Semester Exam Review Name: ____________________________ #______Chemistry Block: __________This packet contains practice problems, but is not meant as a comprehensive review. Please also reference previous test, quizzes, and reviews for more practice problems!Things to KnowUnit 1: Intro to Experimental Practice

Nature of Science Lab Safety Lab Equipment Average Density Percent Error Accuracy vs. Precision Scientific Notation Significant figures (measurement, counting,

calculations) Dimensional Analysis MSDS Graphing

Unit 2: Qualitative Chemistry & Moles Classification of Matter Separation Techniques Phase Changes Endothermic vs. Exothermic Chemical vs. Physical properties/changes Molarity Molar Conversions Molar Mass Percent Composition

Unit 3: Atoms & Molecules History of the Atomic Model Subatomic particles Bohr Models Lewis Dot Diagrams Cations vs. Anions Isotopes Average atomic mass Radioactivity & Nuclear Reactions Half life Empirical vs. Molecular Formulas Ionic Bonding Ionic Nomenclature Diatomic Molecules

Unit 4: Quantum Mechanics & Light The Periodic Table Names of groups (noble gases, halogens, etc.) Metals, Nonmetals, & Metalloids Electron configurations Orbital diagrams

Unit 11. What is the difference between a theory and a law?

Theory –

Law –

2. What is the difference between an observation and an inference?Observation –

Inference –

3. What is the relationship between significant figures and precision?

4. The volume of a gas is known to be 2.20 L. The following data were collected over 5 trials: 5.20 L, 5.20 L, 5.18 L, 5.20 L, and 5.20 L. Is the data precise? Is the data accurate?

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Use the graph to the left to answer Questions 5-7.5. Identify the independent and dependent variables in the graph.

IV –

DV –

6. Does the graph show a direct or an inverse relationship?

7. What is the temperature when the pressure is 100 mmHg? Convert that temperature to the SI unit for temperature.

8. Calculate the density of an object whose mass is 1.6 g and volume is 0.234 mL. Report your answer in g/mL using the proper number of significant figures!

9. A cube has a mass of 3.56g with a length of 33.3cm, width of 10.9cm, and a height of 0.22cm.a. Calculate the density. Report your answer in g/cm3 using the proper number of significant figures!

b. What is the % error if the actual density for the solid is 0.101 g/mL?

10. How many pounds (lbs) are in 12.69 mg of a sample? (1 kg = 2.2 lbs). Report your answer using the proper number of significant figures in scientific notation!

11. If an automobile is able to travel 254 mi on 11.2 gal of gasoline, what is the gas mileage in km/L? (1 in = 2.54 cm) (1 mi = 5280 ft) (1 US gallon = 3.79 L). Report your answer using the proper number of significant figures!

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12. Determine the measurements (using proper # of sig figs) for the following graduated cylinders (in mL):

___________________ ___________________ ___________________

13. Identify the following pieces of lab equipment.

___________________ _________________ _________________ _____________________

___________________ _________________ _________________ _____________________

___________________ _________________ _________________ _____________________

___________________ _________________ _________________ _____________________3

14. Convert the following units:a. 25 mL = _______ L

b. 950 g = _______ kg

c. 25 cm = ______ mm

15. Convert the following to scientific notation:a. 0.005 = _________________

b. 5050 = _________________

c. 0.00081 = _______________

16. Convert the following to standard notation:a. 1.5 x 103 = _______________

b. 3.75 x 10-2 = _____________

c. 6.2 x 105 = _______________

17. Count the number of significant figures in each:

a. 0.02 ____

b. 501 ____

c. 5,000 ____

d. 6051. ____

e. 0.1020 ____

18. Perform the following calculations with your answer in the correct number of significant figures:

a. 1.35 x 2.467 = _________________

b. 12.01 + 35.2 + 6 = _____________

c. 0.021 x 3.1 + 100.1 = ___________

Unit 2

1. Which substance represents a compound?a. C(s)b. Co(s)c. CO(g)d. O2(g)f

2. Two substances, A and Z, are to be identified. Substance A cannot be broken down by a chemical change. Substance Z can be broken down by a chemical change. What can be concluded about these substances?a. Both substances are elements.b. Both substances are compounds.c. Substance A is an element and

Substance Z is a compound.d. Substance A is a compound and

Substance Z is an element.

3. Label the following particle diagrams as an element, compound, or mixture.

4. When KCl(s) is dissolved in water, the resulting solution is classified as a…a. Heterogeneous compoundb. Homogeneous compoundc. Heterogeneous mixtured. Homogeneous mixture

5. Classify as a mixture of pure substance, then as homogeneous, heterogeneous, element, or compound. (only one box will be filled in for each)

Description Mixture (Homo or Hetero) Pure Substance (Element of Compound)1. Water (H2O)

2. Air

3. Trail mix

4. Oxygen (O2)

5. Iron Filings (Fe)

6. Brass (Cu mixed with Zn)

7. Sugar water solution

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6. Classify the following as physical (P) or chemical (C) properties.

a. H2O boils at 100oC _________

b. Gold tarnishing _________

c. Silver is fairly malleable _________

d. Copper has an orange color _________

e. A diamond weighs 0.25 g _________

f. Water has a density of 0.99 g/mL _________

g. Na reacts violently with water _________

h. Sr gives off a red flame in fire _________

7. Identify the following separation techniques or what physical property they take advantage of:Technique Property

Decanting

State / Particle Size

Evaporation

Distillation

8. How many grams are in 2.1 moles of sodium fluoride, NaF?

9. How many moles are in 98.3g of aluminum hydroxide, Al(OH)3?

10. How many molecules are there in 3.16 moles of sodium sulfate, Na2SO4?

11. How many grams are there in 7.50 x 1023 molecules of sulfuric acid, H2SO4?

12. Calculate the molar mass and percent composition of each element for the following compounds:a. Na2CO3 b. NH4NO3

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Technique Property

Polarity

Column Chromatography

Density

13. What is the molarity of 100. mL of solution containing 0.142 mol of NaOH?

14. How many grams of MgCl2 do you need to have 500 mL of solution with a concentration of 2.5 M?

15. You have a stock solution that is 6.0 M. You want to make 100. mL of a 1.25M solution. How much of the stock solution do you need to make your new solution?

