Metallic,+Ionic+and+Covalent+Solids

Intermolecular forces
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Weak interactions
Covalent molecular forces media type="youtube" key="8qfzpJvsp04" height="315" width="560"

Strong interactions
Covalent Netwokrs, Ionic and Metallic Crystals media type="youtube" key="PU9rzTjLyb4" width="560" height="315"

EARTH AND THE NATURE OF SOLIDS
All substances can exist in the solid state, although some require a much colder temperature for this to be possible. The different solids can vary significantly in their physical properties (melting point, solubility in different solvents, conductivity in the solid and liquid states, density, malleability etc). It is important to realise that all substances can be categorised into four major groups and that the properties within each group are very similar. The reason they are similar is that the nature of the particles and the type of bonding within each group is much the same (irrespective of whether a particular substance is in the solid, liquid or gaseous state).

The 4 categories are  **1) ionic**
 *  2) metallic **
 *  3) covalent network **
 *  4) molecular **

The first 3 types of substances are all extended 3-dimensional networks with strong bonding between the particles. This means they all have relatively high melting points and generally exist as solids at room temperature. A molecular substance is quite different - it exists as discrete molecules with only weak forces between the particles so that they have low melting points and often exist as liquids or gases at room temperature. A more detailed description of each category follows:

 H2O NaCl(//s//) ®  Na+(//aq//) + Cl - (//aq//)
 * 1) Ionic - **these substances generally contain metal ions and nonmetal ions e.g. NaCl, MgO, or compounds of polyatomic ions e.g. K2CO3 and NH4NO3. The formula is an empirical formula showing the ratio of ions in the crystal.
 * The compounds are solids at room temperature, made up of positive and negative ions (the particles) held together by strong electrostatic attraction between the ions (ionic bonds). Because this bonding is so strong, the solids have high melting points and boiling points.
 * In the solid stateionic compounds do not conduct electricity as the ions are held in their lattice positions and are not free to move. If the substance is in the molten (liquid) state then it will be able to conduct electricity due to the movement of the ions. Similarly an aqueous solution of an ionic solid will conduct electricity.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Ionic substances are insoluble in non-polar solvents such as cyclohexane. They are soluble in water (a polar solvent) although in some cases the solubility is only small (e.g. CaCO3). When an ionic compound dissolves in water, the ions in the crystal lattice become separated and surrounded by water molecules forming hydrated ions (see p6).
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Ionic solids are hard but brittle as any attempt to distort the crystal results in ions of like
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">charge coming into contact. The strong repulsive forces that result push the crystal apart.


 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">2) Metallic - **<span style="font-family: Arial,sans-serif; font-size: 11pt;">a metallic substance (a metal) is made up of metal atoms held together by strong metallic bonds. Metal atoms only have loosely held valence electrons which readily move between atoms so that more specifically, a metallic bond is described as the electrostatic attraction between the positive metal ions and the “sea of mobile electrons”.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Metals are all solids (except mercury) and have relatively high melting points because the metal atoms are linked by strong metallic bonds.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Metals conduct heat and electricity in the solid and liquid state because the negatively charged valence electrons are free to move.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Metals are generally insoluble although some react with water producing hydrogen gas.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Metalsare malleable and ductile as distortion does not alter the strength of the metallic bonds. (These are sometimes referred to as non-directional bonds).
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Metals have high densities as the atoms are closely packed together.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">3) Covalent network solids - **<span style="font-family: Arial,sans-serif; font-size: 11pt;">this is a small group of substances, which are made up of linear chains, 2-dimensional layers or 3-dimensional networks. They are made up of atoms linked by strong covalent bonds, which are not easily broken.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">The **linear chains** (as in rubber, plastic and some allotropes of sulfur) only have **weak** van der Waals forces between the chains, so the substances are flexible and have low to medium melting points.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Note: **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> It is the weak intermolecular forces **between** the chains that are broken when the substance melts - not the strong covalent intramolecular bonds within the chains themselves.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Graphite is an allotrope of carbon that is made up of **2-dimensional layers**. Within each layer the C atoms are linked by strong covalent bonds but between the layers there are only weak van der Waals forces and delocalised electrons. This means graphite will conduct electricity in the solid state (the only non-metal to do so) and is commonly used to make electrodes for electrical circuits. Because the layers can slide over each other, graphite is commonly used as a lubricant and as the “lead” in pencils. Graphite does have a high melting point and is insoluble in polar solvents such as water.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Carbon (in the form of diamond), silicon and silica (silicon dioxide, SiO2) all exist as
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">3-dimensional covalent network solids. **<span style="font-family: Arial,sans-serif; font-size: 11pt;">All atoms in the lattice are linked by strong covalent bonds so that these substances are rigid, strong and all have very high melting points (as the strong covalent bonds have to be broken before the substance can melt). They are very insoluble. As they have no charged particles that are free to move (either electrons or ions) they do not conduct heat or electricity either in the solid or liquid states.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">4) **Molecular -** A molecular substance is made up of discrete particles (atoms e.g. Ne, or molecules e.g. CO2 and H2O) held together in the solid and liquid states by weak van der Waals forces. These van der Waals forces are the electrostatic attraction between the slightly positive end of 1 molecule and the slightly negative end of another.   d <span style="font-family: Arial,sans-serif; font-size: 11pt;">+ d <span style="font-family: Arial,sans-serif; font-size: 11pt;">-  <span style="font-family: Arial,sans-serif; font-size: 11pt;"> positive H - Cl negative <span style="font-family: Arial,sans-serif; font-size: 11pt;"> end end <span style="font-family: Arial,sans-serif; font-size: 11pt;">Within any molecule the atoms are linked by strong covalent bonds, but the process of melting or boiling does not alter the structure of the molecule or break any covalent bonds. <span style="font-family: Arial,sans-serif; font-size: 11pt;">. <span style="font-family: Arial,sans-serif; font-size: 11pt;">Because the intermolecular van der Waals forces are only weak, molecular substances have low melting and boiling points. The actual m.p.and b.p. values depend on:


