HL Topic3+13 Periodicity

transcript

PERIODICITY

The Periodic Table

• from left to right, in order of increasing proton number (atomic number)

• horizontal rows or periods

3Li 4Be 5B 6C 7N 8O 9F 10Ne

11Na 12Mg

13Al 14Si 15P 16S 17Cl 18Ar

The Periodic Table

1H 2He

11Na 12Mg

• vertical groups

13Al 14Si 15P 16S 17Cl 18Ar

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18

• blocks: ss--blockblock, pp--blockblock, dd--blockblock and ff--blockblock

The Periodic Table

1H 2He

11Na 12Mg

• arrangement is such that elements show periodicity in physical and chemical properties

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18

13Al 14Si 15P 16S 17Cl 18Ar

periodicity

the repeating pattern of physical and chemical properties

Atomic Radii

• Atomic radii describes the relative sizes of atoms.

• Atomic radii increase within a column going from the top to the bottom of the periodic table.

• Atomic radii decrease within a row going from left to right on the periodic table.– This last fact seems contrary to intuition.– How does nature make the elements smaller even though the electron number is increasing?

Periodic Trends in Atomic Radius

Mg Al Si P S

Atomic Radii

• The reason the atomic radii decrease across a period is due to shielding or screening effect.– Effective nuclear charge, Zeff, experienced by an electron is less than the actual nuclear charge, Z.

– The inner electrons block the nuclear charge’s effect on the outer electrons.

• Moving across a period, each element has an increased nuclear charge and the electrons are going into the same shell (2s and 2p or 3s and 3p, etc.).– Consequently, the outer electrons feel a stronger effective nuclear charge.

– For Li, Zeff ~ +1 For Be, Zeff ~ +2

Atomicradius/nm

Nuclearcharge

Atomic Radius Trend across Period 3

+12 +13 +14 +15 +16 +17 +18+11

Mg Al Si P S Cl ArNa

0.145 0.118 0.111 0.098 0.088 0.0790.190 0.071

decreases from left to right

Factors determining size

• Nuclear charge– the greater the nuclear charge, the

smaller the size

• Shielding effect by inner electrons– the greater the shielding effect, the

greater the size

• Number of shells– the greater the

number of shells, the greater the size

+11+17

Reasons for trend across the Period

• Nuclear charge increases across the Period

• Shielding effect is relatively constant

Hence, overall atomic radius decreases across the Period.

• Number of shells remains constant

Element

Atomic Radius/nm

Nuclear Charge

ElectronicStructure

2.8.8.1

2.8.18.8.1

2.8.18.18.8.1

Atomic Radius Trend down Group 1

Atomic Radius vs Proton Number (Group 1)

3 11 19 37 55

AtomicRadius/nm

Proton Number

• The number of shells increases

• The increase in nuclear charge is nullified by the shielding effect of inner electrons

Reasons for trend down a Group

Atomic Radii

• Example 6-1: Arrange these elements based on their atomic radii.– Se, S, O, Te

You do it!You do it!

O < S < Se < Te

Atomic Radii

• Example 6-2: Arrange these elements based on their atomic radii.– P, Cl, S, Si

Cl < S < P < Si

Atomic Radii

• Example 6-3: Arrange these elements based on their atomic radii.– Ga, F, S, As

F < S < As < Ga

Ionic Radii

• Anions (negative ions) are always largerthan their neutral atoms.

1.191.261.71Ionic

Radius(Å

F1-O2-N3-Ion

0.720.730.75Atomic

Radius(Å)

FONElement

Ionic Radii

• Cations (positive ions) are always smallerthan their respective neutral atoms.

0.590.90Ionic

Radius (Å)

Be2+Li+Ion

1.121.52Atomic Radius (Å)

BeLiElement

0.680.851.16Ionic

Radius (Å)

Al3+Mg2+Na+Ion

1.431.601.86Atomic Radius (Å)

AlMgNaElement

A neutral atom is larger than its cation.

A neutral atom is smaller than its anion.

Li+ F-

Mg2+ Al3+ P3- S2-

Ionic Radii

• Cation (positive ions) radii decrease from left to right across a period.– Increasing nuclear charge attracts the electrons and decreases the radius.

0.941.321.66Ionic

Radii(Å)

In3+Sr2+Rb+Ion

Ionic Radii

• Anion (negative ions) radii decrease from left to right across a period.– Increasing electron numbers in highly charged ions cause the electrons to repel and increase the ionic radius.

1.191.261.71Ionic

Radii(Å)

F1-O2-N3-Ion

Ionic Radii

• Cation (positive ions) radii decrease from left to right across a period.– Increasing nuclear charge attracts the electrons and decreases the radius.

0.941.321.66Ionic

Radii(Å)

In3+Sr2+Rb+Ion

Ionic Radii

• Anion (negative ions) radii decrease from left to right across a period.– Increasing electron numbers in highly charged ions cause the electrons to repel and increase the ionic radius.

1.191.261.71Ionic

Radii(Å)

F1-O2-N3-Ion

Na Na+

2.8.1Electronic Structure 2.8

NuclearCharge

+11 +11

• Nuclear charge attracting 1 less electron

• Loss of 1 shell

A neutral atom is larger than its cation.

