1. Chapter 4 Reactions in Aqueous Solution

2. Solutions
Solutions are defined as homogeneous mixtures of two or more pure substances.
The solvent is present in greatest abundance.
All other substances are solutes.
When water is the solvent, the solution is called an aqueous solution.

3. Aqueous Solutions Substances can dissolve in water by different ways:
Ionic compounds dissolve by dissociation, where water surrounds the separated ions.
Molecular compounds interact with water, but most do NOT dissociate.
Some molecular substances react with water when they dissolve.
All substances dissolve by solvation, surrounding of the solute by solvent.

4. Electrolytes and Nonelectrolytes
An electrolyte is a substance that dissociates into ions when dissolved in water. (eg. NaCl in water)
A nonelectrolyte may dissolve in water, but it does not dissociate into ions when it does so. (eg. sucrose in water)
Soluble ionic compounds tend to be electrolytes.
Molecular compounds tend to be nonelectrolytes, except for acids and bases.

5. Electrolytes
A strong electrolyte dissociates completely when dissolved in water.
A weak electrolyte only dissociates partially when dissolved in water.
A nonelectrolyte does NOT dissociate in water. (tend to be molecular compounds)

6. Strong vs. Weak Electrolytes—Equilibrium
A strong electrolyte dissociates completely when dissolved in water. Its equations for the reaction of its dissociation look familiar:
HCl (aq) H+ (aq) + Cl– (aq)
A weak electrolyte only dissociates partially when dissolved in water. Its equations indicate a chemical equilibrium, where a reaction goes both forward and backward:
CH3COOH (aq) H+ (aq) + CH3COO– (aq)

7. Strong Electrolytes Are…
Strong acids
Strong bases
Soluble Ionic Salts

8. Electrolytes
Potassium Iodide, glucose, and hydrofluoric acid are examples of a _________, ________, and _________ electrolytes, respectively.
Weak, non-, strong
Strong, non-, weak
Strong, weak, strong
Strong, weak, weak
Weak, weak, and strong
Answer: B

9. Types of Aqueous Reactions
Precipitation (think Qualitative Analysis lab)
We will use the solubility rules to determine what will form a precipitate
Acid-Base Reactions
We will look at different definitions for acids and bases.
Oxidation Reduction (REDOX) Reactions
These are reactions in which electrons are lost and gained during the reaction itself.

10. Solubility of Ionic Compounds
Not all ionic compounds dissolve in water.
A list of solubility rules is used to decide what combination of ions will dissolve.
Alkali metal cations and NH4+ are usually soluble.

11. Using the Solubility Rules
Which of the following compounds are soluble in water?
PbCl2
Ag2S
Cs3PO4
SrCO3
PbSO4
Answer: D

12. Precipitation Reactions
Precipitation reactions occur when two solutions containing soluble salts are mixed and an insoluble salt is produced. The solid is called a precipitate.

13. How to Predict Whether a Precipitate Forms When Strong Electrolytes are Mixed
Note the ions present in the reactants.
Consider the possible cation-anion combinations.
Use Table 4.1 to determine if any of the combinations is insoluble.

14. Metathesis (Exchange) Reactions
Metathesis comes from a Greek word that means “to transpose.”
It appears as though the ions in the reactant compounds exchange, or transpose, ions, as seen in the given equation.
AX BY AY BX
AgNO3(aq) + KCl(aq)  AgCl(s) + KNO3(aq)

15. Completing and Balancing Metathesis Equations
Steps to follow:
Use the chemical formulas of the reactants to determine which ions are present.
Write formulas for the products: cation from one reactant, anion from the other. Use charges to write proper subscripts.
Check your solubility rules. If either product is insoluble, a precipitate forms.
Balance the equation.