 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">The mass of the molecule – a molecule with a higher mass has a larger cloud of electrons. It is the movement of the electron cloud that can result in instantaneous dipole-dipole attractions. The larger electron cloud results in stronger van der Waals forces and a higher m.p. and b.p. For example, methane, CH4, and propane,C3H8, are both non-polar but propane has a higher b.p. due to the larger electron cloud.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">The polarity of the molecule - the more polar the molecule the stronger the van der Waals forces and the higher the m.p. and b.p. Thus ethanol, CH3OH will have a higher boiling point than ethane CH3CH3 because, although they are similar in molar mass, the ethanol is more polar than ethane because of the presence of the polar O-H bond.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Molecular substances do not conduct electricity as they do not have electrically charged particles (atoms or ions) that are free to move.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Polar molecular substances dissolve in polar solvents (e.g. ethanol dissolves in water) whereas non-polar substances dissolve in non-polar solvents (e.g. wax, a hydrocarbon, and solid iodine both dissolve in cyclohexane and not in water).

<span style="font-family: Arial,sans-serif; font-size: 11pt;">This solubility behaviour is referred to as “like dissolves like”.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">If an electrically charged object is placed near a stream of a polar liquid, the polar molecules will align themselves so that the oppositely charged end of the molecule is attracted to the rod. Non-polar molecules are not charged ans there is no deflection of the liquid stream. This deflection of the liquid can be used to test whether the liquid is polar or not.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Crystalline solids have sharp melting points and a definite characteristic shape because their particles are arranged in a definite order. In contrast amorphous solids such as glass, rubber and plastics do not have a sharply defined melting point or regular shape.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">The elements of Group 17 of the Periodic Table all exist as diatomic molecules with (non-polar) covalent bonding strongly holding the pairs of atoms together as a molecule, but only weak van der Waals forces **between** the molecules. The molar mass (or size of the electron cloud) increases down the group and so therefore does the melting and boiling points as the strength of the inter-molecular forces increases. The effect of this on the physical state at room temperature of each halogen is shown in the table below.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Halogens **

<span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;"> at 25 oC || <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">Pale yellow **gas**  || <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">Greenish-yellow **gas**  || <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">Red **liquid**  || <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">Black shiny **solid**  ||
 * **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Halogen ** ||  **<span style="font-family: Arial,sans-serif; font-size: 11pt;">F2 **  ||  **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Cl2 **  ||  **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Brr **  ||  **<span style="font-family: Arial,sans-serif; font-size: 11pt;">I2 **  ||
 * <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">Physical state

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Because they are all non-polar molecules the halogens are more soluble in non-polar solvents such as cyclohexane. They are however slightly soluble in water.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Chlorine forms a pale green-yellow solution while bromine forms an orange solution. Iodine is special as it forms a purple solution in a non-polar solvent but forms a yellow-brown solution in water. The solubility of iodine in water is increased if iodide ions are present due to the equilibrium below that forms the red triiodide ion (which is more soluble in water).

<span style="font-family: Arial,sans-serif; font-size: 11pt;"> I2(//aq//) + I - <span style="font-family: Arial,sans-serif; font-size: 11pt;">(//aq//)  <span style="font-family: Arial,sans-serif; font-size: 11pt;"> I3 - <span style="font-family: Arial,sans-serif; font-size: 11pt;">(//aq//)

<span style="font-family: Arial,sans-serif; font-size: 11pt;">By considering the oxides of the third row of the periodic table it is possible to examine how the pattern of behaviour varies from the left to right. Most metal oxides are basic oxides. The oxides of Group 1 metals, are soluble in water and form solutions of the metal hydroxides e.g. Na2O dissolves to form a solution of NaOH. Other metal oxides are insoluble in water but are still classed as basic oxides as they dissolve in (react with) acid forming a solution of a metal salt. <span style="font-family: Arial,sans-serif; font-size: 11pt;">For example <span style="font-family: Arial,sans-serif; font-size: 11pt;"> MgO(//s//) + 2HCl(//aq//) ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> Mg2+(//aq//) + 2Cl - <span style="font-family: Arial,sans-serif; font-size: 11pt;">(//aq//) + H2O(//l//) <span style="font-family: Arial,sans-serif; font-size: 11pt;">All the metal oxides are ionic salts; they are solids with high melting points and only conduct electricity when molten or in an aqueous solution.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Oxides of the 3rd Row **

