Structure and Bonding West Midlands Chemistry Teachers Centre November 2009 Presenter: Dr Janice Perkins.

Structure and Bonding

West Midlands Chemistry Teachers Centre

November 2009

Presenter: Dr Janice Perkins

Types of Bonding

1. Ionic Bonding

2. Covalent Bonding

3. Metallic Bonding

Ionic Bonding• Negative Ions (anions)

have a negative charge because of a surplus of e-

• Positive Ions (cations)have a positive charge because of a deficiency of e-

• Ions formed by e- transfer from one atom to anotherthe number of e- lost or gained depends on the elements involved

Ionic Bonding is the attraction between these positive and negative ions

It is called ELECTROSTATIC ATTRACTION

Magnesium and chlorine react together to form the ionic compound magnesium chloride, MgCl2

(i) Explain how each of the ions in this compound is formed

(ii)Explain why compounds with ionic bonding tend to have high melting points

Two e- transferred from Mg to Cl

Question from past paper

one e- to each of two Cl atoms

electrostatic attractionsare strong

Too easy? No, it was very poorly answered

Covalent Bond

• Covalent bond = shared pair of e-

• One electron comes from each atom

Two organic compounds with similar relative

molecular masses are shown below.

Ethanol PropaneState the type of bond present between the C and H atoms in both of these molecules. Explain how this type of bond is formed.

H C C

H

H

H

H

OH

H C C

H

H

H

H

C

H

H

H

Question from June 09

Covalent Shared pair of electrons, one from each atom

Co-ordinate bonding is to do with how a covalent bond is formedOnce formed it is a normal covalent bond

Co-ordinate Bonding

also called Dative covalent bonding

One atom/ion supplies both e- to another atom/ion

Question from Jan 09Phosphorus is in the same group of the Periodic Table as nitrogen.

The molecule PH3 reacts with an H+ ion to form a PH4

+ ion.

Name of the type of bond formed when PH3 reacts with H+ and explain how this bond is formed.

•Coordinate/dative

•Both electrons/ lone pair (on P/PH3)

•Shares/donated from P(H3)/ to H(+)

Metallic bonding

• This type of bonding is often very poorly explained.

• frequently see:ionic bondshydrogen bondsvan der Waals’ forces

in the answer of even the better candidates

• This is wrong!

Metallic bonding

+ + + +

+ + + +

+ + + + X

+ + + +

+ + + +

+ + + +

+ + + +

+ + + +

+ + + +

+ + + +

+ + + +

+ + + +

+ + + +

+ + + +

+ + + +

+ + + +

+ + + +

+ + + +

+ + + +

+ + + +

+ + + +

+ + + +

+ + + +

+ + + +

X

+ + + +

+ + + +

+ + + +

+ + + +

+ + + +

+ + + +

• OUTER electrons are ‘DELOCALISED’ i.e. free to move through metal

• Metal lattice held together by the ATTRACTION between this sea of delocalised e- and the positive ions in the lattice

Question from June 09

State the type of bonding involved in silver.

Draw a diagram to show how the particles are arranged in a silver lattice and show the charges on the particles.

• metallic bonding • regular arrangement of same sized particles • + charge in each ion

Bond Polarity Cl x Cl

• A bond is formed between atoms of the same element

• They both attract the electrons in the bond to the same extent

• This attraction is called electronegativity

ElectronegativityDefinition:• The power of an atom to attract electron density (or a pair of electrons) • from a covalent bond.

Trends:Electronegativity increases across the period• Due to more protons and smaller size, so stronger attraction for the bonding e-

Decreases down a group• Due to larger size/more shells, so weaker attraction for bonding e-

• What’s the sporting Caption?

• ‘Throwing the puppy dog!’??

What’s the sporting Caption?

Wish we were on holiday

What’s the sporting Caption?

Well this one is clear enough

It must be ‘Tug-o-War’

Electronegativity and covalent bonds

• Like a tug-o-war

• The ‘tug’ is provided by the electronegativity of the atom

• A tug-o-war between equal teams is like the pull of the Cl atoms in the covalent bond becauseatoms of the same element have same electronegativity.

• perfectly even sharing of the bonding e-

Cl x Cl

Covalent bonds between atoms with different electronegativities

• The attraction for the bonding e- by the two atoms is different

• This results in unequal sharing of the bonding e- between the two atoms

• The extent of the unequal distribution will depend on how different the electronegativities are- the term we use is ‘electronegativity difference’

• The bigger the difference, the greater the dipole

Polar Bonds

Electron density is low.

Electron density is high.

