1.Definitions Solution - homogeneous mixture of two

or more substancesAqueous - dissolved in water (aq)Anion - negatively charged ionCation - positively charged ion

Electrode - conductor of electricity- electrical conductors (wires, plates) partially immersed in a solution and connected to a source of electricity (e-)

Cathode- connected to the negative pole Anode- connected to the positive pole

2. Arrhenius’ theory of electrolytic dissolution

a. strong electrolyte- solute that is present in solution almost exclusively as

ions; good electrical conductor (ex. KCl, HBr, ) b. weak electrolyte-

solute that is only partly ionized in solution; poor conductor of electricity (ex. Acetic Acid (vinegar))

c. Nonelectrolyte - Solute that is present in solution almost exclusively as

molecules; does not conduct electricity (ex. Sugar)

3. Solubility Rules

Memorize; refer to back of stations lab

4. Dissolution equations Ex5.1Write dissolution equations for: NaCl(s), Al(NO3)3(s), Na2SO4(s), and MgCl2(s)

1. Definitions Double Replacement Reaction:

reaction between two ionic compounds in which the cations and anions are exchanged to form two new ionic compounds

Precipitate: Insoluble ionic solid Symbolized by: (s) or

Full Ionic Equation Equation in which all compounds are written as ions REMEMBER Coefficients, Subscripts, etc.!

Net Ionic Equation “net result” of the chemical reaction Formation of solid, gas, H2O(l), etc.

Spectator Ions Ions that do not take part in the reaction; present before and

after

Double Replacement Rxns The positive ions from one solution interact with

the negative ions from the other solution to form:

A precipitate that will settle out Pb(NO3 )2(aq) + 2 NaCl (aq) ---> PbCl2(s) + 2 NaNO3(aq)

A gas that will bubble out FeS (s) + 2HCl (aq) ---> H2S (g) + FeCl2 (aq)

A molecular compound, usually water HCl (aq) + NaOH (aq) ---> NaCl (aq) + H2O (l)

What is happening in the solution?

The reactants are aqueous, which means that ionic compounds are ionized.

Individual positive and negative ions are present in the aqueous solutions

KCl (s) --H2O---> K+ (aq) + Cl- (aq)

Full (Complete) Ionic Equation

AgNO3 (aq) + NaCl (aq) ---> AgCl (s) + NaNO3(aq)

**Write the Complete Ionic Equation by separating all aqueous compounds

Na+ (aq) + Cl- (aq) + Ag+

(aq) + NO3-(aq) ---> AgCl (s) + Na+

(aq) + NO3-(aq)

Spectator Ions Ions that do not take part in the reaction Ions are present before and after the reaction

Na+ (aq) + Cl- (aq) + Ag+

(aq) + NO3-(aq) ---> AgCl (s) + Na+

(aq) + NO3-(aq)

Spectator Ions

Net Ionic Equation

Remove the Spectator ions from the Complete Ionic Reaction

N.I. = Cl- (aq) + Ag+

(aq) ---> AgCl (s)

2. Ex5.2 Predict whether or not the following pairs of reactants will form precipitates. Then write the double replacement reaction, the full ionic equation, and the net ionic equation and identify spectator ions.

a. CuCl2 and (NH4)2SO4

b. Ba(NO3)2 and Na2CO3

c. MgCl2 and AgNO3

Ex5.3 What two aqueous solutions of soluble compounds could be mixed to produce a precipitate of CuCO3?

Ex5.4 You suspect that a particular unlabeled aqueous solution is one of the following: sodium sulfate, ammonia (NH3), or barium nitrate. Explain how to use precipitation reactions on small samples of the unknown solution to determine its identity.

1. Arrhenius Definition of acids and bases Acids

Produces H+ in soln

Bases Produces OH- in soln

2. Strong acids – strong electrolytes HCl (g) -------> H+

(aq) + Cl- (aq)

Strong Acids to Memorize:

“Br I Cl SO NO ClO 4-3-4” HCl, HBr, HI, HNO3, HClO4, H2SO4**

3. Weak acids – weak electrolytes Any acid that isn’t “strong” Majority of acids

HC2H3O2 (aq) H+ (aq) + C2H3O2-1

(aq)

4. Strong bases – strong electrolytes

NaOH (s) --------> Na+ (aq) + OH- (aq)

Strong Bases to Memorize: (Group I, II Hydroxides, except the 1st one in group I and the 1st two in group II) LiOH, NaOH, KOH, RbOH, CsOH Ca(OH)2, Sr(OH)2, Ba(OH)2

5. Weak bases – weak electrolytes Any base that isn’t strong Watch for Nitrogen-containing compounds

NH3 (aq) + H2O (l) <-----> NH4+ (aq) + OH-

(aq)

6. Oxides of metals and nonmetal Oxides of nonmetals form acids

Ex. CO2 + H2O H2CO3

Oxides of metals form bases Ex. CaO + H2O Ca(OH)2

1. Neutralization reaction and net ionic equation Neutralization Rxn:

Reaction between an acid and a base **Most Rxns: Acid + Base Water + Salt

Net Ionic Eqn:

H+ (aq) + OH-

(aq) H2O (l)

Ex5.5 Write the full neutralization reaction and net ionic equations for the interactions between HNO3 and Ba(OH)2, and between H2SO4 and KOH.