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Non-metal oxides are acidic oxides. They obviously are not Bronsted acids as they do not have a proton to donate. Rather when dissolved in water, they react to form acids that dissociate in water producing H3O+ ions.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">For example SO2(//g//) + H2O(//l//) ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> H2SO3(//aq//) <span style="font-family: Arial,sans-serif; font-size: 11pt;"> and H2SO3(//aq//) + H2O(//l//)  <span style="font-family: Arial,sans-serif; font-size: 11pt;"> H3O+(//aq//) + HSO3 - <span style="font-family: Arial,sans-serif; font-size: 11pt;">(//aq//) <span style="font-family: Arial,sans-serif; font-size: 11pt;">The non-metal oxides (oxides of elements in Groups 14 - 17) are all molecular substances. They are substances with low melting points and are often liquids or gases at room temperature.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">The oxides themselves do not contain ions, but an aqueous solution containing the dissociated acid will act as an electrolyte.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Some oxides (of elements in the middle of the periodic table) can act as both an acid and a base. They are called **amphoteric oxides** e.g. Al2O3 (like most metal oxides) is insoluble in water but will dissolve in both acid and base.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">In acid Al2O3(//s//) + 6HCl(//aq//) ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> 2Al3+(//aq//) + 6Cl - <span style="font-family: Arial,sans-serif; font-size: 11pt;">(//aq//) + 3H2O(//l//) <span style="font-family: Arial,sans-serif; font-size: 11pt;">In base Al2O3(//s//) + 2NaOH(//aq//) + 3H2O(//l//) ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> 2[Al(OH)4] - <span style="font-family: Arial,sans-serif; font-size: 11pt;">(//aq//) + 2Na+(//aq//) <span style="font-family: Arial,sans-serif; font-size: 11pt;">Another amphoteric oxide is ZnO. Aluminium and zinc oxides both have properties that are typical of ionic compounds.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">The melting point of these chlorides reduces from left to right as the structure changes from ionic <span style="font-family: Arial,sans-serif; font-size: 11pt;">through polar covalent to non-polar molecules as shown in the table below.
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Chlorides of the 3rd Row **
 * =====<span style="font-family: Arial,sans-serif; font-size: 11pt;">Formula ===== || **<span style="font-family: Arial,sans-serif; font-size: 11pt;">NaCl **  ||  **<span style="font-family: Arial,sans-serif; font-size: 11pt;">MgCl2 **  ||  **<span style="font-family: Arial,sans-serif; font-size: 11pt;">AlCl3 **  ||  **<span style="font-family: Arial,sans-serif; font-size: 11pt;">SiCl4 **  ||  **<span style="font-family: Arial,sans-serif; font-size: 11pt;">PCl3 **  ||  **<span style="font-family: Arial,sans-serif; font-size: 11pt;">SCl2 **  ||  **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Cl2 **  ||
 * **<span style="font-family: Arial,sans-serif; font-size: 11pt;">M.P. (oC) ** || <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">801  || <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">710  || <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">sub 150  || <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">- 70  || <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">- 112  || <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">- 78  || <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">-101  ||
 * **<span style="font-family: Arial,sans-serif; font-size: 11pt;">State ** || <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">Solid  || <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">Solid  || <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">Solid  || <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">Liquid  || <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">Liquid  || <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">liquid  || <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">(Yellow-green) Gas  ||
 * **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Type of solid ** || <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">Ionic  || <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">Ionic  || <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">Ionic  || <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">Molecular  || <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">Molecular  || <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">Molecular  || <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">Molecular  ||
 * **<span style="font-family: Arial,sans-serif; font-size: 11pt;">Nature of bonding ** |||||| <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">Ionic  |||||||| <span style="display: block; font-family: Arial,sans-serif; font-size: 11pt; text-align: center;">Covalent bonding between atoms in the molecules but weak van der Waals forces between the molecules  ||


 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">NOTES: **
 * 1) <span style="font-family: Arial,sans-serif; font-size: 11pt;">1. In the ionic metal chlorides the formula gives the ratio of positive to negative ions in the crystal lattice. It is an empirical formula. In the non-metal chlorides the formula gives the number of non-metal atoms bonded in each molecule.

<span style="font-family: Arial,sans-serif; font-size: 11pt;">2. Only the ionic chlorides will conduct in the **molten** state although all of the aqueous solutions conduct because of the presence of chloride ions. <span style="font-family: Arial,sans-serif; font-size: 11pt;">In the molecular chlorides these ions arise not from the usual dissolving process but from a chemical reaction with the water (called hydrolysis).

<span style="font-family: Arial,sans-serif; font-size: 11pt;">PCl3 + 3H2O ® <span style="font-family: Arial,sans-serif; font-size: 11pt;"> H3PO4 + 3H+ + 3Cl -