Chap 22 Nonmetals

1
Chap 22 Nonmetals

• Quick review of periodic trends
• Hydrogen
• Nobel gases
• Halogens
• Oxygen sulfur
• Nitrogen phosphorous
• Carbon
• Silicon
• Boron

2
Overview The Periodic Table
3
Remember Effective nuclear charge?
Increasing Nuclear charge
4
Hydrogen Isotopes

• Three isotopes Protium 11H, deuterium 21H, and
  tritium 31H.
• Deuterium (D) 0.0156 of naturally occurring
  H.
• Tritium (T) radioactive t ½ 12.3 yr.
• Deuterium and tritium are substituted for H in
  compounds in order to provide a molecular marker.
  A labeled compound D2O.
• Deuteration (replacement of H for D) results in
  changes in kinetics of reactions (kinetic isotope
  effect).

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Hydrogen Properties

• The electron affinity of H is lower than any
  halogen.
• Colorless, odorless gas at room temperature.
• H2 nonpolar, two electrons, intermolecular
  forces weak (boiling point -253?C, melting point
  -259?C).
• H-H bond enthalpy fairly high (436 kJ/mol).
  Combustion 2H2(g) O2(g) ? 2H2O(l) ?H -571.7
  kJ
• Prepared by reduction of an acid Write a
  reaction for Zn and HCl
• Large quantities of hydrogen can be prepared by
  the reduction of methane in the presence of steam
  at 1100?C
• CH4(g) H2O(g) ? CO(g) 3H2(g)
• CO(g) H2O(g) ? CO2(g) H2(g)

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Hydrogen Uses

• NH3 production hydrogenation of vegetable oils
  to make
• margarine and shortening. (pi bond reduction)
• Manufacture methanol CO(g) 2H2(g) ? CH3OH(g)
• Types of binary hydrogen compounds include
• Ionic hydrides (LiH, metals and H)
• Metallic hydrides (TiH2, transition metals and
  H)
• Molecular hydrides (CH4, nonmetals and
  metalloids and H).

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Noble Gases

• Very unreactive.
• High ionization energies.
• complete valence octet
• He used as a coolant.
• The heavier noble gases slightly more reactive
  than the lighter ones.

8
Halogens

• Outer electron configurations ns2np6.
• Have large electron affinities.
• Most common oxidation state is -1, but oxidation
  states of 1, 3, 5 and 7 are possible.
• Good oxidizing agents.

9
Halogens

• Each halogen is the most electronegative element
  in its row.
• The bond enthalpy of F2 is low, so fluorine very
  reactive.
• Reduction potential of fluorine is very high.
• CFCs used as refrigerants.
• Fluorocarbons are used as lubricants and plastics
  (teflon).
• Chlorine is used in plastics (PVC),
  dichloroethane and other organic chemicals,
  paper and textile industries.
• HF has a high boiling point (strong H-bonds in
  the liquid).
• The ease of oxidation increases F- gt Cl- gt Br- gt
  I-.
• Hydrogen halides prepared by reacting salt with
  sulfuric acid NaCl(s) H2SO4(l) ? HCl(g)
  NaHSO4(s)
• Interhalogen compounds Diatomic molecules
  containing two different halogens have the
  more electronegative halogen in the -1 oxidation
  state and the other in the 1 oxidation state

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Oxygen

• Supports life complements CO2 animal vs plant
  kingdom
• Second most electronegative element
• Two allotropes O2 and O3.
• What is its electron configuration?
• What is the dominant oxidation state and why?
• The OO bond is strong (bond enthalpy 495
  kJ/mol).
• Prepared commercially Fractional distillation of
  air. (Normal boiling point of O2 is -183?C and
  N2 -196?C.)
• Used as an oxidizing agent in the steel industry
  to remove impurities
• Welding with acetylene 2C2H2(g) 5O2(g) ?
  4CO2(g) 2H2O(g)

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Ozone

• Important component of our atmosphere
• Protects us from solar radiation
• Pale blue, poisonous gas
• Dissociates to form oxygen O3(g) ? O2(g)
  O(g), ?H? 107 kJ.
• A stronger oxidizing agent the oxygen
• O3(g) 2H(aq) 2e- ? O2(g) H2O(l), E?
  2.07 V
• O2(g) 4H(aq) 4e- ? 2H2O(l), E? 1.23 V.
• Made by passing an electric current
  through3O2(g) ? 2O3(g)

12
Oxygen oxides

• Oxides compounds with oxygen in the 2- oxidation
  state.
• Nonmetal oxides covalent (carbon monoxide)
• Metal oxides ionic (ferric oxide rust, zinc
  oxide, etc.)
• Basic oxides oxides that react with water to
  form bases.
• BaO(s) H2O(l) ? Ba(OH)2(aq)

13
Oxygen peroxides

• Peroxides have an O-O bond and O in the -1
  oxidation state.
• Hydrogen peroxide is unstable and decomposes into
  water and oxygen
• 2H2O2(l) ? 2H2O(l) O2(g), ?H -196.0 kJ.
• The oxygen produced will kill bacteria.
• Peroxides important in biochemistry it is
  produced when O2 is metabolized.

