1. Reactions in Aqueous Solution
Chapter 4
Reactions in Aqueous Solution
熊同銘

2. Contents
General Properties of Aqueous Solutions
Precipitation Reactions
Acids, Bases, and Neutralization Reactions
Oxidation–Reduction Reactions
Concentrations of Solutions
Solution Stoichiometry and Chemical Analysis

3. 1. General Properties of Aqueous Solutions
Solute: The component in solution other than solvent.
Solvent: The greatest amount component of a solution.
Solution: A homogenous mixture of solute(s) and solvent.
Aqueous solutions: Water is the solvent, and a solid, liquid, or gas is the solute.

4. Electrolytes and Nonelectrolytes
Electrolyte: A substance that provides ions in aqueous solution and this solution conduct electricity.
Nonelectrolyte: A substance that is non-ionized when dissolved in water and this solution does not conduct electricity.

5. How Compounds Dissolve in Water
Ionic compounds and a few molecular compounds dissolves in water caused by hydration.
Non-ionizable molecular substances that dissolve in water caused by mixing.

6. Continued
Solvation: The clustering of solvent molecules around a solute particle which helps stabilize the ions in solution.
Hydration: A solvation when the solvent is water.

7. Strong and Weak Electrolytes
Weak electrolyte: A solute that is only partially ionized in aqueous solution. This solution is a poor conductor of electricity. CH3COOH for example:
*double arrow ( ) indicate that the dissociation does not go to completion (the process is reversible).
Strong electrolyte: A solute that is completely ionized in aqueous solution. This solution is a good conductor of electricity. HCl for example:

8. CH3OH(aq)
MgCl2(aq)
CH3COOH(aq)
Nonelectrolyte
Strong electrolyte
Weak electrolyte

9. Electrolytic Properties of Water-soluble Solutes
Nonelectrolytes include:
Most molecular compounds
Most organic compounds
Weak electrolytes include:
Weak acids (e.g., carboxylic acids, R–COOH) and weak bases (e.g., amines, R–NH2)
A few ionic compounds
Strong electrolytes include:
Strong acids (HCl, HBr, HI, HNO3, H2SO4, HClO4)
Strong bases (Group 1 and 2 hydroxides)
Most water-soluble ionic compounds

10. 2. Precipitation Reactions
Some definition
Precipitation reactions: The reaction of forming insoluble precipitate in a solution
Precipitate: The insoluble solid that deposits from a solution as a result of a precipitation reaction.
Solubility: The concentration of substance in its saturated solution.
Insoluble solid: For the solid with solubility < 0.01 mol/L

12. Solubility Guidelines for Ionic Compounds
common ionic compounds that contain the nitrate anion (NO3-) are soluble.
common ionic compounds of the alkali metal ions (group 1A) and of the ammonium ion (NH4+) are soluble.

13. Eexchange (metathesis) reaction: A reaction in which two substances react through an exchange of their component ions:
AX + BY AY + BX.
* Precipitation and acid–base neutralization reactions are examples of exchange reactions.

14. Eexchange (metathesis) reaction: A reaction in which two substances react through an exchange of their component ions:
AX + BY AY + BX.
* Precipitation and acid–base neutralization reactions are examples of exchange reactions.

15. Symbolic Equations for Precipitation Reaction
Molecular (total) equation:
Ag(NO3)(aq) + NaI(aq) → AgI(s) + NaNO3(aq)
Complete ionic equation:
Ag+(aq) + NO3–(aq) + Na+(aq) + I–(aq)
→ AgI(s) + Na+(aq) + NO3–(aq)
*Na+(aq) and NO3–, called spectator ions
Net ionic equation:
Ag+(aq) + I–(aq) → AgI(s)
spectator ions: 觀眾離子

16. Acids, Bases, and Neutralization Reactions Acids and Bases
Arrhenius:
Acid: A substance that provides H+ in aqueous solution, e.g., HCl
Base: A substance that provides OH– in aqueous solution, e.g., NaOH
Brønsted–Lowry
Acid: A proton donor
Base: A proton acceptor
Acid
Base

18. Common Strong Acids and Bases
* Hydrofluoric acid (HF) is weak acid. *

19. Neutralization Reactions and Salts
Neutralization reaction: A reaction in which
an acid and a base react in stoichiometrically
equivalent amounts; the neutralization reaction
between an acid and a metal hydroxide produces
water and a salt (ionic compound).

