INORGANIC CHEMISTRY

Modern Periodic Table

Features of the Periodic Table• The Periodic Table is an arrangement of elements

in order of increasing atomic number.

• Each element in a horizontal row (Period) differs from the preceding element by addition of an electron to the electron shell and a proton to the nucleus.

• Elements are arranged such that elements in a particular vertical column (Group) have the same number of electrons in its valence shell

8.2

ns1

ns2

ns2

np1

ns2

np2

ns2

np3

ns2

np4

ns2

np5

ns2

np6

d1

d5 d10

4f

5f

Ground State Electron Configurations of the Elements

(1) The chemical and physical properties of the elements are periodic functions of the atomic number (number of protons in the nucleus = number of electrons in the neutral atom).

(2) The elements can be arranged in groups (columns) of elements that possess related chemical and physical properties.

(3) The elements can be arranged in periods (rows) of elements that possess progressively different physical and chemical properties

Orbitals Being Filled

1s

2s

3s

4s

5s

6s

7s

3d

4d

5d

6d

2p

3p

4p

5p

6p

1s

La

Ac

1

3 4 5 6 7

4f

5f

Lanthanide series

Actinide series

Groups 8

Perio

ds

1 2

2

3

4

5

6

7

Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 345

Elements show gradual changes in certain physical properties as one moves across a period or down a group in the periodic table. These properties repeat after certain intervals. In other words they are PERIODIC

Periodic properties Periodic properties include:include: -- Atomic Radius-- Ionization Energy-- Electronegativity-- Electron Affinity-- Ionic Radius

Radius

Atomic Radius = half the distance between 2 nuclei of a diatomic molecule/adjacent atoms.

Atomic Radii

Trends in Atomic Radii • Influenced by three factors:

1. Energy Level–Higher energy level is further away.

2. Charge on nucleus–Higher charge pulls electrons in

closer. 3. Shielding effect

- electron repulsion

Group trends• As we go down a

group...• each atom has

another energy level,

• so the atoms get bigger.

HLi

Na

K

Rb

Periodic Trends• As you go across a period, the radius gets

smaller.• Electrons are in same energy level.• More nuclear charge.• Outermost electrons are closer.

Na Mg Al Si P S Cl Ar

• The radius decreases across a period owing to increase in the positive charge from the protons.

• Each added electron feels a greater and greater + charge because the protons are pulling in the same direction, whereas the electrons are scattered.

LargeLarge SmallSmall

All values are innanometers

.

Atomic Radius

Ionization energy is the minimum energy (kJ/mol) required to remove an electron from a gaseous atom in its ground state.

I1 + X (g) X+

(g) + e-

I2 + X+(g) X2+

(g) + e-

I3 + X2+(g) X3+

(g) + e-

I1 first ionization energy

I2 second ionization energy

I3 third ionization energy

I1 < I2 < I3

Ionization Energy

• Ionization energy is always endothermic, that is energy is added to the atom to remove the electron.

• The larger the atom is, the easier it is to remove its electrons.

• The energy required to remove an electron from an atom reduces as the size of the atom increases

• Ionization energy and atomic radius are inversely proportional.

Ionization energy is the energy required toremove an electron from an isolated gaseous atom

Ionization Energy Cont’d

.15

Factors Affecting Ionization Energy

Nuclear ChargeNuclear ChargeThe larger the nuclear charge, the greater the ionization energy.

Shielding effectShielding effectThe greater the shielding effect, the less the ionization energy.RadiusRadius

The greater the distance between the nucleus and the outer electrons of an atom, the less the ionization energy.

SublevelSublevelAn electron from a full or half-full sublevel requires additional energy to be removed.

Group trends in IE• As you go down a group, first IE

decreases because...

• Atomic radius of the atoms increases

• More shielding.

First Ionization Energies(in kilojoules per mole)

H1312.1

Li520.3

Na495.9

K418.9

Be899.5

Mg737.8

Ca589.9

B800.7

Al577.6

Ga578.6

C1086.5

Si786.5

Ge761.2

N1402.4

P1011.8

As946.5

O1314.0

S999.7

Se940.7

F1681.1

Cl1251.2

Br1142.7

Ne2080.8

Ar1520.6

Kr1350.8

He2372.5

Rb402.9

Sr549.2

In558.2

Sn708.4

Sb833.8

Te869.0

I1008.7

Xe1170.3

Smoot, Price, Smith, Chemistry A Modern Course 1987, page 188

Periodic Trends in IE

• Atoms in the same period have valence electrons in the same energy level.

• Same shielding.• But, nuclear charge increases across the

period• So IE generally increases from left to

right.• Exceptions at full and 1/2 full orbitals.

Firs

t Ion

izat

ion

ener

gy

Atomic number

He

• He has a greater IE than H.