16. You have 250 mL of a stock 3.25 M solution. You added 250 mL to the solution. What is the new concentration?

Use the diagram below to answer questions 17-19:

17. Draw a particle diagram for a sample in the –a. Solid phase (labeled A)b. Liquid phase (labeled B)c. Gas phase (labeled C)

18. Label each arrow with the appropriate phase change (ex: fusion/melting)19. Indicate which phase changes are endothermic and which are exothermic.20. Classify the following as a physical (P) or chemical (C) change.

a. Cutting a piece of ribbon _________

b. Mixing salt and water _________

c. Isopropanol burning _________

d. Hammering metal flat _________

e. Baking a cake _________

f. Crushing a tablet _________

g. Copper oxidizing green _________

h. Mixing egg, milk, and flour _________

i. Baking soda reacts with vinegar to form a

gas _________

CBA

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21. Label the following diagrams as either exothermic or endothermic:

_________________________ _________________________

Unit 31. The Greek philosopher Democritus used the term atomos. This word means _____________________.

2. Antoinne Lavoisier is famous for the Law of ______________________________________________.

3. Dalton’s five major points of his atomic theory are as follows:

a. Elements are made up of tiny particles called _______________.

b. All atoms of a given element are ______________________.

c. Atoms of different elements are ______________________.

d. Atoms of different elements can combine to form ________________________.

e. Atoms cannot be created or _____________________.

4. Thomson knew that cathode rays have a ________________ charge because they were attracted toward

a _______________________ charged electric plate. Thomson had discovered the _______________.

5. In Thomson’s plum pudding model of the atom, the pudding was the ___________________ charge and

the raisins were the ________________________.

6. Rutherford did a famous gold foil experiment using alpha particles. Alpha particles have a ___________

charge. Rutherford discovered the _______________.

7. Draw the two different models of the atom, so that the labels will be points to the correct parts of each model.

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8. Bohr proposed that the electrons surrounding the nucleus travel in _______________ of specific energy.9. What is the difference between a cation and an anion?

10. Complete the following table.Nuclear Notation

Atomic Number

Atomic Mass # Protons # Neutrons # Electrons Cation? Anion?

Neutral?

Li37 3 7 3 4 3 Neutral

K1930 + 18

35 20 Anion (-2)

28 31 Neutral

39 19 19

11. How are the isotopes of carbon-14 and carbon-12 different? How are they the same?

12. Magnesium occurs in nature in three isotopic forms:24Mg 78.7% abundance26Mg 11.2% abundance25Mg 10.1% abundance

Calculate the average atomic mass for magnesium.

13. State the number of valence electrons and draw the Lewis Dot diagrams for:

a. Sulfur b. Calcium c. Chlorine d. Arsenic8

14. List, by number, both the period and group of each of these elements.Element Period Group

Beryllium (Be)

Argon (Ar)

Lead (Pb)15. Complete the following nuclear decay equations and identify the type of radioactive decay

Equation Type of Decay

K1942 → Ca20

42 + γ00 +¿¿

Po84239 → He2

4 +¿¿

U92235 → Th90

231 +¿¿

16. 200 grams of an unknown radioactive element was sealed in a container. After 48 days, only 25 grams remained in the container. What is the half-life of the unknown element?

17. The half-life of polonium-209 is equal to 102 years. If you take 24 grams of this isotope and bury it in a sealed container, how much polonium-209 would remain after 408 years?

22. What is the empirical formula for a compound with the molecular formula C6H12O3? _______________

23. The empirical formula of a compound is NO2. Its molecular mass is 92 g/mol. What is the molecular formula?

24. Name the following ionic compounds from the formula.Formula Name

NaClBa(OH)2

FeBr3

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NH4NO3

25. Write the formulas for the following ionic compounds from the name.Name Formula

Potassium fluoride

Magnesium sulfate

Copper (II) chloride

Calcium nitrideUnit 4

1. Generally, where are metals located on the periodic table? Nonmetals?

2. List the names of each of the following groups:

a. Group 1 _____________________________________

b. Group 2 _____________________________________

c. Group 17 _____________________________________

d. Group 18 _____________________________________

3. Identify the elements with the following electron configurations.

a. ______________________________ 1s22s22p5

b. ______________________________ 1s22s22p63s23p64s23d6

c. ______________________________ 1s22s22p63s23p64s1

4. How many valence electrons are in the element with the electron configuration 1s22s22p3? ___________5. Identify the elements with the following orbital diagrams.

a. ______________________________

b. ______________________________

c. ______________________________

6. Identify whether Hund’s Rule, the Aufbau Principle, or the Pauli Exclusion Principle has been violated in the following orbital diagrams.

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a. __________________________________

b. __________________________________

7. Write the electron configuration and draw the orbital diagram for sodium (Na).

Electron configuration:

Orbital diagram:

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