CHAPTER 1

s-BLOCK ELEMENTS

(GROUP I AND II)

OBJECTIVE

• Group I (Na Cs) : physical & chemical properties usages

• Group II (Mg Ba) : physical & chemical properties usages

1

2

7

4

5

6

3

Period

GroupIntroduction

Classification of the Elementsinto the s, p, d and f Blocks

Periodic Table• Elements are placed in order of their

atomic number / proton number

• Horizontal rows = periods

• Vertical columns = groups

• DIVIDED TO s-block = group 1, 2 & 18 (He only)p-block = group 13 to 18 (except He)d-block = group 3 to 12

(transition elements)

s-Block Elements

• Groups 1 and 2 s-block metals

• Group 1 elements: 1e- in their outer shell

• Group 2 elements: 2e- in their outer shell

• These outer electrons are located in s-orbital (s sub-shell), ns1 & ns2

• Chemistry of these metals is dominated by the loss of s electrons to form a cation.

Electronic Configuration of Group I (ns1, n 2)

Element Symbol Abbreviated Electron Configuration

lithium Li [He]2s1

sodium Na [Ne]3s1

potassium K [Ar]4s1

rubidium Rb [Kr]5s1

cesium Cs [Xe]6s1

francium Fr [Rn]7s1

Electronic Configuration of Group II (ns2, n 2)

Element Symbol Abbreviated Electron Configuration

beryllium Be [He]2s2

magnesium Mg [Ne]3s2

calcium Ca [Ar]4s2

strontium Sr [Kr]5s2

barium Ba [Xe]6s2

radium Ra [Rn]7s2

Characteristic of Group I (Alkali Metals)

Silvery-coloured metal. Soft, easy to cut by knife. Highly reactive metals. React with water to form alkaline

solution. Have 1e- in outer shell that is easily

given off. Form oxidation state of +1 cation

Characteristic of Group II (Alkaline-Earth Metals)

2e- in the outer shell Usually form M2+ ions in compounds

(most compounds are ionic) Very reactive & are powerful reducing

agents Have the oxidation state of +2 The oxides are all basic (except

beryllium oxide, which is amphoteric)

Compared To Group I Elements:

Group II elements are Harder, with higher melting point &

electrical conductivity less reactive as 2e- must be removed to

form M2+ ions The metallic bonds of Group II elements

are stronger as having 2 valence e- for metallic bonding

Physical Trends when go down Group I and II

Melting point (mp) & boiling point (bp) decreases

Atomic radius increases Ionic radius increases First ionization energy decreases As a reducing agent increases Density increases

1. Atomic RadiiAtomic radius are determined by:

1. Nuclear charge

- attraction between the nucleus and the

outer shell e-

2. Screening/shielding effect

- enlarge the atomic radius by mutual

repulsion between the inner shells e- and

the outer shell e-

3. Number of orbital / electronic shells

1. Atomic Radii

Size of atom decreases when going across a period

Effective Nuclear Charge, Z* increase.

The > no. of proton, the nucleus charge but the shielding effect remains constant.

Each added electron feels a greater & greater +ve charge.

e clouds are pulled closer to the nucleus.

Atomic radius decreases.

Size of atom increases when go down Group I and II

increase in the screening effect is greater than the increase in the nuclear charge due

to the increase in number of electronic shells

nuclear charge & screening effect increase due to the increase in the number

of protons & electrons

The attraction between the nucleus & the electron cloud decreases. Atomic radius

2. Ionic Radius

• Ionic radius of cations & anions decreases when going across a period

• For isoelectronic cations, the more positive the ionic charge, the smaller the ionic radius.

• For isoelectronic anions, the more negative charge, the larger the ionic radius.

• Cation is always smaller than atom from which it is formed.

• Anion is always larger than atom from which it is formed.

Comparison of Atomic Radii with Ionic Radii

3. Melting Point & Boiling PointGenerally: depends on STRUCTURE &

BONDING TYPE

• When going down Group 1, 2 & 3, the mp / bp decreases due to the increase in atomic radius.

• When going down group 4, the mp / bp decreases due to the changes of giant molecular structures with strong covalent bonds (carbon, silicon & germanium) to metallic structures with metallic bonds (stanum & plumbum).

• When going down Group 15, the mp / bp increases due to the changes of simple molecules (N2 & P4) to metal (As & Sb).

• When going down Group 16, 17 & 18, the mp / bp increases due to the increase of the molecules size and the strength of the intermolecular bond.