• same shielding

• greater nuclear charge

H

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li has lower IE than H

Outer electron further away

outweighs greater nuclear charge

Li

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Be has higher IE than Li

same shielding

greater nuclear charge

Li

Be

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

O

• Breaks the pattern, because the outer electron is paired in a p orbital and experiences inter-electron repulsion.

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

O

F

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

O

F

Ne• Ne has a lower IE

than He• Both are full,• Ne has more

shielding• Greater distance

Question Arrangement of Elements by First Ionization Energy

PLAN:

SOLUTION:

PROBLEM: Using the periodic table only, rank the elements in each of the following sets in order of decreasing IE1:

(a) Kr, He, Ar (b) Sb, Te, Sn (c) K, Ca, Rb (d) I, Xe, Cs

IE increases as you proceed up in a group; IE increases as you go across a period.

(a) He > Ar > Kr

(b) Te > Sb > Sn

(c) Ca > K > Rb

(d) Xe > I > Cs

Group 8A(18) - IE decreases down a group.

Period 5 elements - IE increases across a period.

Ca is to the right of K; Rb is below K.

I is to the left of Xe; Cs is further to the left and down one period.

Electronegativity

Electronegativity is a measure of the ability of an atom in a molecule to attract electrons to itself.

• This concept was first proposed by Linus Pauling (1901-1994) who later won a Nobel Prize for his efforts.

Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond

Electronegativity - relative,

X (g) + e- X-(g)

Group Trends in Electronegativity• So as you go down a group in the periodic

table, the electrons are further away from the nucleus and better shielded from the nuclear charge and thus not as attracted to the nucleus.

• For that reason the electronegativity decreases as you go down the periodic table.

.29

Periodic Trends in Electronegativity:

• The atoms have same energy levels but size decreases across the period.

• Hence, Electronegativity increases from left to right across a period

• F is highest - or most electronegative element

Summary of Periodic Trends

Periodicity of Period 3 Elements

The Period 3 elements

1234567

Major properties which change across the period are:• Structure and bonding• Elements change from metal through metalloid

to non metals.• Acid-base properties• Redox properties• Solubility and complexing propertiesThe changes are related to Change in size of the atomChange in nuclear chargeIncreasing number of valence electrons

The elements show graduation in properties with exception of argon.

Na Mg Al Si P S ClEo -2.71 -2.37 -1.66 -0.48 +1.36

AR

IEEN 0.9 1.2 1.5 1.8 2.1 2.5 3.0O.S +1 +2 +3 +4 +5,+3 +4,+6 -1to+7

Reactions with oxygen

For example, when sodium is burned in oxygen the oxidation state of the sodium increases from 0 to +1 (oxidation), while the oxidation state of the oxygen decreases from 0 to -2 (reduction).

4Na(s) + O2(g) → 2Na2O(s)

The reactions of the period 3 elements with oxygen are redox reactions. In each reaction, the oxidation state of the elements increases and the oxidation state of the oxygen decreases.

0 0 +1 -2

Reactions with oxygen: summary

S(s) + O2(g) → SO2(g)burns with a blue flame

4P(s) + 5O2(g) → P4O10(s)burns spontaneously with a bright white flame and smoke

Si(s) + O2(g) → SiO2(s)burns with a bright white flame and white smoke

4Al(s) + 3O2(g) → 2Al2O3(s)burns vigorously with a bright white flame

2Mg(s) + O2(g) → 2MgO(s)burns vigorously with a bright white flame

4Na(s) + O2(g) → 2Na2O(s)burns vigorously with a yellow flame

S

P

Si

Al

Mg

Na

EquationDescriptionElement

The metals are oxidised and their oxidation state increases.

2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)

Mg(s) + H2O(g) → MgO(s) + H2(g)

The reactions with water are all redox reactions.

0 +1 -2 +1 -2 +1 0

0 +1 -2 +2 -2 0

The hydrogen is reduced and its oxidation state decreases.

The chlorine is both oxidized and reduced.

Cl2(aq) + H2O(l) HClO(aq) + HCl(aq)

0 +1 -2 +1 -2 +1 -1+1

Reactions with water: summary

–no reaction

dissolves to formchlorine water

–no reaction

–no reaction

–no reaction

–no reaction

Mg(s) + 2H2O(l) → Mg(OH)2(aq) + H2(g)

Mg(s) + H2O(g) → MgO(s) + H2(g)

slow with cold water;vigorous with steam

2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)vigorous

Ar

Cl

S

P

Si

Al

Mg

Na

EquationDescriptionElement

Cl2(aq) + H2O(l) HClO(aq) + HCl(aq)

OXIDES OF PERIOD 3 ELEMENTS

Oxides are binary compounds of an element with oxygen. E + O EOElements can form three types of oxides depending on oxidation state of oxygen:Normal oxide - oxidation state of –II O2-