2.8.7Electronic Structure

NuclearCharge

+17 +17

Nuclear charge attracting 1 more electron

A neutral atom is smaller than its anion.

Cl Cl−−−−

Ionic Radius vs Proton Number for Periods 2 and 3

3 4 5 6 7 8 9 11 12 13 14 15 16 17

Ionic radius/nm

Proton Number

Period 3 (Na+ Cl−)

Ionic radius decreases from Na+ to Si4+ because

2.8 2.8 2.8 2.8 2.8.8 2.8.8 2.8.8

Electronic

Structure

Nuclear

Charge

Na+ Mg2+ Al3+ Si4+ P3−−−− S2−−−− Cl−−−−

+11 +12 +13 +14 +15 +16 +17

Na+, Mg2+, Al3+, Si4+ are isoelectronic but the nuclear charge increases from +11 to +14.

Hence, the electrons are pulled closer to the nucleus.

Ionic radius decreases from P3− to Cl− because

P3−, S2−, Cl− are isoelectronic but the nuclear charge increases from +15 to +17

sharp rise in ionic radius from Si4+ to P3−

2.8 2.8 2.8 2.8 2.8.8 2.8.8 2.8.8

Electronic

Structure

Nuclear

Charge

Na+ Mg2+ Al3+ Si4+ P3−−−− S2−−−− Cl−−−−

+11 +12 +13 +14 +15 +16 +17

Electronegativity

Electronegativity is a measure of how strongly an atom of an element attracts valence electrons in a chemical bond.

Electronegativity

Periodic Trends in Electronegativity

Mg Al Si P S

0.93 1.31 1.61 1.90 2.19 2.58 3.16

Going down a Group, Electronegativity decreases because atomic radius increases.

Across a Period, Electronegativity increases because atomic radius decreases.

Electronegativity

• Example 6-11: Arrange these elements based on their electronegativity.– Se, Ge, Br, As

Ge < As < Se < Br

Electronegativity

• Example 6-12: Arrange these elements based on their electronegativity.– Be, Mg, Ca, Ba

Ba < Ca < Mg < Be

Ionisation Energy

The energy required to remove one mole of electrons from one mole of atoms of the element in the gaseous state to form one mole of positively charged gaseous ions.

e.g. H(g) → H+(g) + e- ∆H = +1310 kJ mol-1

Na(g) → Na+(g) + e- ∆H = +494 kJ mol-1

First Ionization Energies of Some Elements

Factors affecting ionisation energies

• Nuclear charge– the greater the nuclear charge, the

greater the I.E.

• Atomic/ ionic size

– the greater the atomic/ ionic size, the lower the I.E.

• Shielding effect by inner electrons

– the greater the shielding effect, the lower the I.E.

Ionization Energy

• First, second, third, etc. ionization energies exhibit periodicity as well.

• Look at the following table of ionization energies versus third row elements.– Notice that the energy increases enormously when an electron is removed from a completed electron shell.

Ionization Energy

435611,58010,5509540IE4

(kJ/mol)

3232274577336912IE3

(kJ/mol)

1577181714514562IE2

(kJ/mol)

786578738496IE1

(kJ/mol)

Group and

element

Ionization Energy

• The reason Na forms Na+ and not Na2+ is that the energy difference between IE1 and IE2 is so large.– Requires more than 9 times more energy to remove the second electron than the first one.

• The same trend is persistent throughout the series.– Thus Mg forms Mg2+ and not Mg3+.– Al forms Al3+.

Ionization Energy

• Example 6-7: What charge ion would be expected for an element that has these ionization energies?

92040IE8 (kJ/mol)

17870IE7 (kJ/mol)

15160IE6 (kJ/mol)

11020IE5 (kJ/mol)

8410IE4 (kJ/mol)

6050IE3 (kJ/mol)

3370IE2 (kJ/mol)

1680IE1 (kJ/mol)

Notice that the largest increase in ionization energies occurs

between IE7 and IE8. Thus this element would form a 1- ion.

Periodic Trend in First Ionization Energy of Period 3

Na Mg Al Si P S Cl Ar

3rd period

Ionization Energy

• Example 6-4: Arrange these elements based on their first ionization energies.– Sr, Be, Ca, Mg

Sr < Ca < Mg < Be

Ionization Energy

• Example 6-5: Arrange these elements based on their first ionization energies.– Al, Cl, Na, P

Na < Al < P < Cl

Ionization Energy

• Example 6-6: Arrange these elements based on their first ionization energies.– B, O, Be, N

B < Be < O < N

Periodic Trend in First Ionization Energy of Group 7

F Cl Br I

Group 7

Periodic Trend in First Ionization Energy of Group 1

Li Na K Rb Cs

Group 1

Periodic Trend in Melting Points of Period 3

Na Mg Al Si P S Cl Ar

3rd period

melting point depends both on the structure of the element and on the type of attractive forces holding the atoms together

302312336371454M.pt/K

CsRbKNaLiGroup 1

Melting point decreases down the group as the atoms become larger and the strength of the metallic bond decreases.

