Bonding (AS Level) flashcards

Topic 3: Bonding

Ionic bonding

Electrostatic attraction between oppositely charged ions in a lattice.

Co-ordinate (dative covalent) bond

A shared pair of electrons with both electrons supplied by one atom. Drawn as an arrow from the donor atom to the acceptor. Examples: NH₄⁺, H₃O⁺, NH₃BF₃.

Metallic bonding, and what makes it stronger

Attraction between positive metal ions in a lattice and the delocalised electrons around them. Stronger with a bigger ion charge, smaller ions, or more delocalised electrons per atom.

Electronegativity

The ability of an atom to draw the bonding pair of electrons towards itself in a covalent bond. It increases across a period and decreases down a group; fluorine is the most electronegative.

Order of repulsion between electron pairs, and its effect

Lone pair–lone pair > lone pair–bond pair > bond pair–bond pair. Each lone pair squeezes the bond angles.

Shapes and bond angles: CH₄ or NH₄⁺, NH₃, H₂O

NH₄⁺/CH₄: tetrahedral, 109.5°. NH₃: trigonal pyramidal, about 107°. H₂O: bent, about 104.5°. Each lone pair reduces the angle.

Shapes of SF₆, XeF₄, PCl₅ and ClF₃

SF₆: octahedral, 90°. XeF₄: square planar (2 lone pairs opposite each other). PCl₅: trigonal bipyramidal, 120° and 90°. ClF₃: T-shaped (2 lone pairs).

Why is CH₂Cl₂ polar but BF₃ non-polar?

Both have polar bonds. BF₃ is symmetrical (trigonal planar) so the dipoles cancel. CH₂Cl₂ is tetrahedral with two different outer atoms, so its dipoles do not cancel.

The three intermolecular forces, weakest to strongest

Induced dipole–dipole (van der Waals, London) between all molecules; permanent dipole–dipole between polar molecules; hydrogen bonding (H on N, O or F to a lone pair on N, O or F).

Diamond vs graphite: bonding and conductivity

Diamond: each C makes 4 covalent bonds (tetrahedral); no free electrons, so no conduction. Graphite: each C makes 3 bonds in hexagonal layers; one delocalised electron per C, so it conducts. Weak forces between layers let them slide.

What is overcome when a simple molecular solid such as iodine melts?

Only the weak intermolecular forces (van der Waals) between molecules. The covalent bonds inside the molecules do not break.

Why is ice less dense than liquid water?

In ice each H₂O molecule is hydrogen bonded to four others in an open lattice, holding the molecules further apart than in the liquid.