The Periodic Table Note 1

7/29/2019 The Periodic Table Note 1

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SEK. MEN. KEB. SULTAN ISMAIL, JOHOR BAHRU.PHYSICAL CHEMISTRY/ UPPERSIX/ 2013

TOPIC : PERIODIC TABLE : PERIODICITY

ARRANGEMENT OF THE ELEMENTS

1. The Periodic Table is an arrangement of all the elements in order of theirincreasing proton number and also based on the electronicconfiguration of their atoms.

2. The position of an element in the Period Table is identified by the Period andGroup it is in.

3. (a) The periods are the horizontal rows. There are are 7 periods in thePeriodic Table.(b) The elements are arranged across a period in order of increasing

proton number.(c) The valence orbits are being filled across a period.

(d) The total number of main (principal) electron shells an atom hasdetermines the period to which it belongs.

4. (a) The groups are the vertical column. There are 18 groups in thePeriodic Table.

(b) A Group contains elements with similar properties and similar valenceelectronic configuration.

(c) The valence electronic configuration of elements of Groups 1,2 and 13 to18 are shown below :

Group valence electronic configuration1 ns1 (1 valence electron)2 ns2 (2 valence electron)13 ns2np1 (3 valence electron)14 ns2np2 (4 valence electron)15 ns2np3 (5 valence electron)16 ns2np4 (6 valence electron)17 ns2np5 (7 valence electron)

18 ns2np6 (8 valence electron)

5. (a) The periodic table is also divided into blocks. Group 1 and 2 are inthe s-blockas the atoms

outermost electrons are in the valence electrons are in the s-orbitals;while Group 13 to 18

are in the p-blockas the valence electrons are in the p-orbitals.(b) In each of the 4th, 5th, 6th and 7th Period, there is a set of10

elements (located in Group 3 to12) called d-block elements. The elements in the d-block have

electrons filling up the d-orbitals of their atoms. For example, the d-block elements in the 4th

Period have their 3dorbitals filled up progressively after the 4s orbital is fully filled

(according to AufbauPrinciple).

(c) In each of the 6th and 7th Period, there is a series off-block elementselements which have

their f orbitals containing electrons. They are called the lanthanides(4f orbitals are filled) and


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the actinides (5f orbitals are filled).Blocks in the Periodic Table

s d

p

f


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4. Trends in atomic radius down a group of the Periodic TableThe atomic radius increases down a group

(a) Down a group, the nuclear charge increases, but the number of electronsand filled innerelectrons shells also increases. Thus the screening effect of the

inner electrons increases,offsetting the increase in nuclear charge.

(b) The outermost electrons in the group are also further from thenucleus. Hence theattraction of the nucleus on the electrons in the outer shell is not as

strong. Therefore, theatomic radius increases.

Ionic Radius

5. Neutral atoms or ions thatbhave the same number of electrons and the sameelectronic configuration are said to be isoelectronic. Table below shows the

ionic radii of four species (O2-

, F-

, Na+

, and Mg2+

) that are isoelectronic. Theyhave the electronic configuration : 1s2 2s2 2p6

Nuclear charge increases.Inner filled shells increaseShielding increasesOuter electrons further fromnucleusAtomic radius increase


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Species O2- F- Na+ Mg2+

Ionic radius

(nm)

0.140 0.136 0.095 0.065

Nuclear

charge

+8 +9 =11 +12

No.of

electrons

10 10 10 10

The ionic radii of four species that are isoelectronic

6. Table above shows that for a given number of electrons, the higher the

nuclear charge, the higher the force of attraction and the smaller the ionic

radius.

7. A positive ion (cation) is smaller than its neutral atom (refer table below)

Neutral atom Na(11p, 11e-) Mg (12p, 12e-) Al (13p, 13e-)

Atomic radius

(nm)

0.156 0.136 0.125

cation Na+ (11p, 10e-) Mg2+ (12p, 10e-) Al3+ (13p, 10e-)

Ionic radius (nm) 0.095 0.065 0.050

Comparing the size of an atom and its positive ion

8. Conversely, a negative ion (anion) is bigger than its neutral atom.

(refer table below)

Neutral atom P (15p, 15e-) S (16p, 16e-) Cl (17p, 17e-)

Atomic radius(nm)

0.110 0.104 0.099

cation P3- (15p, 18e-) S2- (16p, 18e-) Cl- (17p, 18e-)

Ionic radius (nm) 0.212 0.184 0.181

Comparing the size of an atom and its negative ion

9. The variations of ionic radii for Period 2 (Li+ to F-) and Period 3 (Na+ to Cl-) are

shown in figure below.


