Name _____________________________________ Hour___________

Chemistry Semester 1 Review

PART 1: Introduction to Chemistry DUE TUESDAY 1/17

1. Describe each part of scientific theory

2. Distinguish between a theory and a law

3. Define AND give an example of each of the following

a. Chemical change

b. Physical change

c. Element

d. Compound

e. Homogeneous mixture

f. Heterogeneous mixture

4. What is the formula for density? Describe how you could test if a block of an unknown substance was

more or less dense than water.

PART 2: Dimensional Analysis DUE TUESDAY 1/17

1. What does it mean to measure to “one uncertain digit”?

2. Record the correct measurements of the following pictures

3. Fill in the table below

Scientific Notation Long Hand Number of sig figs Which is the uncertain digit?

0.0008230

9.45x105

100000

0.00780096

4.56x10-3

4. How many millimeters is 0.6 meters?

5. After completing a lab, a student produced 5.5 g of a substance. The true value of the substance

produced should be 6.0 g. What is the percent error in the lab?

PART 3: Atomic Theory DUE TUESDAY 1/17

1. Describe the model of the atom used today

2. Distinguish between protons, neutrons, and electrons (include location in the atom, mass, and charge)

3. What is the distinguishing feature of an atom?

4. Atoms have what charge?

5. Distinguish between an ion and an isotope

6. A nitrogen ion (N3-) has how many protons, neutrons, and electrons?

7. What does each part of this symbol tell you?

X

8. Fill in the following table

Symbol Atomic number

Number of protons

Number of neutrons

Number of electrons

Mass number

Charge

18 18 -1

B

30 35 28

15 31 0

Ag+

9. Draw an atom of carbon and label each part

10. Nitrogen has three naturally occurring isotopes: Nitrogen-14, Nitrogen-15, and Nitrogen-16. Which is

the most abundant? Why?

11. 80.1% of Boron found on earth is Boron-10. The other 19.9% of Boron is Boron-11. What is average

atomic mass of boron?

PART 4: Nuclear Chemistry DUE TUESDAY 1/17

1. Describe how the nucleus of an atom is held together

2. Describe what happens in a nuclear reaction

3. The half life of substance X is 10 days. How much of a 90 gram sample is left after 30 days?

PART 5: Quantum Chemistry/Electron configuration DUE TUESDAY 1/17

1. Describe the relationship between wavelength, frequency, and energy.

2. Draw a wave with high frequency and a wave with low frequency. Draw a wave with a long wavelength

and short wavelength

3. Sketch the electromagnetic spectrum. Label parts of high frequency, high energy, and long wavelength

4. Using the Bohr diagram of a hydrogen atom, describe how an atom emits light. (use the term ground

state, excited state, etc)

5. Draw orbital diagrams for the following elements (the one with the arrows)

a. P

b. C

c. Cl

d. Ca

6. Write the long hand electron configuration for the following elements

a. Sc

b. S

c. Rb

d. Al

7. Write the noble gas configuration for the following elements

a. Ba

b. Au

c. Te

d. U

8. How many electrons are found in:

a. An s orbital?

b. A p ortbital?

c. A d orbital?

d. An f orbital?

e. The third energy level?

f. The 2nd energy level?

9. Label the parts:

2s2

10. In the periodic table above label the s, p, d, and f

block

PART 6: The Periodic Table DUE WEDNESDAY 1/18

1. What is the charge of each group on the periodic table (1A-8A) and why?

2. Circle the correct answer

a. A period goes (from left to right) (up and down)

b. A group goes (from left to right) (up and down)

3. On the periodic table above:

a. Label the alkali metals, alkaline earth metals, transition metals, halogens, and noble

gases

b. Label the metals, nonmetals, metalloids

4. Elements in the same group have similar ___________________________________

5. Elements in the same period have similar ___________________________________

6. Describe trends in atomic radius, ionic radius, electronegativity, and ionization energy

7. Which is larger?

a. Ba or F

b. N or As

c. Al or S

d. S or S2-

e. Na or Na+

8. Which has a greater electronegativity? F or Br? C or F?

9. Which has greater ionization energy? F or Br? C or F?

PART 7: Bonding DUE WEDNESDAY 1/18

1. Draw the Lewis Dot Structure (electron dot structure) for the following elements

Na P Cl I

Mg Xe Ga

2. Compare the types of bonding using the tables below

Types of atoms involved in bonding (ie metals)

