Chapter 19 Summary

This summary must be read alongside the textbook

Representative (Gps 1A to 4A) Elements

• Review the periodic table.

• How size trend causes the properties of the first element in a group to differ form the

remaining element?

• In group 1A H behaves as a non-metal. The rest form +1 cations in ionic compounds. The

electron on the H atom is held strongly.

• In group 2A BeO is amphoteric. The oxides of the remaining elements in group 2A are basic.

The small Be2+

ion strongly polarizes O2- ion and brings about electron sharing.

• In group 3A B behaves as a non-metal (sometimes metalloid), Al as a metal (Al2O3 is

amphoteric) and the rest are active metals.

• In group 4A the chemistry of carbon compounds is dominated by the C-C bond while that of

Si is dominated by the Si-O bond. Compounds with Si-Si bonds are too reactive.

• Overlap between 3p orbitals on Si and the smaller 2p arbitals on O makes the π-bond in

O=Si=O weaker than that in O=C=O. For this reason CO2 exists as single molecules while

SiO2 exists as a network structure consisting of units of SiO4 tetrahedra with single Si-O

bonds.

• The stable triple bond in molecular : N ≡ N : does not exist for phosphorous. The large P

atoms form different structures with single P-P bonds. See figure below for (a) the P4

molecule in white phosphorous, (b) the crystalline network structure of black phosphorous

and (c) the chain structure of red phosphorous.

• Similarly in group 6A O2 is molecular while sulfur is found as S8 and viscous liquid sulfur

may contain chains as long as 10,000 atoms. See figure below

.

• In group 7A, F has a lower electron affinity than Cl. Beyond Cl electron affinity increases.

The smaller 2p orbitals on F result in a large repulsion between lone pairs and the weakening

of the : F – F: bond.

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Abundance & Preparation

• Read the two paragraphs on abundance & study Tables 19.1 & 19.2.

• Most elements (75%) are in the combined state.

• Metallurgy

reduction

Mined Ore � crushing � floatation � Roasting � oxide metal

C or H

heat

2SnO (s) + C (S) 2Sn(s) + CO2(g)

heat

SnO (s) + H2(g) Sn(s) + H2O(g)

• N2 and O2 are obtained by distilling liquefied air. Liquefaction of air involves repeated

cycles of adiabatic expansion (gas cools) and compression (part of the gas is compressed

& the rest carries the heat of compression)

• H2 from electrolysis of water or decomposition of CH4 in natural gas.

Heat, pressure

CH4(g) + H2O(g) CO(g) + 3H2(g)

Catalyst(usually Ni)

• S is found underground in elemental form and recovered By the Frasch process (section

20.6)

• Halogen ion X – in salts oxidized to X. (section 20.7)

The Group 1A elements (excluding H2)

• Many properties of the alkali metals are given in section 7.13. Table 19.3 gives sources

and methods of preparation. More properties in Table 19.4.

• All react with water to be reduced to M+(aq). The small size of Li

+ allows it to

effectively attract H2O molecules & the energy released favors its formation making its

∈o lowest. Because of its higher melting point the heat of reaction does not melt it and

its area of contact with H2O is small thus it reacts more slowly with H2O than Na or K.

• Only Li and Na form(in a limited supply of O2) the regular Li2O & Na2O oxides.

• In excess O2, Na forms Na2O2 (sodium peroxide containing the O2- peroxide ion).

Na2O2 reacts with water to form hydrogen peroxide, H2O2 (hair bleach and disinfectant)

Na2O2(s) + 2H2O(ι) � 2Na+(aq) + H2O2(aq) + 2OH

-(aq)

• K, Rb and Cs form superoxides (MO2) which react with H2O or CO2 and release O2.

2MO2(s) + 2CH2O(ι)� 2M+(aq) + 2OH

- + O2(g) + H2O2(aq)

4MO2(s) + 2CO2(g) � 2M2CO3(s) + 3O2(g)

Superoxides are used in breathing apparatus for protection against toxic fumes.

