Chemistry Final Exam Review Guide
Unit 4: Chemical Bonds
Objectives:
1. Explain the connection between electronegativity and the formation of an ionic or covalent bond.
2. Differentiate between single, double, and triple bonds.
3. Explain how the VSEPR theory impacts how you draw Lewis structures for molecular compounds.
4. Create a summary chart to distinguish between the types of intermolecular forces.
5. Be able to identify a bond or a compound as ionic, covalent or metallic based on a picture, description, difference in electronegativity, or example.
6. Be able to illustrate the formation of ionic and covalent bonds using Lewis structures.
7. Be able to apply nomenclature rules to determine the chemical formula and the name of a covalent or ionic compound, including those with polyatomic ions and/or transition metals.
8. Be able to predict the type of intermolecular force(s) at work when given a picture, description, model or example.
9. Be able to rank substances based on the strength of the intermolecular forces at work when given a description, model, data table, or example.
Practice:
10. Name the following compounds:
1. (NH4)2S
2. CuCO3
11. Write the chemical formula for the following compounds:
1. Magnesium sulfate
2. Gold (II) chloride
12. Draw the Lewis structure for PCl3. Then write its name.
13. Name the following compounds:
1. P2S3
2. BBr3
14. Write the chemical formula for the following compounds:
1. Dinitrogen monosulfide
2. Carbon tetrahydride
15. Draw the Lewis structure for each molecular compound below. Then list the number of electron domains, and identify if any are lone pairs. Then predict the molecule’s shape:
1. H2O
2. HCN
3. PH3
4. CH4
Vocabulary:
Chemical bond
Molecule
Ionic bond
Polyatomic ion
Transition metal
Metallic bond
Covalent bond
Electron domain
Lone pairs
Bonding pairs
Polar molecules
Nonpolar molecules
Intermolecular forces
Unit 5: Chemical Reactions
Objectives:
1. Explain the Law of Conservation of Mass and how it relates to chemical reactions.
2. Be able to label, interpret, and write equations for chemical reactions with correct notation.
3. Be able to balance chemical reactions according to the Law of Conservation of Mass.
4. When given a written description of a reaction, be able to write the balanced chemical reaction for it.
Practice:
5. Write out each chemical reaction described below. Then balance and classify:
1. Solid calcium hydroxide breaks down into solid calcium oxide and water vapor.
2. Solid aluminum and aqueous lead (II) nitrate react to form solid lead and aqueous aluminum nitrate.
6. Predict the products for the chemical reactions below. Then, balance the equations, as needed:
1. Combustion: C3H8 + O2 →
2. Synthesis: Na + Cl2 →
3. Decomposition: Al2O3 →
4. Single: AlBr3 + K →
5. Double: AgNO3 + CaCl2 →
7. Consider the following reaction: N2 (g) + 3H2 (g) 2NH3(g) + heat
1. Is the reaction endothermic or exothermic? Explain how you know.
2. Predict what would happen to the reaction once it has reached equilibrium if each of the following conditions is changed:
1. N2 is added; N2 is removed
2. H2 is added; H2 is removed
3. NH3 is added; NH3 is removed
4. Heat is added
5. Temperature is decreased
6. Pressure is increased; Pressure is decreased
7. Volume is decreased; Volume is increased
Vocabulary:
Chemical equilibrium
Le Chatelier’s Principle
Unit 6: Stoichiometry
Objectives:
1. Explain why the amount of product formed from a reaction is always determined by the limiting reactant.
2. Be able to calculate the amount of moles or mass of a reactant or product when given the amount of moles or mass for another reactant or product (meaning you can do all types of stoichiometry calculations, including mol to mol, mol to g, g to mol, and g to g).
3. Be able to determine the limiting and excess reactant when given starting amounts of reactants.
4. Be able to determine the amount left unused from an excess reactant.
5. Be able to determine how much product can be made when given the amount of moles or masses of the reactants.
6. Be able to calculate percent yield when given the actual yield and quantity of a reactant.
Practice:
7. Answer the following questions about this reaction: Be + HCl → BeCl2 + H2
1. Balance the equation
2. Classify the reaction.
3. Identify the limiting reactant if you begin with 0.275 mol Be and 0.518 mol HCl.
4. How much of the excess reactant, in grams, will be leftover?
5. How much of each product, in grams, will be formed?
8. Write the balanced equation for the double replacement reaction of aqueous calcium chloride with aqueous sodium carbonate to produce solid calcium carbonate and aqueous sodium chloride.
1. Find the theoretical yield of sodium chloride if you start with 12.62 g of sodium carbonate.
2. Calculate the percent yield if 11.38 g are actually produced in the lab.
Vocabulary:
Mole
Molar mass
Limiting reactant
Excess reactant
Actual yield
Theoretical yield
Percent yield
Unit 7: Acids and Bases
Objectives:
1. Explain what an acid-base reaction is. Include a labeled example and an explanation of why it can be referred to as a neutralization reaction.
2. Differentiate between the different definitions of acids and bases covered. Provide examples of each.
3. Differentiate between ionization and dissociation.
4. Explain what makes an acid or base strong versus weak.
5. Summarize the relationship between polarity, percent dissociation, and the strength or weakness of an acid or base.
6. Sketch a graph that shows the concentration of substances at the end of a reaction with a strong base.
7. Sketch a graph that shows the concentration of substances at the end of a reaction with a weak base.
8. Summarize the relationship between pH and ion concentration.
9. Summarize the relationship between pOH and ion concentration.
10. Summarize the relationship between pH and pOH.
11. Be able to identify a substance as a strong or weak acid or base when given a model, description, or example.
Practice:
12. Find the pOH of a substance with a pH of 3.5.
Vocabulary:
Salt
Arrhenius acid
Arrhenius base
Bronsted-Lowry acid
Bronsted-Lowry base
Amphoteric
pH
pOH