Section 2

Periodicity

Bonding in the Elements 1-20 (a)

L.I. To learn about Bonding in the Elements 1-20

S.C. By the end of this lesson you should be able to

•describe the metallic bond•explain what is meant by the term monatomic •explain what London dispersion forces are and how they arise•explain what happens to the strength of LDF as the atom size increases•explain the difference between covalent network and covalent molecular in terms of bpt and mpt•give examples of metallic, covalent molecular, covalent network and monatomic elements

Periodic Pattern

• Johan Wolfgang Dobereiner – triads

the atomic mass of the central element was approximately the mean of the other two.

• What does the periodic table and the sound of music have in common?

• John Newlands – octaves based on atomic mass (musical notes). o

• Every eighth element showed similarities

• The modern Periodic Table is based on the work of Dimtri Mendeleev in 1869• He arranged the elements based on: atomic mass, similar

properties• He left gaps and made predictions for missing elements

fun

Period I II III IV V VI VII VII

1 H

2 Li Be B C N O F

3 NaK

MgCa

Al*

SiTi

PV

SCr

ClMn

Fe Co Ni

4 CuRb

ZnSr

*Y

*Zr

AsNb

SeMo

Br* Ru Rh

Pd

5 Ag Cd In Sn Sb Te I

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Bonding in metals

•Metallic bonding is the electrostatic attraction between the positively charged ions and the delocalised electrons.

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The outer electrons are delocalised and free to move throughout the lattice, making metals good conductors of electricity.

The greater the number of electrons in the outer shell the stronger the metallic bond.

So the melting point of Al>Mg>Na

Bonding in metals

Bonding in Monatomic elements

Noble gases have full outer electron shells

They do not need to combine with other atoms.

The noble gases occur as single atoms, they are said to be monatomic.

He

++

Since they can be liquefied and solidified there must be some weak attraction between the atoms.

Bonding in monatomic elements

The electrons in an atom “wobble” and become unevenly distributed causing one side of the atom becomes slightly negative while the other side becomes slightly positive.

A temporary dipole is therefore formed.

London dispersion forces

These slight charges are given the symbol δ ‘delta’.

A dipole can induce other atoms to form dipoles, resulting in weak attractions between particles.

++

++

δ-

δ+

δ-

δ+

δ-

δ+

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London dispersion forces are a type of van der Waal force. They are very weak attractive forces.

Bonding in monatomic elements

Group 8 element

Electron arrangem

ent

Boiling Point oC

Helium    Argon    

Krypton    Xenon    

Comparing Boiling points

The melting and boiling point of a substance gives an indication of the strength of the forces of attraction holding atoms or molecules together.

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Noble gases b.p.’s

b.p / K

B.p.’s increase as the size of the atom increases

This happens because the London dispersion forces increases with increasing size of atoms.

4 27

87

121

166

0

20

40

60

80

100

120

140

160

180

Helium

NeonArgon

Krypton

Xeon

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Covalent Molecular ElementsMany non-metals exist as discrete covalent molecules held together by covalent bonds.Discrete molecules have a definite number of atoms bonded together.

9+

Fluorine atom

9+ 9+

Fluorine molecule F2

diatomic

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Examples of discrete molecules:

Cl Cl

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weak London dispersion forces

strong covalent bonds

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Melting point low – why?

Comparing Boiling points

Halogen Boiling point (oC)

   

   

   

   

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Halogens b.p.’s

b.p./ K

As the size of the halogen molecule increases the boiling point increases. The bigger the molecule the strongerthe London dispersion forces between the halogen molecules.

0

50

100

150

200

250

300

350

400

450

500

Fluorine

Chlorine

Bromine

Iodine

85

238

332

457

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Fullerenes, molecules of carbon

Fullerenes exists as large covalent molecules with a definite formula.

Fullerenes were discovered in 1985 by Buckminster Fuller. Fullerenes are spherical in shape and usually contain sixty or seventy carbons.

