Periodicity

Period 3 - elements

AS Chemistry

Can you write one fact about each of the following elements?

Na Mg Al Si P S Cl Ar

Starter activity

Learning ObjectivesCandidates should be able to:•describe qualitatively (and indicate the periodicity in) the variations in atomic radius, ionic radius, melting point and electrical conductivity of the elements (see the Data Booklet)•explain qualitatively the variation in atomic radius and ionic radius•interpret the variation in melting point and in electrical conductivity in terms of the presence of simple molecular, giant molecular or metallic bonding in the elements•explain the variation in first ionisation energy.

Na [Ne] 3s1

Mg [Ne] 3s2

Al [Ne] 3s2 3px1

Si [Ne] 3s2 3px1 3py

1

P [Ne] 3s2 3px1 3py

1 3pz1

S [Ne] 3s2 3px2 3py

1 3pz1

Cl [Ne] 3s2 3px2 3py

2 3pz1

Ar [Ne] 3s2 3px2 3py

2 3pz2

Electronic configurations

Atomic radius

0

0.05

0.1

0.15

0.2

0.25

Na Mg Al Si P S Cl ArElement

Atom

ic ra

dius

(nm

)

www.chemguide.co.uk/atoms/properties/atradius.html

Atomic radius – what about argon?

What is atomic radius?

www.rjclarkson.demon.co.uk/found/period3_atomrad.gif

An atomic radius is a measure of the distance from the nucleus to the bonding pair of electrons.

From sodium to chlorine, the bonding electrons are all in the 3rd shell being screened by the electrons in the first and second levels, i.e. the screening remains fairly constant.

The increasing nuclear charge as you go across the period pulls the bonding electrons more tightly towards it.

Explaining the trend

You have to ignore the noble gas at the end of each period. Because neon and argon don't form bonds, you can only measure their van der Waals radius - a case where the atom is pretty well "unsquashed". All the other atoms are being measured where their atomic radius is being lessened by strong attractions. You aren't comparing like with like if you include the noble gases.

Leaving the noble gases out, atoms get smaller as you go across a period.

IONIC RADIUS

Ions aren't the same size as the atoms they come from. Compare the sizes of sodium and chloride ions with the sizes of sodium and chlorine atoms.

Positive ionsPositive ions are smaller than the atoms they come from. Sodium is 2,8,1; Na+ is 2,8. You've lost a whole layer of electrons, and the remaining 10 electrons are being pulled in by the full force of 11 protons.

Negative ionsNegative ions are bigger than the atoms they come from. Chlorine is 2,8,7; Cl- is 2,8,8. Although the electrons are still all in the 3-level, the extra repulsion produced by the incoming electron causes the atom to expand. There are still only 17 protons, but they are now having to hold 18 electrons.

0

0.5

1

1.5

2

2.5

3

3.5

Na Mg Al Si P S Cl Ar

Element

Paul

ing

elec

tron

egati

vity

val

ueElectronegativity

0200400600800

1000120014001600

Na Mg Al Si P S Cl Ar

Element

Firs

t ion

isatio

n en

ergy

(kJ m

ol-1

)First ionisation energy

Electronegativity

Na Mg Al Si P S Cl ArStructure

Giant metallic

Giant metallic

Giant metallic

Giant covalent

Simple molecule

Simple molecule

Simple molecule

Monatomic

Type of element

Metal Metal Metal Non-metal

Non-metal

Non-metal

Non-metal

Non-metal

Bonding Metallic Metallic Metallic Covalent Covalent Covalent CovalentFormula P4 S8 Cl2 ArType of force broken on melting/boiling

Metallic bond

Metallic bond

Metallic bond

Covalent bond

vdW vdW vdW vdW

Does the element conduct electricity?

Yes Yes Yes No No No No No

Bonding, Structure and Properties

Structure and Properties

Electrical conductivity

0

500

1000

1500

2000

2500

3000

Melting and boiling points

Periodicity

Period 3 - oxides

AS Chemistry

Starter activity

Complete these sketches to show how these properties change as you go along Period 3 from Na to Ar.

