Highly detailed notes on Unit 2 - Bonding and Structure covering all the specification points. This allowed me to achieve an A* in chemistry. Covers the different types of bonding and properties of each type. Explains orbitals, shapes of molecules and ions, and polarity in detail. Also covers types...
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2-Bonding and Structure
Bonding
Metallic Bonding
Physical Properties of Metals
Chemical Properties of Metals
Ionic Bonding
Strength of Ionic Bonds
Ionic Radius
Evidence for the Existence of Ions
Covalent Bonding
Expansion of the Octet
Single and Double Bonds
Covalent Bond Strength
Dative Covalent Bonds
Shapes of Molecules and Ions
Predicting Shapes of Molecules
Polarity of Molecules
Shape and Polarity
Forces between Ions, Atoms and Molecules
Chemical Bonds
Forces between Covalently Bonded Molecules
London Forces
Permanent Dipole-Dipole Forces
Hydrogen Bonding
Forces and Physical Properties
2-Bonding and Structure 1
, Melting a Solid
Boiling a Liquid
Structure
Ionic Solids
Melting
Electrical Conductivity
Giant Covalent Substances
Diamond
Graphite
Graphene
Silicon
Simple Molecular Substances
Melting and Boiling
Hydrogen-Bonded Molecular Substances
Boiling
Polymeric Substances
Melting
Solubility
Ionic Solids
Compounds that can form Hydrogen Bonds with the Solvent
Non-Hydrogen-Bonding Subtances
Non-Aqueous Solvents
Bonding
What is Bonding?
A force of attraction between two species. E.g. between two atoms or between
two ions or molecules. Making a bond always releases energy, so it is an
Exothermic reaction. Breaking a bond always absorbs energy, so it is an
Endothermic reaction.
What is Molecule?
Groups of atoms held together by covalent bonds. E.g. H2 and NH3 . The have
no charge - neutral.
3 ways of chemically bonding atoms:
Metallic - between atoms of metallic elements.
Ionic - usually between a non-metal and a metal.
Covalent - usually between two non-metal elements.
2-Bonding and Structure 2
, What is an Ionic Bond?
The force of attraction between positive ions (cations) and negative ions
(anions). Ions form when atoms lose electrons from their outer (valence) shell or
gain electrons into their valence shell:
Usually, metals form cations and non-metals for anions. Except ammonium
ions - NH4 + and hydrogen ions - H+ .
The force of attractions between ions depends on:
The charges on the ions. Higher charge = stronger attraction.
The distance apart of the ions. Further apart = weaker attraction.
The attraction between Mg2+ and O2− in MgO is stronger than between Na
+
and Cl− in NaCl because:
The magnesium and oxygen atoms are more charged.
The magnesium atom is smaller than the sodium ion.
What is a Covalent Bond?
The force of attraction between the nuclei of the two atoms and the shared pair
of electrons.
Double covalent bonds involve sharing two pairs of electrons.
Triple covalent bonds involve the sharing of three pairs of electrons.
Metallic Bonding
How does metallic bonding work?
Metal atoms lose their valence (outer) electrons and form cations (+). These
cations are arranged in a regular lattice, with one layer of ions above another
2-Bonding and Structure 3
, layer, surrounded by a sea of delocalised electrons that can move through the
lattice.
The electrons are:
Delocalised through the structure of metal ions.
The bonding is the attraction between:
Positive ions, which are fixed in position and the negative electrons which are
moving between the ions.
The metallic radius is slightly ______ than the ionic radius:
Larger than the ionic radius, as there are small spaces occupied by delocalised
electrons.
The strength of the metallic bond depends on:
The charge of the metal ion - the same number of delocalised electrons.
The metallic radius.
The structure of the metal lattice - group 1 metals have a body-centred
cubic. Calcium, magnesium and aluminium have a face-centred rubric.
Physical Properties of Metals
Density and Melting Temperature:
2-Bonding and Structure 4
, Group 1 metals - low densities and low melting temperatures. Electron
configuration - ns1 . (n is the orbital number of the outer s orbital, and they
lose the outer s-electron).
Group 2 metals - denser and higher melting temperature. They lose both the
ns2 electrons. Therefore, the ion is smaller than the group 1 ion in the same
period. As there is a smaller radius and higher charge, there is a greater
attraction between the metal's cation and the delocalised electrons.
D-Block metals - they use their (n-1)d electrons and their ns2 electrons in
bonding. So, they are much harder, denser and have higher melting
temperatures.
Across a Period - The melting temperature of metals increases as the
metallic radius decreases and more electrons are released for bonding.
Down a group - the melting temperature decreases - metallic radius
increases and the force of attraction between the metal ions and delocalised
electrons decreases.
Electrical Conductivity:
Electricity is a flow of charge.
All metals conduct electricity when solid and molten - the sea of electrons is
mobile and moves through the lattice of metal ions.
So the electric current in a metal is the flow of electrons.
Thermal Conductivity:
Metals are good conductors of heat - free moving electrons pass kinetic
energy along the metal.
Malleability:
Metals can be hammered or pressed into different shapes.
One layer of metal ions can slide over another layer - electrons between the
layers prevent strong forces of repulsion between the positive ions in one
layer and the positive ions in another layer.
Some metals - lead and gold, are soft. D-block metals are harder as there
are more electrons that bind the layers together.
Chemical Properties of Metals
2-Bonding and Structure 5
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