What happens to electrons

Properties (hardness, conductivity, etc)

Ionic

Covalent

Metallic

3. Draw a Bohr diagram to show how Lithium and Fluorine will bond

4. Use electron dot structures to show how calcium and chlorine will bond

5. Does it create energy or take energy to break a bond?

6. Which is stronger: a single or triple bond? Which is longer: a double or triple bond?

7. Draw Lewis Structures for the following molecules

a. CH4 b. NO3- c. CO2

8. Write the shape for each of the above molecules and state whether the molecule was polar or

nonpolar. How do you know if a molecule is polar or nonpolar?

9. Complete the table below by writing the correct formulas for the compounds found by

combining the positive and negative ions

SO42- OH- NO3

- PO43-

Mg2+

Li+

Al3+

10. Write the prefixes used to name covalent molecules:

1-

2-

3-

4-

5-

6-

7-

8-

9-

10-

11. Remind yourself how to name each type of compound

How to know this is the type of compound

Naming rules

Ionic

Ionic with a transition metal

Ionic with a polyatomic ion

Covalent

Organic

12. Name or write the formula for each compound. Write I for ionic, C for covalent, and O organic

____LiCl

____BaSO4

____CO2

____MgCl2

____IF3

____CuOH

____P2O5

____SnO2

____SI4

____Cu3N

____AgCN

____N2O5

____Al2O3

____Cu(NO3)2

____Mg3(PO4)2

____N2O4

____Ammonium phosphate

____Sodium fluoride

____Zinc hydroxide

____Diphosphorus pentoxide

____Nickel (II) nitride

____Calcium phosphide

____Chlorine dioxide

____Silicon hexafluoride

____Dihydrogen dioxide

____Aluminum sulfide

____Aluminum sulfate

____Calcium phosphate

____Calcium phosphide

____Boron triiodide

____Iron (II) oxide

____Dinitrogen pentoxide

____Cobalt (I) carbonate

____Beryllium ioidide

PART 8: Organic Chemistry DUE THURSDAY 1/19

1. Draw the following molecules

Heptane

Ethane

2-ethyl, 3-methylpentane

2,2,-dimethylbutane

Pentane

3,3-diethyl, 2,4-dimethylheptane

2. Write the names of the following molecules

CH3(CH2)5CH3 CH3CH3

3. What is an isomer?

4. Draw 2 isomers for octane

PART 9: The Mole DUE THURSDAY 1/19

1. 1 mole = ___________________________ representative particles

2. What is the molar mass of each of the following:

a. Al2(SO4)3

b. C6H12O6

c. AgNO3

d. N2O5

e. O2

f. MgCl2

3. How many grams of CO2 are in 0.56 moles of carbon dioxide?

4. What is the mass, in grams, of 6.7 x1023 atoms of aluminum (Al)?

5. How many atoms are in 1.4 moles of sodium (Na)?

6. What is the percent composition of:

a. AgNO3 b. Al2(SO4)3

7. What is the empirical formula of

a. C6H12O6 b. N4O6

8. What is the molecular formula of a compound whose molar mass is 90 g/mol and empirical

formula is CH2O?

9. Find the empirical formula of a compound containing 25.9% nitrogen and 74.1% oxygen

PART 10: Stoichiometry DUE FRIDAY 1/20

Use the equation above for the next 3 Stoichiometry problems: 2H2S + 3O2 2SO2 + 2H2O

1. If we start with 60g H2S and excess O2, how many grams of H2O will be produced?

2. In the lab, we conducted an experiment starting with 60g H2S (see the question above). We

produced 25g H2O. What is our percent yield?

3. If we start with 4 moles O2, how many moles of SO2 will be produced?

Pb(NO3)2 (aq) + 2 NaI (aq) PbI2 (s) + 2 NaNO3 (aq)

4. If I start with 125.0 grams of lead (II) nitrate and 115.0 grams of sodium iodide, what is the limiting reagent? ________________ How many grams of sodium nitrate can be formed?

5. What if I start with 150.0 grams of lead (II) nitrate and 100.0 grams of sodium iodide, what is the limiting reagent? ________________ How many grams of sodium nitrate can be formed?