PHILIP DUTTON UNIVERSITY OF WINDSOR DEPARTMENT OF CHEMISTRY AND BIOCHEMISTRY

1. GENERAL CHEMISTRY Principles and Modern Applications TENTH EDITION PETRUCCI HERRING MADURA BISSONNETTE 5 Introduction to Reactions in Aqueous Solutions PHILIP DUTTON UNIVERSITY OF WINDSOR DEPARTMENT OF CHEMISTRY AND BIOCHEMISTRY

2. Introduction to Reactionsin Aqueous Solutions

3. 5.1 The Nature of Aqueous Solutions • Water • Inexpensive • Can dissolve a vast number of substances • Many substances dissociate into ions • Aqueous solutions are found everywhere • Seawater • Living systems

4. Types of Electrolytes • A strong electrolyte dissociates completely. • A strong electrolyte is present in solution almost exclusively as ions. • Strong electrolyte solutions are good conductors. • A nonelectrolyte does not dissociate. • A nonelectrolyte is present in solution almost exclusively as molecules. • Nonelectrolyte solutions do not conduct electricity. • A weak electrolyte dissociates partially. • Weak electrolyte solutions are poor conductors. • Different weak electrolytes dissociate to different extents.

5. Conduction Illustrated • Electric current is a flow of charged particles. • One type of current is electrons flowing through a wire, from cathode (negative electrode) to anode (positive electrode). • Another type of current: anions and cations moving through a solution as shown here. Cations move to the cathode, anions move to the anode. • Of course, an external source of potential (voltage) is required in either case!

6. 1M MgCl2 1M CH3OH 1M CH3COOH • Electrolytesdissociate to produce ions. • The more the electrolyte dissociates, the more ions it produces.

7. A weak electrolyte: CH3CO2H(aq)← CH3CO2-(aq) + H+(aq) → A non-electrolyte: CH3OH(aq) • Essentially all soluble ionic compounds and only a relatively few molecular compounds are strong electrolytes. • Most molecular compounds are either nonelectrolytes or weak electrolytes. MgCl2(s) → Mg2+(aq) + 2 Cl-(aq) A strong electrolyte: General Chemistry: Chapter 5

8. Is it a strong electrolyte, a weak electrolyte, or a nonelectrolyte? • Strong electrolytes include: • Strong acids (HCl, HBr, HI, HNO3, H2SO4, HClO4) • Strong bases (IA and IIA hydroxides) • Most water-soluble ionic compounds • Weak electrolytes include: • Weak acids and weak bases • A few ionic compounds How do we tell whether an acid (or base) is weak? • Nonelectrolytes include: • Most molecular compounds • Most organic compounds (most of them are molecular)

9. Figure 5-5 The hydrated proton

10. Ion Concentrations in Solution • “Trick” question … • What is the concentration of Na2SO4 in a solution prepared by diluting 0.010 mol Na2SO4 to 1.00 L? • The answer is: • … zero … • WHY?? • And … how do we describe the concentration of this solution?

11. Calculating Ion Concentrations in Solution • In 0.010 M Na2SO4: • two moles of Na+ ions are formed for each mole of Na2SO4 in solution, so [Na+] = 0.020 M. • one mole of SO42– ion is formed for each mole of Na2SO4 in solution, so [SO42–] = 0.010 M. • An ion can have only one concentration in a solution, even if the ion has two or more sources.

12. Relative Concentrations in Solution MgCl2(s) → Mg2+(aq) + 2 Cl-(aq) MgCl2(s) → Mg2+(aq) + 2 Cl-(aq) In 0.0050 MMgCl2: Stoichiometry is important. [Mg2+] = 0.0050 M [Cl-] = 0.0100 M [MgCl2] = 0 M [Mg2+] = 0.0050 M [Cl-] = 0.0100 M [MgCl2] = 0 M General Chemistry: Chapter 5

14. Reactions that Form Precipitates • There are limits to the amount of a solute that will dissolve in a given amount of water. • If the maximum concentration of solute is less than about 0.01 M, we refer to the solute as insoluble in water. • When a chemical reaction forms such a solute, the insoluble solute comes out of solution and is called a precipitate.