20. Neutralization between HCl and KOH
Molecular (total) equation
HCl(aq) + NaOH(aq)  H2O(l) + NaCl(aq)
Complete ionic equation
H+(aq) + Cl–(aq) + Na+(aq) + OH–(aq)
 H2O(l) + Na+(aq) + Cl–(aq)
Net ionic equation
H+(aq) + OH–(aq)  H2O(l)
* In this example, Na+ and Cl– are spectator ions

21. Neutralization Reactions with Gas Formation
Some acid-base reactions form gases as products
Examples:
* NH4+(aq) + OH–(aq)  NH3(g) + H2O(l)

22. 4. Oxidation-Reduction (Redox) Reactions
A chemical reaction contains the electron transfer from one reactant to another, in which the oxidation nunbers of certain atoms change. For examples: corrosion, combustion, and respiration.
Oxidation and Reduction
Oxidation: A process in which a substance loses one or more electrons and the oxidation number of certain atoms increases.
Reduction: A process in which a substance gains one or more electrons and the oxidation number of certain atoms decreases.
* Whenever one substance is oxidized, another substance must be reduced.

23. Continued
Past definition
Oxidation: a substance gains O atoms
Reduction: a substance loses O atoms
Oxidant and Reductant
Oxidant (oxidizing agent): Causes another substance to be oxidized while itself is reduced.
Reductant (reducing agent): Causes another substance to be reduced while itself is oxidized.

24. Familiar corrosion products
(a) A green coating forms (CuCO3 and Cu(OH)2) when copper is oxidized.
(b) Rust forms (Fe2O3 and Fe(OH)3) when iron corrodes
(c) A black tarnish forms (AgS) as silver corrodes.

25. Oxidation of calcium metal by molecular oxygen

26. Oxidation Numbers
A positive or negative whole number assigned to an element in a molecule or ion on the basis of a set of formal rules.
The number of actual charges of the atom on a monoatomic ion or ionic compound. or
The number of apparent charge of the atom (depend on electronegativity) in molecule.

27. Rules for Assigning Oxidation Numbers
For an atom in its elemental form, the oxidation number is always zero.
For a monatomic ion the oxidation number equals the ionic charge.
Metals:
Alkali metal (group 1A) is always +1
Alkaline earth metals (group 2A) is always +2
aluminum (Al, group 3A) is always +3

28. Continued
Nonmetals:
In binary compounds with metals, usually:
group 7A have an oxidation number of –1
group 6A have an oxidation number of –2
group 5A have an oxidation number of –3
The oxidation number of oxygen is usually 2
Exception:
Peroxide, example: H2O2, the oxidation number of oxygen is –1
Superoxide, example: KO2, the oxidation number of oxygen is –1/2.

29. The oxidation number of hydrogen:
usually +1 when bonded to nonmetals
usually 1 when bonded to metals (for example, sodium hydride (NaH))
The oxidation number of fluorine is always 1.
Other halogens (Cl, Br, I) usually have an oxidation number of -1.
Exception: When combined with oxygen (oxyanions), those halogens have positive oxidation numbers.

30. Continued
The sum of the oxidation numbers of all atoms:
In a neutral compound, the sum of all the oxidation numbers is 0.
In a polyatomic ion, the sum of the oxidation numbers is equal to the charge on the ion.
Other elements’ oxidation numbers are determined when they form compounds, including the rules are listed above.

31. Sample Exercise 4.8 Determining Oxidation Numbers
Determine the oxidation number of sulfur in (a) H2S, (b) S8, (c) SCl2, (d) Na2SO3, (e) SO42–.
Solution
(a) 2(+1) + x = 0, x= –2, S has an oxidation number of –2.
(b) S8 is an elemental form, the oxidation number of S is 0.
(c) x + 2(–1) = 0, x= +2, the oxidation number of S is +2.
(d) 2(+1) + x + 3(–2) = 0, x=+4, the oxidation number of S in compound (Na2SO3) is +4.
(e) x + 4 (–2) = –2, x=+6, the oxidation number of S in this ion is +6.

32. Activity Series of Metals
A list of metals in order of decreasing ease of oxidation. Any metal on the list can be oxidized by the ions of elements below it.
* Elements higher on the activity series are more reactive. They are more likely to exist as ions.
Examples:
Mg(s) + Cu2+(aq)  Mg2+(aq) + Cuo(s)
Ag(s) + Cu2+(aq)  No reaction
Zn(s) + HCl(aq)  ZnCl2(aq) + H2(g)
Ag(s) + H+(aq)  No reaction

33. Displacement Reactions
A reaction in which an element reacts with a compound, displacing an element from it.
Oxidation of metals by acids, for example:
Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g) (molecular equation)
Mg(s) + 2H+ (aq)  Mg2+(aq) + H2(g) (net ionic equation)
Oxidation of metals by salts, for example:
Fe(s) + Ni(NO3)2(aq)  Fe(NO3)2(aq) + Ni(s) (molecular equation)
Fe(s) + Ni2+(aq)  Fe2+(aq) + Ni(s) (net ionic equation)
A salt metathesis reaction sometimes called double displacement reaction:
Metathesis reaction: AB + CD = AC + BD
Displacement reaction: AB + C = AC + B