3. Melting Point & Boiling Point

MP and BP (when across a period)

Period 3 Elements

Structure & Bonding Type of Elements in Period 3:

Element Na Mg Al Si P S Cl Ar

Crystal structure

Metallic Giant molecular

Simple molecular

Bonding Metallic Covalent Van der Waals

Explaination: the trend of MP and BP across a Period 3

Na, Mg and Al• Metals with metallic bonding.• Thus, relatively high mp and bp.• Going from Na to Al, No. of delocalize e- increases Metallic bond stronger More heat energy needed Bp & mp increase in order of :

Na < Mg < Al

Silicon• Has a giant covalent

structure with strong covalent bond

• More energy needed to break the bond

• Relatively high mp and bpSilicon

Phosphorus, Sulphur, Chlorine, Argon

• Van der Waals attraction.

• Mp and bp will be lower than those 1st four elements (Na, Mg, Al and Si).

• Phosphorus exists as P4 molecules

• Sulphur exists as S8 molecules

• Chlorine exists as Cl2 molecules

• Argon exists as individual Ar atoms

• The strength of the Van der Waals forces decreases as the size of the molecule decreases

• So the mp and bp decrease in the order

S8 > P4 > Cl2 > Ar 

4. Ionization Energy (IE)1st IE = min energy, E (kJ/mol) required to remove an e- from a gaseous atom to form 1 mol of gaseous ions under standard condition.

1st IE: X(g) X+(g) + e-

2nd IE: X+(g) X2+

(g) + e-

3rd IE: X2+(g) X3+

(g) + e-

Generally,e- easier to be removed = IEe- difficult to be removed = IE

General Trend in First Ionization EnergiesHigh Energy

Low Energy

Incr

easi

ng 1

st IE

Increasing 1st IE

IE when go down a group

Atomic radius increases down the group

Outer e- further from the nucleus

Outer e- become less strongly attracted by the nucleus

Outer e- easier to be removed

IE decreases when go down a group

IE when across a period

When across period, nucleus charge increases

Atomic radius decreases

Outer e- closer to the nucleus and more strongly attracted

e- more difficult to be removed

Thus, IE increases when across period

Anomalous Cases: IE when across a period

• 2 exception:

Boron and Oxygen has relatively lower IE compared to general trend.

1st exception: Why B has lower IE than Be?

• Boron = 1 e- being lost from the 2p subshell which is further from the nucleus than 2s subshell.

• Thus, easier to be removed, IE decreases.

2nd exception: Why O has lower IE than N?

• The rather low IE for O is due to increased repulsion between the 2e- occupying the same 2p orbital.

• So 1e- is more easily lost.

5. Reducing StrengthWhen going down Group I & II, atomic radius increases as extra shell is added

The valence e- are located further away from the nucleus.

The valence e- are easily to be discarded

Metals become increasingly electropositive & easily to form positively charged ions by losing e-

Reducing strengths increase

Chemical Properties of Group I and II Elements

1. Reactivity trend when go down group I and II.

2. Reaction with oxygen (O2) to form oxide compound.

3. Reaction with water (H2O) to form hydroxide compound.

4. Reaction with halogen (X2) to form halide salt.

Reactivity of Group I Elements

Reactivity of Group I :

• Tendency of Group I elements to lose 1e- forming a singly positive charge ions

Example: Na Na+ + 1 e-

Reactivity of Group II Elements

Reactivity of Group II :

• Tendency of Group II elements to lose

2e- forming a doubly positive charge ions

Example: Mg Mg2+ + 2 e-

Trend of Reactivity when go down Group I and II

Atomic radius increases as extra shell is added

Outer e- become further from the positive nuclear charge

Thus, outer e- easier to be removed & reactivity increases

Attraction becomes weaker

Reaction with O2 to form Oxide Compounds (Group I)

• Alkaline metals burn when heated in oxygen or air.

• They form white oxide powders.

• This is redox reaction in which M undergoes oxidation (0 +1) and O2 undergoes reduction (0 -2).

4M(s) + O2(g) 2M2O(s)

Reaction of Oxide Compound (Group I)

Reaction with water:M2O(s) + H2O(l) 2MOH(aq)

Reaction with acid:(a) M2O(s) + 2HCl(aq) 2MCl(aq) + H2O(l)

(b) M2O(s) + 2HNO3(aq) 2MNO3(aq) + H2O(l)

(c) M2O(s) + H2SO4(aq) M2SO4(aq) + H2O(l)

Reaction with O2 to form Oxide Compounds (Group II)

• Burn in oxygen (air) when heated

• They form white oxide powders.

