Chemistry Study Guide 2005



Chemistry Final Study Guide

25 Multiple choice and one free response question. You will be given the Periodic Table of ions and the formula sheet.

Chapter – 7/8 [omit from pg 258, section 8.4 and pg. 268]

1. Define: cation, chemical bond, anion, ionic bond, formula unit, polyatomic ion, noble gas, covalent bond, endothermic, exothermic, oxyacid, diatomic molecule

2. When naming compounds what ion comes first?

3. When are Roman numerals used in naming a compound?

4. When naming a compound that has a transition element in the formula, how does one know what number will represent the Roman numeral in the formula?

5. When is –ide ending used?

6. Where do subscripts go?

7. When are –ate and –ite used as endings?

8. Be able to be given two elements and then from their oxidations numbers, create a formula. Go back into your notes to find a couple of examples to insert into your study guide.

9. Can you write a formula using polyatomic ions?

10. What is the criss cross method and what is it used for?

11. Why must the positive charges be equal to the negative charges in an ionic compound?

12. The octet rule says that all elements want to be like what group?

13. Atoms are __ in a covalent bond.

14. List the 7 diatomic molecules. Memorize them

15. How many electrons are in a single, double and triple bond?

16. Write down the three rules for naming acids, and then memorize them. Give an example of each.

17. Memorize the prefixes that go with naming covalent compounds. Given at least 5 examples from your notes.

18. Know all properties of metals.

19. Know the concept of the electron sea model, what it is made up of, and the charges of it’s components.

20. What are the steps to drawing a lewis structure for a polyatomic ion?

21. How can I tell the difference between a molecular compound and an ionic compound just by looking at the formula or name?

22. What is the smallest particle which represents each type of compound?

23. How do I identify an acid from the chemical formula?

24. Properties of ionic, covalent(molecular) and metallic solids.

Chapter 7/8 practice:

Name the following:

1) C2H6 8) MgCrO4 15) CCl4

2) Fe2(CO3)3 9) Sr(OH)2 16) H2CO3

3) HClO 10) LiC2H3O2 17) Cu3N2

4) CaSO4 11) N2O3 18) Fe3(PO4)2

5) NaI 12) HI 19) H2SO3

6) (NH4)3P 13) N5S7 20) SiBr4

7) KHCO3 14) KMNO4

Write formulas for the following:

21) Carbonic Acid 31) Lithium Phosphate

22) nitrogen trichloride 32) silicon dioxide

23) Copper (II) phosphide 33) sodium oxalate

24) Iron (II) acetate 34) copper (I) fluoride

25) dicarbon pentoxide 35) calcium nitrite

26) aluminum sulfide 36) potassium perchlorate

27) Potassium perchlorate 37) tetraphosphorus pentabromide

28) Barium sulfate 38) sodium bicarbonate

29) ammonium cyanide 39) sulfur tetrachloride

30) hydrobromic Acid 40) phosphoric acid

Chapter – 9

1. Define: Reactant, coefficient, product, precipitate, aqueous solution, Combustion, single-replacement, synthesis, decomposition, word equation, skeleton equation

2. After the definition of each reaction type, give an example that demonstrates this.

3. When is it important to know which elements are diatomic molecules when writing a chemical equation?

4. Why are coefficients used?

5. Be familiar with the activity series of metals and what chemical reaction this series goes with.

6. Be able to predict the products when given the type of reaction.

7. Solubility rules: Which four never make solids?

8. Products of a carbonate and acid

9. Products of acid + base (what is this type of reaction?)

Chapter – 9 Practice:

Write and Balance the following. Identify the type of reaction if not stated.

1. If calcium carbonate is heated strongly, carbon dioxide gas and solid calcium oxide is formed.

2. Liquid propane gas (C3H8) is used for cooking is burned.

3. Solid ammonium carbonate is used as an active ingredient for smelling salts. When solid ammonium carbonate is added to a solution of hydrochloric acid it undergoes a specific type of double replacement reaction.

4. Phosphorus trichloride is used in the manufacture of certain pesticides, and may be synthesized by the direct combination of elemental phosphorus solid and chlorine gas.

5. When a solution of copper (II) oxide is boiled in an aqueous solution of sulfuric acid it undergoes a double replacement reaction.

6. Check to see if the equation is balanced. Write the type of reaction

for each of the equations listed below. If more than one type

applies, please put all types.

1. H2SO4 + KOH ( K2SO4 + H2O

2. Rb + Br2 ( RbBr

3. CaO + SiO2 ( CaSiO3

4. CH4 + O2 ( CO2 + H2O

5. Pb + HBr ( PbBr2 + H2

6. N2O5 + H2O ( HNO3

7. KClO3 ( KCl + O2

Chapter – 10

1. Define: mole, Avogadro’s number, Empirical formula, hydrate, molecular formula, conversion factor, formula unit.

2. Be able to calculate the molar mass of a compound.

3. What is the SI base unit for molar mass?

4. How do chemists’ use the mole?

5. How do I go from moles of an element to grams of an element? What are the conversion factors used?

6. How do I go from particles of an element to moles of an element and back again? What are the conversion factors used?

7. What is Avogadro’s number and how does it relate to the mole?

8. What are the 3 types of particles and what do they represent?

9. Show a math problem of each type from Chapter 10. Show all steps and make side notes on how to do each problem. This is a very important section in your study guide.

