AQA A Level Chemistry - Bonding

• Metallic bonding is the strong electrostatic attraction of positive metal cations surrounded by a sea of delocalised electrons

• Lithium forms 1+ ions • When lithium bonds with other lithium atoms, electrons are delocalised (given up) into a cloud of electrons • There is a strong electrostatic attraction between the positive Li+ cations and the delocalised electrons • The more delocalised electrons there are and the higher the charge density, the stronger the electrostatic attraction • Al and My have higher melting points than Na because they have more delocalised electrons and a higher charge density - so the electrostatic attraction between their positive metal cations and delocalised electrons is stronger • The smaller the ion, the stronger the bond

• Mg has stronger metallic bonding than Na and hence a higher melting point • The metallic bonding gets stronger because in Mg there are more electrons in the outer shell that are released to the sea of electrons • The Mg ion is also smaller and has one more proton • There is therefore a stronger electrostatic attraction between the positive metal ions and the delocalised electrons and higher energy is needed to break bonds • Magnesium ion (Mg2+) has a higher charger than sodium ion (Na+) • This attracts the delocalised sea of electrons more strongly

Examples: - Sodium chloride - Magnesium oxide

• Diatomic molecules eg Cl2, F2 • Molecules that only contain C and H • Molecules where all the outer atoms are the same eg PCl5 • Noble gases

• A covalent bond is a shared pair of electrons

• Ionic bonding is the strong electrostatic attraction between oppositely charged ions

• Physical properties: • High melting point and boiling point • Soluble/dissolves well • Electrical conductivity when molten or in solution • Structural properties: • Brittle - they shatter easily

• K2S has a giant ionic lattice structure so is solid at room temperature because a large amount of energy is needed to break the strong electrostatic attraction between oppositely charged ions • SF2 has a simple covalent/molecular structure so is a gas at room temperature because a small amount of energy is needed to break the weak intermolecular forces between SF2 molecules

• Covalent: non metal + non metal • Ionic: metal + non metal

• Electronegativity is the power of an atom to attract the pair of electrons in a covalent bond

• The nuclear charge (number of protons) • The atomic radius (ie distance from nucleus to outer e-) • Shielding (number of electrons between nucleus and outer e-)

• The nuclear charge increases • The atomic radius decreases • The shielding remains the same • Therefore, the ability to attract electrons in a covalent bond increases

• Molecules made of atoms with no difference in electronegativity have their electrons distributed evenly • When there is a difference in electronegativity within a molecule, electron distribution is asymmetrical • This produces a polar covalent bond • Therefore the atom with a greater electron density is given a delta- charge and the atom with a lower electron density is given a delta+ charge

• Hydrogen bonding is the strongest intermolecular forces • Occurs between H (bonded to N, O or F) and lone pairs on N, O or F atom on another molecule

• Very large different in electronegativity between the oxygen and hydrogen (state which atoms) • creates a dipole on the O-H bond (state the bond) • Lone pair on oxygen atom in one molecule strongly attracts a partially positive hydrogen atom on a different molecule

• Generally weaker than hydrogen bonding • Occurs between polar molecules

• Difference in electronegativity leads to bond polarity • Dipoles do not cancel therefore the molecule has an overall permanent dipole • There is an attraction between delta positive on one molecule and delta negative on another

• Generally the weakest force but can be stronger than both hydrogen bonding and pdd if the molecule is large • This is due to the number of electrons increasing within the molecule • Occurs between all molecules (and atoms of noble gases) but is the most important force for non polar molecules (molecule that do not have the other IMF’s)

• Random movement of electrons in one molecule (atom) • Leads to a uneven distribution of electrons, creating a temporary dipole in one molecule (atom) • This induces a dipole in a neighbouring molecule (atom) • Dipoles attract