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Oxygen super oxides

• Superoxides have an O-O bond and O in an
  oxidation state of -½ (superoxide ion is O2-)
  With active metals (KO2, RbO2 and CsO2).
• Super oxide dismutase (SOD) An antioxidant
  enzyme which helps protect cells from
  free-radical damage
• Cu2 SOD O2- ? Cu1SOD O2
• Cu1 SOD O2- 2H ? Cu2SOD H2O2.

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Sulfur

• Frasch process recovery of underground
• S deposits sulfate and sulfide minerals.
• Superheated water is forced into
• the deposit to melt the S.
• Compressed air then injected,
• forces the S(l) to the surface.
• Used in manufacture of sulfuric acid, and
  rubber.
• SO2 toxic to fungi, used to sterilize dried
  fruit
• Na2SO3 and NaHSO3 are used as preservatives
  cleaners
• When sulfur burns in air both SO2 (major
  product) and SO3formed.
• Oxidation of SO2 to SO3 requires a catalyst
  (usually V2O5 or Pt).
• SO3 is used to produce H2SO4
• SO3(g) H2SO4(l) ? H2S2O7(l) polysulfuric
  acid
• H2S2O7(l) H2O(l) ? 2H2SO4(l)

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Nitrogen

• Large component of our atmosphere (78 N2 vs 21
  O2)
• Unreactive because of the strong triple bond.
• Exception burning Mg or Li in air (78
  nitrogen)
• 3Mg(s) N2(g) ? Mg3N2(s)
• 6Li(s) N2(g) ? 2Li3N(s)
• N3- is a strong Lewis base (forms NH3 in water)
• Mg3N2(s) 6H2O(l) ? 2NH3(aq) 3Mg(OH)3(s)

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Nitrogen
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Nitrogen Preparation uses

• N2 produced by fractional distillation of air.
• Used as an inert gas to exclude oxygen from
  packaged foods, manufacture of chemicals,
  fabrication of metals, and production of
  electronics.
• Nitrogen is fixed by forming NH3 (Haber Process).
• NH3 converted to nitrites (preservative),
  nitrates, etc.
• Plants must secure their nitrogen in "fixed" form
  (incorporated in compounds such as nitrate ions,
  ammonia, urea
• Animals secure their nitrogen (and all other)
  compounds from plants or animals that have fed on
  plants.

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Nitrogen Perparation uses
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Nitrogen Ammonia

• Ammonia, very important nitrogen compound.
• Colorless toxic gas with a pungent aroma.
• In the laboratory, ammonia is produced by the
  reaction between NaOH and an ammonium salt
• NH4Cl(aq) NH3(aq) ? NH3(g) H2O(l) NaCl(aq)
• Commercially ammonia is prepared by the Haber
  process.

See handouts, Reaction trends
21
Carbon

• Constituents about 0.027 of the earths crust.
• The main constituent of living matter.
• Organic chemistry.
• Three crystalline forms of carbon
• graphite (soft and black),
• diamond (clear, hard and forms a covalent
  network), and
• buckminsterfullerene (molecular form of carbon,
  C60, the molecules look something like soccer
  balls).

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Carbon

• Oxides of Carbon
• CO is a good reducing agent (e.g. Fe3O4(s)
  4CO(g) ? 3Fe(s) 4CO2(g)).
• CO2 is produced when organic compounds are burned
  in oxygen
• C(s) O2(g) ? CO2(g)
• CH4(g) 2O2(g) ? CO2(g) 2H2O(l)
• C2H5OH(l) 3O2(g) ? 2CO2(g) 3H2O(l)
• CO2 is produced by treating carbonates with acid.

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Carbon

• Oxides of Carbon
• Fermentation of sugar to produce alcohol also
  produces CO2
• C6H12O6(aq) ? 2C2H5OH(aq) 2CO2(g)
• At atmospheric pressure, CO2 condenses to form
  CO2(s) or dry ice.
• CO2 is used as dry ice (refrigeration),
  carbonation of beverages, washing soda
  (Na2CO3.10H2O) and baking soda (NaHCO3.10H2O).

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Carbon

• Carbonic Acid and Carbonates
• When CO2 dissolves in water (somewhat soluble)
  carbonic acid forms
• CO2(aq) H2O(l) H2CO3(aq)
• Partial neutralization of H2CO3 gives hydrogen
  carbonates (bicarbonates), full neutralization
  gives carbonates.
• Many minerals contain CO32-.

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Silicon Oxygen

• Silicates
• 90 of the earths crust is composed of
  compounds of Si and O.
• Silicates are compounds where Si has four O atoms
  surrounding it in a tetrahedral arrangement.
• The oxidation state of Si is 4.
• Other minerals like zircon, ZrO4 have a similar
  structure.
• The silicate tetrahedra are building blocks for
  more complicated structures.

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Carbon Oxygen

• Silicates
• If two SiO42- link together one O atom is shared.
• This structure is the disilicate ion, Si2O76-.
• To determine the charge on the ion, we need to
  look at oxidation states (4 for Si and -2 for
  O) 2?(4) 7?(-20 -6.

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