Transcript
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CHAPTER 7
Periodic Properties of theElements
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CONTENT
7.1 Development of the Periodic Table7.2 Effective nuclear charge and the Sizes
of Atoms7.3 Ionization Energy7.4 Electron Affinities7.5 Metals, Nonmetals, and Metalloids7.6 Group Trends for the Active Metals7.7 Group Trends for Selected Nonmetals
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Learning outcomes
Able to relate effective nuclearchargeto
size, ionization energy and electron affinity
of elements in periodic table
To differentiate metal, non-metal and
metalloid
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7.1 Development of ThePeriodic Table
The periodic table was first developed byMendeleev and Meyer on the basis of the
similarity in properties and reactivitiesexhibited by certain elements.
Elements in the same column of the periodic
table have the same number of electrons intheir valence orbitals.
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Cont: 7.1 Development of ThePeriodic Table
Group 1A: Group 6A:
3Li - [He] 2s 1 8O - [He] 2s 22p 411Na - [Ne] 3s 1 16S - [Ne] 3s 23p 419K - [Ar] 4s 137
Rb - [Kr] 5s1
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7.2 Effective Nuclear Chargeand Sizes of Atoms
7.2.1 Effective Nuclear Charge
The net positive charge experienced by anelectron on a many-electron atom.
Not the same as the charge on the nucleus
because of the effect of the inner electrons.
The electron is attracted to the nucleus, but
repelled by the inner-shell electrons that shield
or screen it from the full nuclear charge.
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The principal quantum number, n, ofthe valence orbitals of the atoms
changes from top to bottom of thePeriodic Table.
All orbitals with the same value of nare referred to as a shell.
7.2.2 Electron Shells in Atoms
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Cont:
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The maxima(peaks) appear based on electronshaving the same quantum number, n.
n= 1 - 1stpeak (1s)n= 2 - 2ndpeak (2s2p)
n= 3 - 3rdpeak (3s3p3d)
Cont: 7.2.2 Electron Shells in Atoms
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From the graph:
1selectrons for Helium show a maximum inradial electron density - 0.3
1selectrons for Argon - 0.05 only
Cont: 7.2.2 Electron Shells in Atoms
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Question
Why is the 1sshell in Argon so much
closer to the nucleus than the 1sshell inHelium?
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Cont: 7.2.3 Atomic Sizes
1. Within each column (group) the atomicradius (and also the size of orbitals) tends
to increasewhen nincreases.(a) n , orbital size(b) n , Zeffremains relatively constant
n ,Atomic radius
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Cont: 7.2.3 Atomic Sizes
2. Within each row(period), the atomicradius decreaseas we move from left to
right.(a) n constant, orbital sizeconstant(b) number of core electrons stay the
same , nuclear charge (Z) acting on theelectron valence increaseZeff , Atomic radius
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Cont: 7.2.3 Atomic Sizes
same nZeffincreases
B C N O F
0.88 0.77 0.75 0.73 0.71
radius decreases
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Cont:
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Cont: 7.2.3 Atomic Sizes
(Z = 3) Li 1s 22s 1 (1.52 )(Z = 4) Be 1s 22s 2 (1.13 )
For Li, 1s2electrons shield the outer 2s1electron from the 3+ (Z value) charge nucleus.The outer electron 2s1electron experiencesZeffof slightly more than 1+.
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7.3 Ionization Energy
Ionization is a process of removing an electronfrom an atom or ion.
Ionization energy of an atom or ion - theminimumenergy required to remove an electronfrom the ground state of the isolated gaseousatom or ion.
The ease in removing electrons from an atom isan important indicator of the atoms chemicalbehaviour.
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Cont: 7.3 Ionization Energy
First ionization energy, I1(or IE1)-energy needed to remove the firstelectron from a neutral gaseous atom.
Na (g) Na+ (g) + e-
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Cont: 7.3 Ionization Energy
Second ionization energy, I2(or IE2) - energy
required to remove the secondelectron from agaseous ion.Na+(g) Na2+(g) + e-
Ionization energies I1, I2are always positive(endothermic) where energy is absorbed fromthe surrounding.
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Cont:
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Cont: 7.3 Ionization Energy
The greaterthe ionization energy, the moredifficultto remove an electron.
