Structure and Bonding

• The three main types of chemical bonds are Ionic, Covalent, and Metallic.
• Ionic bonding is the electrostatic attraction between positive and negative ions.
• An example of an ionically bonded substance is NaCl (Sodium Chloride - salt).
• Covalent bonding is the electrostatic attraction between a shared pair of electrons and the nuclei.
• Metallic bonding is the electrostatic attraction between the positive metal ions and the sea of delocalised electrons.
• Electrons in the outer shell are represented in a dot and cross diagram.
• Giant ionic lattices conduct electricity when liquid but not when solid because in solid state the ions are in fixed positions and thus cannot move, while when they are in liquid state the ions are mobile and thus can freely carry the charge.
• Giant ionic lattices have high melting and boiling point because a large amount of energy is required to overcome the electrostatic bonds.
• Ionic lattices dissolve in polar solvents such as water.
• Dative covalent bond is a bond where both of the shared electrons are supplied by one atom.
• Expansion of the octet means when a bonded atom has more than 8 electrons in the outer shell.
• Simple molecular structures have low melting and boiling point because a small amount of energy is enough to overcome the intermolecular forces.
• Triple bond is formed when atoms share three pairs of electrons.
• Oxygen forms 2 covalent bonds.
• Single bond is formed when atoms are bonded by a single pair of shared electrons.
• Carbon forms 4 covalent bonds.
• These charges are able to attract charged ions.
• Types of covalent structure include simple molecular lattice and giant covalent lattice.
• Lone pair refers to electrons in the outer shell that are not involved in the bonding.
• In simple molecular structures, atoms within the same molecule are held by strong covalent bonds and different molecules are held by weak intermolecular forces.
• Average bond enthalpy is a measure of average energy needed to break the bond.
• Oxonium ions are formed when acid is added to water, H3O+.
• Simple molecular structures are non conductors.
• Double bond is formed when atoms share two pairs of electrons.
• The most electronegative element is Fluorine.
• London forces are greater in smaller molecules due to more electrons.
• Permanent dipole-induced dipole interactions and permanent dipole-permanent dipole interactions are two interactions that can be referred to as Van der Waals’ forces.
• There is a strong force of attraction between the H nucleus and the lone pair of electrons on O, N, F.
• Hydrogen bonding is the strongest type of intermolecular force.
• London forces are caused by random movements of electrons and lead to instantaneous dipoles which induce a dipole in nearby molecules, causing induced dipoles to attract one another.
• Conditions needed for hydrogen bonding to occur include an O-H, N-H or F-H bond, a lone pair of electrons on O, F, N, and because O, N and F are highly electronegative, the H nucleus is left exposed.
• In ice, the water molecules are arranged in a orderly pattern, forming an open lattice with hydrogen bonds.
• A polar bond is formed when bonding atoms have different electronegativities.
• Water has a melting/ boiling point higher than expected because hydrogen bonds are stronger than other intermolecular forces, requiring extra strength to overcome the forces.
• In water, the lattice is collapsed and the molecules are closer together.
• Hydrogen bonding and Van der Waals’ forces are the two types of intermolecular forces.
• Boiling point increases because the number of electrons increases and hence the strength of London forces also increases.
• In a pyramidal shape there are 3 bonded pairs and 1 lone pairs, the bond angle is 107
• In a tetrahedral shape there are 4 bonded pairs and 0 lone pairs, the bond angle is 109.5
• In a non-linear shape there are 2 bond pairs and 2 lone pairs, the bond angle is 104.5