Chapter 5 Reactions in Aqueous Solution

1. Chapter 5 Reactions in Aqueous Solution Classes of some chemical reactions in solution Combustion reactions Gas forming reactions Dissolution reactions Precipitation reactions Acid-base reactions Oxidation-reduction reactions

2. CombustionReactions Burning in the presence of oxygen is combustion Ex) Combustion of propane In complete combustion all C in the reagents becomes CO2, all H becomes H2O, all N becomes N2 Remaining elements usually combine with oxygen to give the element oxide In incomplete combustion, C compounds can produce CO or even elemental C The formation of soot is the result of incomplete combustion

3. Gas Forming Reactions Gas can be produced Ex) The production of carbon dioxide when washing soda and vinegar are mixed Ex) Decomposition of Hydrogen peroxide Gases also be consumed. Ex) Formation of limestone from lime Ex) Formation of hydrochloric acid

4. Note : The inorganic gas produced forms acids when combined with water SO2(g) + H2O (l)→ H2SO3(aq) CO2 (g) + H2O (l)→ H2CO3 (aq) Sulfurous acid Carbonic acid

5. Solubility The amount of a substance that can combine with another to give a single phase. The major component of a single-phase mixture is called the solvent, while the minor component is called the solute It is governed by the polarity of the two substances involved. “like dissolves like” When the interactions between the solute and solvent molecules are similar to those between solvent molecules and between solute molecules they will mix. A polar solvent will dissolve polar molecules Ex) Salt or sugar in Water A non-polar solvent will dissolve non-polar molecules Ex) Spices in cooking oil

6. Solubility In water these interactions cause water molecules to solvate, i.e. surround, the solute molecule or ion. The partial negative charge on oxygen in water will interact with cations The partial positive charge on hydrogen in water will interact with anions When two substances mix, and dissolve into each other they are referred to as miscible When two substance don’t have compatible polarities they will not mix. They are referred to as immiscible Ex) Oil and water Interactions between the different molecules will be weaker than between the same molecules for at least one component.

7. Solubility The solubility of a substance is usually expressed in the number of grams that dissolves in 100 ml of the solvent at a specified temperature A dissolution reaction can be represented by a chemical equation Dissolution can be accompanied by dissociation, either partial or complete: Solubility is expressed using its solubility product (Ksp): AnBm (s) + H2O(l)→ n Am+(aq)+ m Bn-(aq) Ksp = [Am+]n[Bn-]m Small Ksp → insoluble Large Ksp → Soluble Generally Ksp below 10-6 indicates that a salt is insoluble in water

8. Electrolyte Any compound that generates ions when dissolved in water. strong electrolyte is one which dissociates completely into cation(s) and anion(s) when dissolved in water. Large Ksp Ex) Ionic Salts AnBm(s) + H2O (l)→ nAm+ (aq) + mBn- (aq) Table salt NaCl (s) + H2O (l)→ Na+(aq) + Cl-(aq) weak electrolyte is one which dissociates only partially into cation(s) and anion(s) when dissolved in water. Small “Ksp” Ex) Ammonia NH3(g) + H2O (l)→ NH4+(aq) + OH-(aq) Ex) Vinegar CH3COOH (l) + H2O(l) → CH3COO-(aq) + H3O+(aq) Ksp ≈ 0 The undissociated molecules of compound remain neutral. nonelectrolyteis a compound that dissolves in water but does not dissociate into ions.

9. Solubility of Ionic Compounds Alkaline (gr. 1) and ammonium salts are soluble Nitrates, acetates, chlorates and perchlorates are soluble Most Halides are soluble Hydroxides, phosphates, Sulfites and carbonates tend to be insoluble except for their alkali salts Note that these anions occur in minerals (rocks). The higher the charge of ion the less likely its soluble Which of the following ionic compounds are likely to be soluble in water? (a) LiNO3 (b) Cu3PO4 (c) MgCO3 (d) NiSO4 (e) (NH4)2CO3 (f) Fe(ClO3)3 N Y N Y Y Y

10. Precipitation Reactions Compounds that exceed their solubility commonly precipitate from solution The product of a reaction between two soluble salts can be insoluble, resulting in the product to precipitate from solution. Ionicequation Spectator ions, K+ and Cl-, are ignored giving the Net Ionic Equation

