1. Reactions in Aqueous Solution
Chapter 4
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2. The solute is(are) the substance(s) present in the smaller amount(s)
Reactions Between Ions in Aqueous Solutions
General properties of aqueous solution
A solution is a homogenous mixture of 2 or more substances = solute dissovled in solvent
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
4.1

3. Solute
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
4.1

4. Strong and Weak Electrolytes
Conduct electricity in solution?
Cations (+) and Anions (-)
Strong Electrolyte – 100% dissociation
NaCl (s) Na+ (aq) + Cl- (aq)
H2O
Weak Electrolyte – not completely dissociated
CH3COOH CH3COO- (aq) + H+ (aq)
Acetic acid is a weak electrolyte because its ionization in water is incomplete.
A reversible reaction. The reaction can occur in both directions.
4.1

5. Nonelectrolyte does not conduct electricity?
No cations (+) and anions (-) in solution
C6H12O6 (s) C6H12O6 (aq)
H2O

6. How to Predict Electrolytes
• Water soluble and ionic = strong electrolyte (probably)
• Water soluble and molecular, and a weak acid or
weak base = weak electrolyte
• Otherwise, the compound is probably a nonelectrolyte.
4.1

7. Metathesis Reactions
Double Displacement Reactions
• cations(A and C) & anions (B and D) change partners in the reaction
AB + CD –> AD + CB
Example:
• Pb(NO3)2(aq) + 2KI(aq) –> 2KNO3(aq) + PbI2(s)

8. Precipitation of Lead Iodide
PbI2
Precipitation Reactions
Precipitate – insoluble solid that separates from solution
4.2

9. Solubility is the maximum amount of solute that will dissolve in a given quantity of solvent at a specific temperature.
4.2

10. Molecular equation, Ionic equation and net Ionic equation
precipitate
Pb(NO3)2 (aq) + 2NaI (aq) PbI2 (s) + 2NaNO3 (aq)
ionic equation
Pb2+ + 2NO3- + 2Na+ + 2I PbI2 (s) + 2Na+ + 2NO3-
net ionic equation
Pb2+ + 2I PbI2 (s)
Na+ and NO3- are spectator ions: ions that are not involved in the overall reaction
4.2

11. Writing Net Ionic Equations
Write the balanced molecular equation.
Write the ionic equation showing the strong electrolytes completely dissociated into cations and anions. Weak and non electrolytes are written as molecules
Cancel the spectator ions on both sides of the ionic equation
Check that charges and number of atoms are balanced in the net ionic equation
Write the net ionic equation for the reaction of silver
nitrate with sodium chloride.
AgNO3 (aq) + NaCl (aq) AgCl (s) + NaNO3 (aq)
Ag+ + NO3- + Na+ + Cl AgCl (s) + Na+ + NO3-
Ag+ + Cl AgCl (s)
4.2

12. The net ionic equation tells us:
1. changes in ionic strength
(more ions present before the rxn than after)
2. what actually changed during a reaction
Example: Cd2+ (aq) + S2-(aq) –> CdS (s)
Writing ionic equations, ask:
1. is substance soluble ?
2. is substance a strong electrolyte?
**If yes to both questions, write substance as ions.
3. Weak and non electrolytes are written as molecules.
Conditions that favor product
formation:
• formation of a precipitate
• formation of a soluble weak electrolyte
• formation of a nonelectrolyte
• formation of a gas
• oxidation-reduction reactions

13. Acids-Base Reactions Acids
*Have a sour taste. Vinegar owes its taste to acetic acid.
*Cause color changes in plant dyes.
*React with certain metals to produce hydrogen gas.
2HCl (aq) + Mg (s) MgCl2 (aq) + H2 (g)
*React with carbonates and bicarbonates to produce carbon
dioxide gas
2HCl (aq) + CaCO3 (s) CaCl2 (aq) + CO2 (g) + H2O (l)
*Aqueous acid solutions conduct electricity.
4.3

14. Bases Have a bitter taste. Feel slippery. Many soaps contain bases.
Cause color changes in plant dyes.
Aqueous base solutions conduct electricity.
4.3

15. Arrhennius’s definition can be applied only to aqueous solution.
Arrhenius acid is a substance that produces H+ (H3O+) in water
Arrhenius base is a substance that produces OH- in water
Arrhennius’s definition can be applied only to aqueous solution.
4.3

16. A Brønsted acid is a proton donor A Brønsted base is a proton acceptor
A Brønsted acid must contain at least one
ionizable proton!
4.3

18. Monoprotic acids Diprotic acids Triprotic acids HCl H+ + Cl-
Strong electrolyte, strong acid
HNO H+ + NO3-
Strong electrolyte, strong acid
CH3COOH H+ + CH3COO-
Weak electrolyte, weak acid
Diprotic acids
H2SO H+ + HSO4-
Strong electrolyte, strong acid
HSO H+ + SO42-
Weak electrolyte, weak acid
Triprotic acids
H3PO H+ + H2PO4-
Weak electrolyte, weak acid
H2PO H+ + HPO42-
Weak electrolyte, weak acid
HPO H+ + PO43-
Weak electrolyte, weak acid
4.3

