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UNIT 5

UNIT 5. Aqueous Reactions and Solution Stoichiometry Redox Reactions. Redox Reactions. Redox is short for “ oxidation-reduction .” Plating and combustion are redox reactions. Consider the following redox reaction: Fe (s) + Ni(NO 3 ) 2 (aq)  Fe(NO 3 ) 2 (aq) + Ni (s). O xidation

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UNIT 5

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  1. UNIT 5 Aqueous Reactions and Solution Stoichiometry Redox Reactions

  2. Redox Reactions • Redox is short for “oxidation-reduction.” • Plating and combustion are redox reactions. • Consider the following redox reaction: • Fe (s) + Ni(NO3)2 (aq) Fe(NO3)2(aq) + Ni (s) Oxidation Is Loss of electrons Reduction Is Gain of electrons • Fe goes from an element to an ion: Fe  Fe2+ • Fe got oxidized • Ni goes from an ion to an element: Ni2+  Ni • Ni2+ gotreduced

  3. Oxidation Numbers We determine whether a chemical species has been oxidized or reduced by looking at the change in its oxidation number. Fe (s) + Ni(NO3)2 (aq) Fe(NO3)2(aq) + Ni (s) ox # 0 +2 +2 0 • The oxidation number of an element is 0. • The oxidation number of a monatomic ion is its charge. • The ox # for Fe went up. It was oxidized. • The ox # for Ni2+ went down. It was reduced (reduced = down). Oxidation Is Loss of electrons Reduction Is Gain of electrons

  4. Oxidation Numbers - Rules • The oxidation number of an element is 0. • The oxidation number of a monatomic ion is its charge. • The oxidation number of oxygen in a compound is usually -2 (exception is peroxide, O22-, where the oxidation number is -1). • The oxidation number of hydrogen in a compound is 1 when bonded to nonmetals and -1 when bonded to metals. • The oxidation number of fluoride is always -1. • The sum of the oxidation numbers in a neutral compound is zero. • The sum of the oxidation numbers in a polyatomic ion is the charge of the ion.

  5. Oxidation Numbers - Examples Species oxidation Species oxidation number number Fe Na+ H2 H+ O2 O2- Al3+ Cl- Fe2+ N3- S in SO42- : Cr in Cr2O72- N in NO3- S in SO32- N in NO2- P in PO43- C in CO32- 0 0 0 +3 +2 +1 +1 -2 -1 -3 overall charge of ion is -2 = ox # of S + 4(ox # of O) -2 = ox # of S + 4(-2) -2 = ox # of S – 8 -2+8 = 6 = ox # of S +6 +4 +5 +5 +3 +4

  6. Redox Reactions To see if a redox reaction has occurred, check to see if the oxidation numbers of any element involved in the reaction have changed. Mg (s) + Fe(NO3)2 (aq) Mg(NO3)2(aq) + Fe (s) ox #0 +2 +5 -2 +2 +5 -2 0 • Mg goes from an element to an ion: Mg  Mg2+ • Mg gotoxidized. Fe2+ was the oxidizing agent. • Fe goes from an ion to an element: Fe2+  Fe • Fe2+ gotreduced. Mg was the reducing agent. • This type of redox reaction, where the ion in solution is replaced through the oxidation of an element, is called a single replacement reaction. Oxidation Is Loss of electrons Reduction Is Gain of electrons

  7. Redox Reactions • Single replacement reactions have the form: • A + BX (aq)  AX (aq) + B • Single replacement reactions will occur if the element A is more likely to get oxidized than the ion B. • The activity series gives this information. See the table in the text. Oxidation Is Loss of electrons Reduction Is Gain of electrons • Notice that H2 (g) is included in this series. The series can tell us whether a metal will dissolve in acid.

  8. Will the following pairs react? Zn + Cu(NO3)2 (aq) Cu + Zn(NO3)2 (aq) Zn + NaNO3 (aq) Au + AgNO3 (aq) Ag + Au(NO3)3 (aq) Will the following metals dissolve in acid? Li Al Ni Cu Au What gas will be produced? H2 y n n n y y y y n n

  9. Redox Reactions - Examples Molecular equation: Mg (s) + Fe(NO3)2 (aq) Mg(NO3)2(aq) + Fe (s) Net Ionic equation: Mg (s) + Fe2+ (aq) Mg2+(aq) + Fe (s) Mg got oxidized. Fe2+ was the oxidizing agent. Fe2+ got reduced. Mg was the reducing agent.

  10. Redox Reactions - Examples When Zn dissolves in sulfuric acid, themolecular equationis Zn (s) + H2SO4 (aq) ZnSO4(aq) + H2 (g), thetotal ionic equationis Zn(s) + 2H+(aq) + SO42-(aq) Zn2+(aq) + SO42-(aq) + H2 (g), and thenet ionic equationis: Zn(s) + 2H+(aq) Zn2+(aq) + H2 (g) Zn got oxidized. H+ was the oxidizing agent. H+ got reduced. Zn was the reducing agent.

  11. Summary of Reactions • Combination A + B  AB • Decomposition AB  A + B • Metathesis AX + BY  AY + BX • acid-base neutralization • (example) HCl + NaOH  H2O + NaCl • Redox (example) Fe(s) + Ni2+ Fe2+ + Ni(s) • combustion • (example) CH2O + O2 CO2 +H2O • single replacement • (example) 2Na(s) + 2H2O  2NaOH + H2(g)

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