+ -

+A – B -

AB is a polar molecule it has a dipole

Atom B is more electronegative than atom A

Bonding e- are attracted more strongly to B

A x B

Forces acting between molecules

Types of Intermolecular Forces (IMF)• van der Waals’ forces

(temporary induced dipole - dipole attractions)• Permanent dipole - dipole attractions• Hydrogen bondingAll result from attractions between the partial charges in the dipolesThe stronger the dipole – the stronger the IMF

Van der Waals’ forces• It is the weakest of the three IMFs• It result from temporary unequal

distributions of e- density• The bigger the atom or molecule, the more

electrons there will be and the larger will be the surface area – this increases the strength of the van der Waals’ attractions

• It is always present but is often swamped by stronger IMFs

• It is the only IMF present between non-polar molecules

Permanent dipole-dipole forces

• Always refer to them as ‘dipole-dipole’, not just ‘dipole’, attractions

• They result from attractions between the and of polar molecules

• The greater the electronegativity difference between the two atoms in the bond, the greater the dipole

+

-

Hydrogen Bonding• This is the strongest IMF. It is an extreme

example of dipole-dipole attractions

• It only occurs in molecules with very large electronegativity differences.

• N–H, O–H and F–H bonds

• It results in an attraction that is about 10% as strong as a covalent bond

Hydrogen bonding in HFMark 1 = 3 lone pairs on the fluorine atoms

H F H Fδ+

H F δ- δ+

H F δ-δ+

H F δ- δ+

H F δ-

H F H Fδ+

H F δ- δ+

H F δ-

Mark 2 = dipole correctly shown

Mark 3 = hydrogen bond between lp and H

Question from Jan 09

State the strongest type of intermolecular force in the following compounds.

• Methane (CH4) van der Waals

• Ammonia (NH3) Hydrogen bonding

Use the values in the table to explain how the strongest type of intermolecular force arises between two molecules of ammonia

Large electronegativity difference between N + H Forms N - / H + Lone pair on N attracts H (+)

H C N O

Electronegativity 2.1 2.5 3.0 3.5

States of Matter

• The properties shown by materials are the result of their structure and bonding

• You need to know about crystalline materials which are:

IonicMetallic

Giant covalent Molecular

Ionic Crystal Lattices

• Each ion is surrounded by a number of ions of the opposite charge

• In sodium chloride, each ion is surrounded by 6 of the oppositely charged ions.

• This gives a cubic shape.

Remember. Oppositely charged ions attract each other by

Electrostatic attraction

• Each Na+ ion is surrounded by a total of 6 Cl- ions One more is be in front of the ion; the other is behind it

2D diagram of NaCl

Na+ Cl- Na+ Cl- Cl- Na+ Cl- Na+ Na+ Cl- Na+ Cl-

Na+

Na+

Na+

Na+

Cl-

Cl-

Cl-

Cl-

3D diagram of NaCl

Na+

Na+

Na+

Na+

Cl-

Cl-

Cl-

Cl- Cl

Cl-

Cl -

Cl-

Each Na+ ion is surrounded by six Cl- ions and vice versa

Why are ionic solids brittle animation?

+ + - - + - + + - + -

- + - + + - + -

+ - + -

+ - + - - + - + + - + -

- + - + + - + -

+ - + -

- + - + + - + -

+ - + -

- + - + + - + -

+ - + -

- + - + + - + -

+ - + -

- + - + + - + -

+ - + -

- + - + + - + -

+ - + -

- + - + + - + -

+ - + -

- + - + + - + -

Ionic Lattice

OK – so what was happening there?

Here it is again with sub-captions for the hard-of-thinking!

Why are ionic solids brittle animation?

+ + - - + - + + - + -

- + - + + - + -

+ - + -

+ - + - - + - + + - + -

- + - + + - + -

+ - + -

- + - + + - + -

+ - + -

- + - + + - + -

+ - + -

- + - + + - + -

+ - + -

- + - + + - + -

+ - + -

- + - + + - + -

+ - + -

- + - + + - + -

+ - + -

- + - + + - + -

Like charges repel

Metals In Mg there is an attraction between the

delocalised e and the lattice of metal ions

2+ 2+ 2 + 2 +

2+ 2+ 2 + 2 +

2+ 2+ 2 + 2 +

Most metals have high m.p. / b.p. as a lot of energy is needed to remove an atom from this attraction

The delocalised e- flow thorough the metal (a current) so metals conduct electricity

So how can metals change their shape?

+ + + + + + + + + + + +

+ + + +

+ + + + + + + +

Malleable since the layers can slide over one another

• All ions have the same electron configuration (isoelectronic)• As the proton number increases from Na to Al, so does the

attraction for the outer shell electrons. Size decreases

Na+ Mg2+ Al3+

Comparison of Na, Mg and Al

+ + + + + + + + + + + +

Mg2+

Al3+

Na+

Al3+

Na+

Mg2+

delocalised e-

+ + + + + + + + + + + +

2+ 2+ 2+ 2+ 2+ 2+ 2+ 2+ 2+ 2+ 2+ 2+

3+ 3+ 3+ 3+ 3+ 3+ 3+ 3+ 3+ 3+ 3+ 3+

• Higher ionic charge, more delocalised e-, smaller ionic radius

• Metallic bonding increases; so m.p. / b.p. increase Na Al

Diamond Graphite

Silica

Giant molecular crystals

• Physical strength explain in terms of bond breaking (diamond/Si) and layers sliding (graphite)

Many strong covalent bonds need to be broken in the structure of diamond and silica so strong substances. In graphite the layers can slide over one another so graphite is softer and is often used as a lubricant.