3. Acid Base Titration Titration

Lab procedure where acid and base are made to combine in their stoichiometric proportions

Follows the same stoich process as in Chap 4

1. Definition Reaction in which a gas is formed, sometimes

as the result of a decomposition of one of the products Ex. H2CO3

2. Metal and acid PURE metals react with acid to produce H2

gas Ex. Mg + HCl MgCl2 + H2

3. Carbonate & acidMgCO3 + HCl MgCl2 + H2CO3

H2CO3 H2O + CO2

1. General Classification:Synthesis, Decomposition, Single and Double Rep.

2. Redox Reaction Species lose/gain electrons as reaction occurs COMING SOON Section 5.7

3. Product favored reactions All of the preceding have been product-favored

reactions Removing ions from solution to form a new product

(solid, gas, water, etc.)

1. Examples and categories Uncombined Elements

Fe(s) + Cu+2(aq) Fe+2

(aq) + Cu(s)

Oxygen or Halogens PbO + V+3 + H2O PbO2 + VO + H+

2. Species which loses electrons is said to be: Oxidized

3. Species which gains electrons is said to be: Reduced

Remember… LEO the Lion Says---GER! -or- OIL RIG

“LEO” the Lion

GER!!!!!!

“OIL RIG”

4. Half-reaction examples

Fe(s) + Cu+2(aq) Fe+2

(aq) + Cu(s)

Fe Fe+2 + 2e- OXIDATION

Cu+2 + 2e- Cu REDUCTION

5. Rules for determining oxidation numbers a. Neutral species (isolated atom, molecule, formula

unit) The total of all oxidation numbers is 0. Ex. Fe = 0 Cl2 = 0 C6H12O6 = 0 NaCl = 0

b. Monatomic ions, Grp 1A, 2A, 7A … +1, +2, +3, skip -3, -2, -1 (usually)

c. Oxygen Most of the time, oxidation # = -2 Exception-Peroxides Oxidation # = -1 (O bonded to O

O2-2)

5.7 OXIDATION-REDUCTION REACTIONS

d. Sum of oxidation numbers for neutral and charged species: Neutral Species- total of all oxidation #’s = 0

Charged Species- total of all oxidation #’s = total charge

Ex. NH4+ = +1 NO3

-1 = -1 PO4-3 = -3 Cr+3 = +3

Ex5.6 Assign oxidation numbers for each element in the following species:

CaC2O4 Cr2O72-

N2ON2O4

ClO1- ClO41-

HAsO42- K2O

HCO31- MgSO4

HIO3 Na2MoO4

NO21- NH4

1+

S2O32- TeF8

2-

6. REDUCTION IS…

7. OXIDATION IS…

1. Write all reactants and products EXCEPT H+ and H2O in the form of separate oxidation and reduction half-reactions.

2. For each half-rxn, adjust coefficients for all atoms except H and O.

3. For each half-rxn, if the number of oxygen atoms differs between the left and right sides of the equation, add water molecules to the side needing more oxygen atoms.

4. For each half-rxn, if the number of hydrogen atoms differs between the left and right sides, at H+ to the side need hydrogen atoms.

5. For each half-rxn, add the number of e- need to balance the charge. (Oxidation rxn will have e- on the right and the Reduction will have e- on the left).

6. Multiply each half-rxn by the minimum factor required to equalize the number of e- in each half-rxn.

7.Add the equations for the half-rxns together, canceling e- and excess water molecules or hydrogen ions.

8. IF DONE IN BASIC SOLUTION: Add OH- ions to BOTH sides of the final equation to convert

ALL of the H+ to water Cancel any newly formed H2O molecules if necessary

See Steps on Previous Slide for Explanations!! Ex. PbO + V+3 + H2O PbO2 + VO + H+

