1. Reactions in Aqueous Solutions
CH 4 AP
Reactions in
Aqueous Solutions

2. Water Aqueous means dissolved in H2O
Moderates the Earth’s temperature because of high specific heat
H-bonds cause strong cohesive and adhesive properties
Polar, therefore dissolves many substances both ionic and polar
Solutions: with solvent (dissolving medium) and solute (being dissolved)

3. Electrolyte is a solution which will conduct electricity (solute makes ions in solution allow the flow of e-) Ion Dissociation
Solvated is when an ionic compound dissolves in H2O and the ions become surrounded by the H2O molecules.
Non-electrolyte is a solutions which will NOT conduct electricity (no ions present to permit flow of e-)
See page 126

4. Soluble salts NaCl(s)  Na+1(aq) + Cl-1(aq)
Strong Electrolytes are made with substances which completely ionize in H2O, they have large equilibrium constants.
Soluble salts NaCl(s)  Na+1(aq) + Cl-1(aq)
Strong acids HCl  H+(aq) + Cl-1(aq) Arrhenius acids product H+1 when dissolved in H2O
Strong bases NaOH(s)  Na+1(aq) + OH-1(aq)

5. Non-electrolytes produce NO ions upon dissolving in H2O
Weak electrolytes exhibit small amounts of ionization when dissolved in H2O, have small equilibrium constants.
Weak acids HC2H3O2(aq)  H+1(aq) + C2H3O2-1 (aq)
Weak bases NH3 (aq) + H2O (L)  NH4+1(aq) + OH-1(aq)
Non-electrolytes produce NO ions upon dissolving in H2O

6. Writing half reactions
Ag  Ag+1
Ba  Ba+2
F  F-1
S  S-2
Fe+2  Fe+3

7. Do you know the subatomic particles?Ag
Ag+1 Ba
Ba+2 F
F-1 S
S-2
Fe+2
Fe+3

8. Do you remember how to write formulas?
Oxide
O-2
Iodide
I-1
Sulfate
SO4-2
Phosphate
PO4-3
Cyanide
CN-1
Hydrogen
H+1
Barium
Ba+2
Chromium
Cr+3

9. Molarity = moles/liter
Molarity (M): concentration of solution expressed as moles of solute per volume of solution in liters.
Molarity = moles/liter
Calculate the molarity of a solution prepared by dissolving 17.5 g of solid NaOH in enough H2O to make 1.75 L of solution.

10. Volumetric flask To use: Fill ½ full of distilled H2O Add solute
Cap and invert to mix
Add H2O until close to line on neck
Carefully, add H2O until meniscus is at the line on neck

11. Calculate the concentration of each type of ion in the following:
0.50 M AlCl3
2 M (NH4)2CO3
Calculate the moles of Cl-1 ions in 1.3 L of 0.05 M MgCl2.

12. Typical blood serum is about 0. 14 M NaCl
Typical blood serum is about 0.14 M NaCl. What volume of blood contains 1.0 mg NaCl?
How many grams of AgCl must be used to make 75 mL of 0.1 M AgCl solution?

13. Dilutions: moles = moles M1 V1 = M2 V2
What volume of 0.25 M K2CrO4 must be used to prepare 1.0 L of 0.05 M K2CrO4?
What volume of 12.0 M HCl must be used to prepare 2.5 L of 0.1 M HCl?