38726617253M.pt/K

I2Br2Cl2F2Group 17

The melting points increase as the van der Waals’

forces of attraction between the diatomic molecules

increases down the group.

CsRbKNaLiAlkali

metals

2,8,8,18,18,12,8,8,18,12,8,8,12,8,12,1Electronic

structure

increasing atomic and ionic radius

decreasing ionization energy

increasing reactivity

decreasing electronegativity

ALKALI METALS

increasing atomic and ionic radius

decreasing ionization energy

increasing reactivity

increasing electronegativity

HALOGENS

solidliquidgasgasState at r.t.p.

blackreddish-

yellowish

pale yellowColour

2,8,8,18,72,8,8,72,8,72,7Electronic

structure

IBrClFHalogens

decreasing oxidising strength

No reactionNo reactionNo reactionNo reactionIodine

Solution turns

yellow, then

brown, then

precipitate

No reactionNo reactionNo reactionBromine

Solution turns

yellow, then

brown, then

precipitate

Solution turns

yellow, then

No reactionNo reactionChlorine

Pale yellow

precipitate

(rapidly

darkens to

black in

sunlight)

No reactionAqueous

I-Br-Cl-F-reagent

summary of reactions of the halide ions

Anomalous Behaviour

Bond Dissociation Energy of halogens

F2: 155

Cl2: 242

Br2: 193

I2: 151

Fluorine Chlorine Bromine Iodine

Anomalous Behaviour

Reduction Potentials of alkali metal ions

Li+: -3.05

Na+: -2.71

K+: -2.92

Rb+:-2.93

Cs+: -2.92

872387185532528274013801156Boiling

point /K

2,8,82,8,72,8,62,8,52,8,42,8,32,8,22,8,1Electronic

structure

ArClSPSiAlMgNaPeriod 3

decreasing atomic radius

increasing ionization energy

increasing electronegativity

decreasing metallic character

increasing hydrolysis of chlorides

increasingly acidic oxides

PERIOD 3

acidicweakly

acidic

neutralNature of

solution

dissolvefumes of HCl produceddissolve easilyReaction

with water

simple molecularionicionicStructure

nonenonenonenonepoorgoodgoodElectrical

conductivity

molten state

-351367658-14121413boiling pt/ oC

-101-80-112-70178 (sublimes)

714801Melting pt/ oC

gasliquidliquid(solid)

liquidsolidsolidsolidState at 25oC

Cl2S2Cl2PCl3(PCl5)

SiCl4Al2Cl6MgCl2NaClformula

Chlorides of Period 3

acidicacidicacidicacidicamphotericbasicbasicNatureof oxide

forms HClO4 anacidic

solution

SO3 forms H2SO4 a

strong acid

forms H3PO4 an acidic solution

Does notreact

Does not react

Forms Mg(OH)2(aq)

weakly alkaline

Forms NaOH(aq)

an alkaline solution

Reaction with water

simple moleculargiant covalent

ionicionicionicStructure

nonenonenonevery

goodgoodgoodElectrical conductivity inmolten state

8045175223029803600-boiling pt/ oC

-9217241610202728521275Melting pt/ oC

liquid (gas)

solid (solid)

solidsolidsolidsolidState at 25oC

(Cl2O)

(P4O6)

SiO2Al2O3MgONa2Oformula

Oxides of Period 3

• exhibit variable oxidation states

• ability to form complex ions

• ability to form coloured compounds

• catalytic property

+2+1,+2+2+2+2,+3+2,+4,

+2,+3,

+3Common

oxidation

states

[Ar]3d104s2[Ar]3d104s1[Ar]3d84s2[Ar]3d74s2[Ar]3d64s2[Ar]3d54s2[Ar]3d54s1[Ar]3d34s2[Ar]3d24s2[Ar]3d14s2Electronic

structure

ZnCuNiCoFeMnCrVTiSc1st row of

d-block

decreasing stability of maximum oxidation state

increasing stability of +2 oxidation state

Variable oxidation states

Formation of Complex Ions

Low energy unfilled d-orbitals or p-orbitals can accept a lone pair of electrons from LIGANDS to form a dative bond.

Ligands : species that can donate a pair of non-bonding electrons to the central metal ion eg water, ammonia, chloride ion.

Result:complex ion

↑↓↑↓

↑↓

↑↑↓ ↑↓

↑↓ ↑↓

white light

Energy corresponds to particular frequency in visible spectrum

∆EEnergy is absorbed as electron is promoted.

Reflected light is complementary colour to

absorbed light.

Formation of Complex Ions

•Stabilises certain oxidation states

•Affects the solubility of compounds

•Major effect on the colour of the solution of a metal ion

Stereochemistry of Complex Ions

H3NNH3

tetrahedral

H2O H2O

octahedral

Square planar

Complex ions are capable of exhibiting cis-trans & optical isomerism

Coordination Number: the number of particles around the central metal

Stereochemistry:cis-trans

Stereochemistry: Chirality

Catalytic properties of Transition Elements

HL Topic3+13 Periodicity

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