Step 1: PbO PbO2 V+3 VO

Step 2: (all atoms except O & H are already balanced) PbO PbO2 V+3 VO

Step 3: H2O + PbO PbO2 H2O + V+3 VO

Step 4: H2O + PbO PbO2 + 2H+

H2O + V+3 VO + 2H+

PbO + V+3 + H2O PbO2 + VO + H+

Step 5: H2O + PbO PbO2 + 2H+ + 2e- 1 e- + H2O + V+3 VO + 2H+

Step 6: H2O + PbO PbO2 + 2H+ + 2e- 2e- + 2H2O + 2V+3 2VO + 4H+

Step 7: H2O + PbO + 2e- + 2H2O + 2V+3 PbO2 + 2H+ + 2e- + 2VO + 4H+

FINAL: PbO + 3H2O + 2V+3 PbO2 + 2VO + 6H+

9. Ex5.7 Balance in acidic solution:

MnO41- + Fe2+ Mn2+ + Fe3+

Cr2O72- + H2SO3 Cr3+ + SO4

2-

10. Ex5.8 Balance in Basic Solution

I1- + MnO41- I2 + MnO2

S2O32- + I2 SO4

2- + I1-

1. Molarity – definition, equation, and notation (M and [ ]) Unit of concentration for solutions

Molarity = M = mol-solute

L-soln

[ ] = molar concentrationEx. [Mg+2] [NaCl]

Ex5.9 What is the molarity of a solution made by dissolving 20.0g of NaCl(s) in enough water to make 1.00L of solution? In enough water to make 300.mL of solution?

Ex5.10 How many liters of 0.430M solution of Na2SO4 can be made starting with 10.0g of solid?

Ex5.11 How many grams of solid solute can be recrystallized from 10.0mL of 0.020M CuCl2?

5. Dissolution equations and molarity Concentration of compound is related to

concentration of ions via equation stoich!

Ex. 0.500 M MgCl2

Ex5.12 State the concentration of each ion in a. 0.250M Na2CO3 b. 0.023M barium phosphate c. 0.380M glucose

7. Dilution problems and equation Start with conc. soln. dilute with H2O to lower

conc.

M1V1 = M2V2

8. Ex5.13 What is the new concentration of a solution made by adding 150.mL of water to 23.0mL of a 2.50M solution of NaCl?

9. Ex5.14 A student starts with100mL of a 0.500M solution of HCl. How much water needs to be added to obtain a solution of HCl with a concentration of 0.0250M. What is the final volume of the solution?

1. pH scale – definitions and diagram pH scale: ranges from 0 (acid) to 14 (base); measures

“the Power of Hydrogen”

pH = -log [H+]

[H3O+] = 10-pH

Ex5.15 State the pH of the following solutions with given hydrogen ion concentrations:

a. [H+] = 0.0230M b. [H+] = 0.000560M c. [H+] = 3.00M

Ex5.16 What is the hydrogen ion concentration for each of the following solutions with given pH?

a. pH = 4.30 b. pH = 11.2 c. pH = 6.30

Precipitation Reaction Ex5.17 A precipitate forms when solutions of silver nitrate and scandium (III) chloride are mixed. What volume of 0.0385M scandium (III) chloride is needed to react completely with 22.00mL of 0.130M silver nitrate? What is the mass of the precipitate formed?

Precipitation Reaction Ex5.18 What is the mass percent of NaCl in a mixture of sodium chloride and sodium nitrate if a 0.9056g sample of the mixture yields 0.9372g of AgCl(s) when allowed to react with excess AgNO3(aq)?

Acid Base Reaction Ex5.19 How many milliliters of 0.0195M HCl are required to titrate 10.00mL of 0.0116M Ca(OH)2?

Acid Base Reaction Ex5.20 A Vitamin C capsule (ascorbic acid) is analyzed by titrating is with 0.250M sodium hydroxide. It is found that 10.30mL of the base is required to react with a capsule with a mass of 0.518g. What is the percentage of vitamin C, C8H8O6, in the capsule if the acid and base react in a 1 to 1 molar ratio?

Gas Forming Reaction Ex5.21 If 38.55mL of HCl are used to react with 2.150g of Na2CO3, what is the molarity of the HCl solution?

Redox Reaction Ex5.22 To analyze an iron containing compound, you convert all the iron to Fe2+ in aqueous solution and then titrate the solution with a known concentration of KMnO4 according to the following balanced equation:

MnO41-(aq) + 5Fe2+(aq) + 8H1+(aq) Mn2+(aq) + 5Fe3+(aq) + 4H2O(l)

A 0.598g sample of the iron containing compound requires 22.25mL of 0.0123M KMnO4 for titration to the equivalence point. What is the mass percent of iron in the sample?

#1 – 21, 23, 25, 27, 29, 105, 107, 117 Predicting Solubilities

#2 – 31, 33, 35, 37, 39, 41, 43 Net Ionic Equations & Acid-Base

#3 – 45, 47, 49, 95, 97, 53, 55, 57 Gas forming and Predicting Products

#4 – Worksheet – Balancing Redox Equations #5 – 59, 61, 63, 64, 67, 69, 71, 113

Molar concentrations, preparing solutions #6 – 73, 75, 77, 109, 87, 89, 91, 93, 101

pH & Acid-Base Titration #7 – 79, 83, 99, 1-3, 111, 115, 125

Reaction Stoichiometry