Magnitude I1< I2 < I3Reason:The positive nuclear charge remainsthe same, the number of electrons (producerepulsive interactions) decreases.
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Cont:
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Cont: 7.3 Ionization Energy
Ionization energy increases from:I1 I2 I3 I4 I5(kJ/mol)
786 - 3230 4360 161003s 2 3p 2 3s 2 3p 1 3s 2 3s 1 1s 2 2s 2 2p 6
I1I
4: 786 kJ/mol 4360 kJ/mol
Loss of the four electrons in the 3sand 3psubshells.
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Cont: 7.3 Ionization Energy
I4I5: 4360 kJ/mol 16100 kJ/molThe inner shell 2p electron (core
electron) is much closer to the nucleus -greater Zeff.
Large increase in ionization energy whenelectrons are removed from its noble gascore.
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Cont: 7.3 Ionization Energy
The outermost electrons take part in:a) chemical bonding
b) reaction
The core (noble-gas core) tightly boundto the nucleus.
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Cont: 7.3.1 Periodic Trends inIonization Energies
3. Ionization energy of the transition
elements increase slowly from left to right.
4. The f -block elements show only a small
variation in the values of I1.
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Cont: 7.3.1 Periodic Trends inIonization Energies
Attraction of electrons to the nucleusdepends on:
The effective nuclear charge (Zeff)The average distance of the electron from
the nucleus (atomic radius).
Attraction , Ionization energy
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Cont: 7.3.1 Periodic Trends inIonization Energies
Move across a row:Increase in Zeffand decrease in
atomic radius.
Attraction , Ionization energy
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Cont: 7.3.1 Periodic Trends inIonization Energies
Irregularities:
a) Decrease in ionization energy from:
Beryllium [He] 2s 2 Boron [He] 2s 22p1
o The electrons in the filled 2sorbital
shielding the electrons in 2p.o Zeff decreases from 2+ 1+.
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Cont: 7.3.1 Periodic Trends inIonization Energies
b) Decrease in ionization energy from:Nitrogen [He] 2s 22p 3Oxygen [He] 2s 22p 4
o Due to repulsion of paired electron in thep4configuration.
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Example 1
Arrange the following atoms in order of
increasing first ionization energy:
Ne, Na, P, Ar, K
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Example 1 (Answer)
Use trends to predict.1. Na, P and Ar are in the same row.
exhibit the order of Na
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Example 1 (Answer)
3. K below Na.
I1K is less than Na.
K < Na < P < Ar < Ne
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7.4 Electron Affinities
Measure of attraction towards electron or theease of an atom to gain electron.
The energy change(E) that occurs when an
electron is added to a gaseous atom is calledthe electron affinity.
Cl(g)+ e-
Cl-(g)
[Ne] 3s23p5 [Ne] 3s23p6
E(energy) = EA1= -349 kJ/mol
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Cont: 7.4 Electron Affinities
Energy is released when an electron is added.The electron affinity of Cl is -349 kJ/mol.The greater the attraction, the more negative
the electron affinity will be.EA1= -x
the greater the affinitythe more negative the valueeasier to gain electron
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Cont: 7.4 Electron Affinities
Halogen (F, Cl, Br, I) - by gaining an electron, ahalogen atom forms a stable negative ion that has anoble gas configuration.
Noble gas configuration: Ne : 1s22s22p6Ar : 1s22s22p63s23p6
Cl(g) + e- Cl-(g)[Ne]3s23p5 [Ne] 3s23p6 or [Ar]
psubshell (orbital 3p) is filled to form a stablenegative ion similar to Ar.
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Cont: 7.4 Electron Affinities
Positive Value of the Electron Affinity
The noble gases posses positive values of
electron affinity.The anion is higher in energy than the atom.
Ar(g) + e- Ar-(g) EA1 > 0
[Ne] 3s2
3p6
[Ne] 3s2
3p6
4s1
As EA1 > 0, the Ar-is not stable and will not form.
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Cont: 7.4 Electron Affinities
The addition of an electron to a noble gas
requires the electron to reside in a new, higherenergy subshell or in 4sorbital.
Occupying a higher-energy subshell is
unfavourable.
