AP Chemistry: Course Introduction Sheet



AP Chemistry Unit I Targets Chapters 1-3

By the end of this unit you should be able to . . .

chapter 1

1. define and provide examples for each of the following terms: physical property, chemical property, physical change, chemical change, intensive property, extensive property element, compound, mixture

2. differentiate between the three states of matter.

3. list the commonly used metric prefixes and their meanings.

4. determine the number of significant digits in a measured quantity and determine the appropriate number significant digits in a calculation.

5. convert temperatures between Celsius and Kelvin.

6. perform calculations involving density.

7. convert between units by using dimensional analysis.

chapter 2

1. differentiate between protons, neutrons, and electrons in terms of charge, mass and location in an atom.

2. I can determine the number of protons, neutrons and electrons in isotopes and in ions.

3. describe the works of John Dalton, J.J. Thomson (cathode ray tube), Robert Millikan (Oil Drop Experiment) and Ernst Rutherford (Gold Foil Experiment).

4. use the periodic table to predict the charges of monatomic ions.

5. use the periodic table to predict whether an element is a metal, nonmetal or a metalloid.

6. write the names and formulas of ionic compounds, molecular/covalent compounds and acids.

7. calculate the atomic weight of an element given the abundances and masses of its isotopes.

8. distinguish between empirical formulas, molecular formulas and structural formulas.

chapter 3

1. Predict the products for and write balanced equations for the following types of reactions: combustion, decomposition, sysnthesis (called combination reactions), single displacement and double displacement reactions.

2. interconvert between the number of moles and mass of a substance. I can also use Avogadro’s number and molar mass to calculate the number of particles (atoms, molecules or formula units) making up a substance.

3. calculate the percentage composition of a compound by mass.

4. calculate the empirical formula of a compound, having been given either:

a) mass or % composition, or

b) the mass of CO2 and H2O produced by combustion.

5. calculate the molecular formula, having been given the empirical formula and the molecular weight.

6. use stoichiometry to solve problems involving chemical reactions.

7. determine the limiting reactant in a reaction and determine the amount of excess reactant left over from a reaction.

8. calculate the theoretical and actual yields of a chemical reaction when given the appropriate data.

Target 1. I can define and provide examples for each of the following terms: physical property, chemical property, physical change, chemical change, intensive property, extensive property element, compound, mixture

physical property –

chemical property –

physical change –

chemical change –

intensive property –

extensive property -

element –

compound –

mixture –

Describe the following mixture separation techniques:

Decant - Filtration -

Distillation –

Chromatography –

Target 2: I can differentiate between the three states of matter.

|State |Picture |Movement |Shape |Volume |Compression |

| | | | | | |

|Gas | | | | | |

| | | | | | |

|Liquid | | | | | |

| | | | | | |

|Solid | | | | | |

Target 3: I can list the commonly used metric prefixes and their meanings.

Memorize highlighted

|Unit |Giga |Mega |

|Mg-24 |78.70 |23.985 |

|Mg-25 |10.13 |24.985 |

|Mg-26 |11.17 |25.983 |

Example 2: Gallium has two isotopes: Ga-69 (mass of 68.9255 amu) and Ga-70 (mass of 70.9247). If the average atomic mass of gallium is 69.723 amu, what is the relative abundance (in %) for each of the two isotopes of gallium?

Practice:

1. What is an isotope?

2. What is the atomic weight of Lithium if Lithium has 2 isotopes: Li-6 (7.5%) and Li-7 (92.5%)?

3. Why doesn’t the periodic table use a normal average (mean)? (In other words why doesn’t C-12, C-13 and C-14 just average out to 13?)

4. What is the difference between a normal average (mean) and a weighted average?

5. Chlorine has 2 major isotopes, Cl-35 and Cl-37. What is the %abundance of each isotope?

6. Silver has 2 major isotopes, Ag-107 and Ag- 109. What is the % abundance of each isotope?

7. Copper has 2 major isotopes, Cu-63 and Cu-65. Without doing any major calculations, which isotope is more abundant and why?

