Topic 3
Chemical Bonding

describe covalent bonds in terms of orbital overlap giving σ and π bonds

describe that σ bonds are formed by direct overlap of orbitals between the bonding atoms

describe that π bonds are formed by the sideways overlap of adjacent p orbitals above and below the σ bond

describe how the σ and π bonds form in H2

describe how the σ and π bonds form in C2H6

describe how the σ and π bonds form in C2H4

describe how the σ and π bonds form in HCN

describe how the σ and π bonds form in N2

use the concept of hybridisation to describe sp orbitals

use the concept of hybridisation to describe sp2 orbitals

use the concept of hybridisation to describe sp3 orbitals

define the term bond energy as the energy required to break one mole of a particular covalent bond in the gaseous state

define the term bond length as the internuclear distance of two covalently bonded atoms

use bond energy values and the concept of bond length to compare the reactivity of covalent molecules

state and explain the shape of, and bond angles in, BF3 (trigonal planar, 120°) by using VSEPR theory

state and explain the shape of, and bond angles in, CO2 (linear, 180°) by using VSEPR theory

state and explain the shape of, and bond angles in, CH4 (tetrahedral, 109.5°) by using VSEPR theory

state and explain the shape of, and bond angles in, NH3 (pyramidal, 107°) by using VSEPR theory

state and explain the shape of, and bond angles in, H2O (non-linear, 104.5°) by using VSEPR theory

state and explain the shape of, and bond angles in, SF6 (octahedral, 90°) by using VSEPR theory

state and explain the shape of, and bond angles in, PF5 (trigonal bipyramidal, 120° and 90°) by using VSEPR theory

predict the shapes of, and bond angles in, molecules and ions analogous to those specified in 3.5.1

describe hydrogen bonding, limited to molecules containing N–H and O–H groups, including ammonia and water as simple examples

use the concept of hydrogen bonding to explain the anomalous property of H2O (ice and water): its relatively high melting and boiling points

use the concept of hydrogen bonding to explain the anomalous property of H2O (ice and water): its relatively high surface tension

use the concept of hydrogen bonding to explain the anomalous property of H2O (ice and water): the density of the solid ice compared with the liquid water

use the concept of electronegativity to explain bond polarity

use the concept of electronegativity to explain dipole moments of molecules

describe van der Waals' forces as the intermolecular forces between molecular entities other than those due to bond formation

use the term van der Waals' forces as a generic term to describe all intermolecular forces

describe instantaneous dipole–induced dipole (id–id) forces, also called London dispersion forces, as a type of van der Waals' force

describe permanent dipole–permanent dipole (pd–pd) forces, including hydrogen bonding, as a type of van der Waals' force

describe hydrogen bonding

understand that hydrogen bonding is a special case of permanent dipole–permanent dipole forces between molecules where hydrogen is bonded to a highly electronegative atom

state that, in general, ionic, covalent and metallic bonding are stronger than intermolecular forces

use dot-and-cross diagrams to illustrate ionic bonding including the representation of any compounds stated in 3.4 and 3.5

use dot-and-cross diagrams to illustrate covalent bonding including the representation of any compounds stated in 3.4 and 3.5 (dot-and-cross diagrams may include species with atoms which have an expanded octet or species with an odd number of electrons)

use dot-and-cross diagrams to illustrate coordinate bonding including the representation of any compounds stated in 3.4 and 3.5