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Cont: 7.4 Electron Affinities
Generally, electron affinity becomesincreasingly negative from left to right i.e fromalkali metals to halogens.
Electron affinities of Be : [He] 2s2andMg: [Ne] 3s2 are positive. The added electron
would reside in an emptypsubshell that ishigher in energy (Hunds first rule).
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Cont: 7.4 Electron Affinities
Group 5A
N, P, As, Sb
N [He] 2s 22p 3 0
P [Ne] 3s2
3p3
-72 kJ/molAs [Ar] 4s 23d 104p 3 -78 kJ/molSb [Kr] 5s 24d 105p 3 -103 kJ/mol
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Cont: 7.4 Electron Affinities
These elements have half-filledpsubshells.The added electron must be placed in an orbitalthat is already occupied - resulting largerelectron-electron repulsion (Hunds first rule).
As we proceed from top to bottom, the average
distance of the added electron from thenucleus increases. The electron-nucleusattraction decreases.
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Cont: 7.4 Electron Affinities
N: [He] 2s22p3 Sb : [Kr] 5s 24d 105p 3
strong electron-electron repulsion
lower electron-electron repulsion
will not accept electron
EA1> 0 EA1= -103 kJ/mol
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Cont:
g
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Cont: 7.4 Electron Affinities
You can explain:F [He] 2s 22p 5 - 328 kJ/mol
Br [Ar] 4s 23d1 04p 5 - 325 kJ/mol
I [Kr] 5s 24d 105p 5 - 295 kJ/mol
5p 5- away from nucleus, therefore less nucleusattraction.
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7.5 Metals, Nonmetals andMetalloids
Properties of individual atoms:
1. Atomic radii2. Ionization energies3. Electron affinities
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7.5.1 Metals
Properties:Metals conduct heat and electricity.
They are malleable - can be pounded into thinsheets.Ductile - can be drawn into wire.
Shiny lusterAll are solids (except Hg)High melting point (except Cs, Ga)
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Cont: 7.5.1 Metals
Charge:
Alkali metals 1+.Alkali earth metal 2+.Transition metals ions 2+, 1+, 3+.
Able to form more than one positive ion.
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Cont: 7.5.1 Metals
Basic oxide:
Most metal oxides are basic oxides. Theydissolve in water, react to form metalhydroxides.Metal Oxide + Water Metal HydroxideNa2O(s) + H2O(l) 2NaOH(aq)CaO(s) + H2O(l) Ca(OH)2(aq)
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Cont: 7.5.1 Metals
React with acids to form salts and water:
Metal Oxide + Acid
Salt + Water
MgO(s) + 2HCl(aq) MgCl2(aq) + H2O(l)
NiO(s) + H2SO4(aq)
NiSO4(aq) + H2O(l)Na2O(s)+ H2SO4(aq) Na2SO4(aq) + H2O(l)
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Example 2 (Answer)
a) Aluminium has a 3+ charge,Al3+.The oxide ion is O2-
Al2O3
b) Metal oxides react with acids to form salts
and water.Al2O3(s) + 6HNO3(aq)2Al(NO3)3(aq)+ 3H2O(l)
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7.5.2 Nonmetals
Not lustrous (not shiny).Poor conductors of heat and electricity.Melting points - generally lower than those of
metals (except diamond : 3570 C).Seven nonmetals exist as diatomic molecules.
o H2,N2, O2, F2, Cl2- gases.
o Br2 - liquido I2 - volatile solid
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Cont: 7.5.2 Nonmetals
o Nonmetals, reacting with metals, gain electronsand become anions.
Metal + Nonmetal Salt
2Al(s)+ 3Br2(l) 2AlBr3(s)
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Cont: 7.5.2 Nonmetals
Nonmetals gain electrons to fill their outer psubshell giving a noble-gas electronconfiguration.
Most nonmetal oxidesare acidic oxides(molecular substance).Those dissolve in waterreact to formacids.
Example:Selenium dioxide - SeO2Tetraphosphorus hexoxide - P4O6(s)Tetraphosphorus decoxide - P4O10(s)
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7.5.3 Metalloids
Metalloids have properties intermediatebetween those of metals and nonmetals.
