Chemistry Topic 4: Chemical Bonding and Structure

Ionic bonding and covalent bonding

An ionic bond is formed when electrons are transferred between one atom (metal) to another atom (non metal)

the electrostatic attraction between oppositely charged ions

Ionic compounds for example a crystal of sodium chloride has strong electrostatic forces of attraction between the oppositely charged ions and a lot of energy must be supplied to overcome such strong forces and break the ionic bonds.

the electrostatic forces of attraction between +2 and -2 ions in the magnesium oxide lattice are much stronger than between +1 and -1 ions in the sodium chloride lattice.

in solid form ionic compounds can not conduct electricity as the ions are not free to move, however in molten state in a solution ionic compounds can conduct electricity as the ions can move and carry a current.

the electrostatic attraction between the shared pair of electrons and the nuclei of the atoms making up the bond, each non metal shares an electron with each other to create a pair and therefore full outer shell.

1, 2, 3

single bond < double bonds < triple bonds

single bond > double bonds > triple bonds

the attraction of the two nuclei for two electron pairs is stronger

a type of covalent bond where both electrons come from the same atom

1. make sure the outer atoms have 8 electrons (except hydrogen which has two) 2. if the central atom is from period 2 it should have no more than 8 electrons on its outer shell (generally) 3. if the central atom is from period 3 it may have up to 18 electrons in the outer shell

Lewis drawings for the same compound but with the double bond in different places and lone pairs accordingly

it depends on the number of electron domains in the outer shell of the central atom

1. a lone pair 2. an electron pair 3. the electron pairs that make up a multiple bond i.e. a double or triple bond

the repulsion between electron domains on the valence shell of the central atom. Bonding pairs of electrons repel as far apart as possible however repulsion between bonding pairs of electrons with lone electron pairs is greater.

1. draw a lewis structure 2. count the number of electron domains in the outer shell of the central atom 3. work out the basic shape (Electron domain geometry) 4. work out the actual shape (MDG) by looking at lone pairs and angles.

2 electron pairs, 2 bonding pairs, 0 lone pairs, bond angle 180 e.g. CO2

3 electron pairs, 3 bonding pairs, 0 lone pairs, bond angle 120, e.g. SO3

3 electron pairs, 2 bonding pairs, 1 lone pair, bond angle 117, e.g. SO2

4 electron pairs, 4 bonding pairs, 0 lone pairs, bond angle 109.5 e.g. NH4+

4 electron pairs, 3 bonding pairs, 1 lone pair, bond angle 107 e.g. PCl3

4 electron pairs, 2 bonding pairs, 2 lone pairs, bond angle 105 e.g H2O

polarity is the term used to describe a molecule where two atoms have different electronegativities, causing one atom to have a slightly negative charge and one atom to have a slightly positive charge. This creates a dipole. Additionally, the shape of a molecule has to be asymmetrical for the molecule to be polar.

the more polar the bond

even though C -- O is polar the molecule is symmetrical so the dipoles cancel out.

intermolecular forces are broken when covalent substances are melted or boiled - not the covalent bonds.

intermolecular forces that exist between non-polar molecules. They occur when there is an uneven distribution of electrons in an atom, causing one side to be more negative than the other, attracting oppositely charged areas of other atoms so a temporary dipole is established.

1. the more electrons present in the atom , the stronger the force 2. the greater the relative molecular mass the stronger the forces are.

a type of intermolecular force that exists between polar molecules - a positive dipole in one molecule attracts a negative dipole in another molecule.

true - polar molecules with dipole-dipole IM forces have higher melting and boiling points because the force is stronger.

hydrogen bonding is a type of intermolecular force that occurs between a hydrogen atom and a very electronegative atom (N, O, F) with a lone pair. The force occurs due to the attraction between the negative dipole of the oxygen, nitrogen or fluorine atom in one molecule and a positive hydrogen atom in another molecule.