Peroxide- oxidation state –I ion is O22-

Super oxide = oxidation state -½ O2-

NORMAL OXIDES NORMAL OXIDES

The normal oxides of Periods 3 elements

can be grouped into 3 types according to

the nature of their bonding:

1. Ionic oxides; Na and Mg

2. Ionic oxides with high covalent character; Al

3. Covalent oxides Si-Cl

42

Na2O MgO Al2O3 SiO2

P4O6

P4O10

SO2

SO3

Cl2OCl2O7

Bond Ionic

Ionic with

covalent character

Covalent

Periodicity in nature of bonding in the oxides of Periods 3 elements

Group 1 and 2 Oxides• Na and Mg are metals (form cations) ; they

bond with O2- to form ionic oxides.• The oxide ion can bond with H+ ions and they

act as bases dissolving in water to give alkaline solutions.

• Na2O(s) + H2O(l) 2Na+(aq) + 2OH-

(aq)• They will also neutralize acids to produce salt

and water.• MgO(s) + 2HCl(aq) Mg2+(aq) + 2Cl-(aq)

Group 1 and 2 Oxides are BASIC

Aluminum Oxide• Aluminum oxide does not dissolve in water easily .• It is AMPHOTERIC which means it will react with

(and dissolve in) acids and bases.• Acting like a base:• Al2O3(s) + 6H+(aq) 2Al3+(aq) + 3H2O(l)

• Al2O3(s) + 3H2SO4(aq) Al2(SO4)3(aq) + 3H2O(l)

• Acting like an acid:• Al2O3 (s) + 3H2O(l) + 2OH-(aq) 2Al(OH)4

-(aq)

• Al2O3(s) + 2OH-(aq) 3H2O(l) + 2Al(OH)4-(aq)

Acidic Oxides

• The remaining oxides of period 3 (Si – Cl) form acidic solutions.

• Silicon dioxide has little acid-base activity but it shows weakly acidic properties by slowly dissolving in hot concentrated alkalis to form silicates.

• SiO2(s) + 2OH-(aq) → SiO32-(aq) + H2O(l)

Acidic Oxides

• Phosphorus (V) oxide reacts to form a solution of phosphoric (V) acid, a weak acid

• P4O10(s) + 6H2O(l) → 4H+(aq) + 4H2PO4-(aq)

• Phosporus (III) oxide reacts with water to produce phosphoric (III) acid:

• P4O6(s) + H2O(l) 4H3PO3(aq)

Acidic Oxides• Sulphur (VI)oxide reacts with water to make sulphuric

acid:• SO3(l) + H2O(l) H2SO4(aq)• Suphur (IV) oxide reacts with water to produce

sulphurous acid:• SO2(g) + H2O(l) H2SO3(aq)

• Cl2O7 reacts with water to produce perchloric acid:• Cl2O7(l) + H2O(l) 2HClO4(aq)• Cl2O reacts with water to produce chlorous acid:• Cl2O(l) + H2O(l) 2HClO(aq)

Period 3 oxides and water: summary

covalent

H+(aq), HSO3

-(aq)covalent

H+(aq), H2PO4

-(aq)covalent

– (insoluble)covalent

– (insoluble)ionic/covalent

Mg2+(aq), OH-

(aq)ionic

Na+(aq), OH-

(aq)ionic

H+(aq), HSO4

-(aq)SO3

SO2

P4O10

SiO2

Al2O3

MgO

Na2O

Ions after H2O reaction

BondingOxide Type of solution pH

strongly alkaline

moderately alkaline

–

–

strongly acidic

weakly acidic

strongly acidic

13–14

10

7

7

0–1

2–3

0–1

Properties of the Third Period Oxides

CHLORIDESGroup 1 and 2

• NaCl and MgCl2, are ionic crystalline solids with high melting points.

• NaCl dissolves in water to form a neutral solutionNaCl(s) Na+(aq) + Cl-(aq)

• MgCl2 dissolves to form a slightly acidic solution:

MgCl2(s) Mg2+(aq) + 2Cl-(aq)

• The resulting solutions can conduct electricity due to the free moving ions.

Aluminum chloride• AlCl3 sublimes at 178°C to form Al2Cl6

• AlCl3 dissociates into ions when added to water:AlCl3(s) Al3+(aq) + 3Cl-(aq)

• The aluminum ion is small and has a high charge (3+) thus it has a high charge density.