Starter activity

Starter activity

Learning ObjectivesCandidates should be able to:describe the reactions, if any, of the elements with

oxygen to give Na2O, MgO, Al2O3, P4O10, SO2 and SO3.

state and explain the variation in oxidation number of the oxides.

describe the reactions of the oxides with water.

describe and explain the acid/base behaviour of oxides and hydroxides, including, where relevant, amphoteric behaviour in reaction with NaOH and acids.

Group number 1 2 3 4 5 6Element in Period 3

Na Mg Al Si P S

Nuclear charge

11+ 12+ 13+ 14+ 15+ 16+

[Ne] electronic configuration

3s1 3s2 3s2 3p1 3s2 3p2 3s2 3p3 3s2 3p4

Trend in Atomic radius decreasesTrend in 1st ionisation energy

increases

Trend in electronegativity

increases

Formula of oxide/s

Na2O MgO Al2O3 SiO2 P4O10 SO2/SO3

Oxidation state

+1 +2 +3 +4 +5 +4/+6

Reactions with oxygen

SodiumSodium burns in oxygen with an orange flame to produce the white solid sodium oxide.

MagnesiumMagnesium burns in oxygen with an intense white flame to give white solid magnesium oxide.

Reactions with oxygen

AluminiumAluminium will burn in oxygen if it is powdered, otherwise the strong oxide layer on the aluminium tends to inhibit the reaction. If you sprinkle aluminium powder into a Bunsen flame, you get white sparkles. White aluminium oxide is formed.

SiliconSilicon will burn in oxygen if heated strongly enough. Silicon dioxide is produced.

Reactions with oxygen

PhosphorusWhite phosphorus catches fire spontaneously in air, burning with a white flame and producing clouds of white smoke - a mixture of phosphorus(III) oxide and phosphorus(V) oxide. The proportions of these depend on the amount of oxygen available. In an excess of oxygen, the product will be almost entirely phosphorus(V) oxide:

Reactions with oxygen

Reactions with oxygen

SulphurSulphur burns in air or oxygen on gentle heating with a pale blue flame. It produces colourless sulphur dioxide gas.

In an excess of pure oxygen, some SO3 is also formed. This utilises the highest oxidation state of sulphur.

Reactions with oxygen

Which is which?

Melting points of oxides

Na2O MgO Al2O3 SiO2 P4O10 SO2

Tm/K 1548 3125 2345 1883 573 200

Bonding Ionic Ionic Ionic Covalent

Covalent

Covalent

Structure

Giant lattice

Giant lattice

Giant lattice

Giant lattice

Simple molecul

ar

Simple molecul

ar

Giant ionic and covalent solids contain only strong bonds and have high melting points. Simple molecules have weak vdW forces between molecules and have low melting points.

Reaction with water pHNa2O + H2O 2Na+ + 2OH- 14MgO + H2O Mg2+ + 2OH- 9Al2O3 insoluble – no reaction 7SiO2 insoluble – no reaction 7P4O10 + 6H2O 4H3PO4 0SO2 + H2O H2SO3 3SO3 + H2O H2SO4 0

Acid/base properties of the oxides

Acid/base behaviour of aluminium oxide

BASE:

Al2O3 + 3H2SO4 → Al2(SO4)3 + 3H2O

ACID:

Al2O3 + 2NaOH + 3H2O → 2NaAl(OH)4

Periodicity

Period 3 - chlorides

AS Chemistry

Learning ObjectivesCandidates should be able to:

• describe the reactions, if any, of the elements with chlorine to give NaCl, MgCl2, Al2Cl6, SiCl4, and PCl5.

• describe and explain the reactions of the chlorides with water.

Element Na Mg Al Si PDescription of reaction with chlorine

Very vigorous Vigorous Vigorous Slow Slow

Formula of chloride/s NaCl MgCl2 Al2Cl6 SiCl4 PCl3 / PCl5Oxidation state of period 3 element

+1 +2 +3 +4 +3 / +5

State of chloride at r.t.p.