London Forces < dipole-dipole forces < hydrogen bonding

Normally, a substance will dissolve in a solvent if the intermolecular forces in the solute and solvent and similar e.g. both have dipole-dipole forces. Substances that have hydrogen bonding will dissolve in water as they are able to hydrogen bond to the water.

individual atoms covalently bonded together throughout a lattice structure.

diamond, graphite and buckminsterfullerene

giant covalent structure, covalent bonds between all atoms, tetrahedral arrangement of bonds, hexagonal rings of C atoms. Every carbon atom is bonded to four other carbon atoms.

high melting and boiling point, extremely hard and durable due to strong covalent bonds, does not conduct electricity as there are no delocalised electrons.

giant covalent structure, layers of atoms joined by covalent bonds, London forces between layers. Each Carbon forms three bonds, planar hexagonal rings of atoms. Individual layers of graphite are called graphene.

high melting point because of strong covalent bonds, free electrons as C only forms three bonds, therefore conducts electricity - 1 electron per C atom free to move along the layers. Layers slide over each other so graphite rubs off onto surfaces.

C60 molecular structure, hexagons and pentagons of Carbon atoms. Each Carbon forms 3 bonds, London forces between molecules. Cage structure. Not giant covalent structure as only involves 60 atoms and is instead a covalent molecular structure.

lower melting point than diamond and graphite because only London forces must be broken. Used as catalysts and lubricants.

CO2 - covalent molecular and SiO2 is giant covalent. Therefore, when silicon dioxide melts, the strong covalent bonds require more energy to be broken in comparison to CO2, where less energy is needed to overcome weak London forces when it sublimes.

metals are made up of a regular lattice arrangement of positive ions surrounded by a ‘sea’ of delocalised electrons. Metals have a giant structure.

the electrostatic attraction between the positive ions in the lattice and the delocalised electrons.

Magnesium forms a 2+ ion compared to sodium, which forms a +1 ion. This means that the electrostatic attraction between the ions and the delocalised electrons is stronger in magnesium. The Mg2+ ion has a smaller atomic radius than the Na+ ion and therefore the delocalised electrons are closer to the nucleus of the positive ion in magnesium and are more strongly attracted.

1. conductors of electricity (free electrons) 2. malleable (layers of metal ions slide over each other without affecting bonding as the delocalised electrons are attracted by the ions in all directions) 3. can form alloys which are useful in construction

alloys are homogenous mixtures of two or more metals, or of a metal with a non-metal e.g. steel (Fe and C)

a different sized atom will prevent planes of smaller metal atoms sliding over each other, making the overall metal less bendy/soft.

to decide which Lewis structure is most appropriate for the molecule/compound

FC = (no. of valence electrons in the uncombined atom) - 1/2(no. of bonding electrons) - (no. of non-bonding electrons)

the value closest to 0.

from the axial (head-on) overlap of atomic orbitals.

mostly along the axis joining the two nuclei.

the sideways overlap of parallel p orbitals.

the electron density lies above and below the internuclear axis.

5 electron domains, 5 bonding pairs, 0 lone pairs, 90 and 120 degrees, e,g, PF5

5 electron domains, 4 bonding pairs, 1 lone pair, <90 <120, e.g. SF4

5 electron domains, 3 bonding pairs, 2 lone pairs, <90, BrF3

5 electron domains, 2 bonding pairs, 3 lone pairs, 180, e.g. I3-

6 electron domains, 6 bonding pairs, 0 lone pairs, 90, e.g. SF6

6 electron domains, 5 bonding pairs, 1 lone pair, <90, e.g. SF5-

6 electron domains, 4 bonding pairs, 2 lone pairs, 90, e.g. XeF4

the mixing of atomic orbitals (s and p) on a particular atom to produce a new set of orbitals that have both s and p character and are better arranged in space for covalent bonding.

sp

sp2

sp3

mixing of 1s and 1p orbital on each Carbon

mixing of 1s and 2p orbitals on each Carbon

mixing of 1s and 3p orbitals

the sharing of a pair of electrons between three or more atoms

The Lewis strucutre of O3 would suggest one short double O bond and one longer O-O bond, however the bond lengths are equal as the two electrons in the pi bond are delocalised over the three O atoms.

2 sigma bonds and 1 pi bond shared across 2 electron domains = 3/2 = 1.5

UV radiation is absorbed by O2 and O3 as they undergo dissociation (where their bonds are broken). O2→ O and O3→ O + O2, O2 absorb higher frequnecy radiation (more damaging) as it has a pi bond which makes the molecule stronger.

chlorofluorocarbons (CFCs) such as CCl2F2 are broken down by absorbing UV radiation in the upper atmosphere, forming Cl2.

Cl2 has unpaired electrons causing it to take place in a free radical chain reaction which uses up ozone and regenerates the free radical