C60 is known as Buckminster fullerene

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Covalent Network Elements

Carbon - diamond

m.p.’s C > 3642oC

It is high because many covalent bonds have to be broken.

Diamond has acovalent networkstructure

Each of the outer electrons in a carbon atom canform a covalent bond with another carbon atom. So every C bonds to 4 others.

Carbon - Graphite

Van der Waals forces between the layers allows layers to slide overeach other.

Carbon bonded to only 3 other Carbons

The spare (4th) electron is delocalised and so free to move. Graphite is a conductor of electricity.

Graphite can be used as a lubricant

Properties of graphite and diamond

Property Diamond Graphite

Appearance Colourless transparent solid

Conduction No

Feel Smooth, not slippery

Hardness Very hard

Other Network Structures

In the first 20 elements, only Boron, Carbon and Silicon have covalent network structures.

m.p.’s B 2300oC, C > 3642oC and Si 1410oC

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BONDING IN ELEMENTS - A SUMMARY

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Bonding patterns of the 1st 20 elements

CovalentMolecular

Metalliclattice

Monatomic

CovalentNetwork

C , in the form of fullerenes, is covalent molecular

ArClSPSi

NeFONCB

He

Si

CB

ClSP

FON

CaK

MgNa

BeLi

CaK

AlMgNa

BeLi

H

http://www.ltscotland.org.uk/highersciences/chemistry/animations/bonding_structure.asp

This interactive animation provides a visual representation of the bonding and structure of the first twenty elements in the periodic table, taking into account both the intra- and inter-molecular forces involved.

Questions on elements – bonding and structure

1. Explain why the covalent network elements have high melting and boiling points.

2. Explain why the discrete molecular and monatomic elements have low melting and boiling points.

3. Does diamond conduct electricity? Explain.4. Does graphite conduct electricity? Explain.5. How does the hardness of diamond compare

with graphite? Explain.6. Give a use for both diamond and graphite.7. Complete the following table:

Questions on elements – bonding and structure

7. Complete the following table:

Type of bonding and structure

Properties

Metallic solids ……………. of electricity

Covalent network solids ……….. …. melting points……………. of electricityexception ……………….

Covalent molecular solids ………….. melting points…………… of electricity

Covalent molecular (diatomic) gases

and monatomic gases

…………… boiling points

Section 2

Periodicity

Patterns in the Periodic Table (b)

L.I. To learn about covalent radius

S.C. By the end of this lesson you should be able to

•describe the term covalent radius

•explain the changes in covalent radius down a group

•explain the changes in covalent radius across a period

•explain why there is no stated covalent radius for the noble gases

Covalent Radius

There is no definite edge to an atom.

However, bond lengths can be worked out.

Covalent radius – picometres (pm) 1pm =1 X 10 – 12 m

266pm

The size of an atom is indicated by its covalent radius. (Page 7 of data booklet).

The covalent radius of an element is half the distance between the nuclei of 2 of its bonded atoms

From above the covalentradius would be 133 pm.

Na 154 pm, Mg 145 pm, Al 130 pm, Si 117 pm, P 110 pm, S 102 pm

Trends in covalent radius - Across a period

Why? Going across a period the nuclear charge increases. The attraction between the outer electrons and the positive nucleus increases. Thus the outer electrons are more strongly attracted and so the atom size is smaller.

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Going across a period the covalent radius (atomic size)decreases.

Li 134 pm

Na 154 pm

K 196 pm

Rb 216 pm

Trends in covalent radius – Down a groupactivity in workbook

Going down a group the covalent radius (atomic size) increases.On moving down a group from one element to the next the number of electron shells increases.

So the outer electrons are further from the nucleus and the atom size increases.

Why is there no covalent radius value for the noble gases?

