Chemical Equations Preparation for College Chemistry Columbia University Department of Chemistry

The Chemical Equation Writing and Balancing Equations Information in an Equation Types of Chemical Equations Heat in Chemical Equations The Greenhouse Effect Chapter Outline

The Chemical Equation Shorthand Expression for a Chemical Change Fe + Al 2 O 3   Products Al + Fe 2 O 3   Reactants   Conditions (s) (l) (s)   Physical State    Stoichiometric Coefficients 

Writing Chemical Equations Mg 3 (PO 4 ) 2 + H 2 O  Write the skeleton equation Mg(OH) 2 + H 3 PO 4  Identify the Reaction   magnesium hydroxide + phosphoric acid magnesium phosphate + water 2  3  2  Find the Stoichiometric Coefficients (Balance) 3Mg PO 4 Mg PO 4 R P 12 H 14 O 12 H

Types of Chemical Equations   Combination:   Decomposition   Single -Displacement   Double -Displacement A + B AB AB A + B A + BC AB + C AB + CD AD + CB

Combination Reactions 2Mg(s) + O 2 (g) 2MgO(s) m metal + Oxygen metal oxide m nonmetal + Oxygen non metal oxide 2S (s) + 3O 2 (g) 2SO 3 (g) m metal + nonmetal Salt 2Na (s) + Cl 2 (g) 2NaCl(s) m metal oxide + water Metal Hydroxide m nonmetal oxide + water Oxy-acid MgO (s) + H 2 O(l) Mg(OH) 2 (s) SO 3 (g) + H 2 O(g) H 2 SO 4 (s)

Decomposition Reactions m Metal oxides m Carbonates and Hydrogen carbonates m Other decomposition reactions 2HgO(s) 2Hg (l) + O 2 (g) 2PbO 2 (g) 2PbO (g) + O 2 (g) CaCO 3 (s)CaO (s) + CO 2 (g) 2NaHCO 3 (s)Na 2 CO 3 (s) + H 2 O(l) + CO 2 (g) KClO 3 (s)2KCl (s) + 3O 2 (g) NaNO 3 (s)NaNO 2 (s) + O 2 (g) 2H 2 O 2 (l)2H 2 O (l) + O 2 (g) 2NaN 3 (s)2Na (s) + 3N 2 (g)

Single-Displacement Reactions Zn(s) + 2HCl(g) H 2 (g) + ZnCl 2 (s) m metal + acid Hydrogen + Salt m metal + waterHydrogen + metal hydroxide or oxide 2Na(s) + 2H 2 O(l) H 2 (g) + 2NaOH(aq) m metal + Salt Salt + metal Zn(s) + CuSO 4 (aq) ZnSO 4 (aq) + Cu(s) m halogen + halide salt Halide salt + Halogen Cl 2 (g) + 2NaBr(aq) 2NaCl(aq) + Br 2 (l)

Double-Displacement Reactions NaCl(aq) + KNO 3 (aq) NaNO 3 (aq) + KCl(aq) AB + CD AD + CB   Evolution of Heat (Neutralization Reactions)   Formation of an Insoluble precipitate   Gas Formation Physical Evidences for double-displacement

+ Ionic Dissolution

Precipitation Reactions NO 3 - All nitrates are soluble Appendix V p. A19 Solubility Rules Cl - AgCl, Hg 2 Cl 2, Pb 2 Cl 2 All chlorides are soluble, except SO 4 2- Most sulfates are soluble, except SrSO 4, PbSO 4 and BaSO 4 CaSO 4 is slightly soluble CO 3 2- All carbonates are insoluble, except Group I and NH 4 + OH - All hydroxides are insoluble, except group I Sr(OH) 2 and Ba(OH) 2. Ca(OH) 2 is slightly soluble S 2- All sulfides except Groups I and II and NH 4 + are insoluble

Solubility Rules Used to predict results of precipitation reactions Example 1 What happens when solutions of Ba(NO 3 ) 2 and Na 2 CO 3 are mixed? Ions present: Ba 2+ (aq), NO 3 - (aq), Na + (aq), CO 3 2- (aq) Possible precipitates: BaCO 3, NaNO 3 According to solubility rules, BaCO 3 is insoluble Ba 2+ (aq) + CO 3 2- (aq) BaCO 3 (s)

Solubility Rules Mix solutions of BaCl 2, NaOH Example 2 ions present: Ba 2+ (aq), Cl - (aq), Na + (aq), OH - (aq) possible precipitates: Ba(OH) 2, NaCl both are soluble; no reaction

Net Ionic Equations (Spectator ions do not appear) Mix solutions of Cu(NO 3 ) 2, NaOH ions present: Cu 2+ (aq), NO 3 - (aq), Na + (aq), OH - (aq) possible precipitates: Cu(OH) 2, NaNO 3 NaNO 3 is soluble; Cu(OH) 2 is not. Net Ionic Equation: Cu 2+ (aq) + 2 OH - (aq) Cu(OH) 2 (s) Example Spectator ions: Na + (aq), NO 3 - (aq)

Heat in Chemical Reactions Endothermic Reaction Potential Energy Reactants Products Net Energy absorbed Activation Energy Time

Heat in Chemical Reactions Exothermic Reaction Potential Energy Reactants Products Net Energy released Activation Energy Time