16. Ways to Write Equations for Metathesis Reactions
Molecular equation
Complete ionic equation
Net ionic equation

17. Step 1: Molecular Equation
Begin with the molecular equation, which lists the reactants and products in their molecular form.
AgNO3(aq) + KCl(aq)  AgCl(s) + KNO3(aq)

18. Ag+(aq) + NO3−(aq) + K+(aq) + Cl−(aq)  AgCl(s) + K+(aq) + NO3−(aq)
Step 2: Ionic Equation
In the ionic equation all strong electrolytes (strong acids, strong bases, and soluble ionic salts) are dissociated into their ions.
This more accurately reflects the species that are found in the reaction mixture.
Recall the definitions and examples of strong, weak, and non-electrolytes.
Ag+(aq) + NO3−(aq) + K+(aq) + Cl−(aq)  AgCl(s) + K+(aq) + NO3−(aq)

19. Step 3: Net Ionic Equation
To form the net ionic equation, cross out anything that does not change from the left side of the equation to the right:
Ag+(aq) + NO3−(aq) + K+(aq) + Cl−(aq)  AgCl(s) + K+(aq) + NO3−(aq)
The only things left in the equation are those things that change (i.e., react) during the course of the reaction:
Ag+(aq) + Cl−(aq)  AgCl(s)

20. Ag+(aq) + NO3−(aq) + K+(aq) + Cl−(aq)  AgCl(s) + K+(aq) + NO3−(aq)
Spectator Ions
Those things that didn’t change (and were deleted from the net ionic equation) are called spectator ions:
Ag+(aq) + NO3−(aq) + K+(aq) + Cl−(aq)  AgCl(s) + K+(aq) + NO3−(aq)

21. Writing Net Ionic Equations Stepwise
Write a balanced molecular equation.
Dissociate all strong electrolytes.
Cross out anything that remains unchanged from the left side to the right side of the equation (spectator ions)
Write the net ionic equation with the species that remain.

22. Examples - Precipitation Reactions
Write the net ionic equation for any precipitation reaction when solutions of the following ionic compounds are mixed:
a. KCl and Ag2SO4
b. NaOH and Cu(NO3)2

23. Writing Net Ionic Equations - Example 2
Write a net ionic equation for each of the following reactions in dilute water solution
Nitrous acid with sodium hydroxide

24. Writing Net Ionic Equations - Example 3
Write a net ionic equation for each of the following reactions in dilute water solution
Hydrobromic acid with potassium hydroxide

25. Solubility Rules *Memorize this*
Alkali metal cations and NH4+ are usually soluble.

26. Acids Acids are substances that ionize in aqueous solution to form H+
Because H+ consists of ONLY a proton, acids are often called proton donors.

27. Bases
Bases are substances that react with, or accept, H+ ions; they increase the concentration of OH– when dissolved in water.
Substances do NOT have to contain OH– to be a base.

28. Strong or Weak?
Strong acids completely dissociate in water; weak acids only partially dissociate.
There are only seven strong acids:
Hydrochloric (HCl)
Hydrobromic (HBr)
Hydroiodic (HI)
Nitric (HNO3)
Sulfuric (H2SO4)
Chloric (HClO3)
Perchloric (HClO4)

29. Strong or Weak?
Strong bases dissociate to metal cations and hydroxide anions in water; weak bases only partially react to produce hydroxide anions.
The strong bases are the soluble metal salts of hydroxide ion (OH-):
Alkali metals
Calcium
Strontium
Barium

30. Strong or Weak?
Memorize these!

31. Polyprotic Acids
Acids can form different numbers of H+ ions per molecule.
Monoprotic – produce 1 H+ ion per acid molecule
Ex. HCl, HNO3
Diprotic - produce 2 H+ ions per acid molecule
Ex. H2SO4
Triprotic - produce 3 H+ ions per acid molecule
Ex. H3PO4

32. Strong or Weak Electrolyte?
Is the substance ionic or molecular? If it is ionic, it is a strong electrolyte. (Solubility?)
If molecular: is it an acid or a base?
If it starts with H or ends in COOH, it is an acid. If it is NOT on the list of strong acids, it is a weak acid. (Strong acid = strong electrolyte; weak acid = weak electrolyte.)
Strong bases (Table 4.2) are strong electrolytes; NH3 is a weak base (weak electrolyte).