L.I. To learn about ionisation energies

S.C. By the end of this lesson you should be able to

•describe the term 1st ionisation energy

•write equations for the 1st ionisation energy

•explain the trend in 1st ionisation energy down a group

•explain the trend in 1st ionisation energy across a period

•describe the term 2nd ionisation energy

•carry out calculations involving ionisation energy

Ionisation energies

The first ionisation energy of an element is the energy required to remove one mole of electrons from one mole of atoms in the gaseous state.

Units are kJmol-1.

This is an endothermic process. (Page 11 of the data booklet.)

Na(g) Na+ (g) + e

Cl(g) Cl+ (g) + e

Trends in 1st ionisation energy – Across a periodactivity in workbook

Going across a period the ionisation energy increases.

Going across a period the nuclear charge increases. The attraction between the negative electrons and the positive nucleus increases. Thus the electrons are more tightly held and so more energy is needed to remove the outer electrons.

Trends in 1st ionisation energy – Down a group

activity in workbook

Going down a group the ionisation energy decreases. The explanation for this is

(i) on moving down a group from one element to the next the number of electron shells increases and so the outer electron is further from the nucleus and less tightly held. (ii) the inner shells provide a screening effect which also decreases the attractive forces between the outer electrons and nucleus.

The 2nd Ionisation Energy 

The second ionisation energy of an element is the energy required to remove the second mole of electrons.

ΔH = +738 kJ mol-1

First Ionisation

Mg(g) Mg+(g) + e-

Second Ionisation

Mg+(g) Mg2+ + e-

Third Ionisation

Mg2+(g) Mg3+ + e-

ΔH = +1451 kJ mol-1

ΔH = +7733 kJ mol-1

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L.I. To learn about electronegativity

S.C. By the end of this lesson you should be able to

•describe the term electronegativity

•explain the trend in electronegativity down a group

•explain the trend in electronegativity across a period

•explain why there are no quote values of electronegativity for the noble gases

Electronegativity

The electronegativity is a measure of the attraction an atom involved in a bond has for the shared pair of electrons.

Electronegativity values are based on the Pauling Scale, devised by Linus Pauling an American Chemist. Values on the Pauling Scale range from 0 to 4. A list of these values can be found in the data booklet on page 11.

The higher the number on the Pauling scale is, the greater the attraction an atom has for the bonding electrons.

Electronegativity values can be useful in predicting which type of bonding is most likely between two elements. (More about this later)

Electronegativity – Across a Periodactivity in workbook

On crossing a period, electronegativity values increase. This is caused by an increase in nuclear charge as you move across a period from left to right.

Electronegativity – Down a Groupactivity in workbook

As you go down a group, electronegativity values decrease.This is caused by the addition of another energy level of electrons as you go down a group which shields the bonded electrons from the nucleus; therefore they are not attracted as strongly.

Electronegativity - The Monatomic Gases

Why no values for group 8 elements?

Section 3

Structure and Bonding

Bonding in Compounds

L.I. To learn about bonding in compounds (a)

S.C. By the end of this lesson you should be able to

•describe the bonding and structure in ionic compounds

•explain the melting point of ionic compounds

•describe the bonding and structure in covalent network compounds

•explain the melting point of covalent network compounds

•describe the bonding and structure in covalent molecular compounds

•explain the melting point of covalent molecular compounds

Ionic Bonding

In ionic compounds atoms achieve a full outer shell by either losing or gaining electrons and so form charged particles called ions.

Three different types of compound - ionic, covalent molecular or covalent network.

Na Cl Na+ + Cl-

 2)8)1 2)8)7 2)8 2)8)8

Element Atom electron arrangement

Ion electron arrangement

Ion symbol

Mg 2)8)2 2)8 Mg2+

Complete for sodium, chlorine, bromine, oxygen, aluminium and nitrogen.

Glow: ionic bondingionic compounds

Metal atoms always lose electrons to form positive ions e.g Na+

Non-metal atoms always gain electrons to form negative ions e.g F-

Sodium chloride

Lithium fluoride

Magnesium oxide

Aluminium nitride

Calcium chloride

Now write ionic formula for the above.

