Covalent bonding. Learning Objectives Candidates should be able to: describe, including the use of...

AS Chemistry

Covalent bonding

Learning Objectives

Candidates should be able to:

describe, including the use of ‘dot-and-cross’ diagrams, covalent bonding, as in hydrogen; oxygen; chlorine; hydrogen chloride; carbon dioxide; methane; ethene.

describe covalent bonding in terms of orbital overlap.

Starter Activity

Forming a bond

Shared pair of electrons

Electron density maps for a hydrogen molecule

Electron density map

This single σ covalent bond can be simply represented as:

or H – H

Representing a covalent bond

Overlap of atomic orbitals

Overlap of one s and one p orbital – e.g. HF

Overlap of two p-orbitals – e.g. F2

Dot-cross diagrams

Oxygen

EtheneHydrogen chloride

Carbon dioxide

MethaneChlorine

Unexpected structures !!

Breaking the octet rule

Bonding in CH4 – promotion of an electron

C 1s2 2s2 2p2

Bonding in CH4 – hybridisation

Hybridisation in PCl5

AS Chemistry

Co-ordinate bonding

Learning ObjectivesCandidates should be able to describe, including the use of ‘dot-and-cross diagrams, co-ordinate (dative covalent) bonding, as in the formation of the ammonium ion and in the Al2Cl6 molecule.

A co-ordinate bond (also called a dative covalent bond) is a covalent bond (a shared pair of electrons) in which both electrons come from the same atom.

Starter Activity

Dot-cross diagrams

H2S SF6SiCl4

CH3OH CO

Reaction between NH3(g) and HCl(g)

Reaction between NH3(g) and HCl(g)

Reaction between H2O and HCl

Electron-deficient BF3

Aluminium chloride vapour

Aluminium chloride vapour

AS Chemistry

Electronegativity

Learning Objectives

Candidates should be able to explain the origin of polar bonds, with reference to electronegativity differences between atoms.

Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons.

Electronegativity values

Starter Activity

Why are bonds like bears?

…’cos some of them are POLAR!

What happens if two atoms of equal electronegativity bond together?

What happens if B is slightly more electronegative than A?

Electronegativity differences

Representing polar bonds

Molecule Electronegativity difference

Dipole/debye

HCl 0.9 1.03

HBr 0.7 0.78

HI 0.4 0.38

Representing polar bonds

What happens if B is a lot more electronegative than A?

An ionic bond is formed!!!

Large electronegativity difference

Type of bond

Type of bond Electronegativity difference

Non-polar Covalent 0.5

Polar covalent Between 0.5 and 1.7

Ionic 1.7

As a rough guide:

Ionic with increasing covalent character

Positive cation Negative anion

Polarisation of Anions

Truly ionic

Ionic with some covalent character

Property Cation is the most powerful polarising

agent when....

Anode is most easily polarised

when...

Charge

RadiusHigh

Small

High

Large

Polarisation of Anions

Do polar bonds make polar molecules?

CCl4 molecule is tetrahedral - the partial negative charges on the Cl atoms are distributed pretty symmetrically around the molecule. The partial positive charge on the C is buried in the center of the molecule.

The most electronegative element is fluorine.If you remember that fact, everything becomes easy, because electronegativity must always increase towards fluorine in the Periodic Table.

Trends in electronegativity

AS Chemistry

Ionic bonding

Learning Objectives

Candidates should be able to describe ionic (electrovalent) bonding, as in sodium chloride and magnesium oxide, including the use of ‘dot-and-cross’ diagrams.

“The name’s Bond, Ionic Bond – taken, not shared!!!

Starter Activity

Approaching atoms

Ionic bonding

Ionic bonding

MgO CaCl2

Al2O3

K2O

Why CaCl2 and not CaCl or CaCl3?

Why CaCl2 and not CaCl or CaCl3?

Positive cation Negative anion

Polarisation of Anions

Truly ionic

Ionic with some covalent character

Property Cation is the most powerful polarising

agent when....

Anode is most easily polarised

when...

Charge

RadiusHigh

Small

High

Large

Polarisation of Anions

AS Chemistry

Shapes of molecules

Learning Objectives

Candidates should be able to explain the shapes of, and bond angles in, molecules by using the model of electron-pair repulsion.

Starter Activity

180o Linear BeCl2

Starter Activity

Balloon molecules

2 Balloons give a linear geometry

3 Balloons give a trigonal planar

geometry

4 Balloons give a tetrahedral geometry

5 Balloons give a trigonal bipyramidal

geometry

6 Balloons give an octahedral geometry

Electrons in outer shell of central atom

Electrons added from other atoms (and any charge on an ion)

No. of pairs of electrons

No. of bonding pairs

Diagram of molecule (including

bond angles)

Description of shape

BF3 3 3 3 3Trigonal planar

CCl4 4 4 4 4 Tetrahedral

NH3 5 3 4 3Trigonal

pyramidal

H2O 6 2 4 2Bent (or V-

shaped)

Shapes of molecules

SF6 6 6 6 6 Octahedral

CO2 4 4 2 2 Linear

C2H6 4 4 4 4 Tetrahedral

C2H4 4 4 3 3Trigonal planar

ClF4- 7 5 6 4

Square planar

Shapes of molecules

AS Chemistry

Metallic bonding

Lesson Objectives

Candidates should be able to describe metallic bonding in terms of a lattice of positive ions surrounded by mobile electrons.

