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Structure & Bonding
Properties & Structure
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Keywords

Stable

A full outer shell of electrons

Shell

A ‘layer’ of electrons

Ion

An atom that has lost or gained electrons

Delocalised

Free to move around the system

Oxidation

Loss of electrons

Reduction

Gain of electrons

Electrostatic force

The force between positive and negative charges

Delocalised

Free to move around the system

Lattice

3D arrangement of atoms/ions

Forming Bonds

All atoms want a full outer shell of electrons. This is a stable state. They achieve this by forming chemical bonds with other atoms. (all need 8, except H & He that want 2).

You are only expected to be able to work with the first 20 atoms of the periodic table. An element’s group number tells you how many electrons are in its outer shell, and you can work out how many it needs to fill the shell.

Ionic Bonding

Strong, electrostatic forces hold ions of opposing charges together.

The structure is described as an ionic lattice. The compounds have very high melting points because of this.

Image showing an ionic lattice

OILRIG (Ionisation)

Oxidation is Loss, Reduction is Gain

Metals always form positive ions (oxidised/lose electrons).

A group 2 metal will form a 2+ ion (loses two electrons).

Non-metals form negative ions (reduced/gain electrons).

An atom gaining two electrons

Covalent Bonding

Non-metal atoms bond by sharing electrons to form a very strong covalent bond.

A covalent bond forming between two fluorine atoms

Both fluorine atoms originally have 7 electrons on their outer shell, so need one more each.

By sharing one electron each, they form a single covalent bond. Each atom now has a full outer shell (8 electrons).

Metallic Bonding

Positive ions held together by a “sea of delocalised electrons” from the outer shells of the metal atoms.

The strong electrostatic forces between the ions and electrons mean metals have very high melting points.

A 'sea of delocalised electrons' holds positive metal ions together

The free electrons are able to move so metals are good conductors of electricity and heat.

Metals are malleable. This means the regular layers can slide over each other if they are hammered.

Metals have a regular layered structure which means they can be bent into shape

States of Matter

How each state of matter can be transformed into each other
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Ionic Compounds

Ionic compounds have regular structures. They make giant ionic lattices of oppositely charged ions. These ions experience strong, electrostatic forces of attraction in all directions. This means they have high melting points, as high energy is needed to break the bonds.

Ionic compounds can conduct electricity when molten or dissolved in water, as the ions are free to move = charge can flow.

Small Molecules

These have strong covalent bonds within the molecule, but weak intermolecular forces of attraction. This means they are either gas or liquid at room temperature.

They do not conduct electricity as there are no free electrons.

Water molecules displaying hydrogen bonding

Metals & Alloys

Pure metals are too soft for many uses. An alloy is a mixture of a metal with another element(s), and these have more desirable properties. They have different properties to the metals that are in them.

Alloys are less malleable than pure metals as they have irregular layers, and so they cannot slide over each other as easily. The atoms are still held together by metallic bonding.

Alloys do not have a regular arrangement of atoms, so cannot be easily bent into shape

Giant Covalent

These compounds are solid at room temperature. All of the atoms in a giant covalent structure are held together by strong covalent bonds. These bonds have to be broken, by large amounts of energy, to melt or boil these substances.

Diamond is made up of carbon forming four covalent bonds.

Graphite is made of hexagonal rings of carbon, each atom forming three bonds. Each atom contributes an electron to the “sea of delocalised electrons”.

Polymers

Polymers are large molecules, made of ‘repeating units’ called monomers. All atoms are bonded to other atoms, and they make a chain of strong covalent bonds.

The intermolecular forces are much weaker, and they chains can stretch and move over each other. There are many forces over the long chains, so polymers are solid at room temperature.

To increase the melting point, cross-links between chains can be added.

Graphite

Weak forces hold the layers together, so they can slide over each other, thus making graphite a great lubricant.

Only 3 electrons from each carbon atom form strong covalent bonds, and one delocalised electron.

High melting point, and a good thermal/electrical conductor.

Graphene

Graphene is just a single layer of graphite, and is one atom thick.

Fullerenes can be thought of as graphene sheets rolled into a ball, however graphene is made of 6 sided rings, and fullerenes are made of 5 and 6 sided rings.

Fullerenes can take the shapes of balls, or other shapes like tubes (nanotubes)

Diamond

Every carbon atom is bonded to four others, and because of this it forms a 3D lattice, called a tetrahedron.

No free electrons exist in this structure, so it does not conduct electricity.

Diamonds are used as cutting tools.

They are not ‘shiny’, as they do not reflect light - they refract it.

Nanoparticles

Between 1 and 100 nm in size. Fullerenes and nanotubes are small enough to be nanoparticles.

Common nanoparticles:

· Silver - Clothes/deodorant as antibacterial

· Titanium dioxide - Sunscreens as anti-UV

· Medicine - Drugs designed to work on one type of cell only

Nanoparticles have a large surface area, because of how small their individual volumes are.

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Last updated: 12/08/2017