Greenhouse Effect session/61/39/ /8!xrn_29_0_A &bkm_8_29 session/61/39/ /12!xrn_39_0_A session/61/39/ /25!xrn_4_0_A session/61/39/ /27!xrn_2_0_A

Redox Reactions (electron-transfer reactions) Oxidation Number Oxidation & ReductionBalancing Redox Reactions

Oxidation number(oxidation state) 1. ON of an element in an elementary substance is zero 2. H ON = +1, except in metal hydrides NaH, CaH 2 What is it? 3. O ON = -2 in most compounds, -1 in peroxides Na 2 O 2, +2 in OF 2 4. ON of metallic elements in ionic compounds is positive. 5. Negative ON is assigned to the most electronegative element in a covalent compound. “ pseudocharge” assigned according to arbitrary rules. (rules p.436) # of e - lost, gained or unequally shared by the atom

Oxidation number. Calculation 6. In a compound: Determine the ON of Cr in Cr 2 O In a PAI: Cr 2 O 7 2- Determine the ON of As in K 3 AsO 4 K 3 AsO As + (-2) x 4 +(+1) x 3 = 0 -2 ? ? 2Cr + (-2) x 7 = -2 2Cr = +12 Cr = +6 As = +5

Oxidation & Reduction oxid. # H increases from 0 to +1 (oxidizes) oxid. # O decreases from 0 to -2 (reduces) Oxidation (lost of electrons) Reduction (gain of electrons) ON O 2 (g) + H 2 (g) 2H 2 O(l) REDUCING AGENT OXIDIZING AGENT

Balancing Redox Equations Two Methods q Oxidation number method q Ion-electron method (Molecular redox equations) (Ionic redox equations)

Oxidation number method KMnO 4 + HCl + H 2 S KCl + MnCl 2 + S + H 2 O Reduction Oxidation Oxidation: Reduction: Mn +7 +5e- Mn +2 S -2 S 0 + 2e- x 5 2Mn S -2 2Mn S x 2 2Mn S e-2Mn S0+ 5 S0 + 10e-

Ion-electron method (rules p ) q Acidic Medium H + (aq) q Basic Medium OH - (aq) Mass and charge must balance Neutralization: H + (aq) + OH - (aq) H 2 O(l)

Ion-electron method (Acidic Medium) K + (aq) + MnO 4 - (aq) + H + (aq) + Cl - (aq ) + 2H + (aq) + S 2- (aq) = write the molecular equation in ionic form KMnO 4 + HCl + H 2 S KCl + MnCl 2 + S + H 2 O Oxidation Eliminating spectator ions (appear in both sides of the equation) K + (aq) + Cl - (aq ) + Mn 2+ (aq) + 2Cl - (aq ) ) + S 0 (s) + H 2 O MnO 4 - (aq) + H + (aq) + S 2- (aq)Mn 2+ (aq) + S 0 (s) Reduction Net ionic Equation

Oxidation: Reduction: x 2 Write the two half reactions Balance elements other than O and H Balance O and H, acidic medium: + 8H + + 4H 2 O Balance each half reaction electrically with electrons: MnO 4 - Mn +2 x 5S -2 S 0 + 5e- + 2e- Equalize loss and gain of e- Add the half equations 2MnO H + +5 S -2 2Mn S 0 + 8H 2 O

Ion-electron method (Basic Medium) Oxidation Reduction SbO 2 - (aq) + ClO 2 (aq) Sb(OH) 6 - (aq) + ClO 2 - (aq) Oxidation: Reduction: Write the two half reactions SbO 2 Sb(OH) 6 - ClO 2 ClO 2 -

Balance elements other than O and H Balance O and H, ACIDIC medium, + 2H + + 4H 2 O NEUTRALIZE: add OH- in both sides of the equation Oxidation:SbO 2 - Sb(OH) OH - + 4H 2 OSbO 2 - Sb(OH) H 2 O+ 2OH - Balance each half reaction electrically with electrons: SbO 2 - Sb(OH) H 2 O+ 2OH - + 2e- Reduction:ClO 2 ClO e- Equalize loss and gain of e- x 2 SbO OH - + 2H 2 O + 2ClO 2 2ClO Sb(OH) 6 -

Ba Ba e- Zn Zn e- Cr Cr e- Pb Pb e- Fe Fe e- Ag Ag + + e- As As e- H 2 2H + + 2e- Cu Cu e- Ni Ni e- Sn Sn e- Ca Ca e- Na Na + + e- Mg Mg e- Al Al e- K K + + e- Hg Hg e- Au Au e- Activity Series of Metals (table 17.3) Ease of oxidation

Activity Series of Metals Useful to Predict the Course of Chemical Reactions � Na(s) + HCl(aq) � Hg + AgNO 3 NaCl(aq) + H 2 ? ? No Reaction Na(s) + 2H + (aq) 2Na + (aq) + H 2 Net Ionic Reaction: � Cr(s) + Sn(SO 4 )(aq) Sn + Cr 2 (SO 4 ) 3 ? Cr(s) + 3Sn 2+ (aq) 2Cr 3+ (aq) + Sn Net Ionic Reaction:

 Electrolytic Cells Anode: the OXIDATION SITE Cathode: the REDUCTION SITE Use electrical energy to produce a chemical reaction  Voltaic (Galvanic) Cells Use chemical reactions to produce electrical energy Applications  Corrosion