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GCE AS Chemistry

• Course details• Total 5 units (2 x AS, 3 x A2 units)

• AS units 1 and 2 • 90 min each worth 80 marks (20%)

• A2 units 3 and 4• 1hr45 exam worth 25% each

• Practical unit (2 tasks) 10% qualification

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AS UNIT 1

1.1 Formulae and equations

1.2 Basic ideas about atoms

1.3 Chemical Calculations

1.4 Bonding

1.5 Solid structures

1.6 Periodic Table1.7 Simple

equilibria and acid base reactions

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Unit 11.4 Chemical Bonding

Lesson 7

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Contents

• Ionic bonding • Covalent bonding

ØSimple moleculesØCo-ordinate bondingØElectronegativity &

bond polarity• Metallic bonding

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Ionic Bonding

a) ionic bonding in terms of ion formation and the interaction between positive and negative ions in an ionic crystal

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Recap: Ion Formation

Positive ions

• also known as cations, are smaller than the original atom.

• formed when electrons are removed from atoms.

• the energy associated with the process is known as the ionisation energy

• Na(g) —> Na+(g) + e¯

Negative ions

• known as anions

• are larger than the original atom due to electron repulsion in outer shell

• formed when electrons are added to atoms

• energy is released as the nucleus pulls in an electron

Metals lose electrons to form positive ions (CATIONS)Non-metals gain electrons to form negative ions (ANIONS)The charge on the ion is equal to the number of ions lost or gained

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Ionic BondDefinition: a bond formed by the electrostatic force of attraction between positive and negative ions (cations and anions) as a result of electron transfer.

Ionic bonding usually occurs when a METAL bonds with a NON – METAL.

The transfer of electrons occurs to allow the atoms to achieve a full outer shell of electrons. (noble gas electronic configuration)

This increases stability.

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electron transfer

Sodium ion, Na+, 2.8 Chloride ion, Cl-, 2.8.8

Sodium atom. 2.8.1 Chlorine atom 2.8.7

Formation of sodium chloride (NaCl)

1s2 2s2 2p6 3s1

1s2 2s2 2p6

1s2 2s2 2p6 3s2 3p5

1s2 2s2 2p6 3s2 3p6

Na ClXX

X

X

XX XX

XX

XX

XXX

XX

2.8.1 [2.8]+

+

Na ClXX

X

X

XX XX

XX

XX

XX X

XX

-

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Activity 1: Show the formation of magnesium fluoride (MgF2)

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Magnesium ion, Mg2+, 2.82+ Fluoride ion, F-, 2.8− Fluoride ion, F-, 2.8−

Formation of magnesium fluoride (MgF2)

2+- -

Activity 1: Answer

M g F- F-

Magnesium atom

2.8.2

Fluorine atom 2.7

Fluorine atom 2.7

M g

F

F

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Attractive and repulsive forces

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

Normally ions form a lattice where each cation (+) is surrounded by anions and vice versa to maximise attraction and minimise

repulsion.

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GIANT IONIC CRYSTAL LATTICE

• Oppositely charged ions held in a regular• 3-dimensional lattice by electrostatic attraction• The arrangement of ions in a crystal lattice

depends on the relative sizes of the ions

• Each Na+ is surrounded by 6 Cl¯(co-ordination number = 6)

• each Cl¯ is surrounded by 6 Na+

(co-ordination number = 6).

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Activity 2Explain why caesium chloride is 8:8 co-ordinated

and sodium chloride is 6:6 co-ordinated.

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Activity 2 Answer

Sodium ions are smaller than caesium ions (fewer electron shells).

Therefore a sodium ion can only fit 6 chloride ions around it before the chloride ions come into contact increasing repulsion.

Caesium is larger and can fit more chloride ions around it without increasing repulsion.

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Physical properties of ionic compounds

Melting point - very high

A large amount of energy must be put in to overcome the strong electrostatic attractions and separate the ions.

Strength - Very brittleAny dislocation leads to the layers moving and similar ions being adjacent. The repulsion splits the crystal.Electrical - don’t conduct when solid

Ions are held strongly in the lattice conduct when molten or in aqueous solution - the ions become mobile and conduction takes place.Solubility - Insoluble in non-polar solvents but soluble in water

Water is a polar solvent and stabilises the separated ions.

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Ionic charge and melting & boiling points

Because the electrostatic force of attraction is strong, ionic compounds are solids with high melting and boiling points.