They may have some characteristic metallic
properties but lack others.Example: silicon looks like a metal but it is
brittle rather than malleable and a muchpoorer conductor of heat and electricity than
metals.Several metalloids are semiconductors ( most
notably silicon).
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Cont: 7.6.1 Group 1A: TheAlkali Metals
As a result, the alkali metals are all veryreactive, readily losing one electron to formions with a 1+ charge:
M M+ + e-
*M represents any one of the alkali metals
The alkali metals are the most active metalsand thus exist in nature only as compounds.
C 7 6 1 G 1A Th
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Cont: 7.6.1 Group 1A: The
Alkali Metals
Electrolysis is the process used to obtain metalsfrom compounds.
The chemistry of the alkali metals is dominated by
the formation of 1+ cations.The metals combine directly with most nonmetals.Examples:
2M(s)+ H2(g)2MH(s)2M(s)+ S(s)M2S(s)2M (s) + Cl2(g)2MCl(s)
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Cont: 7.6.1 Group 1A: TheAlkali Metals
The hydrides of the alkali metals (LiH, NaH,and so forth): hydrogen is present as H-, calledthe hydride ion.
Note the difference between hydride ion H-and hydrogen ion H+.
The alkali metals react vigorously with water to
produce hydrogen gas and solutions of alkalimetal hydroxides.2M(s) + 2H2O(l) 2MOH(aq) + H2(g)
h
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Cont: 7.6.1 Group 1A: TheAlkali Metals
This reaction is very exothermic (heat isreleased).
This reaction is most violent for the heaviermembers of the group - weaker hold on the
single outer-shell electron.
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C 7 6 1 G 1A Th
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Cont: 7.6.1 Group 1A: TheAlkali Metals
Potassium, rubidium, and cesium also formcompounds that contain the O2-ion, calledsuperoxides.
K(s)+ O2(g)KO2(s) potassium superoxide
As the alkali metals are extremely reactive
toward water and oxygen, they are usuallystored in hydrocarbon, such as kerosene ormineral oil.
C t 7 6 1 G 1A Th
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Cont: 7.6.1 Group 1A: TheAlkali Metals
Alkali metals salts and their aqueous solutionsare colourless unless they have a coloured anioneg yellow CrO42-.
When alkali metal compounds are placed in a
flame, they emit characteristic colours.(Li: red,Na: yellow,K: blue)
7 6 2 G 2A Th Alk li
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7.6.2 Group 2A: The AlkalineEarth Metals
The elements are all solids with typical metallicproperties.
Compare to elements in Group 1A (alkalinemetals), alkaline earth metals are harder, moredense and melt at higher temperatures.
Their I1
are low, but not as low as those ofalkali metals.
Less reactive than alkali metals.
C t 7 6 2 G 2A Th
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Cont: 7.6.2 Group 2A: TheAlkaline Earth Metals
Be and Mg are the least reactive.Be does not react with water or steam.
Mg does not react with water but does reactwith steam to form Magnesium oxide andhydrogen.Mg(s)+ H2O (g)MgO (s)+ H2(g)
The other elements react readily with water(less reactive than alkali metals).Ca (s)+ 2H2O (s) Ca(OH) 2(aq)+ H2(g)
C t 7 6 2 G 2A Th
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Cont: 7.6.2 Group 2A: TheAlkaline Earth Metals
Pattern in the reactivity of the alkaline earthmetals - the tendency to lose their two outer selectrons and form 2+ ions.
Example:Mg(s)+ Cl2(g) MgCl2(s)2Mg(s)+ O2(g) 2MgO(s)
Like the 1+ ions of the alkali metals, the 2+ ionsof the alkaline earth elements have a noble gasconfiguration.
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cont
r
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7 7 G T d f S l t d
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7.7 Group Trends for SelectedNonmetals
7.7.1 HydrogenThe first element in the periodic table. 1s1electron configuration. Placed above alkali
metals.Unique element, nonmetal, exists as diatomic gas,
H2(g), under most conditions.The ionization energy of hydrogen, 1312 kJ/mol, is
markedly higher than that of the active metals.Reason: the complete absence of nuclear shielding
of its sole electron.
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Cont: 7.7.1 Hydrogen
React with other nonmetals to form molecularcompounds.