• Na+ & K

+ present in all body cells and body fluids and are transported through cell

membranes.

• Li+ (in the form of Li2 CO3) and in the proper concentration affects the levels of

neurotransmitters (molecules that assist transmission of messages along nerve networks)

and is used to treat manic-depressive patients.

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Hydrogen (H2)

• Colorless. Odorless. Nonpolar & of low molar mass (B.P – 253 oC , M.P. – 260

oC).

Flammable. 18% to 60% mixtures by volume with air are explosive.

High T,P

Industrial Sources (1) CH4(g) + H2O(g) -----------� CO(g) + 3H2(g)

Catalyst

(2) Cracking of high molar mass hydrocarbons.

(3) In very pure form by electrolysis of water.

• Major Industrial Use

H H H H

| | H2, Pd | |

-C=C- -C- C-

| |

H H

Unsaturated vegetable oil saturated margarine

Ionic hydrides e.g. LiH & CaH2 contain the hydride (H–) ion. Large electron-electron

repulsion and a small +1 nuclear charge make H– a strong reducing agent (i.e. It is easily

oxidized to H2)

Li H(s) + H2O(ι) --� H2(g) + Li+(aq) + OH

– (aq)

Oxd. No. of H: -1 0

Covalent Hydrides: H2O – a very important substance. Relative to its molar mass it has a high B.P. Its large

∆H vap and heat capacity make it an excellent coolant. Its ice structure which results from

hydrogen bonding leads to a lower density than water. Excellent solvent for ionic & polar

substances. Effective medium for life processes. Other covalent hydrides include HCl,

CH4, NH3, etc.

Metallic Hydrides (Also called interstitial hydrides)

On a metal surface H2 molecules dissociate to H atoms which migrate to occupy interstices

within the metal crystal structure. Pd absorbs 900 times its own volume of H2 gas. Diffusion

through a thin Pd wall is used to purify H2. Most interstitial hydrides are nonstoichiometric.

Pd is used to store H2 which in turn is used as a fuel.

Group 2 A elements

• n s2 valence – electron configuration

• lose 2e- to form M2+

ions

• called alkaline earth metals because their oxides are basic.

(Beo exhibits some acidic properties)

MO(s) + H2O(ι) � M2+

(aq) + 2OH-(aq)

BeO(s) + 2OH-(aq) + H2O(l) � Be (OH)4

2-(aq)

• For the reaction

M(s) + 2H2O(ι) � M2+

(aq) + 2OH-(aq) + H2(g)

Element Be Mg Ca, Sr and Ba

Reaction with H2O No reaction Reacts with boiling H2O Reacts with H2O at 25 ° C

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• Be bonds are more covalent than the rest of the group because of its small size and

relatively high electronegativity. :Cl - Be - Cl: is a gas.

• Ca & Mg are important for human life. Mg which is of low density and moderate

strength is useful as a structural material when alloyed with Al. Important reactions are

given in Table 19.8

• Ca2+

and Mg2+

ions responsible for the hardness of water. They interfere with detergent

action and react with soap to form precipitates.

• In large scale water softening Ca+ is removed as CaCO3(s) by heating the water.

• Cation-exchange resins (see figure below) remove Ca2+

& Mg2+

from water & replace

them by Na+.

• Mg is electrolytically produced.

Ex 19.2 What time is needed to produce 1.00 x 10

2 kg Mg from electrolysis of molten MgCl2 using a

1.00 x 102 A current

Mg2+

+ 2e- � Mg

1.00 x 105g x 1mol Mg x 2mole

- x 96,485 C x 1s x 1day = 918 days

24.3g Mg 1mol Mg mol e- 10

2C 24x3600 s

Group 3A Elements

• ns2 np

1 valence – electron configuration.

• Increase in metallic character down a group.

• Table 19.9 gives physical properties, sources & methods of preparation.