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10. Explanation for the variation of ionic size for Period 2 and 3

elements

(a) The cation Li+ to B3+ are isoelectronic (containing 2 electrons). Theionic radii decrease from

Li+ to B3+ because the number of electrons remains constant but thenumber of protons

increases from 3 protons to 5 protons. The same basic explanationapplies to cations Na+ to

Al3+ which are isoelectronic to one another (containing 10 elecrons).

Species Li+ Be2+ B3+

Nuclear charge +3 +4 +5

Number ofelectrons

2 2 2

Ionic radius (pm) 60 31 20

(b) The anions N3- to F- are isoelectronic (containing 10 electrons). The ionocsize decreasesfrom N3- to F- because the number of electrons remains constant but thenumber of proton increases from 7 protons to 9 protons.

This basic explanation is also true for the anions P3- to Cl- which areisoelectronic to each other (containing 18 electrons).

Species N3- O2- F-

Nuclear charge +7 +8 +9Number ofelectrons

10 10 10

Ionic radius (pm) 171 140 136

(c) The anions are larger than the cations because they have one extraquantum shell filledwith electrons.

Table below shows the cations and anions of period 2 and 3

Cations/Anion

s

Electronic

Configuration

Li+ , Be2+ , B3+ 1s2


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N3- , O2- , F- 1s2 2s2 2p6

Na+ , Mg2+ ,

Al3+1s2 2s2 2p6

P3- , S2- , Cl- 1s2 2s2 2p6 3s2 3p6

Melting Point, Boiling Point, and Enthalpy of Vaporisation

1. The melting point of an element is the temperature at which the element inthe solid state changes into a liquid (i.e. the solid and liquid state are inequilibrium) at constant pressure.The melting point of an element depends on

(a) The strength of the forces holding the particles together in the solid state;(b) The structure of the element in the solid state.

2. Types of structure and bonding in elements

Variation of Boiling Point, Melting Point And Enthalpy of VaporisationAcross The Second and

Third Period1. Boiling point is defined as the temperature at which the saturated vapour

pressure of a liquidis equal to the external (usually 1 atm)

2. Boiling point is also defined as the temperature at which a liquid is inequilibrium with itsvapour at 1 atm.

3. Melting point is the temperature where a solid is in equilibrium with itsliquid at 1 atm.

4. Enthalpy of vaporisation is the heat energy required to change one mole of asubstance from the liquid to the vapour state, at the boiling point of the


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substance.5. Boiling point, melting point and enthalpy of vaporisation are the measure of

the strength of the attractive forces that holds the particles together in thesolid or liquid state.

6. Generally, the stronger the attractive force the higher the boiling point,melting point and enthalpy of vaporisation.

7. The table below shows the variation of the melting point, boiling point andenthalpy of vaporisation of the Second period elements from lithium to neon.

Element Li Be B C(diamond)

N O F Ne

M.p./0C 181 1278 2300 3022 -210 -218 -220 -249B.p./0C 1330 2480 3930 4827 -200 -183 -190 -245v/kJmol-1

135 294 539 717 2.8 3.4 3.2 1.8

Structureof the

element

Giant metallic Giantcovalent

Simple covalent

8. The table below shows the variation of the melting point, boiling point andenthalpy ofvaporisation of the Third Period elements from sodium to argon.

Element Na Mg Al Si P S Cl ArM.p./0C 98 650 660 1423 44 120 -101 -189B.p./0C 890 1120 2450 2680 280 445 -34 -186v/kJmol

-

1

89 129 294 377 12.4 9.7 10.2 6.5

Structureof theelement

Giant metallic Giantcovalent Simple molecule

9. Consider the Third period:(a) The melting/ boiling points of sodium, magnesium and aluminium are

high because the presence of strong metallic bonds in their giant metallicstructures.

(b) The melting/boiling point increases in the order Na


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Molecules P4 S8 Cl2 ArNo. of

electrons60 128 34 18

(j) This causes the strength of the van der Waals forces and themelting/boiling point to increase in the order:

Ar < Cl2 < P4


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2. Going across a period from left to right, the atomic radius decreases but thenuclear charge

increases. Hence, the attraction for bonding electrons increases.

