Copyright McGraw-Hill 20091 Chapter 24 Nonmetallic Elements and Their Compounds Insert picture from...

Copyright McGraw-Hill 2009 1

Chapter 24

Nonmetallic Elements and

Their Compounds

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24.1 General Properties of Nonmetals

• Properties of nonmetals - more varied than those of metals

• Physical state– Gases: hydrogen, oxygen, nitrogen, fluorine,

chlorine, and the noble gases– Liquid : bromine– Solids: All the remaining nonmetals

• Poor conductors of heat and electricity• Exhibit either positive or negative oxidation

numbers.


• Metalloids - small group of elements have properties characteristic of both metals and nonmetals.

• More electronegative than metals• Electronegativity increases from left to right

across any period and from bottom to top in any group in the periodic table

• With the exception of hydrogen, the nonmetals are concentrated in the upper right-hand corner of the periodic table

• Compounds formed by a combination of metals with nonmetals tend to be ionic, having a metallic cation and a nonmetallic anion.


Nonmetals and Metalloids on the Periodic Table

Nonmetals coded in blue and metalloids in orange..


24.2 Hydrogen

• Simplest known element

• Exists as a diatomic molecule

• H2 is a colorless, odorless, and nonpoisonous gas.

• At 1 atm, boiling point is −252.9°C (20.3 K).

• Most abundant element in the universe


• Ground-state electron configuration: 1s1. – Resembles the alkali metals (Group 1A) in

that it can be oxidized to the H+ ion, which exists in aqueous solutions in the hydrated form.

– Resembles the halogens (Group 7A) in that it forms the hydride

• H− (hydride ion) - isoelectronic with helium (1s2)

• Found in a large number of covalent compounds.

• Unique capacity to form hydrogen bonds


• Preparation– Industrial scale

– Laboratory scale


Laboratory Generation of Hydrogen


• Binary hydrides - compounds containing hydrogen and another element, either a metal or a nonmetal.

• Types of hydrides– Ionic hydrides direct combination of molecular

hydrogen and any alkali or alkaline earth metal

– Solids with high melting points– Contain the strong BrØnsted base, H−


– covalent hydrides - the hydrogen atom is covalently bonded to the atom of another element• Types of covalent hydrides

–Discrete unit structure – NH3

–Polymeric structure – (BeH2)x

– Interstitial hydrides – compounds of hydrogen and transition metal in which the atomic ratio is not constant – titanium hydride ranges from TiH1.8 to TiH2.


Binary Hydrides of Representative Elements


• Isotopes of hydrogen– Hydrogen has three naturally occurring

isotopes– , hydrogen, (99.985%) – , deuterium, symbol D, (0.015%)

– , tritium, symbol T, (radioactive, t1/2 =12.5 years.

– Deuterium containing water, D2O

• Called heavy water or deuterated water• Toxic• Affects reaction rates – isotopic effect

H11

H31

H21


• Hydrogenation - addition of hydrogen to compounds containing multiple bonds, usually carbon to carbon double or triple bonds.

– Catalyzed by metals (Pt or Cd)– Important in food industry


• Hydrogen Economy– Hydrogen an alternative fuel source to

petroleum fuels• For automobiles• Electrical power generation

– Pollution free fuel

– Present dilemma – how to obtain sufficient amounts of H2

– Splitting water using solar energy – one possible source for the needed H2.


24.3 Carbon• 0.09 % by mass of Earth’s crust• An essential element of living matter• A component of natural gas, petroleum

and coal. • Combines with oxygen to form carbon

dioxide in the atmosphere• Occur as carbonates in limestone and

chalk.• Found free in allotropic forms of diamond

and graphite


Phase Diagram for Allotropic Forms of Carbon


• catenation – carbon has the unique ability to form long chains stable rings– Responsible for the millions of carbon-

containing compounds

• Reacts with – Metals to form carbides (strong bases), CaC2

– Silicon to form carborundum, SiC– Nitrogen to form cyanides,

• Toxic• Readily complexes metals

C N


Cyanide Pond for Extracting Gold


• Important oxides– Carbon monoxide (CO)

• Formed during incomplete combustion

• Colorless, odorless gas• Used in metallurgical processes• Used in organic synthesis• Not acidic• Only slightly soluble in water• Burns to produce carbon dioxide


– Carbon dioxide (CO2)

• Colorless and odorless gas• Nontoxic—although it is a simple

asphyxiant• Acidic oxide – forms carbonic acid• Uses

– “carbonated” beverages–Fire extinguishers

–Manufacture of baking soda (NaHCO3)

–Manufacture of soda ash (Na2CO3)

–“Dry ice” as a refrigerant


24.4 Nitrogen and Phosphorous

• Nitrogen– Mineral sources of nitrogen: saltpeter (KNO3)

and Chile saltpeter (NaNO3)