Target 8: I can distinguish between empirical formulas, molecular formulas and structural formulas.

Empirical Formula - the simplest whole number ratio of atoms of each element present in a compound.

Lets look at ethene, C2H4…

Ethene can undergo polymerization and become a long strand of linked ethene molecules. The empirical formula would be C2H4 but the polymer may have a molecular formula several hundred times larger than the “building block” of the substance.

Molecular Formula - indicates the number of atoms of each element found in each discrete molecule of that compound.

An example of polyethene: C2H4 x 150 = C300H600

Structural Formula - a graphical representation of the molecular structure, showing how the atoms are arranged. Used extensivily in organic chemistry.

Ethene would look like this:

[pic]

Polyethene would look like this

[pic]

AP Chemistry Name ________________________

Review Chapter 1 and Chapter 2

1. _______ is the state of matter which has a definite volume but no definite shape.

2. Another name for a homogeneous mixture is a(n) ________________.

3. Can the elements in a compound be separated by chemical means? by physical means?

4. List 2 intensive properties and 2 extensive properties.

5. The SI unit for mass is the _______________; the SI unit for length is the _________.

6. Which of the following is the longest distance? 1 km, 106 (m, or 10-6 Mm?

7. 32°F = ________________°C = ___________________ K

8. 17.2 cm + 204.8 mm = ________________ mm

9. Which of John Dalton’s postulates was incorrect about atomic theory? Why?

10. 1 meter = _________________ cm

11. Name an element which is:

a. a gas and also a halogen (assume room Temp)

b. a metalloid

c. an alkaline earth metal

12. Write the empirical formula for glucose.

13. Write the formula for:

a. a polyatomic anion

b. a monatomic cation

c. any molecular compound

d. magnesium cyanide

e. nickel (II) nitride

f. hyponitrous acid

14. Name each of the following:

a. MnO4-

b. Cr(OH)3

c. NBr3

d. NH4Cl

Target 1: I can predict the products for and write balanced equations for the following types of reactions: combustion, decomposition, sysnthesis (called combination reactions), single displacement and double displacement reactions.

1- ALKALI METALS WITH WATER

M (s) + H2O (l) ----------> MOH (aq) + H2(g)

Na(s) + H2O (l) ----------> NaOH (aq) + H2(g)

K(s) + H2O (l) ----------> KOH (aq) + H2(g)

2- COMBUSTION IN AIR

Hydrocarbon (fuel) + oxygen -------> CO2 + H2O

CH4(g) + O2(g) -----------> CO2(g) + H2O(l)

C3H8(g) + O2(g) -----------> CO2(g) + H2O(l)

CH3OH(l) + O2(g) -----------> CO2(g) + H2O(l)

3- COMBINATION REACTIONS - 2 or more reactants form one product

A + B -------> C (generic equation)

2 Mg(s) + O2(g) --------------> 2 MgO(s)

N2(g) + 3 H2(g) -----------> 2 NH3(g)

4- DECOMPOSITION REACTIONS - one reactant forms 2 or more products

A ----------> B + C (generic equation)

2 KClO3(s) -----------> 2 KCl(s) + 3 O2(g) ; chlorates decompose to form chloride salts and oxygen gas

2 NaN3(s) ----------> 2 Na(s) + 3 N2(g)

CaCO3(s) ---------> CaO(s) + CO2(g) ; metal carbonates decompose to form metal oxides and carbon dioxide gas

Target 2: I can interconvert between the number of moles and mass of a substance. I can also use Avogadro’s number and molar mass to calculate the number of particles (atoms, molecules or formula units) making up a substance.

Avogadro’s Number: A number equal to the number of atoms in exactly 12 grams of C-12. Experimentally we have found this number to be 6.022 X 1023.

Molar mass Avogadro’s #

GRAMS MOLES PARTICLES (atoms, ions, molecules)

Example: How many moles of CO2 are in 10.0 grams of CO2?