• This is redox reaction in which M undergoes oxidation (0 +2) and O2

undergoes reduction (0 -2).

2M(s) + O2(g) 2MO(s)  

Reaction of Oxide Compound (Group II)

Reaction with water:MO(s) + H2O(l) M(OH)2(aq or s)

Reaction with acid:(a) MO(s) + 2HCl(aq) MCl2(aq) + H2O(l)

(b) MO(s) + 2HNO3(aq) M(NO3)2(aq) + H2O(l)

(c) MO(s) + H2SO4(aq) MSO4(aq) + H2O(l)

Reaction of Group I Elements with H2O to form Hydroxide

Compounds• The reaction with water is very

exothermic, fast and violent.

• Formation of alkaline solution (hydroxide compounds) together with hydrogen gas

2M(s) + 2H2O(l) 2MOH(aq) + H2(g)

Reaction of Hydroxide Compound (Group I)

Reaction with acid:(a) MOH(aq) + HCl(aq) MCl(aq) + H2O(l)

(b)MOH(aq) + HNO3(aq) MNO3(aq) + H2O(l)

(c) 2MOH(aq) + H2SO4(aq) M2SO4(aq) + 2H2O(l)  

Reaction with halogen:At room temperature,2MOH(aq) + X2(g) MX + MOX + H2O(l)

( X2 = Cl2 , Br2 , I2 )

Reaction of Group II Elements with H2O to form Hydroxide

Compounds• Alkaline earth metals react with water

to form metal hydroxide with the liberation of hydrogen gas

• Beryllium has no reaction with water or steam

• Reactivity of elements with water become more reactive as go down the Group II

M(s) + 2H2O(l) M(OH)2 (s or aq) + H2(g)

(M = Mg, Ca, Sr, Ba)

• However, Mg burns in steam to produce oxide and hydrogen

Mg(s) + H2O(g) MgO(s) + H2(g)

steam

Reaction of Hydroxide Compound (Group II)

Reaction with acid:(a) M(OH)2(aq) + 2HCl(aq) MCl2(aq) + 2H2O(l)

(b) M(OH)2(aq) + 2HNO3(aq)

M(NO3)2(s) + 2H2O(l)

(c) M(OH)2(aq) + H2SO4(aq) MSO4(s) + 2H2O(l)

Reaction of Group I with Halogen, X2

• Heating the metal in chlorine will cause it to burn forming the chloride .

• Salt products, M+X-, are white-colourless crystalline ionic solids that dissolve in water to give neutral solutions of about pH 7.

2M(s) + Cl2(g) 2MCl(s)

Reaction of Group II with Halogen, X2

• Heating the metal in chlorine will cause it to burn forming the chloride.

M(s) + Cl2(g) MCl2(s)

Oxides Properties of Period 3 Elements

Reaction of Oxides with H2OElement Equation

Na2O Na2O(s) + H2O(l) 2NaOH(aq) (soluble)

MgO MgO(s) + H2O(l) Mg(OH)2(aq)

(Slightly soluble)

Al2O3 - insoluble in H2O

- amphoteric as it can react as a base & as an acid As a base:

Al2O3(s) + 6HCl(aq) 2AlCl3(aq) + 3H2O(l)

As an acid: Al2O3(s) + 2NaOH(aq) + 3H2O(l)

2NaAl(OH)4(aq)

Element Equation

SiO2- insoluble in H2O

P4O6

P4O10

P4O6(s) + 6H2O(l) 4H3PO3(aq)

dissolved Phosphonic acid

P4O10(s) + 6H2O(l) 4H3PO4(aq)

dissolved Phosphoric acid

SO2

SO3

SO2(g) + H2O(l) H2SO3(aq)

dissolved Sulphurous acid

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

dissolved Sulphuric acid

Cl2O

Cl2O7

Cl2O(g) + H2O(l) 2HOCl(aq)

Hypochlorous acid

Cl2O7(l) + H2O(l) 2HClO4(aq)

Perchloric acid

Compound of Group I elements

Reaction

Oxide

M2O

• 4M(s) + O2(g) 2M2O(s)

• white ionic solids, very soluble in

water to form the metal hydroxide

Hydroxide

MOH

• 2M(s) + 2H2O(l) 2MOH(aq) + H2(g)

• white ionic solids which very

soluble in water to form strongly

alkaline solutions (pH 13-14).