– Nitrogen is an essential element of life• A component of proteins and nucleic acids

– N2 is obtained by the fractional distillation of air

– N2 contains a triple bond and is stable

– Forms variable oxidation states


• Common (important) forms of nitrogen– Nitride ion, N3−, a strong BrØnsted base

– Ammonia, NH3

• Undergoes autoionization to produce the highly basic amide ion, NH2

-

– Hydrazine, N2H4

• Basic• Reducing Agent


• Important oxides– Nitrous oxide (N2O)

• Supports combustion

• Used as dental anesthetic– Nitric oxide (NO)

• Produced in atmosphere (form of nitrogen fixation)

• Colorless gas• Produced in auto exhaust• Paramagnetic• Resonance stabilized


– Nitrogen Dioxide (NO2)

• Toxic• Paramagnetic

• Dimerizes to N2O4 in the liquid and gas phases

• Acidic oxide–Shown in a disproportionation reaction

with cold water


• Nitric acid (HNO3)– Powerful oxidizing agent– Can be reduced to NH4

+

– Aqua regia – 1:3 mixture of concentrated HCl and concentrated HNO3

• Even oxidizes gold

– Oxidizes nonmetals to oxoacids– Used in manufacture of

• Fertilizers• Drugs• Explosives• Dyes


• Phosphorus– Occurs most commonly in nature as

phosphate rocks

• calcium phosphate [Ca3(PO4)2]

• fluoroapatite [Ca5(PO4)3F]

– Elemental phosphorus produced by

– Allotropic forms of phosphorus• Red phosphorus• White phosphorus

Ca5(PO4)3F


Allotropes of Phosphorus


– Reactions of phosphorus

• Formation of phosphine (PH3)

• Formation of phosphoric acid

• Reaction with the halogens


• Acid production from halides

• Reaction with oxygen to produce acidic oxides


Structure of P4O6 and P4O10


• Oxoacids of phosphorus


24.5 Oxygen and Sulfur

• Oxygen– Most abundant element in Earth’s crust (46%

by mass)– Atmosphere contains about 21% by volume

(23% by mass)

– Diatomic molecule (O2) in the free state

– Essential for human life

– Alloptropic forms: O2 and O3 (ozone)

– Strong oxidizing and bleaching agent


– Oxides• Types of oxides

–Normal oxide, O22−

–Peroxide, O22−

–Superoxide, O2−

• All are strong BrØnsted bases


• Bonding in oxides – ionic to covalent left to right on the periodic table

• Acid-base character of oxides –

–Basicity increases down a group

– Peroxides

• H2O2 (hydrogen peroxide) – most common example

Basic Amphoteric Acidic


Structure of H2O2


• Polar• Miscible with water• Decomposes spontaneously

• Used as mild antiseptic (3% solution) or bleach agent (higher concentrations)

• Used as rocket fuel due to high heat of decompostion

• Serves as an oxidzing agent

• Serves as a reducing agent


– Ozone• Toxic, light-blue gas• Pungent odor• Essential component of the atmosphere• Structure

• Powerful oxidizing agent

• Preparation


Preparation of O3


• Sulfur– Constitutes about 0.06 % of Earth’s crust by

mass– Occurs commonly in nature in the elemental

form• Sedimentary deposits

• Gypsum (CaSO4. 2H2O) and various sulfide

minerals such as pyrite (FeS2)

– Most common allotropic forms• Monoclinic

• Rhombic – most stable form – S8 FeS2


Puckered Ring of S8


– Extracted by the Fasch process


– Forms wide variety of oxidation numbers


– Hydrogen sulfide – H2S

• Used in qualitative analysis• Preparation

• Colorless gas with odor of rotten eggs• Toxic• Weak diprotic acid

• Reducing agent in basic solution


– Oxides of sulfur• Sulfur dioxide (SO2)

–Pungent colorless gas–Toxic–Preparation

–Acidic oxide

–Oxidation


• Sulfur trioxide (SO3)

–Involved in acid rain

–Used in the production of sulfuric acid (H2SO4) in the contact process*

*Vanadium(V) oxide (V2O5) is the catalyst used for the key second step.


• Sulfuric acid–Diprotic acid–Colorless, viscous liquid (m.p. 10.4°C)–Concentrated sulfuric acid is 98 %

H2SO4 by mass (density 1.84 g/cm3), 18 M.

–Oxidizing strength of sulfuric acid depends on temperature and concentration.

–Cold dilute sulfuric acid reacts with active metals


–Hot concetrated sulfuric acid reacts with less active metals

–Depending on the reducing agent, sulfate may be reduced

–Oxidizes nonmetals


• Carbon disulfide (CS2)

–Colorless, flammable liquid (b.p. 46°C)–Preparation

–SIightly soluble in water–Solvent for nonpolar substances

• Sulfur hexafluoride (SF6)– Preparation

– Colorless, nontoxic, inert gas


24.6 The Halogens• The halogens—fluorine, chlorine, bromine,

and iodine—are reactive nonmetals.