Trend Of Electronegativity Down A Group

1. The electronegativity of the Group 2 and Group 17 elements are given below.

Group 1Element Be Mg Ca Sr BeElectronegativity

1.5 1.2 1.0 - -

Group 17Element F Cl Br IElectronegati

vity

4.0 3.0 2.8 2.5

2. Going down a group, the atomic radius increases while the effective nuclearcharge remains

almost constant. Hence, the attraction of the atoms for bonding electronsdecreases causing the

electronegativity to decrease.

3. Summary : increasing

Decreasing

Trend Of Electrical Conductivity

1. The graph below shows the trend of electrical conductivity of the second andthird period

elements.


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2. Li, Be, Na, Mg and Al are metals with delocalised electrons in their giantmetallic structures.

Hence, they are good conductors of electricity.3. B, C(diamond), N, O, F, Ne, P, S, Cl and Ar are all non-metals. They do nothave delocalised

electrons in their solid structure. They are non-conductor of electricity.4. Silicon is metalloid. It is essentially a non-metal but with a certain amount ofmetallic properties.

Its conductivity is higher than the non-metals but lower than the metals. It is

a semiconductor. Itsconductivity can be increased by increasing temperature or by the addition

of a measured ofimpurities.

Ionisation Energy

1. The ionisation energy of an element is a measure of the tendency of theatom of the element to

lose electrons to form positive ions.2. The higher the ionisation energy, the more difficult it is for the atom to loseelectrons.

3. The first ionisation is the minimum energy needed to remove one electronfrom every atom in

one mole of gaseous atom to form one mole of gaseous unipositive ionsunder standard

conditions.

Mg(g) M+(g) + e H = 1st I.E.4. The magnitude of the ionisation energy depends on:

(a) atomic radius(b) nuclear charge(c) screening effect


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Trend Of The First Ionisation Energy Across A Period1. The tables and graphs below show the first ionisation energies of the secondAnd third period

elements.Second Period

Element Li Be B C N O F NeProtonNumber

3 4 5 6 7 8 9 10

AtomicRadius

0.152 0.111 0.088 0.077 0.070 0.066 0.064 0.062

1st I.E./KJmol-1

520 900 801 1086 1402 1314 1681 2081

Third periodElement Na Mg Al Si P S Cl ArProton

Number

11 12 13 14 15 16 17 18

AtomicRadius

0.186 0.160 0.143 0.117 0.110 0.104 0.099 0.094

1st I.E./KJmol-1

496 738 578 798 1012 1000 1251 1520

2. Going across a period, the atomic radius decreases while the nuclear chargeincreases. This

causes a general increase in the first ionisation energy of the second andthird period elements

with increasing proton number.

3. However, the increase is not a smooth one. There is a reversal in trendbetween Be and B, and


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between N and O in the second period.4. The same reversal in trend is also evident in the third period between Mg andAl, and between P

and S.5. The first ionisation energies of Be and Mg are higher than expected becausethe first electron

removed is from a stable s2 configuration. This make the removal of theelectron more difficult

than expected.Be : 1s2 2s2

Mg : 1s2 2s2 2p6 3s2

6. An alternative explanation is that the first ionisation energies of boron andaluminium are lower

than expected. This arises because the first electron removed from the twoatoms are from

higher p orbitals. Take boron as example.

2p

2s

1s

Be B

This electron in the 2p orbital in B is at a higher energy level and is alsobeing shielded from the

nucleus by two inner 2s electrons. This makes the electron in B easier to beremoved than

expected. This same applied to aluminium where the first electron removedis from a higher 3p

orbital.

7. The first ionisation energies of nitrogen and phosphorous are higher thanexpected because the

first electron removed is from a stable p3 configuration where all the three

orbitals are singlyfilled.N : 1s2 2s2 2p3

P : 1s2 2s2 2p6 3s23p3

8. An alternative explanation is that the first ionisation energies of oxygen orsulphur are lower

than expected. This arises because the first electron removed from the twoatoms are formed a p

orbital that is occupied by a pair of electrons. Take sulphur as example :

3p


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3sP s

The two electrons in the 3p orbital in sulphur experience mutual repulsion.This makes the

electron easier to be removed than expected.

9. The first ionisation energies of the second period elements are higher thantheir corresponding

third period elements. This is because, atoms of the second period elementsare smaller and with

higher nuclear charge than their counterparts in Period 3.

