chapter 7 jan13

8/12/2019 Chapter 7 Jan13

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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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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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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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(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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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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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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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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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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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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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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Move across a row:Increase in Zeffand decrease in

atomic radius.

Attraction , Ionization energy


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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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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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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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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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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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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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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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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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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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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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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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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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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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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)



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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)



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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)



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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.



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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.



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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-



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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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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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chapter 7 jan13

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