the periodic table note 1
TRANSCRIPT
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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)