1. Chapter 4 Reactions in Aqueous Solution

2. A solution is a homogenous mixture of 2 or more substances
The solute is(are) the substance(s) present in the smaller amount(s)
The solvent is the substance present in the larger amount
Solution
Solvent
Solute
Soft drink (l)
H2O
Sugar, CO2
Air (g) N2
O2, Ar, CH4
Soft Solder (s) Pb Sn

3. Dissociation
When an ionic substance dissolves in water, the solvent pulls the individual ions from the crystal and solvates them.
This process is called dissociation.

4. Water’s power as an ionizing solvent results from the distribution of its electrons and its overall shape.
A polar molecule is one that has a + end and a - end. Water is one of the most polar molecules. d+ d-
H2O

5. Hydration is the process in which an ion is surrounded by water molecules arranged in a specific manner. d+ d-
H2O

6. Dissociation
An electrolyte is a substances that dissociates into ions when dissolved in water.

7. Solutions An electrolyte is
a substance that dissociates into ions when dissolved in water.
A nonelectrolyte may dissolve in water, but it does not dissociate into ions when it does so.

8. An electrolyte is a substance that, when dissolved in water, results in a solution that can conduct electricity.
A nonelectrolyte is a substance that, when dissolved, results in a solution that does not conduct electricity.
nonelectrolyte
weak electrolyte
strong electrolyte

9. Electrolytes and Nonelectrolytes
Soluble ionic compounds tend to be electrolytes.

10. Electrolytes and Nonelectrolytes
Molecular compounds tend to be nonelectrolytes, except for acids and bases.

11. Electrolytes
A strong electrolyte dissociates completely when dissolved in water.
A weak electrolyte only dissociates partially when dissolved in water.
Strong Electrolyte – 100% dissociation
NaCl (s) Na+ (aq) + Cl- (aq)
H2O
CH3COOH CH3COO- (aq) + H+ (aq)
Reversible reaction - Chemical equilibrium

12. Strong Electrolytes Are…
Strong acids: HCl, HBr, HI, HNO3, HClO4,
HClO3, H2SO4 .
Strong bases: Group 1A metal hydroxide
[LiOH, KOH, NaOH, RbOH, CsOH]
Heavy group 2A metal hydroxides
[Ca(OH)2, Sr(OH)2, Ba(OH)2]

13. Strong Electrolytes Are…
Strong acids
Strong bases
Soluble ionic salts

14. Precipitation Reactions
When one mixes ions that form compounds that are insoluble (as could be predicted by the solubility guidelines),
a precipitate is formed.

15. Metathesis (Exchange) Reactions
Metathesis comes from a Greek word that means “to transpose.”
AgNO3(aq) + KCl(aq)  AgCl(s) + KNO3(aq)

16. 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:
AgNO3(aq) + KCl(aq)  AgCl(s) + KNO3(aq)

17. Solution Chemistry
It is helpful to pay attention to exactly what species are present in a reaction mixture (i.e., solid, liquid, gas, aqueous solution).
If we are to understand reactivity, we must be aware of just what is changing during the course of a reaction.

18. Precipitation Reactions
Precipitate – insoluble solid that separates from solution
PbI2
precipitate
Pb(NO3)2 (aq) + 2NaI (aq) PbI2 (s) + 2NaNO3 (aq)
molecular equation
Pb2+ + 2NO3- + 2Na+ + 2I PbI2 (s) + 2Na+ + 2NO3-
ionic equation
Pb2+ + 2I PbI2 (s)
net ionic equation
Na+ and NO3- are spectator ions

20. Writing Net Ionic Equations
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.
Write the net ionic equation with the species that remain.

21. CaCl2(aq) + Na2CO3(aq) CaCO3(s) + 2 NaCl(aq)
According to the solubility guidelines in Table 4.1, CaCO3 is insoluble and NaCl is soluble. The balanced molecular equation is
CaCl2(aq) + Na2CO3(aq) CaCO3(s) + 2 NaCl(aq)
Thus, the complete ionic equation is
Ca2+(aq) + 2 Cl–(aq) + 2 Na+(aq) + CO32–(aq)
CaCO3(s) + 2 Na+(aq) + 2 Cl–(aq)
Cl– and Na+ are spectator ions. Canceling them gives the
following net ionic equation:
Ca2+(aq) + CO32–(aq) CaCO3(s)

22. Acids
The Swedish physicist and chemist S. A. Arrhenius defined acids as substances that increase the concentration of H+ when dissolved in water.
Both the Danish chemist J. N. Brønsted and the British chemist T. M. Lowry defined them as proton donors.