Aluminum Chloride

• This means it attracts water molecules when in solution and forms the complex ion: [Al(H2O)6]3+

The ion is said to be hydrated

Aluminum Chloride Cont’d

• The ion behaves as an acid be releasing H+ from one of the H2O molecules:

• [Al(H2O)6]3+(aq) [Al(H2O)5OH]2+(aq) + H+(aq)

• Further proton loss can occur:• [Al(H2O)5OH]2+(aq) [Al(H2O)4OH2+(aq) + H+(aq)

• The solution is acidic enough to react with a weak base and produce CO2(g):

• 2AlCl3(aq) + 3Na2CO3(s) 3CO2(g) + Al2O3(s) + 6NaCl(aq)

Chlorides of Silicon- Sulphur

• Chlorides have simple covalent structures.

• The chlorides of non-metals have low mp’s due to weak intermolecular forces between the molecules.

• They react with water to form an acidic solution containing H+, Cl-, O2- or an oxyacid of the element (hydrolysis reaction):

• SiCl4(l) + 2H2O(l) SiO2(s) + 4HCl(aq)

Phosphorus Chlorides

• PCl3 produces phosporous acid and hydrochloric acid:

• PCl3(l) + 3H2O(l) H3PO3(aq) + 3HCl

• PCl5 produces phosphoric acid and hydrochloric acid:

• PCl5(s) + 4H2O(l) H3PO4(aq) + 5HCl(aq)

Chlorine and Water

• In water, Cl2 reacts slowly in a reversible reaction to make a mixture of HCl and HOCl acids:Cl2(aq) + H2O(l) HCl(aq) + HOCl(aq)

• This is a disproportionation reaction where Cl2 is reduced to HCl and oxidized to HOCl.

Properties of the Third Period Chlorides

The s-Block Elements

Li Be

Na

K

Rb

Cs

Fr

Mg

Ca

Sr

Ra

Ba

I II

● Metallic character● Low electronegativity● Basic oxides, hydroxides● Ionic bond with fixed oxidation

states● Characteristic flame colours● Weak tendency to from complex

Characteristic properties

GROUP 1 & 2 ELEMENTS

Atomic radii (nm)

Li 0.152 Be 0.112

Na 0.186 Mg 0.160

K 0.231 Ca 0.197

Rb 0.244 Sr 0.215

Cs 0.262 Ba 0.217

Fr 0.270 Ra 0.220

Atomic radii increases down the group

Group 1 elements are larger than group 2 elements in the same period

Ionization Enthapy

Group I 1st I.E. 2nd I.E.

Li 519 7300

Na 494 4560

K 418 3070

Rb 402 2370

Cs 376 2420

Group I 1st I.E. 2nd I.E. 3rd I.E.

Be 900 1760 14800

Mg 736 1450 7740

Ca 590 1150 4940

Sr 548 1060 4120

Ba 502 966 3390

5.3 Group 2

Element Color Element Color

Li Scarlet Be －Na Yellow Mg －K Lilac Ca Brick-red

Rb Red Sr Crimson

Cs Blue Ba Apple-green

The Flame Color:

Properties of the Alkali Metals• Alkali metals are the largest elements in their

respective periods and their valence electron configuration is ns1.

– The valence e- is relatively far from the nucleus, resulting in weak metallic bonding.

• Alkali metals are unusually soft metals. They can be cut easily with a knife.

• Alkali metals have lower melting and boiling points than any other group of metals.

• Alkali metals have lower densities than most metals.

GROUP 1 ELEMENTS: REACTIONS• Alkali metals are powerful reducing agents.

– They always occur in nature as +1 cations rather than as free metals.

• Alkali metals react vigorously with H2O:

– 2E(s) + H2O(l) → 2E+(aq) + 2OH-(aq) + H2(g)

– Reaction becomes more vigorous down group

(E= ALKALI METAL- Li, Na, K, Rb, Cs)

• Alkali metals reduce halogens to form ionic solids:

– 2E(s) + X2 → 3EX(s) (X = F, Cl, Br, I).

– Alkali metals reduce oxygen in air but product depends on the metal

4Li(s) + O2(g) 2Li2O(s) oxide K(s) + O2(g) KO2(s) superoxide

– Alkali metals reduce Hydrogen to form ionic hydrides

2E(s) + H2(g) 2EH(s)

04/21/23 65

GROUP 2 ELEMENTS: REACTIONS - Metals reduce Oxygen (O2) to form Oxides

2E(s) + O2(g) 2EO(s) E = Mg, Ca, Sr)

Ba + O2 BaO2 (Barium Peroxide)

– Larger metals reduce water to form hydrogen gasE(s) + 2H2O(l) E2+aq) + 2OH- (aq) + H2(g)

(E = Ca, Sr, Ba)

– Metals reduce Halogens to form ionic halidesE(s) + X2 EX2(s) (X = F(not with Be), Cl, Br, I)

– Metals (Be exception) reduce Hydrogen to form ionic hydrides

E(s) + H2(g) EH2 (s) (except Be)

– Elements reduce Nitrogen to form ionic Nitrides3E(s) + N2(g) E3N2(s) (except Be)