15. 5-2 Precipitation Reactions • Soluble ions can combine to form an insoluble compound. • Precipitation occurs. • A test for the presence of chloride ion in water. Ag+(aq) + Cl-(aq) → AgCl(s) Figure 5-6 • Qualitative test for Cl- in tap water

16. Spectator ions Ag+(aq) + NO3-(aq) + Na+(aq) + I-(aq) → AgI(s) + Na+(aq) + NO3-(aq) Net Ionic Equations “whole formula form”: AgNO3(aq) +NaI(aq) → AgI(s) + NaNO3(aq) ionic form: Ag+(aq) + NO3-(aq) + Na+(aq) + I-(aq) → AgI(s) + Na+(aq) + NO3-(aq) Net ionic equation: Ag+(aq) + I-(aq) → AgI(s) General Chemistry: Chapter 5

17. Predicting Precipitation Reactions AgNO3(aq) NaI(aq) AgI(s) Na+(aq) NO3-(aq) Figure 5-7 • A precipitate of silver iodide General Chemistry: Chapter 5

19. With these guidelines we can predict precipitation reactions. • When solutions of sodium carbonate and iron(III) nitrate are mixed, a precipitate will form. • When solutions of lead acetate and calcium chloride are mixed, a precipitate will form.

21. 5-3 Acid-Base Reactions • Latin acidus • Sour taste • Arabic al-qali • Bitter taste • Acid-Base theory • Svante Arrhenius 1884 • Brønsted and Lowry 1923 Figure 5-9 • An acid, a base, and an acid–base indicator

22. Arrhenius: acids provide H+ in aqueous solution. Strong acids completely ionize: Weak acid ionization is not complete: Brønsted Lowry acids are proton donors. → HCl(aq) H+(aq) + Cl-(aq) ← → CH3CO2H(aq) H+(aq) + CH3CO2-(aq) Acids

23. Bases provide OH- in aqueous solution. Strong bases: Weak bases: Brønsted Lowry bases are proton acceptors. → H2O NaOH(aq) Na+(aq) + OH-(aq) ← → Bases NH3(aq) + H2O(l) OH-(aq) + NH4+(aq)

24. Acids have ionizable hydrogen atoms. CH3CO2H or HC2H3O2 Bases areoften indicated by combination of hydroxide ion with various metal cations. KOH or can be identified by chemical equations Na2CO3(s) + H2O(l)→ HCO3-(aq) + 2 Na+(aq) + OH-(aq) Recognizing Acids and Bases

25. ← ← → → Reactions of Acids and Bases:Strong and Weak Acids • Strong acids are strong electrolytes; completely ionized in water. • In water: HCl(g) → H+(aq) + Cl–(aq) No HCl in solution, only H+ and Cl– ions. • Weak acidsare weak electrolytes. Some of the dissolved molecules ionize; the rest remain as molecules. • In water: CH3COOH(l) H+(aq) + CH3COO–(aq) Just a little H+ forms. Some acids have more than one ionizable hydrogen atom. They ionize in “steps” (more in Chapter 15). H2SO4→ H+ + HSO4– HSO4– H+ + SO42–

26. ← ← → → Just a little OH– forms. Most of the weak base remains in the molecular form. Strong and Weak Bases • Strong bases: Most are ionic hydroxides (Group IA and IIA, though some IIA hydroxides aren’t very soluble). • Weak bases: Like weak acids, they ionize partially. Ionization process is different. • Weak bases formOH– by accepting H+ from water … • NH3 + H2O NH4++ OH– • CH3NH2 + H2O CH3NH3++ OH– • methylamine methylammonium ion

28. More Acid-Base Reactions • Milk of magnesia Mg(OH)2 Mg(OH)2(s) + 2 H+(aq)→ Mg2+(aq) + 2 H2O(l) Mg(OH)2(s) + 2 CH3CO2H(aq) → Mg2+(aq) + 2 CH3CO2-(aq) + 2 H2O(l)

29. CaCO3(s) + 2 H+(aq) → Ca2+(aq) + H2CO3(aq) CaCO3(s) + 2 H+(aq) → Ca2+(aq) + H2O(l) + CO2(g) More Acid-Base Reactions • Limestone and marble. But: H2CO3(aq)→ H2O(l) + CO2(g)

30. Acid–Base Reactions:Neutralization • In the reaction of an acid with a base, the identifying characteristics of each “cancel out.” • Neutralizationis the (usually complete) reaction of an acid with a base. • The products of this neutralization are water and a salt.