22. CH3COO- (aq) + H+ (aq) CH3COOH (aq) Brønsted base
Identify each of the following species as a Brønsted acid, base, or both. (a) HI, (b) CH3COO-, (c) H2PO4-
HI (aq) H+ (aq) + I- (aq)
Brønsted acid
CH3COO- (aq) + H+ (aq) CH3COOH (aq)
Brønsted base
H2PO4- (aq) H+ (aq) + HPO42- (aq)
Brønsted acid
H2PO4- (aq) + H+ (aq) H3PO4 (aq)
Brønsted base
Amphoteric species
4.3

23. Neutralization Reaction
Neutralization occurs when a solution of an acid reacts with a base. The products are salt and water.
acid + base salt + water
Salt = ionic compound
cation from a base + anion from an acid
HCl (aq) + NaOH (aq) NaCl (aq) + H2O
H+ + Cl- + Na+ + OH Na+ + Cl- + H2O
H+ + OH H2O
However, a weak base and an acid gives only a salt
HCl +NH3 NH4Cl
4.3

25. Oxidation-Reduction Reactions
(electron transfer reactions)
2Mg Mg2+ + 4e-
Oxidation half-reaction (lose e-)
O2 + 4e O2-
Reduction half-reaction (gain e-)
2Mg + O2 + 4e Mg2+ + 2O2- + 4e-
2Mg + O MgO
4.4

26. 4.4

27. Oxidation and Reduction always occur together
Zn (s) + CuSO4 (aq) ZnSO4 (aq) + Cu (s)
Zn Zn2+ + 2e-
Zn is oxidized
Zn is the reducing agent
Cu2+ + 2e Cu
Cu2+ is reduced
Cu2+ is the oxidizing agent
Copper wire reacts with silver nitrate to form silver metal.
What is the oxidizing agent in the reaction?
Cu (s) + 2AgNO3 (aq) Cu(NO3)2 (aq) + 2Ag (s)
Cu Cu2+ + 2e-
Ag+ + 1e Ag
Ag+ is reduced
Ag+ is the oxidizing agent
4.4

28. Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2
Oxidation number
The charge the atom would have in a molecule (or an
ionic compound) if electrons were completely transferred.
Free elements (uncombined state) have an oxidation number of zero.
Na, Be, K, Pb, H2, O2, P4 = 0
In monatomic ions, the oxidation number is equal to the charge on the ion.
Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2
The oxidation number of oxygen is usually –2. In H2O2 and O22- it is –1.
4.4

29. O = -2 H = +1 3x(-2) + 1 + ? = -1 HCO3- C = +4
The oxidation number of hydrogen is +1 except when it is bonded to metals in binary compounds (e.g. LiH, CaH2). In these cases, its oxidation number is –1.
Group IA metals are +1, IIA metals are +2 and fluorine is always –1.
The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion. For example,NH4+ , the sum of oxidation numbers is
-3+4(+1)=+1
7. Oxidation numbers do not have to be integers. Oxidation number of oxygen in the superoxide ion, O2-, is -½.
Oxidation numbers of all the elements in HCO3- ?
O = -2
H = +1
3x(-2) ? = -1
HCO3-
C = +4
4.4

30. The oxidation numbers of elements in their compounds
*Metallic elements: only positive oxidation number; nonmetallic elements: positive or negative oxidation number;
*The highest oxidation number of an element in group1A-7A is its group number.
4.4

32. IF7 F = -1 7x(-1) + ? = 0 I = +7 K2Cr2O7 NaIO3 Na = +1 O = -2
Oxidation numbers of all the elements in the following ?
F = -1
7x(-1) + ? = 0
I = +7
K2Cr2O7
NaIO3
Na = +1
O = -2
3x(-2) ? = 0
I = +5
O = -2
K = +1
7x(-2) + 2x(+1) + 2x(?) = 0
Cr = +6
4.4

33. 5. Disproportionation reaction

34. Types of Oxidation-Reduction Reactions
Combination Reaction
A + B C +3 -1
2Al + 3Br AlBr3
Decomposition Reaction
C A + B +1 +5 -2 +1 -1
2KClO KCl + 3O2
4.4

35. Types of Oxidation-Reduction Reactions
Combustion Reaction
A + O B +4 -2
S + O SO2 +2 -2
2Mg + O MgO
4.4

36. Types of Oxidation-Reduction Reactions
Displacement Reaction
A + BC AC + B +1 +2
Sr + 2H2O Sr(OH)2 + H2
Hydrogen Displacement +4 +2
TiCl4 + 2Mg Ti + 2MgCl2
Metal Displacement -1 -1
Cl2 + 2KBr KCl + Br2
Halogen Displacement
4.4