– Element Oxides are Basic (except for amphoteric BeO)

EO(s) + H2O(l) E2+(aq) + 2OH-(aq)

Reactions of oxides and hydroxides

2. All group I oxides/hydroxides are basic and the basicity increases down the group.

3. Group II oxides/hydroxides are generally less basic than Group I. Beryllium oxide/hydroxide are amphoteric.

All group I oxides reacts with water to form hydroxides

Oxide: O2- + H2O 2OH-

Peroxide: O22- + 2H2O H2O2 + 2OH-

Superoxide: 2O2- + 2H2O 2OH- + H2O2 + O2

Reactions of chlorides●All group I chlorides are ionic and readily

soluble in water. No hydrolysis occurs. • Group II chlorides show some degree of

covalent character.Beryllium chloride is covalent and hydrolysis

to form Be(OH)2(s) and HCl(aq).BeCl2 + 2 H2O Be(OH)2(s) + 2 HCl(aq) Magnesium chloride is intermediate, it

dissolves and hydrolysis slightly. Other group II chlorides just dissolve

without hydrolysis.

Reactions of hydrides

Hydrides react readily with water to give the metal hydroxide and hydrogen due to the strong basic property of the hydride ion, H:-

H:-(s)+ H2O(l) H2(g)+ OH-(aq)

Thermal StabilityThermal stability describes how easily a compound will decompose on heating. Increased thermal stability means a higher temperature is needed to decompose the compound.

Li2CO3 Li2O + CO2 ( at 700oC)All other group I carbonates are stable at ~800oC

BeCO3 (at 100oC); MgCO3 ( at 540oC); CaCO3 ( at 900oC) SrCO3 ( at 1290oC); BaCO3 ( at 1360oC)

All Carbonates undergo thermal decomposition to the oxideECO3(s) EO(s) + CO2(g)

72

• Certain Period 2 elements exhibit properties that are very similar to those of the Period 3 elements immediately below and to the right.

DIAGONAL RELATIONSHIPS

Lithium and Magnesium reflect similar atomic and ionic size

Both elements form: Nitrides, Hydroxides and Carbonates that

decompose with heat, Organic compounds with polar

covalent metal-carbon bonds Salts with similar solubilities

Uses of s-block compounds● Sodium carbonate

● Manufacture of glass● Water softening● Paper industry

● Sodium hydrocarbonate● Baking powder● Soft drinkBeryl (Be3Al2Si6O18) - Gemstone, source of BeMagnesium oxide (MgO) – Refractory material for furnace bricks

Uses of s-block compounds

● Sodium hydroxide● Manufacture of soaps, dyes, paper and drugs● To make rayon and important chemicals

● Magnesium hydroxide● Milk of magnesia, an antacid

● Calcium hydroxide● To neutralize acids in waste water treatment

● Strontium compound● Fireworks, persistent intense red flame

●

outermost shell electronic configuration of ns2np2

Group 14 Elements Group 14 Elements

carbon

silicon

germanium

tin

lead

carbon

silicon

germanium

tin

lead

Moving down the group

non-metal

metalloids

metals

exhibit a marked change (dissimilarity) among the elements in the same group

Group 14 elements exhibit Allotropy

• Allotropes are different crystalline or molecular forms of the same element.

• One allotrope of a particular element is usually more stable than another at a particular temperature and pressure.

Carbon has several allotropes, including graphite, diamond, and fullerenes.

two important allotropic forms

diamond and graphite

two allotropes

white tin and grey tin

TinTin

White tin Grey tincold

heatMetallic non metallic

Examples

carbon (diamond)

silicon

germanium

grey tin (an allotrope of tin)

Structure and BondingStructure and Bonding

Most common structure : giant covalent

structure

each carbon atom is bonded to four

other C atoms sp3hybridization

Structure of diamond

extremely hard and chemically inertAll electrons are localized non-conductor

Structure of graphite

layered structure Covalent bonds

van der Waals’ forces

Electrons between layers are delocalized

conducts electricity along the layers

The layers slide over each other easily

brittle and soft

network lattice

the atoms are covalently bonded

to one another

2. 2. Silicon and GermaniumSilicon and Germanium

White tin

stable form

metallic lattice structure

atoms are held together by

metallic bonding

3. 3. TinTin

conducts electricity

shows the properties of a typical

metal

Grey tin

network lattice structure

similar to that of diamond

White tin Grey tincold

heatmore dense less

dense

White tin expands and crumbles on cooling

Napoleon’s retreat from Russia

3. 3. Tin and LeadTin and Lead

typical metallic lattice

atoms are held together by

metallic bonding

4. Lead4. Lead

C Si Ge

Electronegativi

ty value

2.5 1.74 2.0

Electronic

configuration

1s22s22p

2

[Ne]

3s23p2

[Ar]

