Complete the following equations by writing a balanced molecular equation and then a net ionic equation.

1.___ RbOH + ___ HBr

2.___ Pb(NO3)2 + ___ NaCl

3.___ Na2CO3 + ___ HCl

4.___ AgNO3 + ___ CuSO4

5.___ AgF + ___ NiCl2

6.____HCl + ____K2SO3

7. ____K2S + ____HNO3

8. ___ NaOH + ___ CaBr2

Write balanced net ionic equations for the following reactions.

9. Aqueous solution of potassium chloride and silver nitrate are mixed

10. A solution of perchloric acid is added to a solution of potassium hydroxide

11. A solution of lead (II) nitrate is added to a solution of silver iodide

12. A solution of ammonium chloride is added to a solution of sodium hydroxide

13. Solutions of nitric acid and sodium sulfide are mixed.

Answer the following question

14. In an AP Chemistry laboratory, students were given two unlabeled beakers and told that one of the beakers contained 1.0 g of solid CaCO3 and the other contained 1.0 g AgNO3. They were told to devise an experiment to identify which compound was which. A student did so be adding 50 mL of deionized water to each beaker. Describe the student’s observations that allowed her to identify each compound.

AP Chemistry Unit 2 Types of Chemical Reactions

Problem Set 2: Oxidation Numbers and Single Replacement Reactions

1. Determine the oxidation numbers on each atom.

  1. CO2
  2. CO
  3. SO3
  4. Ca2+
  5. O3
  6. BF3
  7. PbCl2
  8. H2SO4
  9. HCO3-
  10. N2O5

2. Identify the element that is oxidized and the element that is reduced in each reaction. Justify your answers by identifying the elements that gainedand lost electrons.

  1. 2AgNO3(aq) + Cu(s)  Cu(NO3)2(aq) + Ag(s)
  1. Ca(s) + Cl2(g)  CaCl2(s)
  1. CuO(s) + H2(g)  Cu(s) + H2O(l)
  1. Cu(s) + 4HNO3(aq)  Cu(NO3)2 + 2NO2(s) + 2H2O(l)
  1. 3MnO(s) + 4Al(s)  3Mn(s) + 2Al2O3(s)

3. Complete the following equations by writing a balanced molecular equation and then a net ionic equation.

  1. ___MgO(s) + ___H2(g) 
  1. ___Cu(s) + ___AgNO3(aq) 
  1. ___Cl2(g) + ___KBr(aq) 
  1. ___Ni(s) + ___HCl(aq) 
  1. ___Mn(s) + ___CoCl2(aq) 

AP Chemistry Unit 2 Types of Chemical Reactions

Problem Set 3: Molarity

  1. A 9.98 g sample of glucose, C6H12O6, is dissolved in enough water to produce a 1395 mL solution. What is the molarity of the solution?
  1. How many grams of MgSO4 . 9H2O are needed to prepare 125 mL of 0.200 M magnesium sulfate?
  1. 251 mL of 0.450 M HCl is added to 455 mL water. What is the molarity of the final solution? (Assume the volumes are additive).
  1. Household ammonia used for cleaning contains about 10. g of NH3 in 100.ml of solution. What is the molarity of theNH3 in solution?
  1. The average female adult has about 16 g of sodium ions in her blood. Assuming a total blood volume of 5.0 L, what is the molarity of the Na+ in blood
  1. What is the molarity of each ion present in aqueous solutions prepared by dissolving 20.00 grams of the following compounds in water to make 4.50 L of solution?
  1. Cobalt(II)chloride
  2. Nickel(III)sulfate
  3. Sodium permanganate
  4. Iron(II)bromide
  1. How would you prepare from the solid and pure water
  1. 0.400 L of 0.155 M Sr(OH)3?
  2. 1.75 L of 0.333 M (NH4)2CO3?
  1. A reagent bottle is labeled 0.255 M K2SO4
  1. How many moles of K2SO4 are present in 25.0 mL of this solution?
  2. How many milliliters of this solution of required to supply 0.0600 mol of K2SO4?
  3. Assuming no volume change, how many grams of K2SO4 do you need to add to 1.50 L of this solution to obtain a 0.800 M solution of K2SO4?
  4. If 40.0 mL of the original solution are added to enough water to make 135 mL of solution, what is the molarity of the diluted solution?
  1. A student combines two solutions of KOH and determines the molarity of the resulting solution. He records the following data:

Solution I:30.00 mL of 0.125 M KOH

Solution II:40.00 mL of KOH

Solution I + Solution II70.0 mL of 0.203 M KOH

What is the molarity of solution II?