14. Types of Chemical Reactions
Precipitation Reaction is when 2 solutions are mixed resulting in an insoluble substance.
(Insoluble & slightly soluble are used interchangeably here. Any substance that is less soluble than .01 mol/L is considered insoluble.)
When writing the formula for the precipitate:
a. Precipitate has a net charge of zero
b. Contains a cation (+ ion) and an anion (- ion)
c. Write cation first, anion second

15. Solubility Rules Soluble: the following are always soluble
1. Nitrates (NO3-1) and Acetates (C2H3O2-1)
2. Alkali Metals (Li+1 Na+1 K+1 Cs+1 Rb+1) and
Ammonium (NH4+1)
3. Halides (Cl-1 Br-1 I-1) (except with Ag+1 Pb+2 Hg2+2 are insoluble)
4. Sulfates (SO4-2)
(except with Ba+2 Pb+2 Hg2+2 Sr+2 Ca+2 are insoluble)

16. Mostly Insoluble
1. Hydroxides (OH-1) (except with Alkali metals or ammonium are soluble) (except with Ca+2 Sr+2 Ba+2 are slightly soluble)
2. Sulfides (S-2) (except with Alkali metals, ammonium are soluble ) (except with Alkali Earth metals Be+2 Mg+2 Ca+2 Sr+2 Ba+2 are slightly soluble)
3. Carbonates (CO3-2) and Phosphates (PO4-3) are slightly
soluble (except with alkali metals and ammonium are soluble)
Chromates are insoluble except with alkali metals or ammonium

17. Write the net ionic equation for all precipitates that can be made mixing the following chemicals:
Ba(NO3)2 KCl Na2CO3 Ag2SO4 Mg(C2H3O2)2 NaOH

18. Writing chemical equations
Formula equation: shows everything involved
K2CrO4 (aq) + Ba(NO3)2 (aq)  BaCrO4 (s) KNO3 (aq)
Complete ionic equation: includes all spectator ions, those ions which are not involved in the chemical reaction 2K+1(aq) + CrO4+2(aq) + Ba+2(aq) + 2NO3-1(aq)  BaCrO4 (s) + 2 K+1(aq) + 2 NO3-1(aq)
Net ionic equation: only shows chemicals involved in the reaction, no spectator ions Ba+2(aq) + CrO4-2(aq)  BaCrO4 (s)
All of the above are the same reaction shown in different types of equations.

19. Try writing all 3 types of equations for: Aqueous potassium chloride is added to aqueous silver nitrate
Aqueous potassium hydroxide is mixed with aqueous iron(III) nitrate

20. To calculate ion concentrations after reactions
Identify species present as reactants and products (the complete ionic equation)
Change to moles
Find the Limiting reactant
ICE chart
I = initial concentration
C = change in concentration
E = ending concentration
Use stoichiometric relations
Calculate moles of ions for any soluble chemicals (or excess chemicals)
Change to Molarity using total volume

21. Calculate the concentrations of all ions and mass of precipitate after 100 mL of 1.0 M Mg(NO3)2 reacts with 200 mL of 1.0 M NaOH.
Mg(NO3)2(aq) NaOH  I C E

22. 1. 0 L of 0. 1 M Ba(NO3)2 is mixed with 500 mL of 0. 1 M Na2SO4
1.0 L of 0.1 M Ba(NO3)2 is mixed with 500 mL of 0.1 M Na2SO4. What are the final concentrations of ions in solution and mass of precipitate?
Ba(NO3)2 (aq) + Na2SO4 (aq)  I C E

23. How many grams of silver chloride can be produced by the reaction of 100 mL of .20 M silver nitrate with 100 mL of .15 M calcium chloride. Calculate the concentration of all ion at the end of the reaction.

24. Reaction Type 2 Acid – Base Reaction
Theories: Acid Base
Arrhenius H+ OH-
Bronsted – Lowry proton donor proton acceptor
(proton = H+ , a hydrogen without and e-)
Lewis e- pair acceptor e- pair donor

25. Strong:ionize 100 % Need to memorize strong Acids & Bases both HCl & NaOH strong: H Cl Na OH spectator ions Net ionic equation: H OH-1  H2O(L) Weak: ionize some weak acid (HC2H3O2 stays mainly as molecule) & strong base(KOH ionizes 100 %) Net ionic equation: HC2H3O2 + OH-  H2O + C2H3O2- Non – electrolyte: ionize 0% We will start with only strong acids and bases, therefore ionize 100%

26. strong acid. + strong base  salt + water
strong acid + strong base  salt + water HA BOH  BA H2O (net) H OH-  H2O Neutralization Reactions with gas formation: (Acid) 2 H S-2  H2S(g) H HCO3-1  H2CO3 H2CO3  H2O + CO2(g)

27. When calculating acid base equations, remember that moles A = moles B
When calculating acid base equations, remember that moles A = moles B MA VA = MB VB
What volume of M HCl solution is needed to neutralize 25.0 mL of 0.35 M NaOH?