On show me boards – work out how these elements forman ionic compound

NaCl

The attraction between positive and negative ions holds the compound together.

The electrostatic attraction between positive and negative ions is an ionic bond.

3D lattice – regular repeating pattern

of ions

Ionic Bonding

+ ------- -

ionic bond

Ionic Compounds

Ionic Compounds

NaCl as with all ionic compounds have many strong ionic bonds which are broken on melting thus the melting points are high (801 0C)

Complete workbook

COVALENT BONDING

In covalent bonding the atoms share electrons.

shared pair of electrons

positive nuclei

Covalent bonding is the electrostatic attraction between the shared electrons and the positive nuclei.

Silcion dioxide - SiO2Silcion carbide - SiC

Mpt = 1610oC Mpt = 2700oC

COVALENT NETWORK

Silicon dioxide and silicon carbide exist as a covalent network.

All network structures have very high melting and boiling points.

It is the strong covalent bonds that are broken on melting.

Molecular Compounds

Write formula for the following compounds:

carbon monoxide, sulphur trioxide, carbon tetrafluoride, dinitrogen tetraoxide, phosphorus trifluoride

Draw electron dot cross diagrams for the following molecules and structural formula

1. CH4

2. SCl23. CO2

COVALENT MOLECULAR

Weak intermolecular forces

Strong covalent bonds

Covalent molecules tend to have low melting and boiling points as it is the weak intermolecular forces that are broken on melting.

Complete workbook

Plot a graph of melting points of the carbon tetrahalides against the covalent radius of the halogen in each molecules (see data book)

CF4 = -184oC

CCl4 = -23oC

CBr4 = 90oC

CI4 = 171oC

Temp/ oC

As the molecule size increases the m.pt.s increase. This is because the strength of the London dispersion forces increase, so more energy is needed to separate molecules.

m.p.’s of the carbon halides

increasing size of molecules

CBr4

CF4

CCl4

CI4-183

90

-23

171

COVALENT MOLECULAR

What happens to the melting point as the size of the molecule increase?Why?

Complete workbook

L.I. To learn about polar covalent bonds (b)

S.C. By the end of this lesson you should be able to

•use electronegativities to explain the difference between pure covalent and polar covalent bonds

•explain the term permanent dipole

•use the data book to assign δ+ and δ+ partial charges on atoms

POLAR COVALENT BONDS

Two types of covalent bond can be formed:

Pure covalent (or non-polar covalent)Polar covalent.

Covalent Bonding

Picture a tug-of-war:

If both teams pull with the same force the mid-point of the rope will not move.

Pure Covalent Bond

This even sharing of the rope can be compared to a pure covalent bond, where the bonding pair of electrons are held at the mid-point between the nuclei of the bonding atoms.

H He

e

Covalent Bonding

What if it was an uneven tug-of-war?

The team on the right are far stronger, so will pull the rope harder and the mid-point of the rope will move to the right.

Polar Covalent Bond

A polar covalent bond is a bond formed when the shared pair of electrons in a covalent bond are not shared equally.

This is due to different elements having different electronegativities.

Polar Covalent Bond

e.g. Hydrogen Iodide

If hydrogen iodide contained a pure covalent bond, the electrons would be shared equally as shown

above. However, iodine has a higher electronegativity and pulls the bonding electrons towards itself (winning the tug-of-war)

This makes iodine slightly negative and hydrogen slightly positive. This is known as a dipole.

H Ie

e

δ-δ+

PURE COVALENT (OR NON-POLAR COVALENT)

A pure covalent bond is formed when the atoms involved in the bond have an equal share of the bonding electrons. They have the same electronegativity.

H - H2.2

2.2

Complete workbook

POLAR COVALENT

When atoms with different electronegativity values join together, a polar covalent bond is formed.

The dipole produced is permanent.

A polar covalent bond is a bond where the electrons are not shared equally, one atom in the bond has a greater attraction than the other for the bonded electrons.