34. Note
Only metals above hydrogen in the activity series are able to react with acids to form H2. for example:
Ni(s) + 2HCl(aq)  NiCl2(aq) + H2(g)
The elements below hydrogen in the activity series are not oxidized by H+, for example Cu does not react with HCl(aq).
However, copper react with nitric acid:
Cu(s) + 4HNO3(aq) 
Cu(NO3)2(aq) + 2H2O(l) + 2NO2(g)

35. Sample Exercise 4.9 Writing Equations for Oxidation-Reduction Reactions
Write the balanced molecular and net ionic equations for the reaction of aluminum with hydrobromic acid.
Solution
Al is above H2 in activity series, the reaction occurs.
molecular equation:
complete ionic equation is
Br– is a spectator ion, the net ionic equation:

36. Sample Exercise 4.10 Determining When an Oxidation-Reduction Reaction Can Occur
Will an aqueous solution of iron(II) chloride oxidize magnesium metal? If so, write the balanced molecular and net ionic equations for the reaction.
Solution
Mg is above Fe in the activity series, the reaction occurs.
the balanced molecular equation:
Both FeCl2 and MgCl2 are soluble strong electrolytes and can be written in ionic form, which shows us that Cl– is a spectator ion in the reaction.
The net ionic equation:

37. Concentrations of Solutions
Molarity
Advantages of using molarity:
Substances enter into chemical reactions according to certain molar ratios.
Volumes of solutions are more convenient to measure than their masses.

38. Concentrations of Solutions
Molarity
Advantages of using molarity:
Substances enter into chemical reactions according to certain molar ratios.
Volumes of solutions are more convenient to measure than their masses.

39. Preparing 0.250 L of a 1.00 M solution of CuSO4.

40. Sample Exercise 4.11 Calculating Molarity
Calculate the molarity of a solution made by dissolving 23.4 g of sodium sulfate (Na2SO4) in enough water to form 125 mL of solution.
Solution

41. Expressing the Concentration of an Electrolyte
When ionic compound dissolved in water:
1.0 M NaCl: [Na+] = 1.0 M and [Cl-] = 1.0 M
1.0 M Na2SO4: [Na+] = 2.0 M and [SO42-] = 1.0 M

42. Sample Exercise 4.12 Calculating Molar Concentrations of Ions
What is the molar concentration of each ion present in a M aqueous solution of calcium nitrate?
Solution
0.025 M in Ca(NO3)2:
0.025 M in Ca2+ 2 × M = M in NO3–

43. Sample Exercise 4.13 Using Molarity to Calculate Grams of Solute
How many grams of Na2SO4 are required to make L of M Na2SO4?
Solution

44. Mconc x Vconc = Mdil x Vdil
Dilution
The process of preparing a less concentrated solution from a more concentrated one by adding solvent.
moles solute in conc soln = moles solute in dilute soln
Mconc x Vconc = Mdil x Vdil
* Solutions used routinely in the laboratory are often purchased or prepared in concentrated form, called stock solutions.

45. Preparing 250.0 mL of 0.100 M CuSO4 by dilution of 1.00 M CuSO4

46. Sample Exercise 4.14 Preparing a Solution by Dilution
How many milliliters of 3.0 M H2SO4 are needed to make 450 mL of 0.10 M H2SO4?
Solution

47. 6. Solution Stoichiometry and Chemical Analysis
Procedure for solving stoichiometry problems involving reactions between a pure substance A and a solution containing a known concentration of substance B.

48. Sample Exercise 4.15 Using Mass Relations in a Neutralization Reaction
How many grams of Ca(OH)2 are needed to neutralize 25.0 mL of M HNO3?
Solution

49. Titrations
A procedure for carrying out a chemical reaction between two solutions by the controlled addition (from a buret) of one solution to the other.
Standard solution A solution of known concentration.
Equivalence point: The condition for a titration in which the reactants are in stoichiometric proportions.
End point: The point in a titration where the indicator used changes color.
Indicator: An added substance that changes color at the end point in a titration.

50. Procedure for titrating an acid against a standard solution of NaOH.

51. Procedure for determining the concentration of a solution from titration with a standard solution.

52. Sample Exercise 4.16 Determining Solution Concentration by an Acid–Base Titration
One commercial method used to peel potatoes is to soak them in a NaOH solution for a short time and then remove the potatoes and spray off the peel. The NaOH concentration is normally 3 to 6 M, and the solution must be analyzed periodically. In one such analysis, 45.7 mL of M H2SO4 is required to neutralize 20.0 mL of NaOH solution. What is the concentration of the NaOH solution?
Solution

53. End of Chapter 04