AP Chemistry Unit 2 Types of Chemical Reactions

Problem Set 4 Titrations and Gravimetric Analysis

  1. How many milliliters of 0.250 M KMnO4 are needed to deliver 0.00450 moles of KMnO4 in a titration?
  1. In an AP chemistry laboratory, students were given two unlabeled beakers and told that one of the beakers contained 1.0 g of solid CaCO3 and the other contained 1.0 g of AgNO3. They were told to devise an experiment to identify which compound was which. A student did so by adding 50 mL of distilled water to each beaker. Describe the students observations that allowed him to identify each compound.
  1. Three beakers labeled A, B, and C contain a weak acid H2X. The weak acid is titrated with 0.125 M NaOH. Assume the reaction to be

H2X(aq) + 2OH-(aq)  2H2O(l) + X2-(aq)

  1. Beaker A contains 25.00 mL of 0.316 M H2X. What volume of NaOH is required for complete neutralization?
  2. Beaker B contains 35.00 mL of a solution of H2X and requireds 28.74 mL of NaOH. What is the molarity of the H2X solution?
  3. Beaker C contains 0.124 g of H2X and 25.00 mL of water. To reach the equivalence point, 22.40 mL of NaOH are required. What is the molar mass of H2X?
  1. A 0.0500 M solution of potassium permanganate was used to titrate 250.0 mL of a platinum (II) chloride solution with an unknown concentration. The endpoint was reached after 26.87 mL of 0.0500 M KMnO4 were delivered.
  1. Write the balanced net ionic equation for the chemical reaction that occurred during this titration.
  2. How many moles of KMnO4 were delivered when the endpoint was reached?
  3. How many moles of platinum chloride were contained in the 250.0 mL sample?
  4. Calculate the experimentally determined molar concentration (M) of platinum (II) chloride in solution.
  1. In a gravimetric experiment 200. mL of 2.0 M copper (I) nitrate is mixed with 150. mL of 2.5 M sodium chloride. The mixture produces a precipitate.
  1. Identify the precipitate
  2. What is the limiting reactant? Justify your answer
  3. What is the maximum mass of precipitate that can be formed in this reaction?
  4. What is the percent yield if 31 g of precipitate is formed in the reaction?
  5. The percentage yield increases when the temperature of the solution is reduced. Explain why this is.
  1. A gravimetric experiment was performed to determine the number of moles of sulfate that were present in a solution of sodium sulfate. In this experiment, an excess of barium nitrate was added to the solution and a precipitate formed. The resulting mixture was filtered and the precipitate was dried. The dry mass of the precipitate was measure to be 8.642 g.
  1. Write the balanced complete ionic equation for the precipitation reaction
  2. Write balanced net ionic equation for the precipitation reaction.
  3. Calculate the number of moles of sulfate in the sample.

AP Chemistry Unit 2 Types of Chemical Reactions

Problem Set 5 Identifying Types of Chemical Reactions

Balance the following reactions and indicate which type of chemical reaction is being represented. When appropriate write a net ionic equation

1)____Na3PO4+ ____KOH  ____NaOH + ____K3PO4

2)____MgCl2 + ____Li2CO3____MgCO3 + ____LiCl

3)____C6H12 + ____O2 ____CO2 + ____H2O

4)____Pb + ____FeSO4 ____PbSO4 + ____Fe

5)____CaCO3 ____CaO + ____CO2

6)____P4 + ____O2____P2O3

7) ____RbNO3 + ____BeF2 ____Be(NO3)2 + ____RbF

8) ____AgNO3 + ____Cu  ____Cu(NO3)2 + ____Ag

9)____C3H6O + ____O2 ____CO2 + ____H2O _

10) ____C5H5 + ____Fe  ____Fe(C5H5)2

11)____SeCl6 + ____O2 ____SeO2 + Cl2

12) ____MgI2 + ____Mn(SO3)2 ____MgSO3 + ____MnI4

13)____O3  ____O. + ____O2

14) ____NO2 ____O2 + ____N2

15)____ Pb + ____ H3PO4 ____ H2 + ____ Pb3(PO4)2