Complete workbook

L.I. To learn about the bonding continuum (c)

S.C. By the end of this lesson you should be able to

•explain the relationship between differences in electronegativities and type of bonding

•use data from the properties of compounds to deduce the type of bonding and structure

BONDING CONTINUUM

Electronegativity Difference and Bond Type:Electronegativit

y differenceBond type Example Actual difference

in electronegativity

0.0-0.4 covalent (non polar)

H-H 0.0

0.4-1.0 covalent (polar)

H-Cl 0.9

covalent (polar)

H2O 0.7

1.0-2.0 covalent (very polar)

H-F 1.9

.2.0 ionic NaCl 2.1

The greater the difference in electronegativity the greater the polarity between two bonding atoms and the more ionic in character.

A bonding continuum can be used to help us understand the differences in bonding.

Complete workbook

Bonding Continuum

“Covalent compounds are formed by non-metals only”

Some compounds break this rule….

IS NOT AN ABSOLUTE LAW!

Tin(IV)iodide – covalent or ionic?

Melting point of tin(IV)iodide is 143oC.

Tin electronegativity of 1.8Iodine has electronegativity of 2.6

Molecule contains polar covalent bonds, but the symmetry cancels out the dipoles, therfore only weak London’s forces so low melting an boiling point.

Predict its melting point. Complete workbook

L.I. To learn about intermolecular forces (d)

S.C. By the end of this lesson you should be able to

•explain the difference between intramolecular andintermolecular forces

•name the three types of van der Waals forces

•explain how London dispersion forces arise

INTERMOLECULAR FORCES

Intramolecular bonds are bond between atoms within a molecular – covalent bond.

Intermolecular bonds are bonds which occur between molecules.

Intermolecular bonds are called

van der Waals’ forces. They are named

after the Dutch Chemist Johannes

Diderik van der Waals.

There are three types of van der Waals’ forces:

London dispersion forcesDipole-dipole interactions (permanent dipoles)Hydrogen bonding

1. London Dispersion Forces

Electrons ‘wobble’ and temporary dipole occur. These cause induced dipoles on other atoms.

The attraction between atom resulting from the temporarydipoles are known as London dispersion forces.

London dispersion forces are very weak attractive forces.

London Dispersion Forces

L.I. To learn about intermolecular forces

S.C. By the end of this lesson you should be able to

•explain how dipole-dipole interactions arise

•describe a test that can be used to determine if a molecule is polar

•explain the connection between symmetry and polarity

•describe how most hydrocarbons are classified in terms of polarity

H

O

H

--

+

+

+

+

--

Water has a polar covalent bonding between O and H.

Are all molecules with polar bonds polar?

Polar Molecules

Is water polar?

2. Dipole-Dipole Interactions

see scholar animation on polarity test

Complete activity – testing polarity, and complete the table

Symmetry and polarity

Asymmetrical molecules e.g. H2O are POLAR

In an asymmetrical molecule there is a permanent dipole

workbook activity

Symmetrical molecules e.g. CCl4 are NON-POLAR

In a completely symmetrical molecule the polarities cancel each other out so there is no permanent dipole.

workbook activity

Most hydrocarbons are non-polar

Dipole-Dipole Interactions

Dipole-dipole interactions are intermolecular forces which occur between polar molecules.

Dipole-dipole interactionDipole-dipole interaction

H – H H - H

Dipole – dipole interactions

London dispersion forces

polar molecule

non - polar molecule

Dipole-dipole interactions are stronger than London dispersion forces.

Polar molecules have higher melting and boiling points than non-polar molecules.