Practice 1: How many molecules of H2O are contained in 150. grams of water? SYW!

Practice 2: How many hydrogen atoms are there in 36.0 grams of sulfuric acid? SYW!

Formula weight = sum of atomic weights of the atoms in a formula; we can use the term molecular weight if the substance is a molecular formula or molar mass for either.

Example: Find the formula weight of calcium nitrate.

calcium nitrate is Ca(NO3)2

1 Ca atom = 1(40.1 amu) = 40.1 amu

2 N atoms = 2(14.0 amu) = 28.0 amu

6 O atoms = 6(16.0 amu) = 96.0 amu

___________

164.1 amu

Practice: Find the formula weight of ammonium phosphate.

Target 3: I can calculate the percentage composition of a compound by mass.

Percent Composition: % composition is the % by mass contributed by each element in the substance

(# of atoms of element)(atomic weight) X 100

molar mass of the compound

Example: Find the percentage of nitrogen in calcium nitrate. (Answer is 17.1% N)

Practice: Find the percentage of water in copper sulfate pentahydrate.

Target 4: I can calculate the empirical formula of a compound, having been given either:

a) mass or % composition, or

b) the mass of CO2 and H2O produced by combustion.

mass % ------------> grams of each ---------------> moles of each --------------> empirical element element formula

Example: A compound is 62.58% C, 9.63 % H, and 27.79 % O by mass. Calculate its empirical formula.

Practice: What is the empirical formula of a compound which contains 25.9% nitrogen and 74.1% oxygen? (N2O5)

Determining the empirical formula of a compound from combustion analysis:

Example: Combustion of a 0.2000 gram sample of vitamin C (which contains C, H, and O) yields 0.2998 g of CO2 and 0.0819 g of H2O. Calculate the empirical formula of vitamin C.

Practice: Combustion analysis of toluene, a common organic solvent, gives 3.53mg of CO2 and 0.822mg of H2O. If the compound contains only carbon and hydrogen, what is its empirical formula?

Target 5: I can calculate the molecular formula, having been given the empirical formula and the molecular weight.

Example: A compound is found to contain 16.66 grams of carbon and 3.49 grams of hydrogen.

a. Find the empirical formula. (C2H5)

b. The molar mass is 87. Find the molecular formula: (C6H15)

Practice: Determine the molecular formula for a compound that contains 22.5 % Na, 30.4 % P and 47.1 % O if it has a molar mass of 306 g/mole.

Target 6: I can use stoichiometry to solve problems involving chemical reactions.

Stoichiometry is a process which uses balanced equations to solve problems.

Example: How many grams of O2 can be prepared from the decomposition of 4.50 grams of potassium chlorate?

Target 7: I can determine the limiting reactant in a reaction and determine the amount of excess reactant left over from a reaction.

Example: In order to bake 1 cake you need the following ingredients:

1 box of cake mix

1 egg

2 cups of flour

1/2 cup of sugar

3 tablespoons of water

Assume that you had 1 dozen eggs, 6 boxes of cake mix, 10 cups of flour, 10 cups of sugar, one quart of oil, and an unlimited supply of water. How many cakes can you make? What is the limiting “reagent”?

Example: A 2.00 gram piece of zinc is placed in 2.50 grams of aqueous silver nitrate.

a) Write the balanced equation.

b) Determine the limiting reagent.

c) How many grams of silver will be formed?

d) How many grams of the excess reactant will be left over?

Target 8: I can calculate the theoretical and actual yields of a chemical reaction when given the appropriate data.

Percent Yield Problems

Definitions:

theoretical yield = the quantity of product that should “theoretically” be produced (if all of the reactant is used up)

actual yield = the quantity of product that is actually produced when the reaction is done in the lab

percent yield = the relationship between the actual yield and the theoretical yield

percent yield = actual yield X 100

theoretical yield

Example: Fe2O3(s) + 3 CO(g) -------> 2 Fe(s) + 3 CO2(g)

If 150.0 grams of Fe2O3 reacts with excess CO, 87.9 grams of Fe is produced. Calculate the % yield.

