Bonding and structure

If there are 2 factors eg LDF, pd-pd then say no of electrons outweigh presence / strength of pd-pd interactions

ionic bonding

the electrostatic force of attraction between oppositely charged ions

lattice

the regular three-dimensional arrangement of atoms / ions in a crystal (lattice)

factors affecting strength of ionic bond

• charge of ions

• distance between ions / ionic radii

F ∝ Q1 x Q2 / d²

properties of ionic compounds

• hard, rigid, brittle - layer snaps if it is shifted / displaced so that like charges come on top of one another

• high melting and boiling points

• soluble in water / polar solvents, insoluble in non-polar solvents - hydrating ions is exothermic

• do not conduct electricity when solid; conducts electricity when molten or dissolved in solution

migration of ions

manganate ions attracted to positive anode (purple)

CrO4 2- chromate ions attracted to positive anode (yellow)

covalent bond

the electrostatic force of attraction between the two nuclei and the shared pair of electrons

electron density map

shared electron density = overlapping / joined lines

types of structure

giant or simple molecular

octet rule

tendency to form an outer shell of 8 electrons - associated with stability

sigma bond

• single bond

• the overlap of atomic orbitals between the axis between two nuclei (end-to-end overlap)

• stronger than pi bonds - shared electrons held closer to nuclei + more direct overlap

pi bond

the overlap of singly occupied p orbitals above and below the axis between two nuclei / plane of nuclei (side to side overlap)

lone pair

a pair of electrons in the outer shell not used in bonding

dative covalent bond

a bond between 2 atoms where both of the electrons in the shared pair are donated by the same atom

expansion of the octet

• from P onwards (period 3)

• able to hold the extra electrons in the empty / easily accessibly d subshell / has available d orbitals

• close in energy to p subshell

• endothermic - compensated by making more bonds

bond length

distance between the nuclei of two bonded atoms in a molecule

factors affecting bond length

inversely proportional to bond strength

• atomic radii / no of electron shells

• nuclear charge

• single / double bonds - larger no of electrons shared = stronger force of attraction

properties of simple molecular substances

• low melting and boiling points

• do not conduct heat or electricity

• more soluble in non-polar solvents

electron-pair repulsion theory

electron pairs repel to be as far apart as possible to maximise separation to minimise repulsion (only for when no lone pairs)

If lone pairs say electron pairs arrange themselves to minimise repulsion

shapes of molecules and bond angles

• 2 bonding pairs: linear, 180

• 3 bonding pairs: trigonal planar, 120

• 4 bonding pairs: tetrahedral, 109.5

• 5 bonding pairs: trigonal bipyramidal, 120 / 90 (in triangle one normal line one hatch one wedge)

• 6 bonding pairs: octahedral, 90

• 3 bonding pairs, 1 lone pair: trigonal pyramidal, 107

• 2 bonding pairs, 2 lone pairs: v-shaped / bent, 104.5 (say bent instead / both)

• 4 bonding pairs, 1 lone pair: see-saw, 172, 102 (2 normal lines on the outside 1 hatch 1 wedge)

• 4 bonding pairs, 2 lone pairs: square planar, 90

• 5 bonding pairs, 1 lone pair: square pyramidal, 88

isoelectronic

having the same number and arrangement of electrons

why do lone pairs repel more than bonding pairs

lone pairs: fatter, shorter, more concentrated distribution of negative charge held closer to the nucleus

bonding pairs: thinner, longer electron distribution

exam answer about shapes of molecules

• number of bonding pairs / regions / lone pairs

• electron pairs repel to be as far apart as possible to maximise separation to minimise repulsion

• lone pairs repel more strongly than bonding pairs, compressing the bond angles / bonding pairs repel each other equally

• therefore the molecule adopts the shape / angle

** or if lone pairs, electron pairs repel and arrange themselves to minimise repulsion

electronegativity

the ability of an atom to attract the shared pair of bonding electrons in a covalent bond

Pauling scale - 0 to 4

most electronegative: fluorine (4.0), oxygen, nitrogen

factors affecting electronegativity

• nuclear charge

• atomic radius

• shielding

increases across a period - increased nuclear charge, decreased atomic radius, similar shielding

decreases down a group - increased atomic radius, shielding, outweighs increased nuclear charge

difference in electronegativity

indicates type of bonding

• small difference in electronegativity 0 - 0.4: covalent bonding

• medium difference in electronegativity 0.4 - 1.7: polar covalent bonding

• large difference in electronegativity >= 1.7: ionic bonding - electrons are completely transferred to the more electronegative atom

polar molecules

• contain polar bonds

• asymmetrical molecule (non-identical bonds / presence of lone pairs) - bond dipoles do not cancel out

London (dispersion) forces

1. fluctuations in electron density around the molecule leads to a temporary / instantaneous dipole

2. induce a dipole in a neighbouring molecule

3. weak force of attraction between dipoles

present in all molecules

factors affecting LDFs

• number of electrons in the molecule - larger no. / electron cloud = more polarisable as a greater chance of asymmetry in electron cloud / more easily distorted

• shape of molecule - long, straight molecules have a larger surface area of contact between molecules

permanent dipole - permanent dipole interactions

• attraction between terminal dipoles are in addition to london forces

• electropositive pole of a molecule is attracted to the electronegative pole of another molecule

hydrogen bonding

• H is covalently bonded to an oxygen, fluorine or nitrogen atom - electronegative enough to polarise H so it is very electron deficient

• O, F, N must have an available lone pair of electrons

• 180 degrees around the central H atom

• The bond is polar, X is more electronegative than H, force of attraction from electropositive H to lone pair of electrons on electronegative X, say specifically which bond is involved / the molecule has

low density of ice

• each water molecule forms 4 hydrogen bonds with adjacent molecules

• hydrogen bond lengths are relatively long (and longer than covalent bonds)

• molecules are arranged in an open / hexagonal lattice structure

• held further apart than in water / in water molecules get closer

boiling points of hydrides

H2O > HF > NH3

• H-F is the most polar bond - stronger hydrogen bonds

• H2O able to form 2 hydrogen bonds per molecule vs 1 hydrogen bond per molecule of HF / NH3

large decrease down the group then increase

• no H bonds - weaker intermolecular forces

• number of electrons increase down the group

solubility

‘Like dissolves like’ - solute will only dissolve if:

energy required to break the intermolecular forces within the solute and solvent < energy released when intermolecular forces are made between molecules of solute and solvent

• non-polar substances disrupt the H bonding in water + reforms weak intermolecular forces

hydration of ions

• exothermic

• energy released from hydrating ions > energy required to break ionic bonds + intermolecular forces in water

miscible

two liquids that mix with each other

allotropes

different structural forms of the same element in the same physical state

fullerenes

• C60 - molecular

• each C forms 3 bonds with adjacent C atoms

• cannot conduct electricity as delocalised electrons from C atoms are unable to move between molecules

properties of metals

• dense - regularly arranged in a metallic lattice structure

• high melting and boiling points

• malleable and ductile - layers can slide over each other / slip without breaking metallic bonds

• good conductors of heat and electricity

factors affecting boiling points of metals

• charge on ions - higher charge = stronger force of attraction + more delocalised electrons holding the structure together

• ionic radius / shielding

• nuclear charge