23. Acids Hydrochloric (HCl) Hydrobromic (HBr) Hydroiodic (HI)
There are only seven strong acids:
Hydrochloric (HCl)
Hydrobromic (HBr)
Hydroiodic (HI)
Nitric (HNO3)
Sulfuric (H2SO4)
Chloric (HClO3)
Perchloric (HClO4)

24. Acids that ionize to form one H+ ion are called monoprotic acids
Acids that ionize to form one H+ ion are called monoprotic acids. Common monoprotic acids include HCl, HNO3 and HC2H3O2. Acids that ionize to form two H+ ions are called diprotic acids. A common diprotic acid is H2SO4.

25. Bases
Arrhenius defined bases as substances that increase the concentration of OH− when dissolved in water.
Brønsted and Lowry defined them as proton acceptors.

26. Bases The strong bases are the soluble metal salts of hydroxide ion:
Alkali metals
Calcium
Strontium
Barium

27. Common bases are NaOH, KOH, and Ca(OH)2
Common bases are NaOH, KOH, and Ca(OH)2. Compounds that do not contain OH– ions can also be bases. Proton transfer to NH3 (a weak base) from water (a weak acid) is an example of an acid–base reaction. Since there is a mixture of NH3, H2O, NH4+, and OH– in solution, We write: NH3(aq) + H2O(l)  NH4+(aq) + OH–(aq)

28. Strong and Weak Acids and Bases
• Strong acids and strong bases are strong electrolytes.
• They are completely ionized in solution.
• Strong bases include: Group 1A metal hydroxides, Ca(OH)2, Ba(OH)2, and Sr(OH)2.
• Strong acids include: HCl, HBr, HI, HClO3, HClO4, H2SO4, and HNO3.
• We write the ionization of HCl as:
HCl  H+ + Cl–
• Weak acids and weak bases are weak electrolytes.
• They are partially ionized in aqueous solution.
• HF(aq) is a weak acid; most acids are weak acids.
• We write the ionization of HF as:
HF  H+ + F–

29. Acid-Base Reactions
In an acid–base reaction, the acid donates a proton (H+) to the base.

30. Neutralization Reactions
Generally, when solutions of an acid and a base are combined, the products are a salt and water:
CH3COOH(aq) + NaOH(aq) CH3COONa(aq) + H2O(l)

31. Neutralization Reactions
When a strong acid reacts with a strong base, the net ionic equation is
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)

32. Neutralization Reactions
When a strong acid reacts with a strong base, the net ionic equation is
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
H+(aq) + Cl−(aq) + Na+(aq) + OH−(aq) 
Na+(aq) + Cl−(aq) + H2O(l)

33. Neutralization Reactions
When a strong acid reacts with a strong base, the net ionic equation is
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
H+(aq) + Cl−(aq) + Na+(aq) + OH−(aq) 
Na+(aq) + Cl−(aq) + H2O(l)
H+(aq) + OH−(aq)  H2O(l)

34. The balanced molecular equation is:
The complete ionic equation is:
The net ionic equation is:

35. Gas-Forming Reactions
Some metathesis reactions do not give the product expected.
In this reaction, the expected product (H2CO3) decomposes to give a gaseous product (CO2):
CaCO3(s) + 2HCl(aq)  CaCl2(aq) + CO2(g) + H2O(l)

36. Gas-Forming Reactions
When a carbonate or bicarbonate reacts with an acid, the products are a salt, carbon dioxide, and water:
CaCO3(s) + 2HCl(aq)  CaCl2(aq) + CO2(g) + H2O(l)
NaHCO3(aq) + HBr(aq)  NaBr(aq) + CO2(g) + H2O(l)