Starter Activity

This is sometimes described as "an array of positive ions in a sea of electrons".

Metallic Bonding

Close packed structures?

Dense metals

Group 1 metals

Malleability

Metal grains

AS Chemistry

Intermolecular forces

Lesson Objectives

Candidates should be able to describe intermolecular forces (van der Waals’ forces), based on permanent and induced dipoles, as in CHCl3(l), Br2(l) and the liquid noble gases.

Starter Activity

Particles in solids, liquids and gasesIn a liquid or a solid there must be forces between the molecules causing them to be attracted to one another, otherwise they would move apart from each other and become a gas.

Intermolecular attractions are attractions between one molecule and a neighbouring molecule.

Intermolecular forces

The lozenge-shaped diagram represents a small symmetrical molecule - H2, perhaps, or Br2. The even shading shows that on average there is no electrical distortion (i.e. the molecule is non-polar).

How do intermolecular (or van der Waals) forces arise?

Temporary or instantaneous dipoles

Temporary or instantaneous dipoles

Electrons are mobile. The constant "sloshing around" of the electrons in the molecule causes rapidly fluctuating dipoles even in the most symmetrical molecule.

It even happens in monatomic molecules - molecules of noble gases, like helium, which consist of a single atom.

Temporary dipole - induced dipole interaction

How temporary dipoles give rise to intermolecular attractions

This diagram shows how a whole lattice of molecules could be held together in a solid using van der Waals’ forces.

An instant later, of course, you would have to draw a quite different arrangement of the distribution of the electrons as they shifted around - but always in synchronisation.

van der Waals’ forces

van der Waals lattice

helium -269°C

neon -246°C

argon -186°C

krypton -152°C

xenon -108°C

radon -62°C

How molecular size affects the strength of the dispersion forces

The boiling points of the noble gases are:

How molecular size affects the strength of the dispersion forcesThere is a gradual increase in the very low boiling temperatures of the noble gases with increasing atomic size.

As the size of the atoms increases the number of electrons increases and the magnitude of the van der Waals forces increases.

How molecular shape affects the strength of the temporary dipole interactions

Butane has a higher boiling point because the intermolecular forces are greater. The molecules are longer and can lie closer together than the shorter, fatter 2-methylpropane molecules.

Permanent dipoles

Permanent dipoles

AS Chemistry

Hydrogen bonding

Lesson Objectives

Candidates should be able to describe hydrogen bonding, using ammonia and water as simple examples of molecules containing N-H and O-H groups.

Starter Activity

The increase in boiling point happens because the molecules are getting larger with more electrons, and so van der Waals forces become greater.

Boiling points of the Group 4 hydrides

Boiling points of the hydrides in Groups 5, 6 and 7.

The origin of hydrogen bonding

Hydrogen bonding is a particularly strong intermolecular force that involves three features:

a large dipole between an H atom and the highly electronegative atoms N, O or F;

the small H atom which can get very close to other atoms;

a lone pair of electrons on another N, O or F, with which the positively charge H atom can line up.

The origin of hydrogen bonding

Drawing hydrogen bonds

1 mark for indicating bond polarity

1 mark for showing lone pair

1 mark for showing H-bond

Hydrogen bonding accounts for many of the other unusual properties of water including:

its high specific heat capacity

its very high surface tension

its high viscosity and

the low density of ice compared to water

Hydrogen bonding in water

Which type of intermolecular force?

AS Chemistry

A summary

Bonding, structure and properties

Lesson Objectives

Candidates should be able to:describe, interpret and/or predict the effect of

different types of bonding on the physical properties of substances.

describe, in simple terms, the lattice structure of a crystalline solid which is ionic, simple molecular, giant molecular, hydrogen-bonded and metallic.

suggest from quoted physical data the type of structure and bonding present in a substance.

Starter Activity

The properties of substances are decided by their bonding and structure.

Bonding means the way the particles are held together: ionic, covalent, metallic or weak intermolecular bonds.

Structure means the way the particles are arranged relative to one another. You have already met the major types of structure at IGCSE.

Bonding, structure and properties

Bonding, structure and properties

Bond Average bond enthalpy/kJmol-1

Bond length/nm

C – C +347 0.154

C = C +612 0.134

C ≡ C +838 0.120

C – H +413 0.108

O – H +464 0.096

C – O +358 0.143

C = O +805 0.116

Bond energies

GIANT LATTICE COVALENT MOLECULARIonic Covalent network Metallic Simple molecular Macromolecular

What substances have this type of structure?

Compounds of metals with non-metals.

Some elements in Group 4 and some of their compounds.

Metals Some non-metal elements and usually some non-metal/non-metal compounds.

Polymers

Examples NaCl SiO2 Cu H2O Poly(ethene)What type of particle does it contain?

ions atoms positive ions and

delocalised electrons

molecules molecules

How are the particles bonded together?

