1. Chapter 5 Molecular View of Reactions in Aqueous Solutions
Chemistry: The Molecular Nature of Matter, 6E
Jespersen/Brady/Hyslop

2. Reactions in Solution Solution For reaction to occur
Reactants needs to come into physical contact
Happens best in gas or liquid phase
Movement occurs
Solution
Homogeneous mixture
2 or more components mix freely
Molecules or ions completely intermingled
Contains at least 2 substances
5.1 Describing Solutions

3. Definitions: Solvent Solute Medium that dissolves solutes
Component present in largest amount
Can be gas, liquid, or solid
Liquids most common
Aqueous solution—water is solvent
Solute
Substance dissolved in solvent
Solution is named by solute
Can be gas—CO2 in soda
Liquid—Ethylene glycol in antifreeze
Solid—Sugar in syrup

4. Iodine Molecules in Ethanol
Figure 5.1 | Formation of a solution of iodine molecules in alcohol.
Ethanol = solvent
Iodine = solute
Crystal of solute placed in solvent
Solute molecules dispersed throughout solvent

5. Solutions May be characterized using Concentration
Solute-to-solvent ratio
Percent Concentration or

6. Relative Concentration
Dilute solution
Small solute to solvent ratio
Ex. Eyedrops
Concentrated solution
Large solute to solvent ratio
Ex. Pickle brine
Dilute solution contains less solute per unit volume than more concentrated solution
Figure 5.2
Eyedrops = Low concentration of NaCl in water
Pickle brine = high concentration ofNaCl in water

7. Concentration Solubility Saturated solution Unsaturated solution
Temperature dependent
Saturated solution
Solution in which no more solute can be dissolved at a given temperature
Unsaturated solution
Solution containing less solute than maximum amount
Able to dissolve more solute

8. Solubilities of Some Common Substances
Formula
Solubility
(g/100 g water)
Sodium chloride
NaCl
35.7 at 0°C
39.1 at 100°C
Sodium hydroxide
NaOH
42 at 0°C
347 at 100°C
Calcium carbonate
CaCO3
at 25°C

9. Concentrations Supersaturated Solutions
Contains more solute than required for saturation at a given temperature
Formed by careful cooling of saturated solutions
Unstable
Crystallize out when add seed crystal – results in formation of solid or precipitate (ppt.)
Figure 5.3
WileyPlus video: supersaturated solutions.

10. Preciptates Precipitate Precipitation reaction
Solid product formed when reaction carried out in solutions and one product has low solubility
Insoluble product
Separates out of solution
Precipitation reaction
Reaction that produces precipitate
Pb(NO3)2(aq) + 2KI(aq)  PbI2(s) + 2KNO3(aq)
1 mol Pb(NO3)2  2 mol KI
0.100 mol Pb(NO3)2  mol KI

11. Electrolytes in Aqueous Solution
Ionic compounds conduct electricity
Molecular compounds don’t conduct electricity
Why?
Bright light
No light
Ions present
Figure 5.4 Electrolytes, Weak electrolytes and non-electrolytes
Molecular
CuSO4 & water
Sugar & water

12. Ionic Compounds (Salts) in Water
H2O molecules arrange themselves around ions & remove them from lattice.
Dissociation
Break salts apart into ions when enter solution
Separated ions
Hydrated
Conduct electricity
Note: Polyatomic ions remain intact
Ex. KIO3  K+ + IO3
Fig 5.5
Process of dissolution involves active interaction between specific areas of the water molecules.
Ex. Na+ surrounded by oxygen atoms of water Cl surrounded by hydrogen atoms of water
NaCl(s)  Na+(aq) + Cl–(aq)

13. Molecular Compounds In Water
When molecules dissolve in water
Solute particles are surrounded by water
Molecules are not dissociated
Fig 5.6

14. Electrical Conductivity
Electrolyte
Solutes that yield electrically conducting solutions
Separate into ions when enter into solution
Strong electrolyte
Electrolyte that dissociates 100% in water
Yields aqueous solution that conducts electricity
Good electrical conduction
Ionic compounds
Strong acids and bases
Ex. NaBr, KNO3, HClO4, HCl

15. Electrical Conductivity
Weak electrolyte
Aqueous solution that weakly conducts electricity due to low ionization
Weak acids and bases
Ex. Acetic acid (HC2H3O2), ammonia (NH3)
Non-electrolyte
Aqueous solution that doesn’t conduct electricity
Molecules remain intact in solution
Ex. Sugar, alcohol
Distinguishing between strong electrolytes, weak electrolytes, and non-electrolytes?

16. Your Turn
How many ions form on the dissociation of Na3PO4? 1 2 3 4 8

17. Your Turn How many ions form on the dissociation of Al2(SO4)3? 2 3 5 914

18. Equations for Dissociation Reactions
Ionic compound dissolves to form hydrated ions
Hydrated = surrounded by water molecules
In chemical equations, hydrated ions are indicated by
Symbol (aq) after each ions
Ions are written separately
KBr(s)  K+(aq) + Br(aq)
Mg(HCO3)2(s)  Mg2+(aq) + 2HCO3(aq)

19. Learning Check Na3PO4(aq) → Al2(SO4)3(aq) → 3 Na+(aq) + PO43(aq)
Write the equations that illustrate the dissociation of the following salts:
Na3PO4(aq) →
Al2(SO4)3(aq) →
CaCl2(aq) →
Ca(MnO4)2(aq) →
3 Na+(aq) + PO43(aq)
2 Al3+(aq) SO42(aq)
Ca2+(aq) Cl(aq)
Ca2+(aq) MnO4(aq)

20. Equations of Ionic Reactions
Consider the reaction of Pb(NO3)2 with KI
Fig 5.6 {darkened}
PbI2(s)
Pb2+
NO3– K+ I–

21. Equations of Ionic Reactions
When two soluble ionic solutions are mixed, sometimes an insoluble solid forms.
Three types of equations used to describe
Molecular Equation
Substances listed as complete formulas
Ionic Equation
All soluble substances broken into ions
Net Ionic Equation
Only lists ions that actually take part in reaction

22. Equations of Ionic Reactions
1. Molecular Equation
Complete formulas for all reactants and products
Formulas written with ions together
Does not indicate presence of ions
Gives identities of all compounds
Good for planning experiments
Ex.
Pb(NO3)2(aq) + 2KI(aq)  PbI2(s) + 2KNO3(aq)

23. Equations of Ionic Reactions
2. Ionic Equation
Emphasizes the reaction between ions
All strong electrolytes dissociate into ions
Used to visualize what is actually occurring in solution
Insoluble solids written together as they don’t dissociate to any appreciable extent
Ex.
Pb2+(aq) + 2NO3(aq) + 2K+(aq) + 2I(aq)  PbI2(s) + 2K+(aq) + 2NO3(aq)
Provides a more accurate description of what’s happening.

24. Equations of Ionic Reactions
Spectator Ions
Ions that don’t take part in reaction
They hang around and watch
K+ & NO3 in our example
3. Net Ionic Equation
Eliminate all spectator ions
Emphasizes the actual reaction
Focus on chemical change that occurs
Ex. Pb2+(aq) + 2I(aq)  PbI2(s)
Net Ionic Equation shows that more than one set of reactants can lead to the same net reaction.

25. Net Ionic Equations Many ways to make PbI2
Pb(NO3)2(aq) + 2KI(aq)  PbI2(s) + 2KNO3(aq)
Pb(C2H3O2)2(aq) + 2NH4I(aq)  PbI2(s) NH4C2H3O2(aq)
Different starting reagents
Same net ionic equation
Pb2+(aq) + 2I(aq)  PbI2(s)

26. Converting Molecular Equations to Ionic Equations
Strong electrolytes exist as dissociated ions in solution
Strategy
Identify strong electrolytes
Use subscript coefficients to determine total number of each type of ion
Separate ions in all strong electrolytes
Show states as recorded in molecular equations

27. Learning Check: Convert Molecular to Ionic Equations:
Write the correct ionic equation for each:
Pb(NO3)2(aq) + 2NH4IO3(aq) → Pb(IO3)2(s) + 2NH4NO3(aq)
2NaCl (aq) + Hg2(NO3)2 (aq) → 2NaNO3 (aq) + Hg2Cl2 (s)
Pb2+(aq) + 2NO3–(aq) + 2NH4+(aq) + 2IO3–(aq) → Pb(IO3)2(s) + 2NH4+(aq) + 2NO3–(aq)
Note that Hg22+ is a polyatomic ion.
2Na+(aq) + 2Cl–(aq) + Hg22+(aq) + 2NO3–(aq) → 2Na+(aq) NO3–(aq) + Hg2Cl2(s)

28. Your Turn Consider the following reaction :
Na2SO4(aq) + BaCl2(aq) → 2NaCl(aq) + BaSO4(s)
Which is the correct ionic equation?
2Na+(aq) + SO42–(aq) + Ba2+(aq) + Cl22–(aq) → 2Na+(aq) Cl–(aq) + BaSO4(s)
2Na+(aq) + SO42–(aq) + Ba2+(aq) + 2Cl–(aq) → 2Na+(aq) Cl–(aq) + BaSO4(s)
2Na+(aq) + SO42–(aq) + Ba2+(aq) + Cl22–(aq) → 2Na+(aq) Cl–(aq) + Ba2+(s) + SO42–(s)
Ba2+(aq) + SO42–(aq) → BaSO4(s)
Ba2+(aq) + SO42–(aq) → Ba2+(s) + SO42–(s)

29. Converting Ionic Equations to Net Ionic Equations
Strategy
Identify spectator ions
Eliminate from both sides
Rewrite equation using only ions that actually react.
Show states as recorded in molecular and ionic equations

30. Learning Check: Convert Ionic Equation to Net Ionic Equation
Write the correct net ionic equation for each.
Pb2+(aq) + 2NO3–(aq) + 2K+(aq) + 2IO3–(aq) →Pb(IO3)2(s) K+(aq) + 2NO3–(aq)
2Na+(aq) + 2Cl–(aq) + Hg22+(aq) + 2NO3–(aq) → 2Na+(aq) NO3–(aq) + Hg2Cl2(s)
Pb2+(aq) + 2IO3–(aq) → Pb(IO3)2(s)
Note that Hg22+ is a polyatomic ion.
2Cl–(aq) + Hg22+(aq) → Hg2Cl2(s)

31. Your Turn Consider the following molecular equation:
(NH4)2SO4(aq) + Ba(CH3CO2)2(aq) → NH4CH3CO2(aq) + BaSO4(s)
Which is the correct net ionic equation?
Ba2+(aq) + SO42–(aq) → BaSO4(s)
2NH4+(aq) + 2CH3CO2–(aq) → 2NH4CH3CO2(s)
Ba2+(aq) + SO42–(aq) → BaSO4(aq)
2NH4+(aq) + Ba2+(aq) + SO42–(aq) + 2CH3CO2–(aq) → 2NH4+(aq) + 2CH3CO2–(aq) + BaSO4(s)
2NH4+(aq) + 2CH3CO2–(aq) → 2NH4CH3CO2(aq)

32. Criteria for Balancing Ionic and Net Ionic Equations
Material Balance
There must be the same number of atoms of each kind on both sides of the arrow
Electrical Balance
The net electrical charge on the left must equal the net electrical charge on the right
Charge does not have to be zero

33. Learning Check: Balancing Equations for Mass & Charge
Balance Molecular Eqn. for mass
2Na3PO4(aq) + 3Pb(NO3)2(aq)  6NaNO3(aq) Pb3(PO4)2(s)
Can keep polyatomic ions together when counting
Balance Ionic Eqn. for charge
6Na+(aq) + 2PO43(aq) + 3Pb2+(aq) + 6NO3(aq)  Na+(aq) + 6NO3(aq) + Pb3(PO4)2(s)
Charge must add up to zero on both sides.
Net Ionic Eqn. Balanced for both mass & charge
3Pb2+(aq) + 2PO43(aq)  Pb3(PO4)2(s)
Molecular Equation :
2  3 = 6Na+ on each side
2  1 = 2PO43– on each side
3  1 = 3Pb2+ on each side
2  3 = 6NO3– on each side
Reactants: {6  (+1)} + {2  (–3)} + {3  (+2)} + {6  (–1)} = 6 – – 6 = 0
Products: {6  (+1)} + {6  (–1)} = 6 – 6 = 0
{3(+2)} +{2(-3)} = 6 – 6 = 0

34. Acids & Bases as Electrolytes
Many common laboratory chemicals and household products
Indicators
Dye molecules that change color in presence of acids or bases
Acids
Turn blue litmus red
Lemon juice, vinegar, H2SO4
Bases
Turn red litmus blue
Drano (lye, NaOH), ammonia (NH3)
Section 5.3
Indicators = Useful to detect presence of acids or bases
Acids = sour taste
Acids corrode metal
Some OK to eat, some dangerous
Bases = slippery, soapy feel
Mostly used as cleaning products

35. Neutralization Reaction
Important reaction of acids and bases
Acid reacts with base to form water and salt (ionic compound).
Acid + base  salt + H2O
Ex. HCl(aq) + NaOH(aq)  NaCl(aq) + H2O
HBr(aq) + LiOH(aq)  LiBr(aq) + H2O
1:1 mole ratio of acid:base gives neutral solution
Ionization reactions
Ions form where none have been before
Reactions of acids or bases with water
Neutral = Neither acidic nor basic

36. Arrhenius Acid-base neutralization is H+(aq) + OH–(aq)  H2O
In solution, H+ attaches itself to H2O to form H3O+ or hydronium ion in water
H+ does not ever exist in aqueous solution
When H3O+ reacts, it releases H+
H+ is active ingredient
Often use just H+ for simplicity
Arrhenius acid-base = reaction of H+ (hydrogen ion) with OH– (hydroxide ion) to form H2O (water).

37. Arrhenius Acid
Substance that reacts with water to produce the hydronium ion, H3O+
Acid H2O  Anion + H3O+
HA H2O  A– + H3O+
HC2H3O2(aq) + H2O  H3O+(aq) + C2H3O2−(aq)
HCl(g) + H2O 
Cl–(aq) + H3O+(aq)
Fig 5.10
Note for organic acids
In general, only hydrogen written first in formula transfers to H2O to give H3O+.

38. Acids Categorized by Number of H+s
Monoprotic Acids
Furnish only one H+
HNO3(aq) + H2O  H3O+(aq) + NO3–(aq)
HC2H3O2(aq) + H2O  H3O+(aq) + C2H3O2–(aq)
Polyprotic acids
Furnish more than one H+
Diprotic acids — furnish two H+
H2SO3(aq) + H2O  H3O+(aq) + HSO3–(aq)
HSO3–(aq) + H2O  H3O+(aq) + SO32–(aq)

39. Acids Catagorized by Number of H+s
Polyprotic acids
Triprotic acids — furnish three H+
H3PO4  H2PO4–  HPO42–  PO43–
Stepwise equations
H3PO4(aq) + H2O  H3O+(aq) + H2PO4–(aq)
H2PO4–(aq) + H2O  H3O+(aq) + HPO42–(aq)
HPO42–(aq) + H2O  H3O+(aq) + PO43–(aq)
Net:
H3PO4(aq) + 3H2O  3H3O+(aq) + PO43–(aq)
–H+
–H+
–H+

40. Acidic Anhydrides Nonmetal Oxides SO3(g) + H2O  H2SO4(aq)
Act as Acids
React with water to form molecular acids that contain hydrogen
SO3(g) + H2O  H2SO4(aq)
sulfuric acid
N2O5(g) + H2O  2HNO3(aq)
nitric acid
CO2(g) + H2O  H2CO3(aq)
carbonic acid

41. Arrhenius Bases 1. Ionic compounds containing OH– or O2–
Ionic compounds that contain hydroxide ion, OH–, or oxide ion, O2–. or
Molecular compounds that react with water to give OH–.
1. Ionic compounds containing OH– or O2–
a. Metal Hydroxides
Dissociate into metal & hydroxide ions
NaOH(s)  Na+(aq) + OH–(aq)
Mg(OH)2(s)  Mg2+(aq) + 2OH–(aq)

42. Ionic Oxides b. Basic Anhydrides
Soluble metal oxides
Undergo ionization (hydrolysis) reaction to form hydroxide ions
Oxide reacts with water to form metal hydroxide
CaO(s) + H2O  Ca(OH)2(aq)
Then metal hydroxide dissociates in water
Ca(OH)2(aq)  Ca2+(aq) + 2OH–(aq)
Fig. 5.12
O2–
H2O
2OH–

43. Strong vs. Weak Electrolyte
Fig 5.14
Strong Electrolytes =Ionic compounds that dissociate 100% in water
All strong acids & bases
Weak electrolytes = Electrolytes that only partially dissociate (<100%)
All weak acids & bases.
HCl(aq)
CH3COOH(aq)
NH3(aq)

44. Strong Acids perchloric acid chloric acid hydrochloric acid
HClO4(aq)
perchloric acid
HClO3(aq)
chloric acid
HCl(aq)
hydrochloric acid
HBr(aq)
hydrobromic acid
HI(aq)
hydroiodic acid
HNO3(aq)
nitric acid
H2SO4(aq)
sulfuric acid
Dissociate completely when dissolved in water
Ex. HBr(g) + H2O  H3O+(aq) + Br–(aq)
Good electrical conduction
Any acid not on this list, assume weak

45. Arrhenius Bases  2. Molecular Bases
Undergo ionization (hydrolysis) reaction to form hydroxide ions
Base + H2O  BaseH+(aq) + OH–(aq)
B H2O  BH+(aq) + OH–(aq)
NH3(aq) + H2O  NH4+(aq) + OH–(aq)
Fig 5.14 
NH4+
NH3
OH–
H2O

46. Strong Bases Common strong bases are:
Bases that dissociate completely in water
Soluble metal hydroxides
KOH(aq)  K+(aq) + OH–(aq)
Good electrical conductors
Behave as (aq) ionic compounds
Common strong bases are:
Group IA metal hydroxides
LiOH, NaOH, KOH, RbOH, CsOH
Group IIA metal hydroxides
Ca(OH)2, Sr(OH)2, Ba(OH)2

47. Weak Acids Any acid other than 7 strong acids
Only ionize partially (<100%)
Organic acids
HC2H3O2(aq) + H2O  H3O+(aq) + C2H3O2–(aq)
Ex. HCO2H(aq) + H2O  H3O+(aq) + HCO2–(aq)
Acetic Acid
Molecule, HC2H3O2
Fig 5.11 separated
Acetate ion, C2H3O2–
Only this H comes off as H+

48. Why is Acetic Acid Weak? H2O + C2H3O2–(aq)  HC2H3O2(aq) + H3O+(aq)
Fig 5.15 Acetic acid is a poor electrical conductor. Only a few H3O+ & C2H3O2– ions form. C2H3O2– strong tendency to react with H3O+.
Balance reached when ions form & combine at same rate.
Here corresponds to only a few % dissociated.
H3O+(aq) + C2H3O2–(aq)  HC2H3O2(aq) + H2O

49. Dynamic Equilibrium 2 opposing reactions occurring at same rate
Also called Chemical equilibrium
Equilibrium
Concentrations of substances present in solution do not change with time
Dynamic
Both opposing reactions occur continuously
Represented by double arrow
HC2H3O2(aq) + H2O H3O+(aq) + C2H3O2–(aq)
Forward reaction – Forms ions
Reverse reaction – Removes ions

50. Weak Bases NH3(aq) + H2O  NH4+(aq) + OH(aq)
Molecular bases
Do not dissociate
Accept H+ from water inefficiently
Accept H+ from acids preferentially
NH3(aq) + HCl(aq)  NH4Cl(aq)
Ex.
NH3(aq) + H2O  NH4+(aq) + OH(aq)
Or for general base
B(aq) + H2O  BH+(aq) + OH(aq)

51. Equilibrium for Weak Base
Forward reaction
Reverse reaction
Fig. 5.16
Net is dynamic equilibrium
NH3(aq) + H2O NH4+(aq) + OH(aq)

52. Position of Equilibrium
Extent of completion
Depends on electrolyte
Weak electrolyte
Small % ionizes
 dominant
Mostly reactants
Weak acids and bases
Lots of back reaction
Write eqn. as
Strong electrolyte
Large % ionizes
 dominant
Mostly products
Strong acids & bases
Little back reaction
Write eqn. as 

53. Learning Check
Write the ionization equation for each of the following with water:
Weak base methylamine, CH3NH2.
Weak acid nitrous acid, HNO2.
Strong acid chloric acid, HClO3.
Strong base strontium hydroxide, Sr(OH)2.
CH3NH2(aq) + H2O CH3NH3+(aq) + OH–(aq)
HNO2(aq) + H2O H3O+(aq) + NO2–(aq)
HClO3(aq) + H2O  H3O+(aq) + ClO3–(aq)
Sr(OH)2(aq)  Sr2+(aq) + 2 OH–(aq)

54. Your Turn Which of the following is a weak acid? HCl HNO3 HClO4
HC2H3O2
H2SO4

55. Your Turn Which of the following is not a strong base? NaOH CH3NH2
Cs2O
Ba(OH)2
CaO

56. Your Turn Which of the following is not a product of the reaction:
NH3(aq) +HCN(aq) ?
CN–(aq)
NH4+(aq)
NH3CN(s)
H2O
HCN

57. Acid—Base Nomenclature
System for naming acids and bases
Acids
Hydrogen compounds of non-metals = binary acids
Hydrogen compounds of oxoanions = Oxoacids
Naming acid salts
Bases
Metal Hydroxides and oxides = ionic
Molecular = molecular names
5.4 Acid—Base Nomenclature

58. Naming Acids A. Binary Acids — hydrogen + nonmetal
Take molecular name
Drop –gen from H name
Merge hydro– with nonmetal name
Replace –ide with –ic acid
Name of Molecular compound
Name of Aqueous Binary Acid
HCl(g)
hydrogen chloride
HCl(aq)
hydrochloric acid
H2S(g)
hydrogen sulfide
H2S(aq)
hydrosulfuric acid
Name of (aq) form differs from other states due to ionization that occurs in water

59. Naming Acids B. Oxo Acids To name:
Acids with hydrogen, oxygen and another nonmetal element
Most of the polyatomic ions in Table 3.5
To name:
Based on parent oxoanion name
Take parent ion name
Anion ends in –ate change to –ic (more O's)
Anion ends in –ite change to–ous (less O's)
End name with acid to indicate H+

60. Oxoacids (Aqueous) Named according to the anion suffix
Anion ends in -ite, acid name is -ous acid
Anion ends in -ate, acid name is -ic acid
Name of Parent Oxoanion
Name of Oxoacid
NO3
HNO3
SO42
H2SO4
ClO2
HClO2
PO32
H2PO3
nitrate
nitric acid
sulfate
sulfuric acid
Common mnemonic: something I -ate was icky.
chlorite
chlorous acid
phosphite
phosphorous acid

61. Learning Check: Name Each Aqueous Acid
HNO2
HCN
HClO4 HF
H2CO3
nitrous acid
hydrocyanic acid
perchloric acid
hydrofluoric acid
carbonic acid

62. Your Turn Which of the following is the correct name for HClO4 (aq)?
chloric acid
hydrochloric acid
perchloric acid
hypochlorous acid
chlorous acid

63. Your Turn Which of the following is the correct name for H2SO3(aq)?
sulfuric acid
sulfurous acid
hydrosulfuric acid
hydrosulfurous acid
hydrogen sulfite acid

64. Acid Salts If polyprotic acids are neutralized stepwise Acid salt
Can halt neutralization before all H+’s are removed
Must specify # of H's that remain on salt
Acid salt
Ion containing H+ and anion
Contains anion capable of furnishing additional hydrogen ions
H2SO4(aq) + KOH(aq)  KHSO4(aq) + H2O(ℓ)
acid salt

65. Naming Acid Salts—Polyprotic
Must specify number of hydrogens still attached to the anion
Can be neutralized by additional base
Ex. Na2HPO4
NaH2PO4
KHSO4
Some acid salts have common names
NaHCO3
sodium hydrogen phosphate
sodium dihydrogen phosphate
potassium hydrogen sulfate
sodium hydrogen carbonate or sodium bicarbonate

66. C. Naming Bases Oxides & Hydroxides Molecular Bases Ionic compounds
Named like ionic compounds
Ca(OH)2 calcium hydroxide
Li2O lithium oxide
Molecular Bases
Named like molecules
NH3 ammonia
CH3NH2 methylamine
(CH3)2NH dimethylamine
(CH3)3N trimethylamine