Post on 26-Mar-2015
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)