33. Neutralization Reactions
Reactions between an acid and a base are called neutralization reactions.
When the base is a metal hydroxide, water and a salt (an ionic compound) are produced.
These equations can be written as molecular, complete ionic, or net ionic equations.

34. Neutralization Reactions with Gas Formation
Some metathesis reactions do not give the product expected.
When a carbonate or bicarbonate reacts with an acid, the products are a salt, carbon dioxide, and water.
In these reactions, the expected product (H2CO3) decomposes to give a gaseous product (CO2)
CaCO3(s) + 2 HCl(aq) CaCl2(aq) + CO2(g) + H2O(l)
NaHCO3(aq) + HBr(aq) NaBr(aq) + CO2(g) + H2O(l)

35. Neutralization Reactions with Gas Formation
Similarly, when a sulfite reacts with an acid, the products are a salt, sulfur dioxide, and water:
Example: SrSO3(s) + 2HI(aq) SrI2(aq) + SO2(g) + H2O(l) A rather smelly example:
Na2S(aq) + H2SO4(aq)  Na2SO4(aq) + H2S(g)

36. Application of Neutralization Reactions: Antacids

37. Oxidation-Reduction Reactions
Loss of electrons is oxidation.
Gain of electrons is reduction.
One cannot occur without the other.
The reactions are often called redox reactions.

38. Oxidation Numbers
To determine if an oxidation-reduction reaction has occurred, we assign an oxidation number to each element in a neutral compound or charged entity.
This is a “bookkeeping” method—it does NOT imply that the atoms have these charges!

39. Rules to Assign Oxidation Numbers
Atoms in their elemental form have an oxidation number of zero. (eg. Cl2)
The oxidation number of a monatomic ion is the same as its charge. (eg. H+, Cl-)

40. Rules to Assign Oxidation Numbers
Nonmetals usually have negative oxidation numbers, although they sometimes can be positive:
Oxygen: usually −2, except in the peroxide ion, where it is −1. (eg. H2O, H2O2)
Hydrogen: usually +1 when bonded to a nonmetal and −1 when bonded to a metal. (eg. H2O, LiH)
Fluorine: –1. Other halogens: usually –1, unless combined with oxygen (oxyanions), where they will be positive.

41. Rules to Assign Oxidation Numbers
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 on the ion. (Remember to count EVERY atom, no matter how large the subscript, when assigning oxidation numbers!)

42. Redox Reactions
Zn (s) + 2H+ (aq)  Zn2+ (aq) + H2 (g) Half Reactions: Oxidation: Zn (s)  Zn2+ (aq) + 2e- Reduction: 2H+ (aq) + 2e-  H2 (g) - Oxidation and reduction must occur together - No net change in the number of electrons in a redox reaction

43. More Redox Reaction Terminology
Zn (s) + 2H+ (aq)  Zn2+ (aq) + H2 (g)
Oxidation: Zn (s)  Zn2+ (aq) + 2e-
Reduction: 2H+ (aq) + 2e-  H2 (g)
Reducing agent – species that donates the electrons (Zn)
Oxidizing agent – species that accepts the electrons (H+)

44. Example – Redox Reactions
What is the oxidation number of manganese in MnO4-?
Give the oxidation number of each atom in:
a. HSO4-
b. N2H4
What is the oxidation number of sulfur in Na2SO4?

45. Displacement Reactions
In displacement reactions, ions oxidize an element. (H+ oxidizes Mg below.)
The ion is displaced (replaced) in solution. (Mg replaces H+ below.)

46. Example - Displacement Reactions
In this reaction, silver ions oxidize copper metal:
Cu(s) + 2Ag+(aq)  Cu2+(aq) + 2Ag(s)

47. Example Continued - Displacement Reactions
The reverse reaction, however, does not occur:
Cu2+(aq) + 2Ag(s)  Cu(s) + 2Ag+(aq)\
Question: How can we tell which element is going to be the one that gets oxidized by ions of the other element?
Answer: The Activity Series x

48. Activity Series and Hydrogen
The elements above hydrogen will react with acids to produce hydrogen gas. Elements below will NOT react!
A reactive metal is oxidized to a cation.

