Gen Chem 1

mod 4 practice problems

Gen Chem 1 Module 1 Module 2 Module 3 Module 4 Module 5 Module 6 Module 7

Concentration

Determine the molarity for each of the following solutions:
1. 0.444 mol of CoCl₂ in 0.654 L of solution
2. 98.0 g of phosphoric acid, H₃PO₄, in 1.00 L of solution
3. 0.2074 g of calcium hydroxide, Ca(OH)₂, in 40.00 mL of solution
  1. What is the molarity of Calcium cations in the above solution? What is the molarity of the hydroxide anions in the above solution?
4. 2.76 kg of CuSO₄•5H₂O in 1.45 L of solution
What is the mass of the solute in 0.500 L of 0.30 M glucose, C₆H₁₂O₆, used for intravenous injection?
Calculate the number of moles and the mass of the solute in each of the following solutions:
1. 2.00 L of 18.5 M H₂SO₄, concentrated sulfuric acid
2. 100.0 mL of 3.8×10⁻⁵MNaCN, the minimum lethal concentration of sodium cyanide in blood serum
What is the molarity of KMnO₄ in a solution of 0.0908 g of KMnO₄ in 0.500 L of solution?
What is the molarity of HCl if 35.23 mL of a solution of HCl contain 0.3366 g of HCl?
What volume of a 1.00-M Fe(NO₃)₃ solution can be diluted to prepare 1.00 L of a solution with a concentration of 0.250 M?
If 0.1718 L of a 0.3556-M C₃H₇OH solution is diluted to a concentration of 0.1222 M, what is the volume of the resulting solution?

Types of reactions

Classify the following as acid-base reactions or oxidation-reduction reactions:
1. Na₂O(aq)+2HCl(aq)→2NaCl(aq)+H₂O(g)
2. 2Na(s)+2HCl(aq)→2NaCl(aq)+H₂(g)
3. Mg(s)+Cl₂(g)→MgCl₂(s)
4. MgO(s)+2HCl(aq)→MgCl₂(aq)+H₂O(l)
5. K₃P(s)+2O₂(g)→K₃PO₄(s)
6. 3KOH(aq)+H₃PO₄(aq)→K₃PO₄(aq)+3H₂O(l)
Identify the atoms that are oxidized and reduced, the change in oxidation state for each, and the oxidizing and reducing agents in each of the following equations:
1. Mg(s)+NiCl₂(aq)→MgCl₂(aq)+Ni(s)
2. PCl₃(l)+Cl₂(g)→PCl₅(s)
3. C₂H₄(g)+3O₂(g)→2CO₂(g)+2H₂O(g)
Complete and balance the following acid-base equations:
1. HCl gas reacts with solid Ca(OH)₂(s).
2. A solution of Sr(OH)₂ is added to a solution of HNO₃.
Complete and balance the following oxidation-reduction reactions, which give the highest possible oxidation state for the oxidized atoms.
1. Al(s)+F₂(g)→
2. Al(s)+CuBr₂(aq)→
3. Ca(s)+H₂O(l)→ (products are a strong base and a diatomic gas)
Complete and balance the equations for the following acid-base neutralization reactions. If water is used as a solvent, write the reactants and products as aqueous ions.
1. Mg(OH)₂(s)+HClO₄(aq)→
2. Sr(OH)₂(s)+H₂SO₄(l)→
Write the molecular, total ionic, and net ionic equations for the following reactions:
1. Ca(OH)₂(aq)+HC₂H₃O₂(aq)→
2. H₃PO₄(aq)+CaCl₂(aq)→

Balancing Chemical Equations

Balance the following equations with the smallest, whole numbers possible. Classify each equation as Synthesis, Decomposition, Single Displacement (replacement), Double Displacement (replacement). Also, note if any of the reactions are Acid Base or Redox reactions.

Fe_(s) + S_(s) → Fe₂S_3(s)

Ca(OH)₂ +Na₂CO₃→ CaCO₃ + NaOH

KMnO₄ + HCl → KCl + MnCl₂ + H₂O + Cl₂

K₄Fe(CN)₆+ H₂SO₄ + H₂O → K₂SO₄ + FeSO₄ + (NH₄)₂SO₄ + CO

C₆H₅COOH + O₂→ CO₂ + H₂O (There is an error in the video for this problem - I didn't multiply the product by 2 at the end)

Calcium hydroxide reacts with carbon dioxide to form calcium carbonate and water.

Phosphoric acid reacts with sodium hydroxide to form sodium phosphate and water.

Aqueous mercury (II) oxide reacts with gaseous chlorine to form solid mercury (i) chloride and gaseous oxygen.

Solubility practice

Solubility rules (from chem.libretext.org)

The following are the solubility rules for common ionic solids. If there two rules appear to contradict each other, the preceding rule takes precedence.

Salts containing Group I elements (Li⁺, Na⁺, K⁺, Cs⁺, Rb⁺) are soluble . There are few exceptions to this rule. Salts containing the ammonium ion (NH₄⁺) are also soluble.
Salts containing nitrate ion (NO₃^-) are generally soluble.
Salts containing Cl ^-, Br ^-, or I ^- are generally soluble. Important exceptions to this rule are halide salts of Ag⁺, Pb²⁺, and (Hg₂)²⁺. Thus, AgCl, PbBr₂, and Hg₂Cl₂ are insoluble.
Most silver salts are insoluble. AgNO₃ and Ag(C₂H₃O₂) are common soluble salts of silver; virtually all others are insoluble.
Most sulfate salts are soluble. Important exceptions to this rule include CaSO₄, BaSO₄, PbSO₄, Ag₂SO₄ and SrSO₄ .
Most hydroxide salts are only slightly soluble. Hydroxide salts of Group I elements are soluble. Hydroxide salts of Group II elements (Ca, Sr, and Ba) are slightly soluble. Hydroxide salts of transition metals and Al³⁺ are insoluble. Thus, Fe(OH)₃, Al(OH)₃, Co(OH)₂ are not soluble.
Most sulfides of transition metals are highly insoluble, including CdS, FeS, ZnS, and Ag₂S. Arsenic, antimony, bismuth, and lead sulfides are also insoluble.
Carbonates are frequently insoluble. Group II carbonates (CaCO₃, SrCO₃, and BaCO₃) are insoluble, as are FeCO₃ and PbCO₃.
Chromates are frequently insoluble. Examples include PbCrO₄ and BaCrO₄.
Phosphates such as Ca₃(PO₄)₂and Ag₃PO₄ are frequently insoluble.
Fluorides such as BaF₂, MgF₂, and PbF₂ are frequently insoluble.

Predict the products of the following reactions. Use the solubility rules to determine if a precipitate will be produced. Indicate that precipitate with the appropriate phase tag. If a precipitate is produced, write the net ionic reaction. Be sure the reactions are balanced.