• B is a non-metal. It forms covalent boron hydrides (BORANES). The instability of

BH3 and the high reactivities of B2H4 and B5H9 (see structures below)

are because of their high electron deficiencies. They undergo highly exothermic

reactions with oxygen.

• Al is an important structural material. It has metallic properties but forms highly

covalent bonds with nonmetals and Al2O3 is amphoteric.

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• Ga has a melting point of 29.8 o C and a boiling point of 2400

o C. Its high liquid

range makes it useful for thermometers that measure high temperatures. Ga expands on

freezing. Ga2O3 is amphoteric.

Group 4A elements

• ns2 np

2 valence – electron configuration.

• Includes C and Si; two very important elements.

• Metallic character increases down the group.

• Table 19.11 lists their physical properties, sources and methods of preparation.

• All form covalent bonds to nonmetals. In the tetrahedral structure CH4, SiF4, GeBr4,

SnCl4 and PbCl4 all the group 4 central atoms have sp3 hybridization. All except C

which is too small in size and lacking d-orbital can add two Cl – ions

SnCl4 + 2 Cl – � SnCl6

2–

• We already discussed how the chemistry of C differs from the rest of the group because

of its ability to form π-bonds and how it is dominated by C-C bonds while that of Si is

dominated by Si-O bonds.

• C occurs as graphite and diamond .

Copyright © Houghton Mifflin Company. All rights reserved. 10–22

Figure 10.22 The Structures of

Diamond and Graphite

Also Buckminster fullerene (C60) & other forms were identified.

Oxides of C:

: C ≡ O: (Carbon monoxide). Odorless, colorless & toxic. Formed by the combustion of

C-containing compounds in a limited oxygen supply.

O = C = O carbon dioxide). Linear. sp hybridization on C. Produced by respiration & in

the combustion of fossil fuels. Also produced by fermentation of sugar in fruit.

C6H12O6(aq) Enzymes 2C2H5OH(aq) + 2CO2(g)

(Glucose)

Its solution in water is acidic:

CO2(g) + H2O(ι) ⇔ H+(aq) + HCO3

-(aq)

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O = C = C = C = O (carbon suboxidse) has sp hybridized C atoms.

Silicon (Si). Metalloid. 85% of the earth's crust consists of silica (SiO2, structure shown

below) and silicates (such a Feldspar K2O.Al2O3. 6H2O & Na2O. Al2O3. 6H2O)

Copyright © Houghton Mifflin Company. All rights reserved. 10–25

Figure 10.26

The Structure

of Quartz

(Empirical

Formula

SiO2)

• Si is mainly used in a highly pure form as a semiconductor.

Germanium (Ge): Metalloid. Called Germanium (Ge) because it was first discovered in an

ore in a mine in Freiburg, Germany. Rare. Used in semiconductors.

Tin (Sn): is a soft metal. Used for many centuries in various alloys. Sn has three allotropes

1. Gray tin ; powdery. Stable below 13.2 o

C

2. White tin ; stable at normal temperatures.

3. Brittle tin ; found above 161 o

C

It is electrolytically coated on steel where it forms a protective oxide layer against

corrosion, especially for cans used as food containers.

Lead (Pb) Obtained from its ore, galena (PbS). Probably the first metal discovered & used

by man. It is poisoness. The extensive use of Pb utensils by Roman emperors gave them

dementia (a disease in which the patient suffers from forgetfulness) and this could be one of

the reasons for the downfall of the Roman Empire. Analysis of bones of dead Romans

showed a high lead level. Roman used Pb to make pipes for use in their baths.

In our modern life we are exposed to lead that comes from tetraethyl lead (C2H5)4 Pb that is

an antiknock agent added to gasoline. Tetraethyl lead is now being replaced by less harmful

antiknock agents.

The largest use of Pb (1.3 million tons/yr) is for the electrodes of lead storage batteries used

in cars.

Table 19.13 gives some of the important reactions of group 4A elements.