2H2(g)+ O2(g)2H2O(l)
Hydrogen reacts with other active metals to
form solid metal hydrides, contain the hydrideion, H-.
7 7 2 Group 6A: The Oxygen
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7.7.2 Group 6A: The OxygenGroup
Increase in metallic character as moving downthe group.
Oxygen is colourless gas, the rest are solids.Oxygen, sulfur, and selenium are typical
nonmetals.
Tellurium is a metalloid.Polonium is a metal.
OS
Se
Te
Po
Cont: 7 7 2 Group 6A: The
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Cont: 7.7.2 Group 6A: TheOxygen Group
Oxygen occurs in two forms, O2and O3.O2- dioxygen (normally called as oxygen)
O3ozone
Ozone - toxic and pungent gas. It is alsoformed from O2in electrical discharge, eg
lightning storm.3O2 (g) 2O3(g)
Cont: 7 7 2 Group 6A: The
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Cont: 7.7.2 Group 6A: TheOxygen Group
Oxygen has a great tendency to attractelectrons from other elements ( to oxidisethem).
Form oxide, O2-ion. This ion has a noble gasconfiguration and thus stable.
Peroxide, O22-and superoxide, O2-ions often
react with themselves to produce O2-
and O2.Eg.: 2H2O2(aq) 2H2O(l) + O2(g)
Cont: 7 7 2 Group 6A: The
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Cont: 7.7.2 Group 6A: TheOxygen Group
After oxygen, the most important element issulfur.
The most common and stable is the yellow solid,S8.
Sulfur is written simply as S(s).Sulfur has the tendency to gain electron
forming sulfides, S2-.Eg.: 2Na(s) + S(s) Na2S(s)
Cont: 7 7 2 Group 6A: The
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Con t: 7.7.2 Group 6A: TheOxygen Group
Most sulfur in nature is present as metalsulfides.
The chemistry of sulfur is more complex thanthat of oxygen.
Sulfur can be burned in oxygen to producesulfur dioxide, the main pollutant.
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7.7.3 Group 7A: The Halogens
As we move from group 6A to 7A, thenonmetallic behaviour of the elements
increases.All the halogens are typical nonmetals (exceptAt: metalloid).
Melting points and boiling points increase withincreasing atomic number.
Cont: 7 7 3 Group 7A: The
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Con t: 7.7.3 Group 7A: TheHalogens
Fluorineand chlorine- gas at roomtemperature.
Bromine - liquidIodine solid(sublime easily)Each element consists of diatomic molecules:
F2, Cl2, Br2, and I2.
F2- pale yellow gas; Cl2(g)-yellow-greencolour;Br2(l)- reddish brown; I2(s)- greyish black.
Cont: 7 7 3 Group 7A: The
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Con t: 7.7.3 Group 7A: TheHalogens
The halogens have highly negative electronaffinities.
The chemistry of halogens are dominated bytheir tendency to gain electrons from other
elements to form halide ions:X2 + 2e- 2X-
Cont: 7 7 3 Group 7A: The
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Con t: 7.7.3 Group 7A: TheHalogens
Fluorine and chlorine are more reactive thanbromine and iodine.
2Na(s) + F2(g) 2NaF(s)2H2O(l)+ 2F2(g) 4HF(aq)+ O2(g)
Fluorine gas is difficult and dangerous to beused in laboratory because it is very reactive.
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Cont: 7 7 3 Group 7A: The
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Con t: 7.7.3 Group 7A: TheHalogens
Chlorine react slowly with water to formrelatively stable aqueous solutions of HCl andHOCl (hypochlorous acid):
Cl2(g)+ H2O(l) HCl(aq) + HOCl(aq)
The halogens react directly with most metalsto form ionic halides. Also react with hydrogento form gaseous hydrogen halide compounds.
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cont
7 7 4 Group 8A: The Noble
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7.7.4 Group 8A: The NobleGases
The elements are all nonmetals that are gases atroom temperature.
They are all monoatomic(consist of single atoms
rather than molecules).The noble gases have completely filledsandp
subshells.All elements have high first ionization energies,
the values decrease as moving down the group.The noble gases are exceptionally unreactive.Also called inert gases.
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cont
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