37. Displacement Reactions
Hydrogen Displacement
• Group I and Ca, Sr, & Ba from Group II will
displace hydrogen from cold water.
2 Na(s) + 2H2O(l) --> 2NaOH (aq) + H2(g)
• Many metals will displace hydrogen from acids.
Mg(s) + 2 HCl (aq) --> MgCl2 (aq) + H2 (g)
2. Metal Displacement
More active metals will displace a less active
metal from its compound
Fe (s) + CuSO4 (aq) --> FeSO4(aq) + Cu(s)

38. The Activity Series for Metals
* Arranges metals according to their ease of oxidation
*The higher the metal on the Activity Series, the more active that metal (the easier it is oxidized.)Any metal can be oxidized by the metal ions below it.
*Any metal above hydrogen will displace it from water or from an acid.

39. The Activity Series
• arranges metals according to their ease of oxidation
• can be used to predict reactions
The higher the metal on the Activity Series, Figure 4.16,
the more active that metal (the easier it is oxidized.)
Any metal can be oxidized by the ions of elements below
it on Figure 4.16.
Cu + AgNO3 --> Cu(NO3)2 + Ag
• Cu is above Ag in the activity series; Cu is more active.
Therefore Cu will displace Ag+ from a solution of AgNO3.
• Silver metal will come out of the solution (reduction.)
• The solution will begin to turn blue from the
presence of Cu2+ as the copper metal reacts
(oxidizes.)

40. 3. Halogen Displacement
F2 > Cl2 > Br2 > I2
F2 is the greatest oxidizing halogen
I2 is the least oxidizing halogen
Example:
2 Br- + Cl2 --> 2 Cl- + Br2
Br2 + Cl- --> no reaction (NR)

41. Types of Oxidation-Reduction Reactions
Disproportionation Reaction
Element is simultaneously oxidized and reduced. +1 -1
Cl2 + 2OH ClO- + Cl- + H2O
Chlorine Chemistry
4.4

42. Worked Example 4.5

43. Solution Stoichiometry
The concentration of a solution is the amount of solute present in a given quantity of solvent or solution.
M = molarity =
moles of solute
liters of solution =
mole
liters n V
• Read 2.5 M NaCl as 2.5 molar sodium chloride
Often use square bracket [ ] to indicate the concentration
Example:
What is the molarity of a solution made from 4.00 g of NaOH diluted to a final volume of 250 mL?
First find the number of moles of NaOH.
Then divide by the volume in Liters.
4.5

44. What mass of KI is required to make 500 mL of a 2.80 M KI solution?
volume of KI solution
moles KI
grams KI
500 mL
= 232 g KI
166 g KI
1 mol KI x
2.80 mol KI
1 L soln
1 L
1000 mL

45. Dilution is the procedure for preparing a less concentrated solution from a more concentrated solution.
Dilution
Add Solvent
Moles of solute
before dilution (i)
after dilution (f) =
MiVi
MfVf =
4.5

46. How would you prepare 60.0 mL of 0.200 M
HNO3 from a stock solution of 4.00 M HNO3?
MiVi = MfVf
Mi = 4.00
Mf = 0.200
Vf = 0.06 L
Vi = ? L
Vi =
MfVf Mi =
0.200 x 0.06
4.00
= L = 3 mL
3 mL of acid
+ 57 mL of water
= 60 mL of solution
4.5

47. Gravimetric Analysis - an analytical technique based on the measurement of mass
Dissolve unknown substance in water
React unknown with known substance to form a precipitate
Filter and dry precipitate
Weigh precipitate
Use chemical formula and mass of precipitate to determine amount of unknown ion
AgCl precipitate
NaCl
Add in AgNO3
4.6

50. Equivalence point – the point at which the reaction is complete
Titrations
In a titration a solution of accurately known concentration is added gradually to another solution of unknown concentration until the chemical reaction between the two solutions is complete.
Equivalence point – the point at which the reaction is complete
Indicator – substance that changes color at (or near) the
equivalence point
e.g. phenolphthalein, colorless in acidic and neutral solution but reddish pink in basic solutions
Slowly add base
to unknown acid
UNTIL
the indicator
changes color
4.7

52. What volume of a 1.420 M NaOH solution is
Required to titrate mL of a 4.50 M H2SO4
solution?
WRITE THE CHEMICAL EQUATION!
H2SO4 + 2NaOH H2O + Na2SO4 M
acid rx
coef. M
base
volume acid
moles acid
moles base
volume base
4.50 mol H2SO4
1000 mL soln x
2 mol NaOH
1 mol H2SO4 x
1000 ml soln
1.420 mol NaOH x
25.00 mL
= 158 mL
4.7