10. What are the steps to writing an empirical formula? Write down an example from your notes in your study guide for reference.

11. What are the steps to writing a molecular formula? Write down an example from your notes in your study guide for reference.

12. What is the difference between a molecular formula and an empirical formula?

13. Review hydrates and steps for calculations

Chapter - 10 Practice:

|Substance |Type of Particle |

|1. Potassium carbonate | |

|2. Carbon dioxide | |

|3. Calcium | |

|4. Boron trichloride | |

|5. Sodium nitrate | |

6. How many molecules of sucrose are in 3.6 moles of sucrose? (C12H22O11)

7. Calculate the grams of 6.25 mol of magnesium.

8. How many moles of lithium are in 61.3 grams?

9. Calculate the number of moles that are in 7.89 x 1016 atoms of fluorine.

10.In the formula for aluminum sulfate, Al2(SO4)3:

a) How many moles of aluminum are present? _____

b) How many moles of sulfur are present? _____

c) How many moles of oxygen are present? _____

11.How many moles of each element is in 0.550 mol K2CrO4?

a) Potassium

b) Chromium

c) Oxygen

12.What is the percent composition of beryllium oxide?

%Be __________

%O __________

13. What are the empirical and molecular formula of a compound that is composed of 83.64% tin and 16.36% phosphorus? The compound’s molar mass is 567.80 g/mol.

14. A compound is composed of 82.68% mercury and 17.32% nitrogen. It has a molar mass of 485.20 g/mol. Calculate the empirical and molecular formula of the compound.

15. A 1.628 g sample of a hydrate of magnesium iodide is heated until its mass is reduced to 1.072 g. Assume all the water has been removed. What is the formula of the hydrate?

Chapter – 11

1. Be able to develop mole ratios from a balanced equation and use the fencepost method to convert between species in the equation.

2. Be able to solve for limiting reactant and remaining excess reactant.

3. Be able to calculate the percent yield.

NOTE: Since the average for this test was an A, it is assumed that you don’t need extra practice. However, if you do, you can print a blank copy of the practice packet and check your answers.

Chapter – 12/13

Review basic concepts associated with the gas laws. Know what the variables are involved, whether they are directly related or inversely related and what the constants are. Be able to match the names of the laws with the formulas and use all the formulas. Remember the molar volume at STP and use it to convert via stoichiometry. Be able to use the ideal gas law and stoichiometry to get information from a chemical reaction.

If you want more practice,you can print a blank copy of the practice packet and check your answers.

Chapter – 14 [omit 14.4]

Don’t forget Molarity and how to use the formula to solve for moles or liters. Also, percent composition of solutions.

1. The solubility of a gas is 0.34 g/L at STP. What is its solubility at a pressure of 0.80 atm and the same temperature?

3. 1.56 g of a gas dissolves in 2.00 L of water at a pressure of 1.75 atm. At what pressure will

2.00 g of the gas dissolve in 2.00 L of water if the temperature remains constant?

5. A 500.0 g-sample of aqueous hydrogen peroxide (H2O2) contains 31.50% H2O2 by mass.

a. Find the mass of hydrogen peroxide in the solution.

b. Find the mass of water in the solution.

7. An aqueous solution of methanol is 45.0% methanol by volume.

a. Find the volume of methanol in a 250.0-mL sample of the solution.

b. Find the volume of water in this sample of the solution.

9. Calculate the molarity of 0.205 L of a solution that contains 156.5 g of sucrose (C12H22O11).

11. What mass of ammonium chloride (NH4Cl) would you use to prepare 85.0 mL of a 1.20M solution NH4Cl?

13. A 22.0-mL sample of 12M H2SO4 is diluted to a volume of 1200.0 mL. What is the molarity of the diluted solution?

Chapter 18 Acids & Bases, pH and pOH

Know the vocabulary and how to do the calculations from the homework worksheets. You are only responsible for the parts assigned on the website.

Write balanced chemical equations for each of the following reactions that involve acids and bases.

1. aluminum and hydrochloric acid

2. nitric acid and sodium carbonate

3. potassium hydroxide and sulfuric acid

A solution has a [OH−] of 3.6×10−7M.

9. What is its [H+]?

10. What is its pH?

11. What is its pOH?

A solution has a pH of 5.79.

15. What is its pOH?

16. What is its [H+]?

17. What is its [OH−]?

18. What is the pH of a 0.50M solution of HCl, a strong acid?

19. What is the pH of a 1.5×10−3M solution of NaOH? [pic]

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