11. Exercise Write a net ionic equation for each of the following reactions, and identify the spectator ions: (a) Mixing potassium hydroxide & copper(II) sulfate solutions 2 KOH (aq) + CuSO4(aq)→ Cu(OH)2(s) + K2SO4(aq) X 2 K+(aq) + 2 OH_(aq) X X X + Cu2+(aq) + SO42-(aq) → Cu(OH)2(s) + 2 K+(aq) +SO42-(aq) 2 OH_(aq) + Cu2+(aq) → Cu(OH)2(s) (b) Mixing ammonium phosphate & barium hydroxide solutions (NH4)3PO4 (aq) + Ba(OH)2 (aq) → NH4OH (aq) + Ba3(PO4)2 (s) 6 2 3 6 NH4+(aq) + 2 PO43-(aq) X X Ba3(PO4)2(s) X + 6 NH4+ (aq) + 6 OH-(aq) + 3 Ba2+(aq) + 6 OH_(aq)→ X 2 PO43- (aq) + 3 Ba2+ (aq) → Ba3(PO4)2 (s) (c) Mixing silver fluoride & magnesium iodide solutions AgF (s) + MgI2(aq)→ AgI (s) + MgF2(aq) 2 2 X X 2 Ag+(aq) + 2 F_(aq) + Mg2+(aq) + 2 I_(aq)→ X 2 AgI (s) + Mg2+(aq) + 2 F_(aq) X Ag+(aq) + I_(aq)→ AgI (s)

12. Acids and Bases Three major ways to define acids and bases introduced by Lewis, Brønsted and Arrhenius. They differ in the role of water Arrhenius and Brønsted require water, Lewis does not Acid Donates an H+ Brønsted Accepts an H+ Base Acid HCl + H2O → H3O+ + Cl- Base NaOH + H+→ Na+ + H2O

13. Acids and Bases Acid Produces H3O+ when added to water Arrhenius Produces OH- when added to water Base HCl + H2O → H3O+ + Cl- NH3+ H2O → NH4+ + OH-

14. : H H H H H B H H B N N H H H H H Acids and Bases Acid Accepts electrons Lewis Note: Electrons are not transferred between acids and bases, they are shared. Base Donates electrons BH3 + NH3 →BH3NH3 Base Acid

15. Acids and Bases The Lewis definition is the most general H+ leaves electrons behind Consider a Brønsted Acid It donates a H+ A-H i.e.A accepts the electrons A : _ H+ A is a Lewis Acid All Brønsted acids are Lewis Acids An Arrhenius acid, is a Bronsted Acid, since it produces H3O+ when dissolved in water as it “donates” H+ to H2O.

16. Acids and Bases A strong acid, just like a strong electrolyte, is an acid which dissociates completely when dissolved in water. The concentration of H3O+ is thereby the highest possible, determined exactly by how much acid was added to water Ex) HCl + H2O → H3O+ + Cl- Ex) H2SO4 + H2O → H3O+ + HSO4_ Inorganic acids tend to be strong acids (except HF) A strong base, just like a strong electrolyte, is dissociates completely when dissolved in water. The concentration of OH_ is thereby the highest possible, determined exactly by how much acid was added to water. Ex) NaOH → Na+ + OH_ The hydroxides of alkali metals are strong bases.

17. Acids and Bases A weak acid, just like a weak electrolyte, does not dissociate completely when dissolved in water The concentration of H3O+ is not the highest possible, since much remains in the undissociated form. The concentration of H3O+ is determined from the dissociation constant, similar to Ksp, and the amount of acid added. Ex) CH3COOH (l) + H2O(l) → CH3COO-(aq) + H3O+(aq) Organic acid tend to be weak acids A weak base, just like a weak electrolyte, does not dissociate completely when dissolved in water. The concentration of OH- is not the highest possible, since much remains in the undissociated form. The concentration of OH- is determined from the dissociation constant, similar to Ksp, and the amount of acid added. Ex) NH3+ H2O → NH4+ + OH- Metal Oxides and nitrogen containing organic compounds tend to be weak bases

19. The oxides are anhydrides Non-metal oxides react with water to give oxoacids Therefore non-metal oxides are anhydrides of acids Metal oxides react with water to give hydroxide bases Therefore metal oxides are anhydrides of bases Metalloid oxides are amphoteric: they react with either strong acids or strong bases Amphotericoxides usually do not dissolve with water by themselves, but react with both strong acids and strong bases to give soluble products

20. Metal Metalliod Non-metal 1 2 13 15 16 14 17 The oxides are anhydrides Strength of acids and bases is correlated with the positions of the oxides on the PT

21. pH The acidity (or basicity) of a solution is reported as pH: pH = -log [H3O+] or [H3O+] = 10-pH = concentration of H3O+ in mol./l = molar (M) p = power of For pH < 7 solution is acidic For pH > 7 solution is basic Ex) 0.10 M solution of HCl [H3O+] = 0.10 M pH = - log [0.10] = -(-1.00) = 1.00 # of sig. figs. Increased from 2 to 3? For logarithmic quantities only the decimal numbers are significant. Therefore a pH = 1.00 has only 2 sig. figs, Note: pH does not have units