3d104s24p2

Atomic radius

(nm)

0.077 0.117 0.122

Bond enthalpy

(kJ mol–1)

347 226 188

Melting point

(C)

3527 1414 1211

Boiling point

(C)

4027 3265 2833

Some physical properties of the Group 14 elements

Sn Pb

Electronegativity

value1.7 1.55

Electronic

configuration[Kr]4d10

5s25p2

[Xe] 4f145d10

6s26p2

Atomic radius (nm) 0.140 0.154

Bond enthalpy (kJ

mol–1)150 –

Melting point (C) 232 327

Boiling point (C) 2602 1749

Some physical properties of the Group 14 elements

Special features of Carbon

• C cannot expand it’s octet. It has no empty d orbital to accommodate electrons.

• Carbon can catenate. (can form -C-C-C- C- chains. This ability explains formation of huge number of hydrocarbons. Silicon forms only Si-O-Si-O chains in silica

Special features of Carbon Cont’d

• Carbon is only member of the group that can form p bonds. Explains the formation of >C=C<, >C=O,- C=N bonds.

• Carbon forms gaseous oxides, CO and CO2.

on going down the groupThe very high m.p. of diamond is due to the strong C – C bonds & the giant structure

Going from C to Ge

bond length

bond strength

melting point

boiling point

Variation in Melting Point &Boiling PointVariation in Melting Point &Boiling Point

C 3527C

Si 1414C

Ge 1211C

Sn 232C

Pb 327C

4027C

3265C

2833C

2602C

1749C

M.P. B.P.

Variation in Melting Point & Boiling PointVariation in Melting Point & Boiling Point

Sn and Pb have exceptionally low m.p. because

C 3527C

Si 1414C

Ge 1211C

Sn 232C

Pb 327C

1. metallic structures

extent of bond breaking on melting is small2. only two (ns2) of the

four valence electrons are involved in the sea of electrons

ChloridesChlorides

Two series of chlorides formed by the Group 14 elements

the dichlorides (MCl2)

the tetrachlorides (MCl4)

ChloridesChlorides

All Group IV elements

form tetrachlorides

liquids at room temperature and pressure

all are simple covalent molecules with a tetrahedral shape

CCl4

SiCl4

GeCl4

SnCl4

PbCl4

M – Cl bonds are polar with ionic character

+ -

-

-

-

Molecules as a whole are non-polar

Reactions with waterReactions with water

CCl4 + H2O no reaction

SiCl4 + H2O Si(OH)Cl3 + HCl

H4SiO4, silicic acid

Si(OH)Cl3 + H2O Si(OH)2Cl2 + HCl

Si(OH)2Cl2 + H2O Si(OH)3Cl + HCl

Si(OH)3Cl + H2O Si(OH)4 + HCl

Si in SiCl4 is more positively charged than C in CCl4

More susceptible to nucleophilic attack

Cl

Si

ClCl

Cl

O

H

H

Cl Si

O

Cl

Cl

Cl

H

H

+

Si, unlike C, can expand its octet to accept an additional electron pair

Cl

Si

ClCl

Cl

O

H

H

Cl Si

O

Cl

Cl

Cl

H

HCl

Si

ClOH

Cl+ HCl

+

OH

Si

HOOHOH H4SiO4, silicic acid

ChloridesChlorides

all possess covalent character though they exist as

crystalline solids at room temperature and pressure

tendency to form dichlorides, MCl2

down the group-

-

GeCl2

SnCl2

PbCl2

On moving down the group,

Metallic character of elements

Ionic character of MCl2

-

-

GeCl2

SnCl2

PbCl2

mainly covalent

mainly ionic

ChloridesChlorides

On moving down the group,

the relative stability of +4 oxidation state

the relative stability of +2 oxidation state

ChloridesChlorides

The outermost ns2 electrons are less shielded by the more diffused inner d and/or f electrons.

Tin (Sn) Lead (Pb)

[Kr]4d10 5s25p2 [Xe] 4f145d10

6s26p2

They are attracted more by the positive nucleus

Less available for forming bonds

Form only two bonds using np2

OxidesOxides

Two series of oxides are formed by the Group 14 elements

the monoxides (MO) oxidation state II

the dioxides (MO2) oxidation state IV

monoxides are more basic than the dioxides.

