between nuclear charge and shielding effect –
electrostatic attraction between nucleus and measure of how tightly electrons attracted to nucleus
valence electrons increases
Atomic radius decreases - valence electrons Atomic radius increases - increased Cationic radius smaller than parents - fewer electrons in
attracted closer to nucleus distance reduces electrostatic attraction cation; shielding effect decreases
Ionic radius decreases - increase in nuclear between protons in nucleus and valence Anionic radius larger than parents - increase in
charge electrons repulsion between electrons
Ionic radii increase - more electron energy
levels
Ionisation energy increases - more energy ionisation energy decreases - size of atoms - Less energy required to remove 2p electron than 2s
required to remove valence electrons increases; reduced electrostatic attraction; electron - 2p electron further from nucleus
less energy required to remove - Less energy required to remove paired 2p electron than
unpaired 2p due to repulsion
Electronegativity increases - atoms get Electronegativity decreases - decreased
smaller, electrostatic attraction between electrostatic attraction between bonding
bonding electrons and nuclei increases electrons and nuclei of atoms
Electron affinity becomes more exothermic Electron affinity becomes less exothermic - Fluorine less exothermic than chlorine (electron-electron
- added electron strongly attracted to decrease in attraction between nucleus and repulsion)
nucleus added electron Nitrogen less exothermic (positive first EA)
, Melting points
- Normal melting point is the temperature at which a pure solid is in equilibrium
with its pure liquid at one atmospheric pressure
- Overcome intermolecular force of attraction
- Depends on structure and electrostatic force of attraction between
particles
Trend across period 3
Na to Al Giant metallic Increase in metallic bond strength:
structure - Decrease in metallic radius
- increase in number of electrons donated to sea
of electrons
Increased charge density of cation
Si Giant molecular Covalent bonds throughout giant covalent structure
P to Cl Simple molecular Instantaneous dipole-induced dipole (london
(dispersion)) forces
- Increases when Mr increases [S8 > P4 > Cl2]
Ar Monoatomic Instantaneous dipole-induced dipole (london
(dispersion)) forces
Trends in group 1 (alkali metals)
Physical properties Chemical properties
- Good conductors of electricity and heat - Very reactive
- Low densities - Form ionic compounds with
- Shiny grey surface when cut nonmetals
- Low melting point - Good reducing agents
