Calculate the Concentration (In Mol/L) of Chloride Ions in Each

Ch 4 Problems

Calculate the concentration (in mol/L) of chloride ions in each: a. 1.00 M CaCl2 b. 21.2 g LiCl in 500.0 mL water c. b. 38.3 g Na3PO4 in 500.0 mL water

Calculate the potassium ion concentration when 40.0 mL of 2.0 M potassium chloride is added to 60.0 mL of 3.0 M potassium sulfate.

A stock solution is prepared by dissolving 40.0 g NaOH in 100.0 mL water. A 25.0 mL sample of this stock solution is added to 25.0 mL water. What is the concentration of the new solution?

For each of the following reactions, write the balanced molecular equation, complete ionic equation and the net ionic equation. (use Table 4.1 on p.105 to help you) a. lead nitrate is added to potassium hydroxide: b. silver nitrate is added to barium chloride:

A 1.42 g sample of a pure compound, with formula M2SO4, was dissolved in water and treated with excess aqueous calcium chloride resulting in the precipitation of all the sulfate ions as calcium sulfate. The precipitate was collected, dried and found to weigh 1.36 g. Determine the atomic mass and identity of M.

A 3.48 g sample of a pure compound, with formula M2SO4, was dissolved in water and treated with excess aqueous calcium chloride resulting in the precipitation of all the sulfate ions as calcium sulfate. The precipitate was collected, dried and found to weigh 2.72 g. Determine the atomic mass and identity of M.

Define the terms acid and base and write the balanced reaction of hydrochloric acid with barium hydroxide. Using the above reaction, what volume of 0.120 M HCl is needed to neutralize 50.0 mL of 0.101 M Ba(OH)2? Redox Rxns

**Rules for assigning oxidation states: 1. Elements = 0 2. The sum of oxidation states for a molecule or ion = the charge on molecule or ion 3. Group 1 Metals (Li, Na, K, …) = +1 4. Group 2 Metals (Be, Mg, Ca…) = +2 5. Fluorine = -1 6. Hydrogen = +1 2- 7. Oxygen = -2 (except for peroxides O2 = -1) 8. Halogens (Cl, Br, I) = -1

Assign oxidation states to each atom in the following:

+ NH4 KMnO4 Na2C2O4 Br2 H2SO4

**Definitions: Oxidation- increase in oxidation state (or loss of electrons) Reduction- decrease in oxidation state ( or gain of electrons)

The substance being oxidized is the reducing agent, substance being reduced is the oxidizing agent.

For the following reactions, identify the substance oxidized and reduced, the oxidizing and reducing agents and the number of electrons transferred:

2+ -1 Mg(s) + HCl(aq) → Mg(aq) + Cl(aq) + H2(g)

CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) **Rules for balancing redox reactions by the half-reaction method: 1. Write the oxidation and reduction half-reactions 2. For each half-reaction: a. balance all elements besides O and H b. balance O using H2O c. balance H using H+ d. balance the charge using electrons 3. Equalize the numbers of electrons in both half-reactions 4. Add the half-reactions and simplify by canceling anything present on both sides 5. If the reaction is in basic solution, add the number of OH- ions equal to the number of H+ ions to both sides + - a. H + OH forms H2O b. cancel the H2O that appears on both sides

Balance the following reactions in both acidic and basic conditions:

- 2+ 2+ 3+ MnO4 + Fe → Mn + Fe

2- - - CrI3 + Cl2 → CrO4 + IO4 + Cl

- 2+ MnO4 + CH3OH  Mn + HCO2H

- - Pb(OH)4 + ClO  PbO2 + Cl

- 2- 3+ - NO2 + Cr2O7  Cr + NO3

- H2O2 + Cl2O7  ClO2 + O2