Oxides of C and Si are acidic those of Ge, Sn and Pb are amphoteric

OxidesOxides

Carbon dioxide (CO2)

the only dioxide which consists of simple molecules

exists as a gas at room temperature and pressure

All Group IV elements

form the dioxides

OxidesOxides

The dioxides of other Group IV elements

crystalline solids of high melting points

either giant covalent or giant ionic structures

CO2

SiO2

GeO2

SnO2

PbO2

stability of dioxide

Decrease

down the group

OxidesOxides

CO2 dissolves in water to form an acid

CO2 + H2O HCO3

Si

the monoxides (MO) oxidation state II

the dioxides (MO2) oxidation state IV

MonoxidesMonoxides

All Group IV elements (except silicon)

form the monoxides at normal conditions

CO

-

GeO

SnO

PbO

Stability of MO down the group

Group IV

element

Oxide

s

forme

d

Bond type of

the oxide

Relative

stability

CarbonCO Covalent

Unstable

(reducing)

CO2 Covalent Stable

Silicon(SiO) – Very unstable

SiO2 Covalent Stable

Germaniu

m

GeOPredominantly

ionic

Unstable in the

presence of O2

GeO2

Partly ionic,

partly covalentStable

The bond type and the relative stabilitiy of the monoxides and dioxides formed by the Group IV

elements

Group IV

element

Oxides

formed

Bond type of the

oxide

Relative

stability

Tin

SnO Predominantly ionic Unstable

(reducing)

SnO2

Partly ionic,

partly covalent

Unstable

(oxidizing)

Lead

PbO Ionic Stable

PbO2 Predominantly ionic Unstable

(oxidizing)

The bond type and the relative stabilitiy of the monoxides and dioxides formed by the Group IV

elements

Group 16 Elements

Element EC AR EN Ox. states

O [He] 2s22p4

S [Ne] 3s23p4

Se [Ar] 4s2 3d10 4p4

Te [Kr] 5s2 4d10 5p4

Po [Xe] 6s2 5d10 6p4

Group 16 Elements

• Oxygen, like nitrogen, occurs as a low-boiling diatomic gas, O2.

• Sulfur, like phosphorus, occurs as a polyatomic molecular solid.

• Selenium, like arsenic, commonly occurs as a gray metalloid.

• Tellurium, like antimony, displays network covalent bonding.

• Polonium, like bismuth, has a metallic crystal structure.

GROUP 16 REACTIONS• Reacts with halogens to form halides E(s) + X2(g) → various halides

(E = S, Se, Te; X = F, Cl)

• The other elements in the group are oxidized by O2:E(s) + O2(g) → EO2

(E = S, Se, Te, Po)

• SO2 is oxidized further:2SO2(g) + O2(g) → 2SO3(g)

• The thiosulfate ion is formed when an alkali metal sulfite reacts with sulfur:S8(g) + 8Na2SO3(s) → 8Na2S2O3(aq)

Allotropes in the Oxygen Family

Oxygen has two allotropes:- O2, which is essential to life, and- O3 or ozone, which is poisonous.Sulfur has more than 10 different forms, due to the ability of S to catenate. S–S bond lengths and bond angles may vary greatly.

Selenium has several allotropes, some consisting of crown-shaped Se8 molecules.

Hydrides of the Oxygen Family

• Oxygen forms two hydrides:

– water (H2O)

– hydrogen peroxide (H2O2).

– H2O2 contains oxygen in a -1 oxidation state.

– H2O and H2O2 can form H bonds, and therefore have higher melting and boiling points than other H2E compounds.

The hydrides of the other 16 elements are foul-smelling, poisonous gases.

• H2S forms naturally in swamps from the breakdown of organic matter and is as toxic as HCN.

Hydride bond angles decrease and bond lengths increase down the group.

Halides of the Oxygen Family

Except for O, the Group 16 elements form a wide range of halides.Their structure and reactivity patterns depend on the sizes of the central atom and the surrounding halogens.

As the central atom becomes larger, the halides become more stable.

This pattern is related to the effect of electron repulsions due to crowding of lone pairs and halogen atoms around the central atom.This is opposite to the previously observed bonding patterns, where bond strength decreases as bond length increases.

Highlights of Sulfur Chemistry

• Sulfur forms two important oxides:

– SO2 has S in its +4 oxidation state. It is a colorless, choking gas that forms when S, H2S or a metal sulfide burns in air.

– SO3 has S in the +6 oxidation state.

• Sulfur forms two important oxoacids.

– Sulfurous acid (H2SO3) is a weak acid with two acidic protons.

– Sulfuric acid (H2SO4) is a strong acid, and is an important industrial chemical. It is an excellent dehydrating agent.

GROUP 7A(17) REACTIONS

• The halogens (X2) oxidize many metals and nonmetals. The reaction with H2 is characteristic:

– X2 + H2(g) → 2HX(g)

• The halogens undergo disproportionation in water:

– X2 + H2O(l) HX(aq) + HXO(aq) (X = Cl, Br, I)

• In aqueous base, the reaction goes to completion to form hypohalites and, at higher temperatures, halates:

3Cl2(g) + 6OH-(aq) ClO3-(aq) + 5Cl-(aq) + 3H2O(l)

Δ

Chemical properties• Halogens (Cl2,Br2 and I2) – Group 17 (VII)

General Properties – • diatomics, colored, phase changes as one goes down

the family. Cl2 is gas (green yellow), Br2 is liquid (brown/red) and Iodine is a purple solid

• not soluble in water (non polar substance) (hence use of oil in experiments-non polar to dissolve halogens).

