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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.
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• 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.
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Nonmetals and Metalloids on the Periodic Table
Nonmetals coded in blue and metalloids in orange..
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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
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• 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
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• Preparation
– Industrial scale
– Laboratory scale
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Laboratory Generation of Hydrogen
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• 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−
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– 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.
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Binary Hydrides of Representative Elements
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• Isotopes of hydrogen
– Hydrogen has three naturally occurring
isotopes
– 11H, hydrogen, (99.985%)
– 21H , deuterium, symbol D, (0.015%)
3
– 1H , 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
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• 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
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• 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.
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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
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Phase Diagram for Allotropic Forms of Carbon
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• catenation – carbon has the unique ability
to form long chains stable rings
– Responsible for the millions of carboncontaining compounds
• Reacts with
– Metals to form carbides (strong bases), CaC2
– Silicon to form carborundum, SiC
– Nitrogen to form cyanides, C N
• Toxic
• Readily complexes metals
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Cyanide Pond for Extracting Gold
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• 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
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– 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
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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
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• 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
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• 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
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– Nitrogen Dioxide (NO2)
• Toxic
• Paramagnetic
• Dimerizes to N2O4 in the liquid and gas
phases
• Acidic oxide
–Shown in a disproportionation reaction
with cold water
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• 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
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• Phosphorus
– Occurs most commonly in nature as
phosphate rocks
• calcium phosphate [Ca3(PO4)2]
• fluoroapatite [Ca5(PO4)3F]
– Elemental phosphorus produced by
Ca5(PO4)3F
– Allotropic forms of phosphorus
• Red phosphorus
• White phosphorus
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Allotropes of Phosphorus
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– Reactions of phosphorus
• Formation of phosphine (PH3)
• Formation of phosphoric acid
• Reaction with the halogens
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• Acid production from halides
• Reaction with oxygen to produce acidic
oxides
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Structure of P4O6 and P4O10
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• Oxoacids of phosphorus
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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
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– Oxides
• Types of oxides
–Normal oxide, O22−
–Peroxide, O22−
–Superoxide, O2−
• All are strong BrØnsted bases
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• Bonding in oxides – ionic to covalent left to
right on the periodic table
• Acid-base character of oxides –
Basic
Amphoteric
Acidic
–Basicity increases down a group
– Peroxides
• H2O2 (hydrogen peroxide) – most common
example
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Structure of H2O2
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• 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
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– Ozone
• Toxic, light-blue gas
• Pungent odor
• Essential component of the atmosphere
• Structure
• Powerful oxidizing agent
• Preparation
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Preparation of O3
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• 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
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FeS2
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Puckered Ring of S8
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– Extracted by the Fasch process
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– Forms wide variety of oxidation numbers
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– Hydrogen sulfide – H2S
• Used in qualitative analysis
• Preparation
• Colorless gas with odor of rotten eggs
• Toxic
• Weak diprotic acid
• Reducing agent in basic solution
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– Oxides of sulfur
• Sulfur dioxide (SO2)
–Pungent colorless gas
–Toxic
–Preparation
–Acidic oxide
–Oxidation
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• 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.
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• 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
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–Hot concetrated sulfuric acid reacts with
less active metals
–Depending on the reducing agent,
sulfate may be reduced
–Oxidizes nonmetals
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• 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
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24.6 The Halogens
• The halogens—fluorine, chlorine, bromine,
and iodine—are reactive nonmetals.
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• 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.
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– 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.
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• 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
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• Preparation and Properties of F2 and Cl2 –
determined by their strong oxidizing
capability
– Fluorine
• From liquid HF
• At 70oC
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Electrolytic Preparation of F2
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– Chlorine
• Electrolysis of molten NaCl
• Overall reaction
• Chlor-alkali process
–Designed to prevent side reactions
–Mercury cell
–Diaphragm cell
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Mercury Cell in the Chlor-alkali Process
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Diaphragm Cell in the Chlor-alkali Process
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• 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.
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– Hydrogen Halides
• Preparation from elements – can occur
violently
• Preparation varies with the halogen, for
example
HCl
HF
HBr
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– 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
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– Aqueous solutions of HX
• Acidic
• Variation in acid strength
– Oxoacids – halogens form a series of acids
increasing acid strength
• Only Cl forms the entire series
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• 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
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– Bromine
• Insecticides (BrCH2CH2Br)
• Scavenger for Pb in gasoline
• Photographic films (AgBr)
– Iodine
• Antiseptic (tincture of iodine)
• Thyroxine (thyroid hormone derivative)
• Cloud seeding (AgI)
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Key Points
• General properties of the nonmetals
• Hydrogen
– Properties
– Preparation
– Binary Halides
• Ionic
• Covalent
• Interstitial
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– Isotopes of hydrogen
• Hydrogen (protium)
• Deuterium
• Tritium
– Hydrogenation
– The hydrogen economy
• Carbon
– Properties
– Allotropes
• Diamond
• Graphite
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– Carbides
– Cyanides
– Oxides
• Carbon monoxide
• Carbon dioxide
• Nitrogen and Phosphorus
– Nitrogen
•
•
•
•
Properties
Nitrides
Ammonia
Hydrazine
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• Oxides
–Nitrous oxide
–Nitric oxide
–Nitrogen dioxide
• Nitric Acid
– Phosphorus
• Properties
• Allotropes
–White phosphorus
–Red phosphorus
• Phosphine
• Halogen compounds
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• Oxides
• Oxoacids
• Oxygen and Sulfur
– Oxygen
• Properties
• Allotropes
• Oxides
–Normal oxide
–Peroxide
–Superoxide
• Acidity of oxides
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• Hydrogen peroxide
• Ozone
– Sulfur
• Properties
• Industrial production
• Hydrogen sulfide
• Oxides
–Sulfur dioxide
–Sulfur trioxide
• Sulfuric Acid
–Production
–Uses
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• 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
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• Hydrogen Halides
• Oxoacids
– Uses of the halogens
• Fluorine
• Chlorine
• Bromine
• Iodine
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