Solid Solid Solid Liquid liquid / solid

b.pt. of chloride (oC) 1465 1418 423 57 74 / 164Structure of chloride Giant lattice Simple molecularBonding in chloride Ionic Covalent

Reaction with chlorine

Structure of Al2Cl6

Reaction with water pHNaCl(s) Na+

(aq) + Cl-(aq) 7MgCl2(s) Mg2+

(aq) + 2Cl-(aq) 6/7Al2Cl6(s) + 12H2O(l) 2[Al(H2O)6]3+ 6Cl-(aq) 3SiCl4(l) + 2H2O(l) SiO2(s) + 4HCl(g) 0PCl5(l) + 4H2O(l) H3PO4(aq) + 5HCl(g) 0

Acid/base properties of the chlorides

Periodicity

Group 7 - Halogens

AS Chemistry

Starter activity

Can you complete task 1 – IGCSE revision on the halogens

Learning ObjectivesCandidates should be able to:

• describe the trends in volatility and colour of chlorine, bromine and iodine.

• interpret the volatility of the elements in terms of van der Waals’ forces.

• describe the relative reactivity of the elements as oxidising agents.

• describe and explain the reactions of the elements with hydrogen.

• describe and explain the relative thermal stabilities of the hydrides.

• interpret these relative stabilities in terms of bond energies.

The Halogens

Ionic

Covalent

The Halogens

Trend in colour and physical state

Like dissolves like!

Periodicity

Group 7 – the halides

AS Chemistry

Starter activity

Can you complete the table ‘Reducing power of the halides’?

Trend in reducing abilityNaX Observations Produc

tsType of reaction

NaF steamy fumes HF acid-base (F- acting as a base)NaCl steamy fumes HCl acid-base (Cl- acting as a base)NaBr steamy fumes HBr acid-base (Br- acting as a base)

colourless gas SO2 redox (reduction product of H2SO4)

brown fumes Br2 redox (oxidation product of Br-) NaI steamy fumes HI acid-base (I- acting as a base)

colourless gas SO2 redox (reduction product of H2SO4)

yellow solid S redox (reduction product of H2SO4)

smell of bad eggs

H2S redox (reduction product of H2SO4)

Grey solid, purple fumes

I2 redox (oxidation product of I-)

Learning ObjectivesCandidates should be able to:

• describe and explain the reactions of halide ions with aqueous silver ions followed by aqueous ammoniaconcentrated sulphuric acid.

• describe and interpret in terms of changes of oxidation number the reaction of chlorine with cold, and with hot, aqueous sodium hydroxide.

• explain the use of chlorine in water purification.

• recognise the industrial importance and environmental significance of the halogens and their compounds, (e.g. for bleaches; PVC; halogenated hydrocarbons as solvents, refrigerants and in aerosols).

Fluoride

Chloride

Bromide

Iodide

Relative reducing power

Summary of reducing power

Testing for halide ionsion present observation

F- No precipitateCl- White precipitateBr- Cream precipitateI- Yellow precipitate

Silver fluoride is soluble, and so you don't get a precipitate.

Confirmatory tests

original precipitate observation

AgCl Soluble in dilute NH3(aq)

AgBr Soluble in concentrated NH3(aq)

AgI Insoluble in concentrated NH3(aq)

E.g. AgCl(s) + 2NH3(aq) → [Ag(NH3)2]+(aq) +

Cl-(aq)

Cl2(g) + 2NaOH(aq) → NaClO(aq) + NaCl(aq) + H2O(l)

0 +1 -1

3Cl2(g) + 6NaOH(aq) → NaClO3(aq) + 5NaCl(aq) + 3H2O(l)

Reactions of chlorine with NaOH

Uses of halogens and their compounds

bleachrefrigerants

solvents

Aerosol propellants

PVC

Periodicity

Nitrogen and Sulphur

AS Chemistry

Starter activity

Can you use information in your textbooks to complete task 9?

Learning ObjectivesCandidates should be able to:

• explain the lack of reactivity of nitrogen.• describe the

o formation, and structure of, the ammonium iono the displacement of ammonia from its salts.

