The Solid State

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The Solid State

Five types of solids require investigation in terms of structure and bonding and properties.

1. Metallic

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2. Giant Ionic

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3. Giant Covalent (Macromolecular)

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4. Simple Molecular

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5. Hydrogen Bonded


Metals usually have high melting points and boiling points. These high values suggest that strong forces exist between the separate atoms in the metal.

The physical properties of a metal can be explained by using a model in which the outer shell electrons of the metal move randomly throughout, the lattice of regularly spaced positive ions.

The moving electrons are sometimes described as a 'sea' or 'cloud' of moving and fluctuating negative charge (delocalised).

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In the metal lattice, each positively charged ion is attracted to the 'cloud' of negative electrons and vice versa. These electrostatic attractions bind the entire crystal together as a single unit.

In this model, one particular electron does not belong to one particular metal ion.

Note: these ions are often referred to as pseudo cations due to them 'pretending' or 'acting' like ions.

As you move from Na to Mg to Al this bonding gradually gets stronger. This is due to the increase in outer electrons and charge of ion. Hence, boiling point is seen to increase.

Due to the mobility of their electrons, metals are good conductors of heat and electricity.

Ionic lattices

As with metal atoms, ions may be thought of as solid spheres, which pack together as closely as possible.

The lattice is made up of ions, cations(+) surrounding anions(-) and vice versa. Hence the force holding these ions together and ultimately the lattice is due to electrostatic attraction.

Example: NaCl

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Properties of ionic crystals are:

a) Hard, brittle solids

b) High melting point and boiling point due to large amount of energy required to overcome electrostatic forces. The greater the charge of the ions, the greater the electrostatic attraction, the greater the melting point and boiling point.

Example: MgO has a greater boiling point and melting point than NaCl.

c) Soluble in polar solvents such as water due to charge on ions attracting polar water molecules.

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d) Good conductors of electricity when molten or liquefied, but not as solids. Remember, for electricity to flow there must be a moving charge (delocalised electrons or freely moving ions).

Giant covalent

The metalloids (carbon (graphite) and silicon) and the non-metal carbon (diamond) have giant covalent structures. These elements have medium electronegativities and the ability to form four covalent bonds. In general each atom can be imagined to be situated at the centre of a tetrahedron strongly bound by four other atoms. The covalent linking in these elements extends from one atom to the next through the whole lattice forming a three-dimensional giant molecule. Thus, these kinds of structures have very high melting points and boiling points.

Example: Diamond

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Diamond is an allotrope of carbon and has a different structure and properties to graphite. In this structure all electrons are localised within strong covalent bonds throughout the whole of the lattice. Hence, diamond is very hard, with a high melting point and boiling point. Due to electrons localised in bonds it is a poor conductor of electricity.

Example: Graphite

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The structure of graphite is unique. The carbon atoms are arranged inflat, parallel layers. Each layer contains millions of hexagonally arranged carbon atoms.

Each carbon atom is covalently bonded to three other atoms in its layer. The strong covalent bonds account for a high melting point of 37300C.

Between the layers are weak van der Waal's forces, these allow slippage between the layers, this accounts for the use of graphite as a lubricant.

Since only three electrons from each carbon atom is used in forming covalent bonds, the fourth electron is delocalised around the layer. This leads to graphite being a good conductor of electricity.

Simple Molecular

Example: Iodine

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These are governed by the weak Van der Waal's forces holding the molecules in the lattice.

a) Low melting points and boiling points.

b) Non-conductors of electricity, when solid or in solution, unless they react with the solvent to form ions.

c) Insoluble in polar-solvents such as water, unless the molecules themselves are polar.

Rule: Like dissolves like

Iodine does not dissolve in water, because the water molecules attract each other strongly and prefer to stick together rather than mix with iodine molecules. The attraction between water and iodine molecules would be quite weak.

CCl4 does dissolve iodine because the attractions within the solvent are similar in strength to the attractions within the iodine crystal. Hence, I2 and CCl4 molecules are able to mix freely and form new I2 - CCl4 bonds of similar strength.

Hydrogen bonded

Example: Ice

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a) High melting point due to hydrogen bonding in comparison to similar compounds.

b) Less dense as solid than liquid due to solid structure having much free space between molecules. As it melts some H-bonds are broken and the structure becomes less ordered. Molecules are able to pack closer together.

Note: A hydrogen bond is an electrostatic attraction between the poorly shielded proton of the hydrogen atom bonded to a small highly electronegative atom, such as N, O or F and a lone pair of electrons on a neighbouring molecule.

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