31. Acid–Base Reactions:Net Ionic Equations HCl + NaOH H2O + NaCl • In the reaction above, the HCl, NaOH, and NaCl all are strong electrolytes and dissociate completely. • The actual reaction occurs between ions. Na+ and Cl– are spectator ions. H+ + Cl– + Na+ + OH– H2O + Na++ Cl– H++ OH– H2O A net ionic equation shows the species actually involved in the reaction.

34. Reactions InvolvingOxidation and Reduction • Oxidation: Loss of electrons • Reduction: Gain of electrons • Both oxidation and reduction must occur simultaneously. • A species that loses electrons must lose them to something else (something that gains them). • A species that gains electrons must gain them from something else (something that loses them). • Historical: “oxidation” used to mean “combines with oxygen”; the modern definition is much more general.

35. Oxidation Numbers • An oxidation number is the charge on an ion, or a hypothetical charge assigned to an atom in a molecule or polyatomic ion. • Examples: in NaCl, the oxidation number of Na is +1, that of Cl is –1 (the actual charge). • In CO2 (a molecular compound, no ions) the oxidation number of oxygen is –2, because oxygen as an ion would be expected to have a 2– charge. • The carbon in CO2 has an oxidation number of +4 (Why?)

36. Rules for Assigning Oxidation Numbers • For the atoms in a neutral species—an isolated atom, a molecule, or a formula unit—the sum of all the oxidation numbers is 0. • For the atoms in an ion, the sum of the oxidation numbers is equal to the charge on the ion. • In compounds, the group 1A metals all have an oxidation number of +1and the group 2A metals all have an oxidation number of +2. • In compounds, the oxidation number of fluorine is –1. • In compounds, hydrogen has an oxidation number of +1. • In most compounds, oxygen has an oxidation number of –2. • In binary compounds with metals, group 7A elements have an oxidation number of –1, group 6A elements have an oxidation number of –2, and group 5A elements have an oxidation number of –3.

37. Hematite is converted to iron in a blast furnace. D Fe2O3(s) + 3 CO(g)→ 2 Fe(l) + 3 CO2(g) D Fe2O3(s) + 3 CO(g) → 2 Fe(l) + 3 CO2(g) D Fe2O3(s) + 3 CO(g) → 2 Fe(l) + 3 CO2(g) 5-4 Oxidation-Reduction Reactions: Some General Principles Oxidation and reduction always occur together. Fe3+ is reduced to metallic iron. CO(g) is oxidized to carbon dioxide.

38. Assign oxidation states: D Fe2O3(s) + 3 CO(g) → 2 Fe(l) + 3 CO2(g) Oxidation State Changes 0 3+ 2- 2+ 2- 4+ 2- Fe3+ is reduced to metallic iron. CO(g) is oxidized to carbon dioxide.

39. Fe2O3(s) + 2 Al(s)  Al2O3(s) + 2 Fe(l) Figure 5-11 Thermite Reaction

41. Oxidation and Reduction Half-Reactions Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) Figure 5-12 • An oxidation-reduction reaction

42. The reaction represented by two half-reactions. Zn(s) → Zn2+(aq) + 2 e- Oxidation: Cu2+(aq) + 2 e-→ Cu(s) Reduction: Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq) Overall:

43. Oxidation and Reduction • Oxidation • O.S. of some element increases in the reaction. • Electrons are on the right of the equation • Reduction • O.S. of some element decreases in the reaction. • Electrons are on the left of the equation.

45. 5-5 Balancing Oxidation-Reduction Equations • Few can be balanced by inspection. • Systematic approach required.

46. Write and balance separate half-equations for oxidation and reduction. Adjust coefficients in the two half-equations so that the same number of electrons appear in each half-equation. Add together the two half-equations (canceling out electrons) to obtain the balanced overall equation The Half-Equation Method