Trend Of First Ionisation Energy Down A Group

1. The first ionisation energies of the group 2 elements are as shown below.

Element Be Mg Ca Sr BaProtonNumber

4 12 20 38 56

AtomicRadius

0.112 0.160 0.197 0.215 0.222

1st I.E./KJmol-1

900 740 590 550 500

2. Going down Group 2 the atomic radius increases while the effective nuclearcharge remains

almost constant. The attraction between nucleus and electron gets weaker.3. As a result, the first ionisation energy decreases when going down Group 2.4. The same trend is also observed in other groups such as Group 17.

Element F Cl Br I1st I.E./KJ mol-1 1680 1260 1140 1010


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SekolahMenengah Kebangsaan Sultan Ismail, Johor Bahru.

Inorganic Chemistry/ Upper Six/ 2013

Topic : Period 3 Elements

Chemical properties Of The Period 3 Elements

1. Oxidising / reducing power of Period 3 elemets

Metal are usually reducing agents while non-metals are oxidizing agents.

Going across the period (Na to Cl) the oxidizing power of the elements

increases while the reducing power decreases.


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Sodium, magnesium, aluminium Phosphorous, sulphur, chlorine

(a) relatively low ionization energy (a) relatively small atomic

radius and

andlarge atomic radius. high nuclear chargeresulting in

high electron affinity.

(b) form cations by losing electrons (b) form anions by gaining

electrons.

(c) good reducing agents (c) function as oxidizing agents;

P and S

are weak oxidizing

agents while

chlorine is a

powerful oxidizing

agent.

P + 3e P3-

S + 2e S2-

Cl2 + 2e 2Cl-

2. Reaction of the elements with water

(a) The metals become less reactive across the period.

(i) Sodium reacts vigorously with cold water to form the strong alkali

sodium hydroxide and hydrogen gas

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

(ii) Magnesium very slow reaction with cold water (almost negligible)

Mg(s) + 2H2O(l) Mg(OH)2(s) + H2(g)

Reacts vigorously with steam to form magnesium oxideand


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hydrogen

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

(iii) Aluminium does not react with water or steam because it is covered

with a

layer of impermeable Al2O3

(**but will react with warm water if its layer of Al2O3 is

removed,

forming aluminium oxide and hydrogen :

2Al(s) + 3H2O(l) Al2O3(s) + 3H2(g)

(b) Silicon, phosphorous, and sulphur do not react with water.

(c) Chlorine dissolves in water to form hydrochloric acid and chloric

(I) acid.

Cl2(g) + H2O(l) HCl(aq) + HOCl(aq)

3. Reaction of the elements with oxygen

Generally the reactivity decreases across the Period.

(a) Sodium, magnesium and aluminium must be heated to react directly

with

oxygen to form ionic oxides.


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(b) Silicon powder reacts vigorously to form a covalent giant molecular

oxide.

(c) Dry phosphorous can ignite spontaneously in air to form acidic oxides.

(that is why it is stored under water)

(d) Sulphur burns in air to form an acidic oxide.

(e) Chlorine does not react directly with oxygen.

Oxides Of Period 3 Elements

1. Structure, bonding, and melting point of oxides of Period 3

The bonds between the element and oxygen change from ionic to covalent

across the


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period, as the difference in electronegativity between the element and

oxygen decreases.

(a) Sodium, magnesium, and aluminiumare

metals,so they form

ionicoxides.

They have high melting and boiling due to strong electrovalent

bonds between

the oppositely charged ions.

2. Acid/ base nature of the oxides of Period 3 elements

Across the period the oxides change from basic oxides to amphoteric

oxides to

acidic oxide.

(a) Basic oxides : Na2O and MgO

Basic oxides are ionic metal oxides. Soluble basic oxides dissolve

in water to

form alkalis.All basic oxides will react with acids to form their

respective salt

and water.

(i) Sodium oxide

dissolves in water to form sodium hydroxide, a strong alkali;

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

reacts readily with dilute acid to form salt and water.

Na2O(s) + H2SO4(aq) Na2SO4(aq) + H2O(l)

(ii) Magnesium oxide


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is only sparingly soluble in water

(the magnesium hydroxide formed is also only slightly soluble in

water

forming a weak alkaline solution)

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

reacts readily with dilute acid to form salt and water.

