Types of bonds

1. Metallic Solids2. Ionic Solids3. Molecular Solids4. Covalent network solids

Gallium

What is a metallic bond ?

both atoms have low ionization energies and low electronegativities and will lose electrons easily (i.e. 2 metals)

In metallic bonding, positive metal ions are arranged with valence electrons delocalized around them.

Since the electrons are delocalized, they are mobile and able to move throughout the metal structure.

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Metallic solids

In metals, the valence electrons of neighbouring atoms form a sort of ‘electron soup’

This ‘delocalized electron soup’ can be thought of as a ‘glue’ that holds the positive nuclei of the metal atoms together.

Metal solids can be thought of as being like a rice krispy square

The marshmallow is like the ‘delocalized electron soup’ that acts a glue that holds it together

The rice krispies are like the positive nuclei of the metal atoms that compose the solid.

Properties of Metallic Compounds

1. Malleable and ductile – no fixed bond and the ions can roll past each other.

2. Good conductors of heat and electricity because the valence e- are mobile and can transmit energy rapidly.

3. Shiny because when light strikes a metal, the valence electrons absorb energy, oscillate at the same frequency as the incident light (incoming light) and then emit light as a reflection of the original light.

4. Solids at room temperature (except mercury) due to the strong bonds (intermolecular and intramolecular forces are the same).

5. Hardness – metals with many, small crystals is harder (less uniform layers) but more brittle (more boundaries around crystals)

6. Melting & boiling points –* larger +ve charge across period – higher b.p/m.p.*down group – bigger n – less attraction*the more available e- in electron sea, the stronger the attraction (stronger bond)

Properties of Metallic Compounds

IONIC BONDS

Occur between an element with low ionization energy (a metal) and an element with high ionization energy (a non-metal).

An actual transfer of electrons from the metal (becoming a cation) to the non-metal (becoming an anion) occurs.

This transfer results in the formation of 2 oppositely charged ions.

The electrostatic interaction between the 2 ions holds the compound together.

Atoms are arranged in a highly ordered crystal lattice structure which maximizes the attractive forces between oppositely charged ions and minimizes the repulsion between like charged ions (the crystal structures are determined experimentally using X-ray crystallography).

Different Types of Crystal Lattices

No, you don’t need to know these! They’re just neat lookin’

Arrangement of Ions in a Sodium chloride cystal (cubic)

Properties of Ionic Compounds

Solids at room temperature because the attraction between the ions is very strong.

The intramolecular forces are the same as the intermolecular forces (i.e. the electrostatic attraction of oppositely charged ions)

Mostly soluble in water (the ions will dissociate in water) – can vary from high to very low

Properties of Ionic Compounds

Brittle – when impact makes anions/cations align – repulsion

Will conduct electricity as a liquid or aqueous solution because the ions are free to move to oppositely charged electrodes.

High m.p. and b.p. – the greater the charge on the ions, the higher the m.p./b.p.

Covalent bonds

COVALENT BONDS

Occur between 2 atoms with high ionization energies (i.e. 2 non-metal atoms).

sharing of electrons to obtain a full outer energy level (octet rule).

If the electronegativity difference between the two atoms is less than 0.4, the bond is a true covalent bond and the electrons are shared equally (for example?).

Covalent bonds

If the electronegativity difference between the two atoms is between 0.4 and 1.7, the bond is polar and there is an unequal sharing of electrons.

If a molecule contains polar covalent bonds and is asymmetrical, the molecule will be polar.

If both electrons being shared come from the same atom, the bond is a coordinate covalent bond.

Properties of Non-Polar Covalent Compounds

Have low melting and boiling points and are usually gases at room temperature. This is due to the low intermolecular forces that exist in these molecules (London Dispersion Forces only).

If solid at room temperature, the solid is usually soft and waxy.

Soluble in non-polar solvents such as ethers. Will not conduct electricity in any form due to

the fact that there are no ions present.

Properties of Polar Covalent Molecules

Dipoles - greater intermolecular forces than non-polar covalent compounds (the presence of dipole-dipole forces and possibly H-bonding).

higher melting and boiling points and are more likely to be liquids or solids at room temperature (may even exhibit a crystal lattice like sugar).

Will dissolve in polar solvents if H-bonding is present (sugar in water).

Will not conduct electricity to any appreciable degree (only ionize to a very small degree)

Covalent Network Solid

Consider carbon dioxide (CO2) and silicon dioxide (SiO2).

What would you expect the physical properties of SiO2 to be?

CO2 b.p. = -78.5oC If London forces are the only

intermolecular force, then you might predict the b.p. of SiO2 to be slightly more than CO2.

Covalent Network Solid

However SiO2 has a m.p. of 1650oC, and b.p. of 2230oC!

SiO2 is also known as quartz, or sand, is used in the production of glass.

Clearly, SiO2 is not a molecular solid like CO2 is.

SiO2 is a covalent network solid.

Raw silica (SiO2)

Silica glass (SiO2)

Covalent Network Solid

A covalent network solid is a solid that consists of atoms held together in large networks or chains by covalent bonds.

Every atom is covalently bonded forming a 1,2 or3-dimensional network

Examples include: diamond, graphite, silicon, asbestos

Network Solids

Generally strong bonds - very high melting and boiling points, solids at room temperature, not soluble in polar or nonpolar solvents and they do not conduct electricity.

Can be soft (2-d networks like graphite) or hard (3-d networks like diamonds) or fibres (1-d networks like asbestos)

Allotropes: elements that exist in different physical forms with different physical properties but the same chemical properties.

For example graphite and diamond are allotropes of carbon (as is “Buckey ball”). Both form carbon dioxide and water when undergoing combustion but only graphite conducts electricity and diamond is one of the hardest substances known.

Diamond

Diamond is another allotrope of carbon.

Like graphite, it is a covalent network solid.

However, instead of sheets, it forms a 3-dimensional lattice of carbon atoms.

This is what gives diamond its characteristic hardness.