Topic 4: Chemical Bonding and Structure

atoms interact by transferring/sharing electrons to achieve stability

• aka electrovalent bond - established by the transfer of electrons from a metallic atom to a non-metallic atom

• no of valence electrons: 1st atom should have 1-3 valence e’s while 2nd atom should have 5-7 valence e’s • net lowering of energy: energy released as a result of electron transfer and formation of ionic compounds • higher lattice energy: the greater the energy, the greater the strength of an ionic bond • a big difference in electronegativity

the amount of energy released when one mole of an ionic compound crystal is formed from its cations and anions in gaseous state

• solids at room temp: strong electrostatic forces of attraction between locked ions in crystal lattice • brittle: ions of the same charge are beside each other, so the repulsive forces cause it to split • high m.pt/b.pt: high temps required to overcome attractive forces • soluble in water: water’s polar properties detach ions from the crystal lattice by their electrostatic pull • good electricity conductors in molten state: as ions are free to move about • low volatility

• bond formed by mutual sharing of electrons between the combining atoms

• equal electronegativity so no transferring of electrons occurs • equal electron affinity to equally attract the electron pair • both atoms should have 5-7 valence electrons

• bond length • bond angle • bond energy

• the average distance between centres of the nuclei of 2 bonded atoms • measured in picometer (pm)/angstrom (Å)

• bond multiplicity: single, double, or triple bonds. it’s inversely proportional to bond length. • size of atom: directly proportional to bond length

the average angle between orbitals containing bonding electron pairs around the central atom in a molecule

• the amount of energy required to break one mole of a certain type between atoms in gaseous state • expressed in kJ/mol • measures bond strength

• bond length: inversely proportional relation - size of bonded atom: inversely proportional relation

• a covalent bond in which electrons are shared unequally • due to unequal electronegativity • bonded atoms acquire a partial positive/negative charge

a covalent bond in which both electrons in the shared pair of electrons come from the same atom

Valence Shell Electron Pair Repulsion Theory • the geometry of a molecule is dependent on the no. of valence e-s around the central atom • electron pairs tend to repel each other as e- clouds are negatively charged, so pairs try to stay as far apart to have min. energy/max. stability

when the molecule’s characteristic properties can be described by 2+ structures, the real structure is a resonance hybrid of the possible structures

• Van der Waals/London Dispersion Forces • dipole-dipole forces • hydrogen bonding

ionic bonds > hydrogen bonds > dipole-dipole > London forces

• weak force existing in all covalent compounds • relatively stronger between easily-polarised molecules than between molecules that don’t polarise easily • due to temporary dipole caused by random electron movement • has an inductive effect on neighbouring molecules

• no. of e-s: higher no. of e-s = higher distance between valence e-s and nucleus = lower attraction to nucleus = electron cloud is more easily polarised • size/vol of electron cloud: the bigger the cloud, the higher the polarisability of electrons • shape of molecule: straight chain molecules can interact with each other across the molecule

the ease of distortion of the electron cloud of a molecule by an electric field

• occurs due to electrostatic attraction between molecules with permanent dipoles • exists in all polar molecules with permanent dipole movement • b.pt increases significantly

occurs as a result of interaction between a non-bonding electron pair on one of the atoms with a H atom in a different molecule that carries a high partial positive charge