28. moles A = moles B MA VA = MB VB
This works using M x Liters = moles
mole/Liter = M
This also works using
M x milliliters = millimoles
millimole/milliliter = M

29. 28. 0 mL of 0. 25 M HNO3 reacts with 53. 0 mL of 0. 32 M KOH
28.0 mL of 0.25 M HNO3 reacts with 53.0 mL of 0.32 M KOH. Calculate the amount of H2O formed in the resulting reaction. What concentration of H+1 or OH-1 is in excess?
KOH(aq) HNO3(aq)  I C E

30. Volumetric Analysis: Titration delivers a given amount of known concentration (titrant) to a given amount of unknown concentration (analyte). Equivalence point is the stoichiometric point when the moles = moles. An indicator chemical is added which changes color at the equivalence point (end point). Standardizing is using a solution of a very stable chemical to determine the M of an unknown.

31. A student standardizes an NaOH solution using Potassium Hydrogen Phthalate (KHC8H4O4  KHP) (m=204.22). The student dissolves g KHP in H2O, adds phenolphthalein and titrates with mL NaOH of unknown concentration. Find the concentration H OH-1  H2O HC8H4O OH-1  H2O C8H4O4-2

32. therefore Reducing Agent therefore Oxidizing Agent
Reaction Types #3 Oxidation – Reduction (Redox) A reaction in which one or more e- are transferred.
Oxidation Reduction
Increase in os
therefore Reducing Agent
Loss of e-
Gain of oxygen
Loss of hydrogen
Decrease in os
therefore Oxidizing Agent
Gain of e-
Loss of oxygen
Gain of hydrogen
Reduction
Oxidation

33. Oxidation State (o.s.) (oxidation number)
*Is an imaginary charge.
*It is a way to arbitrarily assign e- to atoms especially in covalent bonds
*The e- are often not shared equally, one atom may have a stronger e- affinity or attraction.

34. Rules for assigning o.s. Whenever 2 rules appear to contradict one another, follow the rule that appears higher on the list!
The oxidation state (os) of an atom in the free (uncombined) element is zero (o).
The total of the os of all the atoms in a molecule is zero (o). For an ion, the total is equal to the charge on the ion.
In their compounds,
Alkali metals (Li, Na, K, Rb, Cs, Fr) have os = +1
Alkali Earth metals (Be, Mg, Ca, Sr, Ba, Ra) have os = +2
In their compounds, Hydrogen has os = Fluorine has os = -1
In it’s compounds, Oxygen has os = -2
In 2 element compounds,
Halogens (Cl, Br, I) have os = -1
(O, S, Se, Te, Po) have os = -2
(N, P, As, Sb, Bi) have os = -3

35. Determine the os for the following underlined elements.
Cr2O7-2
Cl2O
KO2
S2O3-2
KMnO4
H2CO3

36. Write the formulas for the oxides with the following charges.
Cr+3
Cr+4
Cr+6
N+1
N+2
N+3
N+4
N+5

37. Activity Series of Metals see page 144
A list of metals arranged in order of decreasing ease of oxidation. The higher on the chart, the more active the metal. Any metal on the list can be oxidized by the ions of elements below it. Will an aqueous solution of iron (II) chloride oxidize Mg(s)

38. Metal Activity Series – Any metal on the list can be oxidized by the ions of elements below it.
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