Chemistry of Nonmetals
Dr.Riham Hazzaa
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OXIDATION-REDUCTION REACTIONS
• A redox reaction involves the transfer of electrons
between reactants
• Electrons gained by one species must equal electrons
lost by another
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• Both oxidation and reduction must occur
simultaneously.
• Oxidation: removal of electrons
• Reduction: gain of electrons
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• oxidising and reducing agents
• An oxidizing agent is an element which causes
oxidation (and is reduced as a result) by
removing electrons from another species.
oxidizing agent is the electron acceptor
• A reducing agent is an element which causes
reduction (and is oxidized as a result) by giving
electrons to another species. reducing agent
is the electron donor.
Na(s) + Cl (g) → NaCl(s)
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The main group metals are oxidized in
all of their chemical reactions. These
metals are oxidized when they react
with nonmetal elements.
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• Nonmetals can undergo both oxidation and
reduction.
• Phosphorus, is oxidized when it reacts with
oxygen to form P4O10.
• it is reduced when it reacts with calcium to
form calcium phosphide.
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• Phosphorus (EN = 2.19) is less electronegative
than oxygen (EN = 3.44). When these elements
react, the electrons are drawn toward the more
electronegative oxygen atoms. Phosphorus is
therefore oxidized in this reaction, and oxygen is
reduced.
• Calcium (EN = 1.00), on the other hand, is
significantly less electronegative than phosphorus
(EN = 2.19). When these elements react, the
electrons are drawn toward the more
electronegative phosphorus atoms. As a result,
calcium is oxidized and phosphorus is reduced.
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The behavior of the nonmetals
• Nonmetals tend to oxidize metals.
2 Mg(s) + O2(g)→ 2 MgO(s)
• Nonmetals with relatively large electronegativities
(such as oxygen and chlorine) oxidize substances
with which they react.
2 H2S(g) +3 O2(g)→ 2 SO2(g)+2 H2O(g)
• Nonmetals with relatively small electronegativities
(such as carbon and hydrogen) can reduce other
substances.
CuO(s) +H2(g)→ Cu(s) + H2O(g)
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Hydrogen
• Compounds of hydrogen are frequently called
hydrides, even though the name hydride
describes compounds that contain an H- ion.
H+
=
1s0
H
=
1s1
H-
=
1s2
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• Because hydrogen forms compounds with
oxidation numbers of both +1 and -1, many
periodic tables include this element in both Group
IA (with Li, Na, K, Rb, Cs, and Fr) and Group VIIA
(with F, Cl, Br, I).
• The first ionization energy of hydrogen (1312
kJ/mol), is halfway between the elements with
the largest (2372 kJ/mol) and smallest (376
kJ/mol) ionization energies.
• Hydrogen has an electronegativity (EN = 2.20)
halfway between the extremes of the most
electronegative (EN = 3.98) and least
electronegative (EN = 0.7) elements.
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• Hydrogen is oxidized by elements that are
more electronegative to form compounds in
which it has an oxidation number of +1.
• Hydrogen is reduced by elements that are less
electronegative to form compounds in which
its oxidation number is -1.
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Formation of Hydrogen
• By reacting an active metal with water.
2 Na(s) +2 H2O(l)
2 Na+(aq) + 2 OH-(aq) + H2(g)↑
• By reacting a less active metal with a strong
acid.
Zn(s) + 2 HCl(aq)
Zn2+(aq) + 2 Cl-(aq)
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+ H2(g)↑
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• By reacting an ionic metal hydride with water
NaH(s) + H2O(l)
Na+(aq) + OH-(aq) + H2(g)
• By decomposing water into its elements with
an electric current.
electrolysis
2 H2O(l)
2 H2(g) +
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O2(g)
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The Chemistry of Nitrogen (GroupV)
• A neutral nitrogen atom contains five valence
electrons: 2s2 2p3.
• Because the covalent radius of a nitrogen
atom is relatively small (only 0.070 nm),
nitrogen atoms come close enough together
to form very strong bonds.
• The strength of the nitrogen-nitrogen triple
bond makes the N2 molecule very unreactive.
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The Synthesis of Ammonia NH3 Haber process
• The Haber process, a mixture of N2 and H2 gas at
200 to 300 atm and 400 to 600oC is passed over a
catalyst of finely divided iron metal.
Fe
N2(g) + 3 H2(g)
2 NH3(g)
• Two-thirds of the ammonia used for fertilizers is
converted into solids such as ammonium nitrate,
NH4NO3; ammonium phosphate, (NH4)3PO4;
ammonium sulfate, (NH4)2SO4; and urea,
H2NCONH2. The other third is applied directly to
the soil as anhydrous ammonia.
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• The Synthesis of Nitric Acid Ostwald process
4 NH3(g) + 5 O2(g)
4 NO(g)
2 NO(g) + O2(g)
2 NO2(g)
3 NO2(g) + H2O(l)
2 HNO3(aq) + NO(g)
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+ 6 H2O(g)
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The Nitrogen Oxides
1. Dinitrogen oxide, N2O which is also known as
nitrous oxide, can be prepared by carefully
decomposing ammonium nitrate.
170 to 200oC
NH4NO3(s)
N2O(g) + 2 H2O(g)
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2. Nitrogen oxide, or nitric oxide,
N2(g) + O2(g)
2 NO(g)
Nitrogen oxide NO can be prepared in the
laboratory by reacting copper metal with
dilute nitric acid.
3 Cu(s) + 8 HNO3(aq)
3 Cu(NO3)2(aq) + 2 NO(g) + 4 H2O(l)
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3. Nitrogen dioxide NO2
NO reacts rapidly with O2 to form nitrogen
dioxide (once known as nitrogen peroxide),
which is a dark brown gas at room temperature.
2 NO(g)
+
O2(g)
2 NO2(g)
It can also be made by reacting copper metal
with concentrated nitric acid,
Cu(s) + 4 HNO3(aq)
Cu(NO3)2(aq) + 2 NO2(g) + 2 H2O(l)
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4. Dinitrogen pentoxide
By carefully removing water from concentrated
nitric acid at low temperatures with a
dehydrating agent we can form dinitrogen
pentoxide.
4 HNO3(aq) + P4O10(s)
2 N2O5(s) + 4 HPO3(s)
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