37. Gas-Forming Reactions
Similarly, when a sulfite reacts with an acid, the products are a salt, sulfur dioxide, and water:
SrSO3(s) + 2HI(aq)  SrI2(aq) + SO2(g) + H2O(l)

38. Gas-Forming Reactions
This reaction gives the predicted product, but you had better carry it out in the hood, or you will be very unpopular!
But just as in the previous examples, a gas is formed as a product of this reaction:
Na2S(aq) + H2SO4(aq)  Na2SO4(aq) + H2S(g)

39. Oxidation-Reduction Reactions
An oxidation occurs when an atom or ion loses electrons.
A reduction occurs when an atom or ion gains electrons.
One cannot occur without the other.

40. 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.

41. Oxidation Numbers
Elements in their elemental form have an oxidation number of 0.
The oxidation number of a monatomic ion is the same as its charge.
Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions

42. Oxidation Numbers Fluorine always has an oxidation number of −1.
The other halogens have an oxidation number of −1 when they are negative; they can have positive oxidation numbers, however, most notably in oxyanions.
Hydrogen is −1 when bonded to a metal, +1 when bonded to a nonmetal.
Oxygen has an oxidation number of −2, except in the peroxide ion, in which it has an oxidation number of −1.

43. Oxidation Numbers
The sum of the oxidation numbers in a neutral compound is 0.
The sum of the oxidation numbers in a polyatomic ion is the charge on the ion.

44. Displacement Reactions
In displacement reactions, ions oxidize an element.
The ions, then, are reduced.

45. Displacement Reactions
In this reaction, silver ions oxidize copper metal:
Cu(s) + 2Ag+(aq)  Cu2+(aq) + 2Ag(s)
The reverse reaction, however, does not occur:
Cu2+(aq) + 2Ag(s)  Cu(s) + 2Ag+(aq) x

46. The molecular equation: 2 Al(s) + 6 HBr(aq) 2 AlBr3(aq) + 3 H2(g)
Both HBr and AlBr3 are soluble strong electrolytes.
Thus, the complete ionic equation is:
2 Al(s) + 6 H+(aq) + 6 Br–(aq) Al3+(aq) +
6 Br–(aq) + 3 H2(g)
Because Br– is a spectator ion, the net ionic equation is:
2 Al(s) + 6 H+(aq) Al3+(aq) + 3 H2(g)

47. Activity Series

48. The magnesium salt formed in the reaction is MgCl2, meaning
The balanced molecular equation is:
Mg(s) + FeCl2(aq) MgCl2(aq) + Fe(s)
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 is:
Mg(s) + Fe2+(aq) Mg2+(aq) + Fe(s)
The net ionic equation shows that Mg is oxidized and Fe2+ is
reduced in this reaction.

49. volume of solution in liters
Molarity
Two solutions can contain the same compounds but be quite different because the proportions of those compounds are different.
Molarity is one way to measure the concentration of a solution:
moles of solute
volume of solution in liters
Molarity (M) =

50. Mixing a Solution
To create a solution of a known molarity, one weighs out a known mass (and, therefore, number of moles) of the solute.
The solute is added to a volumetric flask, and solvent is added to the line on the neck of the flask.

52. Calcium nitrate is composed of calcium ions (Ca2+) and nitrate
ions (NO3–), so its chemical formula is Ca(NO3)2.
Because there are two NO3– ions for each Ca2+ ion, each mole
of Ca(NO3)2 that dissolves dissociates into 1 mol of Ca2+ and
2 mol of NO3–.
Thus, a solution that is M in Ca(NO3)2 is M in
Ca2+ and 2  M = M in NO3–:

54. Dilution One can also dilute a more concentrated solution by
Using a pipet to deliver a volume of the solution to a new volumetric flask, and
Adding solvent to the line on the neck of the new flask.

55. Dilution
The molarity of the new solution can be determined from the equation
Mc  Vc = Md  Vd
where Mc and Md are the molarity of the concentrated and dilute soln., respectively, and Vc and Vd are the volumes of the two solutions.

57. Using Molarities in Stoichiometric Calculations

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

60. Titration

62. Knowing the number of moles of NaOH in 20.0 mL of
solution allows us to calculate the molarity of this solution: