1. Chapter 4 Reactions in Aqueous Solution

2. Background: 1. Many reactions occur in aqueous solutions:
Three common types of reactions in solution: Precipitation, Acid-base, and Oxidation-reduction
Concentration of solutions is measured in units of Molarity
Water is the universal solvent

3. 2. Definitions
Solution: homogeneous mixture of a solvent and a solute (not always in same phase – solid/liquid/gas)
The solvent is present in the greatest amount,
while all other substances are solutes.
Aqueous: dissolved in water
Anion: a negatively charged ion
Ex O2-, CN-1
Cation: positively charged ion
Ex H+
Electrodes: measure “electron” flow
in a solution

4. 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) =
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5. Solute Concentrations - Molarity
Definition of molarity
Molarity = moles of solute/liters of solution
Symbol is M
Square brackets are used to indicate concentration in M
[Na+] = 1.0 M
Example: 1.5 moles of NaCl are dissolved to make 250mL aqueous solution.

6. Additivity Masses are additive; volumes are not
The total mass of a solution is the sum of the mass of the solute and the solvent
The total volume of a solution is not the sum of the volumes of the solute and solvent
Molarity as a conversion:
Use: # moles = 1 Liter

7. Volumetric Glassware
Volumetric pipets, burets and flasks are made so that they contain an exact volume of liquid at a given temperature
Preparation of molar solutions involves using volumetric glassware

8. Example 4.1

9. Example 4.1

10. 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.
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11. How to Use a Volumetric Flask

12. 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.
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13. 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 solutions, respectively, and Vc and Vd are the volumes of the two solutions.
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14. Using Molarities in Stoichiometric Calculations
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15. Titration
Titration is an analytical technique in which one can calculate the concentration of a solute in a solution.
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16. 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.
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17. Dissociation
An electrolyte is a substances that dissociates into ions when dissolved in water.
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18. Electrolytes
An electrolyte is a substances 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.
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19. Electrolytes and Nonelectrolytes
Soluble ionic compounds tend to be electrolytes.
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20. Electrolytes and Nonelectrolytes
Molecular compounds tend to be nonelectrolytes, except for acids and bases.
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21. Electrolytes
A strong electrolyte dissociates completely when dissolved in water.
A weak electrolyte only dissociates partially when dissolved in water.
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22. Strong Electrolytes Are…
Strong acids
Strong bases
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23. Strong Electrolytes are also….
Soluble ionic salts
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24. Dissolving Ionic Solids
When an ionic solid is dissolved in a solvent, the ions separate from each other
MgCl2 (s) → Mg2+ (aq) + 2 Cl-1 (aq)
The concentrations of ions are related to each other by the formula of the compound:
Molarity of MgCl2 = Molarity of Mg2+
Molarity of Cl-1 = 2 X Molarity of MgCl2
Total number of moles of ions per mole of MgCl2 is 3

25. Example 4.2

26. Example 4.2

27. Precipitation Reactions
Precipitation in chemical reactions is the formation of a solid where no solid existed before reaction
Precipitation is the reverse of solubility, where a solid dissolves in a solvent to produce a solution
When one mixes ions that form compounds that are insoluble (as could be predicted by the solubility guidelines), a precipitate is formed.

28. Figure 4.3 – Precipitation Diagram

29. Solubility Trends Mostly soluble Compounds of Group 1 and NH4+ cations
All nitrates
All chlorides, except for AgCl
All sulfates, except for BaSO4
Mostly insoluble
Carbonates and phosphates, except for the Group I and ammonium
Hydroxides, except for the Group 1, Group 2 and ammonium
SAP (compounds containing sodium, ammonium, and potassium are soluble)
CAN (chlorate, acetate, and nitrate containing compounds are soluble)

30. Example 4.3

31. Possible precipitates include: copper (II) sulfate or ammonium nitrate  both soluble…NO PPT!
Possible precipitates include: iron (III) nitrate or silver chloride PPT: AgCl !!

32. AgNO3 (aq) + KCl (aq)  AgCl (s) + KNO3 (aq)
Molecular Equation
The molecular equation lists the reactants and products in their molecular form.
AgNO3 (aq) + KCl (aq)  AgCl (s) + KNO3 (aq)
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33. 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.
Ag+ (aq) + NO3- (aq) + K+ (aq) + Cl- (aq) 
AgCl (s) + K+ (aq) + NO3- (aq)
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34. 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)
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35. Ag+(aq) + Cl-(aq)  AgCl (s)
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.
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)
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36. 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.
The only things left in the equation are those things that change (i.e., react) during the course of the reaction.
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)
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37. 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.
Atoms must balance
Charges must balance
Show only the ions that react
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38. Example 4.4

39. Example 4.4

40. Another common type of reaction in aqueous solution involves a transfer of electrons between two species.
Such a reaction is called an oxidation-reduction, aka
REDOX REACTION.

41. Redox Reactions and Oxidation Numbers
Oxidation is an increase in oxidation number
This is the same as a loss of electrons (LEO)
Increases charge
Reduction is a decrease in oxidation number
This is the same as a gain of electrons (GER)
Decreases charge
OIL RIG

42. Example: Which element is being oxidized and which is being reduced?
Fe  Fe e-
F + 1e-  F-1

43. Example: Which element is being oxidized and which is being reduced?
Fe  Fe e- OXIDIZED
F + 1e-  F-1 REDUCED

44. REDOX TIME A species is oxidized when it loses electrons.
Here, zinc loses two electrons to go from neutral zinc metal to the Zn2+ ion.
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45. REDOX TIME A species is reduced when it gains electrons.
Here, each of the H+ gains an electron, and they combine to form H2.
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46. REDOX TIME What is reduced is the oxidizing agent.
H+ oxidizes Zn by taking electrons from it.
What is oxidized is the reducing agent.
Zn reduces H+ by giving it electrons.
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47. Oxidation Numbers
In order to keep track of what loses electrons and what gains them, we assign oxidation numbers.

48. Assigning 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.

49. Assigning Oxidation Numbers
Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions.
Oxygen has an oxidation number of −2, except in the peroxide ion, which has an oxidation number of −1.
Hydrogen is −1 when bonded to a metal and +1 when bonded to a nonmetal.

50. Assigning Oxidation Numbers
Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions.
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.

51. Assigning 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.
Always remember and start with the common charges of the main group families on the PT! Then work backwards to figure out the oxidation numbers that can change!

52. Balancing Redox Equations
Assign oxidation numbers to determine what is oxidized and what is reduced.
Split the equation into the oxidation and reduction half-reaction.
Balance each half-reaction with respect to both atoms and charge.
Balance the atoms of the element being oxidized or reduced.
Balance oxidation number by adding electrons. (reduction: electrons on left, oxidation: electrons on right)

53. Balancing Redox Equations Continued…
Balance charge by adding H+ ions in acidic solution OR OH- ions in basic solution
Balance hydrogen by adding H2O molecules.
Check to make sure that oxygen is balanced.
4. Combine the two half-reactions in such a way to eliminate the electrons.
Balancing Redox by Oxidation Number Method