• understand the environmental consequences of the uncontrolled use of nitrate fertilisers.

• understand and explain the occurrence, and catalytic removal, of oxides of nitrogen.

• explain why atmospheric oxides of nitrogen are pollutants, including their catalytic role in the oxidation of atmospheric sulphur dioxide.

• describe the formation of atmospheric sulphur dioxide from the combustion of sulphur contaminated carbonaceous fuels.

• state the role of sulphur dioxide in the formation of acid-rain and describe the main environmental consequences of acid-rain.

• understand the industrial importance of sulphuric acid.• describe the use of sulphur dioxide in food preservation.

Gases in the air

Structure of nitrogen

Bond dissociation enthalpy = +946 kJ/mol

High energy

Enzymes

Nitrogen

The Haber Process

The ammonium ion

Ammonium NH4

+

Add sodium hydroxide solution to a solution of the substance and gently heat.

Ammonia gas is given off.

n. Any of a large number of natural and synthetic materials, including manure and nitrogen, phosphorus, and potassium compounds, spread on or worked into soil to increase its capacity to support plant growth.

Fertiliser

Fertilisers

Ammonium nitrate

Potassium chloride

Ammonium phosphate

Nitric acid

NH3(g) + 2O2(g) HNO3(l) + H2O(l)

Pt/Rh catalyst 900oC

The environment and fertilisersfertilisers applied to farm land

too much used, at the wrong time of year, during wet weather,

if

Excessive growth of aquatic plants. The bacteria which live on dead plants

thrive and use up the oxygen in the water. The lack of oxygen causes death of fish.

This is called eutrophication.

Harm to infants - called‘blue baby’ syndrome

washed intorivers and lakes

causes

excess excess

contaminatesunderground

drinking water supplies

causes

Natural sources of SO2 and NOx

Sulphur dioxide - SO2

Used as a food preservative:

Inhibits growth of moulds, yeasts and aerobic bacteria.

Acts as a reducing agent and retards the oxidation of foodstuffs.

Acid rain

SO2(g) + NO2(g) SO3(g) + NO(g)

SO3(g) + H2O(l) H2SO4(l)

Uses of sulphuric acid

Paints , pigments and dyestuffs

Detergents

Fertilisers

Periodicity

Group 2

AS Chemistry

Learning Objectives

Candidates should be able to interpret, and make predictions from, the trends in physical and chemical properties of the elements and their compounds.

Trends in Atomic Radius

The radius of an atom is governed by:the distance of the outer electrons from the

nucleusthe nuclear charge, and the amount of shielding

Trends in First Ionisation Energy

Ionisation energy is also governed by:

the charge on the nucleus,

the amount of shielding by the inner electrons,

the distance between the outer electrons and the nucleus.

Trends in First Ionisation Energy

Increased nuclear charge ‘off-set’ by increased shielding. Outermost electron increasingly distant from pull of nucleus, less energy needed to remove it.

Trends in Electronegativity

Trends in Electronegativity

Trends in Electronegativity

As the metal atoms get bigger, any bonding pair gets further and further away from the metal nucleus, and so is less strongly attracted towards it.

Increasing nuclear charge ‘off-set’ by more shielding.

In other words, as you go down the Group, the elements become less electronegative.

Trend in Electronegativity

Trend in Melting Point

Going down Group 2:

• the number of delocalised electrons remains the same ...

• the charge on each metal cation stays the same at 2+, but ...

• the ionic radius increases ...

• so the attraction between the delocalised electrons and the metal cations decreases.

Trend in Melting point

Periodicity

Group 2 – chemical properties

AS Chemistry

Learning ObjectivesCandidates should be able to:

• describe the reactions of the elements with oxygen and water.

• describe the behaviour of the oxides with water.• describe the thermal decomposition of the nitrates

and carbonates.• interpret, and make predictions from, the trends in

chemical properties of the elements and their compounds.

• explain the use of magnesium oxide as a refractory lining material and calcium carbonate as a building material.

• describe the use of lime in agriculture.