Compound of Group I elements

Reaction

Chlorides

MCl

• 2M(s) + Cl2(g) 2MCl(s)

• The chlorides are colourless

crystalline solids, soluble in water

to give a neutral solution pH 7

Nitrates MNO3

• MOH(aq) + HNO3(aq)

MNO3(aq) + H2O(l)

• Colourless, soluble, neutral

crystalline salts, formed by

neutralising the alkaline oxide or

hydroxide with nitric acid.

Compound of Group I elements

Reaction

Sulphates

M2SO4

• 2MOH + H2SO4 M2SO4 + 2H2O

• Colourless, soluble, neutral

crystalline salts, formed by

neutralising the alkaline oxide/

hydroxide with sulphuric acid.

Carbonates M2CO3

• 2MOH + CO2 M2CO3 + H2O

• White, soluble, weakly alkaline

solids formed by reacting the

hydroxide with carbon dioxide gas

eg the formation of Na2CO3 & H2O

Group II compounds

Reaction

Oxide

MO

• 2M(s) + O2(g) 2MO(s)

• The oxide, apart from beryllium, is slightly soluble in water forming the alkaline hydroxide, which increases in strength of basic character down the group.

Hydroxide

M(OH)2

• M(s) + 2H2O(l) M(OH)2(aq/s) + H2(aq)

• Magnesium hydroxide and calcium hydroxide (limewater) are sparingly soluble, but the solubility increases down the group, so barium hydroxide is moderately soluble.

Group II compounds

Reaction

Chloride

MCl2

• M(s) + 2HCl(aq) MCl2(aq) + H2(g) 

• M(s) + Cl2(g) MCl2(s) 

• MCO3(aq) + 2HCl(aq)

MCl2(aq) + H2O(l) + CO2(g)

• M(OH)2 + 2HCl(aq) MCl2(aq) + 2H2O(l) 

• Readily react with acids (except be) with increasing vigorous down the group. • Chloride salts are soluble

Nitrate

M(NO3)2

• M + 2HNO3 M(NO3)2 + H2

• MCO3+ 2HNO3 M(NO3)2+ H2O +CO2

• M(OH)2+ 2HNO3 M(NO3)2 + 2H2O

•Nitrate salt are soluble. 

Group II compounds

Reaction

Sulphate

MSO4

• MO + H2SO4 MSO4 + H2O 

• M(OH)2 + H2SO4 MSO4 + 2H2O

• M + H2SO4 MSO4 + H2 

reaction increasingly slower for calcium barium as the sulphate becomes less soluble.

Carbonate

MCO3

• M(OH)2 + CO2 MCO3 + H2O 

• Carbonate salt insoluble in water. • The carbonates decompose on heating to give the oxide and carbon dioxide and exhibit a clear thermal stability trend

Thermal Dissociation of Nitrates, Carbonates & Hydroxides of

Group II Elements

Decompose to form a metal oxide when heated.

e.g. Magnesium compounds

• Mg(NO3)2(s) MgO(s) + 2NO2(g) + 1/2O2(g)

• MgCO3(s) MgO(s) + CO2(g)

• Mg(OH)2(s) MgO(s) + H2O(l)

Uses of Group I Metals

Na vapour is used in the yellow-orange street lamps.

Na liquid is used as a coolant in specialized high-temperature applications

NaCl, 'common salt' is used as a food flavouring and preservative.

Na2CO3 is used in the manufacture of glass and the treatment of hard water.

NaOH is used in the manufacture of soaps, detergents, bleaches, rayon.

KNO3 is used in fertilizers.

Cs, because of its low ionization energy, is used in photosensors in automatic doors, toilets, burglar alarms, and other electronic devices

Uses of Group I Metals

Uses of Elements of Group IICompound Use Reason of use

Be Applications involving radioactivity

Its low atomic noThe lowest tendency to absorb X-rays of metallic elements

Mg Lightweight metal alloys in aircraft, aircraft frames & automobile engine parts

Its low density

Compound Use Reason of use

MgO Refractory (heat resistant) lining of furnaces.

It has a very high mp.

CaO “Quicklime” & Ca(OH)2

Spread onto agricultural land to neutralize excess acidity

CaO and Ca(OH)2 are alkaline.

CaCO3

“Limestone”

Major component of cement

-

Compound Use Reason of use

CaCl2 Road salt to lower the freezing point of water on roads in cold temperatures

-

CaSO4 Plaster casts for broken limbs

It absorbs water and sets to a hard solid

Learning Outcomes

Understand and explain the similarities, variations and trends

in some physical and chemical properties of the elements in

Groups I and II (s-block elements)