49. Activity Series *Ions of a particular metal on the list can oxidize
any metal above it.
* Any metal on the list can be oxidized by ions of the element below it
*Notice the
placement of Cu
and Ag
placement of the
precious metals.
Precious metals include palladium, platinum, and gold

50. Metal/Acid Displacement Reactions
Elements higher on the activity series are more reactive.
They will exist as ions.
The element below will exist as the element.

51. Example
Will an aqueous solution of iron(II) chloride oxidize magnesium metal? If so, write the balanced molecular and net ionic equations for the reaction.
Mg(s) + FeCl2(aq) -> MgCl2(aq) + Fe(s)
Mg(s) + Fe2+(aq) -> Mg2+(aq) + Fe(s)

52. volume of solution in liters
Molarity
The quantity of solute in a solution can matter to a chemist.
We call the amount dissolved its concentration.
Molarity is one way to measure the concentration of a solution:
Molarity can be used as a conversion factor between moles and liters.
moles of solute
volume of solution in liters
Molarity (M) =

53. Mixing a Solution
To create a solution of a known molarity, weigh out a known mass (and, therefore, number of moles) of the solute.
Then add solute to a volumetric flask, and add solvent to the line on the neck of the flask.

54. Example Cu(NO3)2 (aq)  Cu2+(aq) + 2 NO3-(aq)
Suppose that 1 mol of cupric nitrate is dissolved in water to a final volume of 1 L.
What is the molarity of cupric nitrate?
1 L
M = = 1 M Cu(NO3)2
1 mol Cu(NO3)2
What is the molarity of nitrate ion?
M = = 2 M NO3-
1 mol Cu(NO3)2
1 L
2 mol NO3-

55. Dilution
A solution can be diluted by adding ONLY solvent. The concentration is LOWER, but the MOLES don’t change.

56. Dilution
The molarity of the new solution can be determined from the equation
Mc × Vc = Md × Vd,
where Mc = molarity of the concentrated solution
Vc = volume of the concentrated solution
Md = molarity of the dilute solution
Vd = volume of the dilute solution
[Remember: M × V (in L) = moles!]
Liters times molarity = moles

57. Example - Dilution Equation
A bottle labeled “concentrated HCl” in the lab contains 12.0 mol of HCl per liter of solution, so [HCl] = 12.0 M.
A. What would the concentration of a dilute HCl solution be if 15.0 mL of a 12.0 M stock solution were diluted to mL?
C. How many mL of a 12.0 M stock solution are needed to make 1.5 L of a 0.10 M solution?

58. Stoichiometry Applied to Solutions

59. Example - Molarities in Stoichiometry
Find the number of grams of Na2SO4 that are needed to make L of M Na2SO4.
What can I solve for given the information I have?
What am I being asked to find?
How am I going to get there?

60. Example
I want to prepare 0.10 L of 0.25 M NaOH. How many grams of NaOH do I need?
Molar Mass NaOH: g/mol

61. Example
I want to prepare 0.20 L of M NaOH. How many mL of a 0.25 M NaOH stock solution should I use for dilution?

62. Example Sodium carbonate reacts with HCl as follows:
Na2CO3 + 2 HCl  2 NaCl + H2O + CO2
How many liters of 0.15 M HCl are required to
completely react with 2.8 g of Na2CO3?
Molar Mass Na2CO3: g/mol

63. Titration
A titration is an analytical technique in which one can calculate the concentration of a solute in a solution.

64. Titration
A solution of known concentration, called a standard solution, is used to determine the unknown concentration of another solution.
The reaction is complete at the equivalence point, which is based on the seen end point (color change).

65. Indicators
Indicator: phenolphthalein

66. Types of Indicators Indicator Color HIn Color In- Endpoint pH
Methyl red
Red
Yellow
Bromothymol blue
Blue
6-7.6
Phenophthalein
No color
Pink
8.2-10
Bromocresol green
Blue-green

67. Example
In an acid-base titration, HCl is used to neutralize KOH based on the following equation:
HCl + KOH KCl+ H2O
What is the original molarity of a KOH solution if mL of M HCl is required to neutralize mL of KOH solution?