Practice: Nitrogen gas reacts with hydrogen gas to make ammonia. 15.5 L of Nitrogen reacts at STP to make 30 L of ammonia. What is the % yield?

Law of Definite and Multiple Proportions Experimental Demonstration

Purpose

To experiemntally determine the emperical formula of a copper iodide compound. The compound will either be copper (I) iodide, CuI, or copper (II) iodide, CuI2.

Your instructor will demonstrate a reaction between copper and iodine in the hood. The procedure is as follows:

Procedure

1. Mass an empty beaker.

2. Clean the surface of a Cu strip with steel wool. Place the clean Cu strip into the empty beaker and determine the mass of both.

3. Bend the strip into a concave shape, and sprinkle a small amount of iodine onto it. Under a fume hood and using tongs, heat gently to react the iodine with the copper. The unreacted iodine will leave as a dense, purple gas. The copper iodide product will remain on the slip as a creamy white solid.

4. Carefully place the strip into the beaker and find the mass.

5. Remove the strip from the flask and remove the compound from the strip by scraping it off.

6. Place the cleaned Cu strip in the beaker and find the mass.

Data

mass of copper before . . . . . . . . . . . _______________ g

mass of copper & compound . . . . . . _______________ g

mass of copper after . . . . . . . . . . . . . _______________ g

Calculations

Using the data above, show by calculations whether the formula of the compound is CuI or CuI2.

Discussion Questions

Using complete sentences, explain how the ratio of copper to iodine would be altered if:

a) some of the copper iodide had been knocked off before being massed.

b) some of the copper iodide had been knocked off after being massed.

c) In some procedures, after the copper iodiide is scraped off, the copper strip is rinsed with water and then dried off. How would the copper to iodine ratio be altered if the strip were not completely dried?

d) some scientists dry the strip by heating it over a bunsen burner. If heated too long some dark copper (II) oxide can form on the surface. Explain how this would affect the Cu:I ratio?

AP Chemistry Ch. 3 Review Sheet Name_____________________

Directions: Answer the following questions on a separate sheet of paper. Show all of your work, label your final answers with correct units, and circle your final answers!

1. Predict the products for the following reactions and write a balanced equation for each:

a) The combustion of glucose.

b) The synthesis reaction between potassium and chlorine gas.

c) The decomposition of magnesium carbonate.

d) Reacting magnesium oxide and water.

e) Reacting sulfur trioxide and water.

f) The decomposition of sodium chlorate.

2. Below is a chart containing data for the three naturally occurring isotopes of Mg:

Isotope abundance (%) mass (u)

Mg-24 78.70 23.98504

Mg-25 10.13 24.98584

Mg-26 11.17 25.98259

Calculate the atomic mass of magnesium.

3. Calculate the percentage of oxygen (by mass) in nickel (II) acetate.

4. Assume you have 5.0 liters of water. Calculate the number of each of the following:

a) the number of grams of water.

b) the number of moles of water.

c) the number of molecules of water.

d) the number of hydrogen atoms in this sample of water.

5. Antifreeze is composed of 51.6 % oxygen, 9.70% hydrogen, and 38.7% carbon by mass. The molar mass of antifreeze is 62.1 g/mol. Calculate its empirical and molecular formulas.

6. Menthol, the substance we can smell in mentholated cough drops, is composed of C, H, and O. A 1.005-g sample of menthol is combusted, producing 0.2829 g CO2 and 0.1159 g of H2O. What is the empirical formula of menthol? If the compound has a molecular mass of 156 g/mol, what is its molecular formula?

7. When a mixture of 10.0 g of acetylene, C2H2, and 10.0 g of oxygen, O2, is ignited, the resultant combustion produces CO2 and H2O.

a) Write the balanced equation for this reaction.

b) Which reactant is the limiting reactant?

c) How many grams of C2H2, O2, CO2, and H2O are present after the reaction is complete?

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