The strength of the attraction depends on the size and charges of the ions.

The electrostatic force increases when the charge on the ion increases and the size of the ion decreases.

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Example: Why does CaF2 have a higher melting point than CaCl2?

CaF2 has a higher melting point than CaCl2 because:-

1) The charges on the ions are the same

2) The fluoride ion has a smaller radius than the chloride ion

Therefore because the fluoride ion is smaller than the chloride ion then there is a greater force of attraction between the calcium ion and the fluoride ions compared to the calcium ion and the chloride ions.

More energy is needed to overcome the stronger force of attraction and so the melting point for CaF2 is higher than for CaCl2.

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Activity 3 Explain the differences in melting

points of the following:

NaCl 1074KKCl 1043KMgO 3125K

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Activity 3 Answer

• NaCl/KCl – the Na+ ion is smaller than K+ ion

• Charges the same

• Greater force of attraction for smaller ion therefore NaCl requires more energy than KCl

• MgO – in MgO both ions have a higher charge than ions in NaCl or KCl

• In Mg2+ and O2- higher charge – greater forces of attraction therefore higher melting point

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Covalent Bonding

• = a pair of SHARED electrons with opposed spin shared between two atoms with each atom giving one electron.

• We show electrons as dots and crosses but really bonds are formed by the overlap of atomic orbitals – regions of electron density• This forms a region in space where an electron pair can be found; new

molecular orbitals are formed.• the electrons in the covalent bond repel each other - their attractions to both

nuclei overcome this repulsion. • If atoms get too close, nuclei and inner electrons repel those in the other atom,

therefore bonds have a certain length. • Electrons must spin in opposite directions for bonds to form.

+--

+--

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Properties of Simple Molecules

vStrong covalent bonds (intramolecular)vWeak intermolecular forces (between molecules) Therefore it is easy to

separate the molecules.

Overlapping orbitals

Be careful! Covalent bonds are very strong and high temperatures are needed to separate.

It is the bonds between the simple molecules that are easily overcome.

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Activity 4: Draw dot and cross diagrams for:

H2

HCl

H2O

NH3

CH4

PCl5

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Activity 4 Answers

H2

HCl

H2O

NH3

CH4

PCl5

XH H

XH Cl X

X

X X

XXH O

X X

X X

H

XXH C

XX

H

H

H

XXH N

X

H

H

XX

X

ClX

XX

XX

XP

X

ClXX

X X

X X

X

XClX

X

X X

X X ClXX

XX

X X

X

ClX

XX

XXX

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Double and triple bonds

• When atoms share 4 electrons in the overlapped region a double covalent bond is formed.

• When atoms share 6 electrons in the overlapped region a triple covalent bond is formed.

Activity 5:Draw dot and cross diagrams for

O2

CO2

N2

C2H4

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Activity 5 Answers

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Co-ordinate bonding

Co-ordinate bonding occurs when one of the atoms contributes both electrons.

One of the atoms must have a lone pair of electrons (electrons in the outer shell that have not been used in bonding)

The other atom must have an empty orbital (short of their maximum)

Once formed the bond acts as any other covalent bond

New AS Content!

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Example: Ammonium ion, NH4+

• The lone pair of electrons on N is used to share with the hydrogen ion which needs two electrons to complete its outer shell.

The N now has a + charge as it is now sharing two electrons rather than owning them.

X

X

H

NX

H

H

XX

H+

X

X

H

+NX

H

H

XX

H

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Example: Boron trifluoride and ammonia NH3BF3

The B atom has an incomplete shell in BF3 and is electron deficient. There is room for another pair. It can accept a lone pair of electrons donated by ammonia. A co-ordinate bond is formed. The B becomes - as it shares a pair of electrons (it is up one electron) it didn’t have before.

X

X

H

NXH

H

XX

X

X

F

B XF

F

XX

X

XX

XX

XX

XXX

XX

X

XXX

X

X

H

+NXH

H

XX

X

X

F

B- XF

F

XX

X

XX

XX

XX

XXX

XX

X

XXX

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Example: Oxonium ion formation

X

XH

O

XX

XX

H

H+

X

XH

O

XX

XX

HH

H2O + H+ à H3O+

+

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Activity 6• Define a co-ordinate bond

• Draw the bonding in the formation of an ammonium ion

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Activity 6 Answer

• Co-ordinate bonding occurs when one of the atoms contributes both electrons. One of the atoms must have a lone pair of electrons, the other must have an empty orbital (be electron deficient).