C

O

CCH

H

H

H

H

H

b.p. 56 o C b.p. -1 o C

non - polar moleculepolar molecule

L.I. To learn about intermolecular forces

S.C. By the end of this lesson you should be able to

•explain how H-bonds arise

•what is necessary in a molecule to allow H-bonds to arise

Hydrogen Bonding

Hydrogen bonding is a special type of dipole-dipole interaction involving; H-N, H-O or H-F bonds workbook activity

For H-bonds to exist between molecules

1. the molecules must have a strong polar covalent bond

2. the polar covalent bond must be between a hydrogen atom and either nitrogen, fluorine and oxygen (NOF)

Hydrogen bonds are stronger than normal dipole-dipole interactions and London dispersion forces.

Molecules which contain hydrogen bonding have much higher melting and boiling points than those with dipole-dipole interactions or London dispersion forces.

L.I. To learn about relating properties of compounds to intermolecular forces (e)

S.C. By the end of this lesson you should be able to

•explain the connection between size of molecule and strength of London dispersion forces•describe the evidence that proves the existence of permanent dipole-dipole interactions.•describe the evidence that proves the existence of H-bonds•explain why ice is less dense that water•explain how intermolecular forces affect bpts, mpts, •viscosity and solubility•explain the term “like dissolves like”

RELATING PROPERTIES TO INTERMOLECULAR BONDING

1. Melting and boiling points give an indication of the amount of energy needed to overcome the van der Waal’s forces between molecules.

London dispersion forces – between non polar molecules

Alkane name

Alkane formula

boiling point(oC)

pentane    

hexane    heptane    

octane    

workbook activity

From the table we can see that as the molecular size increases (number of electron shells) the boiling point increase and so the strength of the London dispersion forces increase.

Dipole-dipole interactions – between polar molecules

C

O

CCH

H

H

H

H

H

b.p. 56 o C b.p. -1 o Cpolar molecule non - polar molecule

For polar molecules, the melting and boiling points are higher than those of non-polar molecules. More energy is needed to overcome the dipole-dipole interactions between polar molecules. Molecules with a similar mass are used to allow us to ignore the London Dispersion Forces.

workbook activity

Formula mass 58

Formula mass 58

Hydrogen bonding

For polar molecules that contain H-N, H-O or H-F the melting points are related to the strong hydrogen bonds between molecules.

workbook activity

For polar molecules which contain H-N, H-O or H-F bonds the melting point and boiling points are higher than those of other polar and non-polar molecules. More energy is required to overcome the strong hydrogen bond.

NH3, has a higher boiling point than expected.

Group 5

0

100

200

300

NH3 PH3 AsH3 SbH3

Boi

ling

Poin

t (K

)

Group 5

H2O has a higher boiling point than expected.

Group 6

0

100

200

300

400

H2O H2S H2Se H2Te

Boilin

g Po

int (

K)

Group 6

HF has a higher boiling point than expected.

Group 7

0

100

200

300

400

HF HCl HBr HI

Boilin

g Po

int (

K)

Group 7

Evidence of Hydrogen Bonding

Boiling Points of Hydrides

0

100

200

300

400

Series Number

Bo

ilin

g P

oin

t (K

)

Group 4

Group 5

Group 6

Group 7

NH3

H2O

HF

Strength of van der Waals’ forces: Hydrogen > Dipole-dipole > London dispersion forces

2. SOLUBILITY

“like dissolves like”

polar solvents, such as water, will dissolve polar and ionic solutes

Na+Cl- (s) Na+(aq) + Cl-(aq)

Dissolving in Water

3. VISCOSITY

scholar animation

The stronger the van der Waals’ forces between a liquid are, the more viscous it will be.

Liquids containing hydrogen bonding will be more viscous than molecules containing dipole-dipole interactions or London dispersion forces

The most viscous liquid is propane-1,2,3-triol.

Why does the ice float in Mrs Brown’s gin and tonic?

Why do icebergs float?

Why do fish in ponds not die in winter when the water freezes?

DENSITY OF WATER/ICE

look at candle wax and ice

The density of ice is unusual. Normally solids sink in their liquids.

On cooling, water contracts but at 4oC it expands. At freezing point an open structure exists as a result of H-bonds.

So ice floats!!

SUMMARY