General Reactivity- • highly reactive due to need for a single electron to fill

valence shell

Group 7A Elements (ns2np5, n 2)

8.6

Halogens• Reactivity decreases as one goes down the halogen

family.

• Halogens will react by adding an electron to themselves (they behave as oxidizing agents -they are reduced - gain electrons). The smallest and most electronegative element F is the most reactive.

• Valence electrons that are farther from the nucleus will have less attraction and are therefore less reactive.

Group 7A Elements (ns2np5, n 2)

X + 1e- X-1

X2(g) + H2(g) 2HX(g)

Incr

easi

ng r

eact

ivity

8.6

Halide Ions (F-, Cl-, Br- and I-)

• Reactivity oxidizing power of the ions decreases going

down the table (size of atom increases and attraction for electrons decreases) so Cl will oxidize I but I will not oxidize Cl (higher halogen will displace a lower halogen from its salts.)

Halide Ions

• Reactions : assume that the halogen is the one reacting by removing electrons from the ion, therefore if the halogen (diatomic) is higher on the table than the ion , the reaction will take place, but if the ION is higher on the table than the HALOGEN the reaction will not take place.

• Cl2 + 2 I- → I2 + 2 Cl-

• Br2 + 2 I- → I2 + Br-

• I2 + 2 Br- → no rxn

Reactivity of the Halogens

A halogen atom needs only one electron to fill its valence shell. Halogens are therefore very reactive elements.

The halogens display a wide range of electronegativities, but all are electronegative enough to behave as nonmetals.

A halogen will either- gain one electron to form a halide anion or- share an electron pair with a nonmetal atom.

The reactivity of the halogens decreases down the group, reflecting the decrease in electronegativity.

Why does melting point increase going down the halogens?

• The halogens are diatomic molecules, so F2, Cl2, Br2, I2

• As the molecules get bigger there are more electrons that can cause more influential intermolecular attractions between molecules.

• The stronger the I.M. forces, the more difficult it will be to melt. (more energy needed to break the I.M. forces)

Figure 14.22 Bond energies and bond lengths of the halogens.

F2 shows an anomalous bond energy. The F-F bond is weaker than expected since the lone pairs on the small F atom repel each other more than the lone pairs of other halogens.F2 is the most and I2 the least reactive halogen.

Figure 14.23 The relative oxidizing ability of the halogens.

Halogens are strong oxidizing agents. The oxidizing ability of X2 decreases down the group while the reducing ability of X- increases.

Cl2(aq) + 2I-(aq) → 2Cl-(aq) + I2 (in CCl4)

Figure 14.23 continued

Cl2 is a stronger oxidizing agent than I2. Cl2 will therefore displace I- from solution. I2 will not displace Cl- ions.

Interhalogen Compounds

Halogens bond with each other to form interhalogen compounds.The central atom will have the lower electronegativity and a positive oxidation state.

The interhalogens illustrate a general principle of oxidation states: odd-numbered groups exhibit odd-numbered oxidation states while even-numbered groups exhibit even-numbered oxidation states.When bonds form or break, two electrons are involved, so the oxidation states of the atoms involved commonly change by 2.

Odd-numbered oxidation states:

I2 + F2 → 2IF0

0

+1

-1IF + F2 → IF3

0-1

+1 +3

-1

Even-numbered oxidation states:

F and I are both in Group 7A, an odd-numbered group.

S is in Group 6A, an even-numbered group.

S + F2 → SF2

+20SF2 + F2 → SF4

+2 +4

Figure 14.24 Molecular shapes of the main types of interhalogen compounds.

ClF

linear, XY

BrF5

Square pyramidal, XY5

BrF5

Square pyramidal, XY5

86

IF7

Pentagonal bipyramidal, XY7

IF7

Pentagonal bipyramidal, XY7

90

ClF3

T-shaped, XY3

ClF3

T-shaped, XY3

88

Figure 14.25 Chlorine oxides.

dichlorine monoxideCl2O

chlorine dioxideClO2

lone e-

dichlorine heptaoxideCl2O7

Table 14.4 The Known Halogen Oxoacids*

Relative Strength of Halogen Oxoacids

The relative strength of halogen oxoacids depends on both the electronegativity and the oxidation state of the halogen.

For oxoacids with the halogen in the same oxidation state, acid strength decreases as the halogen EN decreases.

HOClO2 > HOBrO2 > HOIO2

For oxoacids of a given halogen, acid strength decreases as the oxidation state of the halogen decreases.

HOClO3 > HOClO2 > HOClO