WJEC AS Chemistry Unit 1.4 and 1.5 - Bonding and Solid Structures

Ionic bond

• Electrostatic attraction between oppositely charged ions

• Metal + non-metal

• Predominantly groups 1 and 2 with 6 and 7

Process of ionic bonding

• metal atom loses electrons and becomes a cation (+ve)

• Non-metal atom gains electrons and becomes an anion (-ve)

Covalent bond

• pair of electrons shared between two atoms, with each atom giving one electron, forming a bond pair in which the electron ships are opposed

• Non-metal + Non-metal

• Includes diatomic molecules; HNFOICB

Forces of attraction and repulsion in covalent molecules

• the electrons in the pair between the atoms repel one another but this is overcome but their attractions to both nuclei

• If atoms get too close together the nuclei and inner electrons will repel those of the other atom, so the bond has a certain length

• The electron spins must be opposite for the bonds to form

Forces of attraction and repulsion in ionic bonding

• Cations and anions are arranged so that each cation is surrounded by several anions + vv to maximise attraction and minimise repulsion

• Repulsions from inner electrons and nuclei prevent the ions from getting too close together

Coordinate bonding

Both electrons are contributed by the same atom

Intermediate character of ionic and covalent bonds

• the degree of ionic and covalent character and properties depends on the difference in electronegativity between the bonded atoms

Electronegativity

• The ability to attract electrons in a covalent bond

• Higher electronegativity = better the element can attract bonding electrons

• No units as relative value

• N, O and F are the 3 most e.n. elements

• F = most as high charge and small atomic radius

• Depends on;

  • nuclear charge (e.n. is higher as proton number increases) therefore increases across a period

  • atomic radius (smaller = higher e.n.) a.r. decreases slightly across a period and increases down a group

  • group 8/0 not considered as does not participate in covalent bonding

Polar bonds

• Polar = one end of bond with a slightly positive charge, other end with a slightly negative charge

• bonding electrons are pulled towards more electronegative atom

• Partial charged written above atom with symbol delta

• Equal electronegativity = electrons equally shared = non-polar bond

• Difference in e.n. < 0.4 = non-polar covalent bonds

• Difference 0.4-1.9 = polar covalent bonds

• Difference 2.0 or more = ionic

Intermolecular bonding

• Weak bonding holding the molecules together

• Governs the physical properties of the substance

• Three types

• Much weaker than covalent and ionic bonds

Intermolecular bonding

• strong bonding between atoms in the molecules

• Governs its chemistry

Dipole-dipole forces

• Polar molecules have dipoles; one end slightly positive charge, one end slightly negative charge due to difference in e.n.s between the atoms in the molecule

• There will be an attraction between them if they arrange them selves to that the negative region of one molecule is close to the positive region of another molecule

• Forms a permanent dipole-dipole force

No dipole-induced dipole forces

• A molecule with no dipole (symmetrical distribution of electron cloud)

• The + end of the molecule can pull the electron cloud of a neighbouring molecules towards it, giving the left side of that molecule a - charge

• Induced a temporary dipole in the neighbouring molecule

• The two dipoles are attracted to each other

Temporary dipole-induced dipole interactions

• A temporary dipole induces an induced dipole in a neighbouring molecule

van der Waals forces

• Weakest form of intermolecular forces

• Includes dipole-dipole and temporary dipole-temporary dipole forces

Induced dipole-Induced dipole and the effect on physical properties

• Strength of forces increases with increasing number of electrons int he molecules

• The more electrons in a molecule, greater the fluctuation in the electron cloud around the nuclei + the larger the temp + induced dipoles created

• —> stronger forces between the molecule

• Shown in the boiling temperatures of the noble gases or the halogens

Hydrogen bonding

• Strongest intermolecular forces

• Occur between molecules containing hydrogen atoms bonded to small, very electronegative elements which have lone pairs - Fluorine, Oxygen or Nitrogen

• F, O + N = highly electronegative elements

• The + charge in the bonded H atom is spread over a small volume + so has a high charge density

• Highly Polarising + H atom then attracts a lone pairs of e-s from a small highly en atom in another molecule

Effect of hydrogen bonding on physical properties

• Solubility;

  • the most significant IMFs between water molecules = H bonds

  • Covalent compounds that can replace these bonds by forming new H bonds with water will dissolve

• Boiling Temperatures;

  • H bonds = strongest IMFs

  • Molecules that form H bonds have a higher boiling temperatures than molecules of a similar size that cannot H bond

VSEPR principle

• The shape of a molecule or ion is governed by the number of pairs of electrons in the outer (valence) shell of the central atom

• The electron pairs arrange themselves around the central atom as far apart as possible from each other to minimise repulsion between them

• lp:lp repulsion > lp:bp repulsion > bp:bp repulsion

Use of the VSEPR Principle in predicting the shapes of simple molecules and ions

Geometry of molecules with two bond pairs around the central atom

• The bps of e-s arrange themselves at 180 degrees to each other

• Geometry; linear

Geometry of molecules with 3 bond pairs around the central atom

• The bps = 120 degrees to each other

• Trigonal planar

Geometry of molecules with 4 bond pairs around the central atom

• 109.5 degrees

• Tetrahedral

Geometry of molecules with 5 bond pairs around the central atom

• Three of the atoms are in a plane at 120 degrees to each other

• The other two atoms are at 90 degrees to this plane

• Trigonal bipyramidal

Geometry of molecules with 6 bond pairs around the central atom

• 90 degrees

• Octahedral

Geometry of molecules with 3 bond pairs and 1 lone pair

• The strong repulsion between an lp and a bp forces the bps together, slightly reducing the bonding angel between them to 107 degrees