Strong ionic bonds; attraction between oppositely charged ions

Strong covalent bonds; attraction of atoms’ nuclei for shared electrons

Strong metallic bonds; attraction of atoms’ nuclei for delocalised electrons

Weak intermolecular bonds between molecules; strong covalent bonds between atoms within each molecule.

What are the typical properties?

M. pt and b.pt.

high very high generally high lowmoderate (often

decompose on heating)

Hardness hard but brittle

very hard (if 3D)

hard but malleable

soft variable

Electrical conductivity

conduct when (l) or (aq)

do not normally conduct

conduct when (s) or (l)

do not conduct do not normally conduct

Solubility in water

often soluble insoluble insoluble (but some react)

usually insoluble (but

some H-bond)

sometimes soluble

Solubility in non-polar solvents (e.g. hexane)

insoluble insoluble insoluble usually solublesometimes

soluble

Structure table

AS Chemistry

The modern use of materials

Lesson ObjectivesCandidates should be able to:Explain the strength, high melting point and

insulating properties of ceramics in terms of their giant molecular structure.

Relate the uses of ceramics to their properties.

Describe and interpret the uses of the metals aluminium and copper (and their alloys) in terms of their physical properties.

Understand that materials are a finite resource and the importance of recycling processes.

Starter Activity

Five most common metals

Aluminium

Copper

Zinc

Steel

Brass

Low density, corrosion resistance and strength make it ideal for construction of aircraft, lightweight vehicles, and ladders.

Malleability, low density, corrosion resistance and good thermal conduction make it a good material for food packaging.

Good electrical conduction, corrosion resistance and low density leads to its use for overhead power cables hung from pylons (low density gives it an advantage over copper).

Uses of aluminium

Uses of copperCopper is an excellent conductor of electricity and heat.

Copper is soft and malleable.

Copper is very unreactive and therefore corrosion resistant.

Copper Alloy

Other metal it contains

Main properties

Uses

BrassZinc

Fairly soft and

malleable

Screws and hinges

BronzeTin Strong

Propellors and

bearings

Alloys of copper

A mineral is a naturally occurring solid formed through geological processes that has a characteristic chemical composition, a highly ordered atomic structure, and specific physical properties.

Ceramics

Ceramic: Any of various hard, brittle, heat-resistant and corrosion-resistant materials made by shaping and then firing a nonmetallic mineral.

Ceramicsfurnace: an enclosed chamber in which heat is produced to heat buildings, destroy refuse, smelt or refine ores, etc.

heat shields glass and crockery

furnace linings

brake pads electrical insulators

Raw materials extracted (removed by chemical means) from the Earth cannot last forever. Although some materials are more (present in great quantity) abundant than others, they are all finite (have a limit) resources. Increasing demand for raw materials (items used to produce something else), coupled with ever growing problems of waste disposal, have led to considerable interest in recycling (processing for reuse) waste.

Recycling has a number of possible advantages (beneficial factors):

It leads to reduced demand for new raw materials;

It leads to a reduction in environmental damage (harm to the surroundings);

It reduces the demand for landfill sites (a place for burying waste) to dump waste;

It reduces the cost of waste disposal;

It may reduce energy costs.

Recycling

AS Chemistry

The kinetic-molecular model of

liquids

Lesson Objectives

Candidates should be able to describe using a kinetic-molecular model, the liquid state; melting; vaporisation and vapour pressure.

Starter Activity

Solid Liquid GasArrangement of particles

very orderly short-range order, longer range disorder

almost complete disorder

Movement of particles

vibrate about fixed positions

some movement from place to place

continuous, rapid, random movement

Proximity of particles

close (~10-10m) close (~10-10m) far apart (~10-

8m)Compressibility of substance

very low very low high

Conduction of heat

poor except metals and graphite

metals very good; others poor

very poor

The Kinetic-Molecular Model of Liquids

The Kinetic-Molecular Model of LiquidsLiquids do not have a fixed __________ because the particles can move about. However, they remain very __________ together. This shows that the inter-particle forces have not been __________ broken. If sufficient __________ is supplied, the particles overcome the inter-particle forces almost completely and __________ from the liquid. This is called __________ or boiling. The energy required to boil a liquid is always __________ than that required to melt the same substance and is a better __________ of the strength of inter-particle forces.

Vapour pressure

Vapour pressure is the pressure of a vapour over a liquid at equilibrium.

Vapour pressure

Even at low temperature there are particles with high energy.

Vapour pressure

At equilibrium, the rate at which molecules leave the liquid equals the rate at which molecules join the liquid.

Measuring vapour pressure

AS Chemistry

More on ideal gases

Lesson Objectives

Candidates should be able to explain qualitatively in terms of intermolecular forces and molecular size the limitations of ideality at very high pressures and very low temperatures.

Starter Activity

More on ideal gases

More on ideal gasesY is hydrogen. It is closest to ideal under all conditions. Hydrogen has the weakest intermolecular forces and is the smallest molecule.

Z is ammonia. It is the least ideal at lower pressures. Ammonia molecules can hydrogen bond.

X is nitrogen. Deviates greatly from ideality at high pressures where its larger molecular volume becomes important.