• All are highly reactive and toxic• Magnitude of reactivity and toxicity generally

decreases from fluorine to iodine.• The chemistry of fluorine differs from that of the

rest of the halogens in the following ways:– Fluorine is the most reactive due to the

relative weakness of the F−F bond.

– The difference in reactivity between fluorine and chlorine is greater than that between chlorine and bromine.


– Hydrogen fluoride (HF) has a relatively high boiling point (19.5°C)

– Hydrofluoric acid is a weak acid, all other hydrohalic acids are strong acids.

– Fluorine uniquely reacts with cold sodium hydroxide solution to produce oxygen difluoride as follows:

– Silver fluoride (AgF) is soluble. All other silver halides (AgCl, AgBr, and AgI) are insoluble.


• Elemental state, halogens form diatomic molecules (X2).

• In nature, always found combined with other elements. – Chlorine, bromine, and iodine occur as

halides in seawater

– Fluorine occurs in the minerals fluorite (CaF2) and cryolite (Na3AlF6).

• All isotopes of astatine (As) are radioactive


• Preparation and Properties of F2 and Cl2 – determined by their strong oxidizing capability– Fluorine

• From liquid HF• At 70oC


Electrolytic Preparation of F2


– Chlorine • Electrolysis of molten NaCl• Overall reaction

• Chlor-alkali process–Designed to prevent side reactions

–Mercury cell–Diaphragm cell


Mercury Cell in the Chlor-alkali Process


Diaphragm Cell in the Chlor-alkali Process


• Compounds of the Halogens– Either ionic or covalent.

• The fluorides and chlorides especially those belonging to the alkali metal and alkaline earth metal are ionic compounds (except halides of Be).

• Most of the halides of nonmetals are covalent compounds.

– Oxidation numbers range from −1 to +7 except F which can only be 0 (in F2) and −1, in all compounds.


– Hydrogen Halides• Preparation from elements – can occur

violently

• Preparation varies with the halogen, for example

HCl

HBr

HF


– Industrial uses of hydrogen fluoride (HF)• Reactive enough to etch glass

• Used in the manufacture of Freons

– Industrial uses of hydrogen chloride (HCl)• Preparation of hydrochloric acid• Inorganic chlorides• Various metallurgical processes


– Aqueous solutions of HX• Acidic• Variation in acid strength

– Oxoacids – halogens form a series of acids

• Only Cl forms the entire series

increasing acid strength


• Uses of the halogens– Fluorine

• UF6 separating isotopes of U

• Production of polytetrafluorethyline (Teflon ©)

– Chlorine• Biological role as Cl−(aq)

• Industrial bleaching – Cl2• Water purification – Cl2, ClO−

• Organic solvents – CHCl3• Polymer production - PVC


– Bromine• Insecticides (BrCH2CH2Br)

• Scavenger for Pb in gasoline• Photographic films (AgBr)

– Iodine• Antiseptic (tincture of iodine)• Thyroxine (thyroid hormone derivative)• Cloud seeding (AgI)


Key Points

• General properties of the nonmetals• Hydrogen

– Properties– Preparation– Binary Halides

• Ionic• Covalent• Interstitial


– Isotopes of hydrogen• Hydrogen (protium)• Deuterium• Tritium

– Hydrogenation– The hydrogen economy

• Carbon– Properties– Allotropes

• Diamond• Graphite


– Carbides– Cyanides– Oxides

• Carbon monoxide• Carbon dioxide

• Nitrogen and Phosphorus– Nitrogen

• Properties• Nitrides• Ammonia• Hydrazine


• Oxides–Nitrous oxide–Nitric oxide–Nitrogen dioxide

• Nitric Acid– Phosphorus

• Properties• Allotropes

–White phosphorus–Red phosphorus

• Phosphine• Halogen compounds


• Oxides• Oxoacids

• Oxygen and Sulfur– Oxygen

• Properties• Allotropes• Oxides

–Normal oxide–Peroxide–Superoxide

• Acidity of oxides


• Hydrogen peroxide• Ozone

– Sulfur• Properties• Industrial production• Hydrogen sulfide• Oxides

–Sulfur dioxide–Sulfur trioxide

• Sulfuric Acid–Production–Uses


• Carbon disulfide• Sulfur hexafluoride

• The Halogens– Properties

• Special properties of fluorine– Preparation and Properites

• Preparation of fluorine• Preparation of chlorine – chlor-alkali

process–Mercury Cell–Diaphragm cell


• Hydrogen Halides• Oxoacids

– Uses of the halogens• Fluorine• Chlorine• Bromine• Iodine

Copyright McGraw-Hill 20091 Chapter 24 Nonmetallic Elements and Their Compounds Insert picture from...

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