• Trigonal pyramidal

Geometry of molecules with 2 bond pairs and 2 lone pairs

• lp:lp repulsion forced the bps closer together and reduce angel between them to 104.5 degrees

• V-shaped

Simple structures

• weak intermolecular forces

• strong intramolecular forces

• Examples; water, carbon dioxide, hydrogen gas

• Gases or volatile (low bp) liquids

• Covalent in groups 1, 2, 3, 5, 6, 7 (not 0 (full os) or 4 (giant covalent))

Giant structures

• strong forces of attraction

• Solids at room temperature

• Ionic compounds adopt a giant structure

• Group 4 covalent —> giant covalent (e.g. diamond and graphite)

• Metallic —> giant structure

Ionic solids

• Giant lattices of positive and negative ions

• Structure of the crystal depends on the relative number of ions + their sizes

• Made of the same base unit repeated

• Ions are arranged so the electrostatic forces of attraction between the oppositely charged ions is greater than the electrostatic repulsion between ions with the same charge

Sodium chloride

• Consists of sodium ions and chloride ions

• Each Na+ ion is surrounded by 6 Cl- ions + vv

• Coordinate number of each ion = 6

• Face centred cubic structure

Caesium chloride

• Caesium ion is larger than the Sodium ion, therefore more Cl- ions can fit around it

• Each Cs+ ion is surrounded by 8 Cl- ions + vv

• Coordination number of each ion = 8

• Adopts a body centred cubic lattice structure

Reason for the difference in structure between sodium chloride and caesium chloride

• Different cation size

• rCs+ > rNa+ and both have a 1+ charge

• Charge in Na+ is more concentrated due to smaller size

• Electrostatic attraction is stronger in NaCl

• Caesium is smaller than Sodium

Physical properties of ionic compounds

• Relatively high melting points (higher than simple covalent, lower than giant covalent);

  • giant lattices are held together by strong esa between opp charged ions. Takes a large amount of energy to overcome these forces of attraction

• Non-Conductors/poor electrical conductivity;

  • conduction requires moving negatively charged particles

  • In the solid state, ions are in fixed positions by the strong ionic bonds

• Electrolytes;

  • When molten/dissolved, ions are free to move and will move conduct electricity

• Most are soluble in water;

  • Water molecules = polar

  • Oxygen ends attracted to negative ions

  • Hydrogen ends attracted to negative ions

  • Like dissolved like; a charged solute dissolves in a charged solvent

Diamond

• Giant tetrahedral structure

• Each C atom is covalently bonded to 4 others

• Bonding forces are uniform throughout the structure

Physical Properties of Diamond

• Very high melting temperature; energy need to break strong c.b.s = very high

• Extremely hard; strength of the c.b.s + the geometrical rigidity of the structure

• Insoluble in water; no ions to attract the polar water molecules

• Poor conductor of electricity; no delocalised electrons or ions present

Graphite

• Hexagonal layer structure

• Each layer; C is joined to 3 others by strong cobalt. Bonds

• 4th electron from each Carbon atom = delocalised within the layer

• Hexagonal layers held together by weak induced dipole-induced dipole forces

Physical properties of Graphite

• Very high melting temperature; strong c.b.s in the hexagonal layers

• Soft + slippery; weak forced between layers easily broken, layers can slide over each other

• Insoluble in water; no ions to attract the polar water molecules

• Good conductors of electricity; delocalised electrons are free to move along the layer so electricity can flow. Delocalised electrons not free to move one layer to the next so it can only conduct parallel to its layers

• Low density; relatively large amount of space between layers so length of c.b.s is much shorter than the length of the vdW forced between the layers

Physical properties of giant molecular substances

• Very high melting temperature due to strong covalent bonds

• Insoluble in water as there are no ions (sometimes no delocalised electrons) to attract the polar water molecules

Iodine

• Solid; molecules held in a lattice by weak intermolecular forces

• Atoms covalently bonded in pairs to form diatomic iodine molecules

• Held together by weak can der Waals forces

Ice

• Giant tetrahedral structure containing stronger intermolecular hydrogen bonds

• Molecule of water arranged in rings of six

• Water molecules = further apart than they are in the liquid state. The structure —> large areas of open space inside the rings + therefore ice = less dense than liquid water

Metallic bonding

• Electrostatic attraction between delocalised electrons and the nucleus of the cations

• Metals = lattice

Physical properties of metals

• High melting temperatures; larger energy needed to overcome stronger forces of attraction between the nuclei of the metal cations and the delocalised electrons. Melting affected by number of dl e-s per cation + the size of the cation

• Hard; metallic bond is very strong

• Insoluble in water; no ions to attract the polar water molecules

• Good conductors of electricity in both solid and molten state; delocalised electrons can carry electricity

• Good thermal conductors; delocalised electrons can pass KE to each other

• Malleable + ductile; when a force is applied to a metal the layers of cations can slide over each other

Physical properties of simple molecular substances

• Low melting and boiling temperatures; intermolecular forces are weak and don’t need much energy to break

• Soft; weak IMFs between the molecules are easily broken

• Normally insoluble in water; no ions to attract the polar water molecules. Compounds that can form H bonds with water will be soluble

• Poor conductors of electricity; do not contain delocalised electrons or ions