MgO(s) + 2HCl(aq) MgCl2(aq) + H2O(l)

(b) Amphoteric oxide : Al2O3

Amphoteric oxides are ionic metal oxides with covalent character.

Amphoteric

oxides have both acidic and basic properties. They react with both

acid and

base to form salts.

(i) Aluminium oxide is ionic but the small ionic radius and highcharge on the

Al3+ ion causes it to polarize the O2- ion. Ad a result the ionic Al2O3

has some

covalent character. Hence Al2O3 is amphoteric, i.e. having both

acidic and

basic properties.

(ii) Aluminium oxide is insoluble in water, but reacts with both acidand alkali

to form soluble salts.

Al2O3(s) + 6H+

(aq) 2Al3+

(aq) + 3H2O(l)

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

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

-(aq)

Al2O3(s) + 2NaOH(aq) + 3H2O(l) 2Na[Al(OH)4](aq) (acidic


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property)

(c) Acidic oxides (oxides of Si, P, S and Cl)

Acidic oxides are covalent oxides ofnon-metals.

Soluble acidic oxides react with water to form acids, and with alkalis to

form salts.

Acidic oxides that are not soluble will react with alkalis to form salts.

Silicon (IV) oxide is not soluble in water but reacts with hot

concentrated

alkalis to form silicates (salts) :

SiO2(s) + 2NaOH(aq) Na2SiO3(aq) + H2O(l)

(sodium silicate)

Oxides of phosphorous, sulphur, and chlorine all dissolve in water to

form

acids.

P4O6(l) + 6H2O(l) 4H3PO3(aq) (phosphoric (III) acid)

P4O10(s) + 6H2O(l) 4H3PO4(aq) (phosphoric (V) acid)

SO2(g) + H2O(l) H2SO3(aq) (sulphuric (IV) acid)

SO3(g) + H2O(l) H2SO4(aq) (sulphuric (VI) acid)

Cl2O(g) + H2O(l) 2HOCl(aq) (chloric (I) acid)

Cl2O7(l) + H2O(l) 2HClO4(aq) (chloric (VII) acid)


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3. Properties of oxides of Period 3 Summary

Formula of

oxide

Na2O MgO Al2O3 SiO2 P4O6

P4O10

SO2

SO3

Cl2O

Cl2O7

Oxidation

number

+1 +2 +3 +4 +3

+5

+4

+6

+1

+7

Structure Ionic lattice structure Giant

covale

nt

simple covalent

structure

Physical

state (at

room temp)

solid liquid

solid

gas gas

liquid

Acid/base

nature

basic amphoter

ic

acidic

Solubility in

water

very

sparingly

pH 14

pH 9

insoluble soluble form acid

Chlorides Of Period 3

1. Properties of chlorides of Period 3- summary

Formula of

chloride

NaCl MgCl2 AlCl3

Al2Cl6

SiCl4 PCl3

PCl5

S2Cl2

Oxidationnumber

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

Structure

(bond)

giant ionic lattice

(ionic)

simple molecular

(covalent)

Physical

state (at

room temp)

solid liquid liquid

solid

liquid

Solubility in

water

soluble

soluble

Hydrolysed by water/white fumes of HCl/


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(neutral)

(slightly

acidic)

forms acidic solution

2. Melting Point

NaCl, MgCl2 : Typical ionic compounds. Giant ionic lattice structure, and

strong

ionic bonds result in high melting points.

AlCl3 , SiCl4 , PCl5, S2Cl2 , Cl2 : simple covalent molecular structure. Weak

intermolecular van der Waals forces

give rise to low

melting points

3. Solubility in water

Across the period- bonding changes from ionic to covalent which is more

likely to be

hydrolysed in water, i.e. react with water, changing from neutral to acidic

solution

(i) NaCl(s) + aq Na+

(aq) + Cl-(aq) pH = 7 / ionic chloride, not

hydrolysed

(ii) MgCl2(s) + aq Mg2+

(aq) + 2Cl-(aq) pH = 6/ ionic with a little

covalent character,

slightly hydrolysed

(iii) Al2Cl6(s) + 6H2O(l) 2Al(OH)3(s) +

covalent chlorides,

hydrolysed by water

forming acidic solutions


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6HCl(aq)

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

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

2S2Cl2(l) + 2H2O(l) 3S(s) + SO2(g) + 4HCl(aq)

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The Periodic Table Note 1