X

X

H

NX

H

H

XX

H+

X

X

H

+NX

H

H

XX

H

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Electronegativity

The ability of an atom to attract the electron pair in a covalent bond to itself

In a covalent bond the electron pair not usually shared evenly between the 2 atoms unless they're the same.

So one atom will have a slightly positive charge and one will have a slightly negative

New AS Content!

H𝛿 +

F𝛿 −

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Bond Polarity

• The overall molecule is neutral, one end is slightly + and one slightly –• This gives a DIPOLE• Small charges are written over the atoms using symbols (delta +/-)• Co-ordinate bonds are always polar• Bond polarity is governed by difference in EN of the two atoms

forming the bond• On Pauling’s EN scale Fluorine is 4.0 and Cs has a value of 0.7

H𝛿 +

F𝛿 − New AS Content!

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Pauling Scale

• a scale for measuring electronegativity• values increase across periods• values decrease down groups• fluorine has the highest value

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Summary:

Non-polar bond -atoms that have the same electronegativity will both pull on the electrons equally so electrons are equally shared H H

Polar bond - different atoms have different electronegativities.a dipole is formed as the electron pair is pulled towards the more electronegative atom. The bond is said to be polar.

greater electronegativity difference = greater polarity

Pauling Scale -a scale for measuring electronegativity

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Metallic Bonding: positively charged

metal ions surrounded by a

sea of delocalisedelectrons

The metallic bond is the force of attraction between the lattice of positive ions and the delocalisedelectrons.

The electrons come from the outer shell of the atom.

Atoms arrange in regular close packed 3-dimensional crystal lattices.

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Properties: can you explain the

properties using the information on the previous slide?

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Electrons are free to move and carry charge and heat through the metal

Ions are the same size and can slide over each other when force applied so metal can be shaped. Electrons act as a lubricant between layers of ions and allow layers to slide over each other making metal ductile.

Metallic bonds are strong and a lot of energy is needed to melt or reshape them so they have HIGH mpt and bpt.

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Activity 7: complete the table of properties in your notes

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Melting point trends:

Melting point is a measure of how easy

it is to separate individual particles

Metals generally have high melting points generally

because they have giant structures.

There is a strong electrostatic force of attraction between

the delocalisedelectrons and the

positive metal ions, therefore a large

amount of energy is required to overcome this and separate the

ions.

In metals melting point is a measure of

how strong the electron cloud holds

the positive ions.

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The strength of the metallic bond increases as:

The charge on the positive ion increases.

The number of delocalised

electrons per atom increases.

The size of the positive ion decreases.

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METALLIC BOND STRENGTH EXPLAINED

Each atom donates one electron to the cloud So the strength of the metallic bonding in sodium is relatively weak

The metallic bonding in potassium is weaker than in sodium because the ion is larger. The delocalised electron cloud has a bigger volume to cover so it is less effective at holding the ions together.

Each atom has donated two electrons to the cloud. The greater the electron density holds the ions together more strongly so themetallic bonding in magnesium is stronger than in sodium

Na

Mg

K

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What happens to melting points as you go across a period?

MELTING POINT INCREASES ACROSS THE PERIOD

• Number of delocalised electrons increases (electron density) which = stronger electrostatic attraction between the positive ions and electrons.

• The size of the ion decreases and the charge increases (electrons held closer to the nucleus). The greater charge is spread over a smaller area.

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What happens to the melting point in metals as you go down a group?

• MELTING POINT decreases DOWN A GROUP

• Size of ion increases down the group (electrons further from nucleus)

• Charge remains the same and is spread over a larger area. (same number of delocalisedelectrons)

• Therefore less electrostatic attraction between the positive ions and the electrons.

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Activity 8: What 3 things does the strength of a metallic bond depend on?

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Activity 8 Answer

the charge on the positive ion

The number of delocalised electrons per atom

The size of the ion

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Activity 9:Which metal will have the highest melting point

and why? Sodium, Magnesium or Aluminium?

Think about metallic bonding, outer shell electrons, forces of attraction.

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Activity 9 answer

Aluminium

Number of delocalised electrons increases (electron density)

The size of the ion decreases

the charge on the ion increases (electrons held closer to the nucleus).

The greater charge is spread over a smaller area which = stronger electrostatic attraction between the positive ions and electrons.

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