22. pOH Basicity of a solution can be reported as pOH: pOH = -log [OH-] Where [OH-] = conc. of OH in mol/l pH and pOH are related by: pH + pOH = 14 at 25 oC Therefore an acidic solution as pOH > 7, and a basic solution has pOH < 7 Determine the pH and pOH of: Exercise a) 0.275 M HNO3 solution [H3O+] = 0.275 M pH = - log (0.275) = -(0.561) = 0.561 pOH = 14.000 – pH = 14.000 -0.561 = 13.439 b) 0.0051 M NaOH solution [OH-] = 0.0051 M pOH = - log (0.0051) = -(-2.29) = 2.29 pH = 14.000 – pOH = 14.000 -2.29 = 11.71

23. Reactivity A reaction between an acid and a base produces water and a salt HCl(aq) + NaOH(aq)→ H2O (l) + NaCl (aq) Acid Base Water Salt Ionic Equation H3O+(aq) + Cl-(aq) + Na+(aq) + OH_(aq) → H2O(l) + Na+(aq) + Cl-(aq) Net equation H3O+(aq) + OH_(aq)→ H2O(l) A strong acid will react completely with any base. A strong base will react completely with any acid.

24. Reactivity A reaction between a weak acid and a weak base will not go to completion unless there is a driving force (e.g. making a gas): Ex) CH3COOH (aq) + NH3 (aq) → CH3COONH4(aq) + H2O(l) Write an ionic equation for each of the following acid-base reactions: (a) KOH(aq) + H2SO4(aq) → (b) H3PO4(aq) + LiOH(aq) →

25. Reduction-Oxidation (REDOX) Reactions In chemical reaction bonds, both covalent and ionic, are made and broken by moving electrons. For all reaction dealt with to date the number of electrons on each atom has been preserved 1+ 2- 0 0 Loss e’s Oxidation = LEO 2 Na + O2→ 2 NaO 2 Na+ + O2- Gain e’s Reduction = GER Na → Na+ + 1 e ½ O2 + 2 e→ O2- Oxidized Reduced In REDOX reaction electrons are transferred between atoms as well has bonds being broken and formed. The balance between ionic/covalency of the bonds in the reactants is different from that in the product. We therefore need to keep track of the number of electrons of each atom. The oxidation state of an atom is used for this purpose, which is related to the idea of the formal charge.

26. Oxidation State The oxidation state of an element is its charge assuming ionic bonding Electrons are not shared they are placed on the more electronegative element Rules for Assigning Oxidation States Pure elements have oxidation states of 0 Ions have oxidation states that add up to the charge of the ion Hydrogen has an oxidation state of +1 unless bonded to a less electronegative atom. When bound to metals or boron it has an oxidation state of -1. Fluorine has an oxidation state of -1 Oxygen has an oxidation state of -2 unless bonded to fluorine or another oxygen. Halogens other than fluorine have oxidation states of -1 unless bonded to oxygen or a more electronegative halogen The rest are determined by the process of elimination, where the oxidation states must add up to the total charge of the molecule or ion.

27. Oxidation State Determine the oxidation state of all the element in the following molecules: a) H2 H = 0 C = 4 O = -2 0 = C + 2(-2) b) CO2 B = 3 0 = B + 3(-1) c) BF3 F= -1 S = 6 H = 1 O = -2 0 = S +(2(+1) + 4(-2)) = S - 6 d) H2SO4 O = -2 e) SO2ClF F = -1 Cl = -1 0 = S +(2(-2) -1 -1) = S - 6 S = 6 f) IO2F2_ F = -1 -1 = I + (2(-2) + 2(-1)) = I - 6 O = -2 I = 5 O = -2 H = 1 g) HPO42- -2 = P + (4(-2) + 1) = P - 7 P = 5

28. Oxidation States of Poly Atomic Ions CO32- NO3- 5 4 NO2- 3 SO42- ClO3- PO43- 5 6 5 PO33- ClO2- SO32- 4 3 3

29. Recognizing Redox Reactions Consider the following reaction: Reactants Products Cu = 0 Cu2+ = 2 HNO3 N = 5 N2O4 N = 4 Cu → Cu2+ + 2 e Oxidation 4 HNO3 + 2 e → 2NO3- +N2O4 + 2 H2O Reduction REDOX equation seem difficult to balance # of e’s gained = # of e’s lost

30. Recognizing Redox Reactions The species that is oxidized in a redox reaction is called the reducing agent In the previous example Cu is the reducing agent The species that is reduced in a redox reaction is called the oxidizing agent In the previous example HNO3is the oxidizing agent

31. Concepts from Chapter 5 Electrolytes Solubility Role of Polarity Ksp Miscible vs. Immiscible Reaction Types Combustion Reactions Gas Forming Reactions Precipitation Reactions Acid-Base Reactions Redox Reactions Net Ionic Equations Lewis vs. Brönsted vs. Arrhenius Definitions of Acids/Bases Strong Acids/Bases vs. Weak Acids/Bases Acidity/Basicity of Oxides pH and pOH Assigning Oxidation States