• Melting points explain in terms of bond breaking

Many strong covalent bonds need to be broken in macromolecules so high melting points.

• Electrical conductivityexplain in terms of presence/absence of delocalised e-

Graphite has delocalised electrons between the layers so it conducts electricity.

Molecular crystals• The example you need to know is Iodine

(but you could be asked to apply your understanding to other elements, such as sulphur)

• There is only one element present

• The bond is non-polar

• The only IMF present is van Der Waals’

• These IMFs are weak so

• The melting point is relatively low

• The IMFs operate between molecules

Shapes of molecules/ions

• The basic shape is determined by the number of electron pairs present

• If all the electron pairs are bonding-pairs then we get a ‘regular’ shape

• If some of the electron pairs are lone-pairs then we get an ‘irregular’ shape

The Electron Pair Repulsion Theory

• The shape results from repulsions between e- pairs – NOT bonds/atoms

• The rules are:

• Lone pairs repel more strongly than bonding pairs

• Order is

l.p/l.p. > l.p/b.p >b.p/b.p

• Remember: When lone pairs are present, the basic shape will be distorted.

B A B

Equal repulsion between

2 bonding pairs

Linear Shape

A B

B

B Equal repulsion between

3 bonding pairs

Trigonal Planar shape

Tetrahedral shape

Equal repulsion between

4 bonding pairs

B

BBB

A

Trigonal Bipyramid shape

B

A B

B

B

B

Equal repulsion between

5 bonding pairs

Octahedral shape

B B

A B

B

B

B

Equal repulsion between

6 bonding pairs

Regular Shapes

e- pairs

Formula Bond angle Name

2 B – A – B 180 linear

3 AB3120 trigonal planar

4 AB4 109½ tetrahedral

5 AB5

120

90trigonal

bipyramid

6 AB6 90 octahedral

Working out the shape

• You can work out the shape using simple maths• Find the group number of the central atom• Add to this number the number of bonds present

• If it is an ion– add one if the charge on the ion is -1– Deduct 1 if the charge is +1

• Divide total by 2 - this gives the number of e- pairs• Number of e- pairs = basic shape• If lone pairs present, basic shape modified

To draw the shape• Work out the basic shape

• Lightly sketch this in pencil

• Put the atoms in

• Draw the lone pairs along the ‘spare’ bonds

• Rub out you working pencil lines

N

H H

– N is in Group 5, so five outer electrons– 3 more electrons (1 from each H)– Makes 8 electrons– 4 pairs of electrons– 3 bonding pairs and 1 lone pair

Basic shape (4 e- pairs) = tetrahedral– One lone pair, so doesn’t have basic

shape– 3 bonding pairs + 1 lone pair

Working out the shape of NH3

Ammonia

Pyramidal shape

Unequal repulsion between

3 bonding pairs and 1 lone pair

N

H H

H

N is in Group 5, so five outer electrons– 2 bonds to H atoms, so 2 extra bonding e-

– 1 charge on molecule, so add extra e-

– = 8 electrons

– Makes 2 bonding pairs and 2 lone pairs

Basic shape (4 e- pairs) = tetrahedral– 2 bonding pairs + 2 lone pair – so doesn’t have basic shape

Working out the shape of the ion NH2-

The shape of the ion NH2-

‘Bent’ or ‘V-shaped’

Unequal repulsion between

2 lone pair and 2 bonding pairs

N

H H

Shapes of molecules / ionsSpecies AlCl3 CCl4 NH2

- NH3 H3O+

Group number

e- from central atom

3 4 5 5 6

Bonds present

e- from other atoms

3 4 2 3 3

Charge+ charge

reduces e-'0 0 1 0 -1

Total e- Sum of above 6 8 8 8 8

e- pairsDivide above

by 2 3 4 4 4 4

Lone pairs

= e- pairs - bonds'

0 0 2 1 1

Shape based on

Closest shape without l.p.

Trig Planar

Tetra Tetr Tetr Tetr

Actual shape

shape with any l.p. present

Trig Planar

Tetra Bent Pyramidal Pyramidal

Question from Jan 09 Arsenic is in the same group as nitrogen. It forms a

compound AsH3

Draw the shape of an AsH3 molecule and name the shape made by its atoms.

3 bonds and 1 lp attached to As

trigonal pyramidal

As

H H

H

The End