Chemistry I Semester I Exam Study Guide
Topics 1 & 2: Chapter(s) 1, 2, 3
All vocabulary from
this semester is fair
game for the final.
Understand the process and stages of scientific (logical) problem solving
Understand the differences between law and theory.
Understand and be able to use standard form (scientific notation) for representing numbers
Recall and use some SI units
Recall, and be able to use the rules for determining significant figures and rounding off
Understand the difference between Precision and Accuracy
Be able to calculate % Error.
Be able to interconvert between units using conversion factors
Learn, and be able to use, formula for the conversion of the three different temperature units studied in
this unit
10. Understand and recall definitions for physical and chemical change and properties
11. Know the difference between elements, mixtures and compounds
1.
2.
3.
4.
5.
6.
7.
8.
9.
Topic 3: Atoms, Elements, Molecules and Ions: Chapter(s) 4 and 24
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
Recall a very brief history of Atomic Theory
Know and understand the main aspects of Dalton's Atomic Theory
Know the three particles that make up the atom and their relative charges, masses and positions in the
atom
Be able to use and understand symbols and formula to represent elements and compounds
Know and be able to apply the Laws of Definite and Multiply Proportions.
Know what the term isotope means and be able to calculate average atomic masses.
Be able to use the Atomic # and Mass # of an isotope to calculate the numbers of protons, neutrons and
electrons
Understand the Band of Stability and how radioactive elements decay into the band.
Know the types of radiation (alpha, beta, positron, electron capture, gamma) and how they affect the
nucleus
Be able to balance decay equations.
Know the concept of Half-Life. Be able to use the half-life equation. (p. 871)
Topic 4 Electron configuration Chapter 5
1. Identify frequency, wavelength, and amplitude of a wave.
2. Calculate frequency, wavelength, and energy of a photon.
3. Explain the scientific contributions of the following people:
Planck
Einstein
de Broglie
Schrödinger
4. Explain how Bohr’s model of the atom explained the line emission spectrum of hydrogen.
5. Be able to determine the number of energy levels, sublevels, orbitals, and electrons present on
the periodic table
6. Know the three rules (Aufbau, Pauli exclusion, and Hunds) for creating electron configurations.
7. Be able to write orbital/arrow, longhand, & shorthand electron configurations.
8. Explain how stability leads to exceptions in the electron configurations of Cu and Cr
Topic 5 Inorganic Nomenclature - Sections 7.1 & 7.3
1.
2.
3.
4.
5.
Know how to combine those anions and cations in the correct proportions to form ionic compounds
with no net charge
Be able to write and name Monatomic Ions
Be able to name binary/ternary compounds of metals and non-metals
Understand and be able to use the system of Roman numerals indicating charge on metals of varying
oxidation state
Know what is meant by the term polyatomic ion
Semester I Review Questions
Matter
1. Classify the following substances as solid, liquid, gas, or plasma based on their properties.
a. flexible volume, high KE, particles can disperse freely.
b. flexible volume, very high KE, particles are charged.
c. fixed volume, very low KE, orderly particles.
d. fixed volume, low KE, particles can move past each other.
2. Compare and contrast a solution, colloid, and suspension.
3. Classify the following as element, compound, heterogeneous mixture, or solution.
a. graphite (carbon)
b. grape juice
c. table salt (NaCl)
d. pepper
4. Classify the following as chemical or physical changes.
a. cutting wire
b. ripening tomato
c. apple slices turning brown
d. compressing a gas
5. Classify the following properties as physical or chemical.
a. melts at 68.0C
b. corrosive
c. reacts violently with water
d. decomposes in air
e. magnetic
Measurement
6. In a lab, the average measured density for Pre-1982 pennies was 7.98 g/cm3. Given that the
literature value for the density is 8.92 g/cm3, calculate the percent error.
7. How many sig figs are in the following numbers? Round each value to 3 sig figs.
a. 2.2500
c. 89.70157
b. 34,0100
d. 0.0052001
8. Convert the following numbers into or out of scientific notation. How many sig figs are in each
value?
a. 548,000
c. 1.200 × 10-3
b. 0.0000770
d. 9.25 × 107
9. Osmium is the densest element with a density of 22.57 g/cm3. Find the mass of a 56.2 cm3 sample
of osmium.
10. Perform the following SI prefix conversions.
a. 65.2 mm = ? dm c. 65,000 L = ?
mL
b. 2.3 kg = ?g
d. 0.502 km = ? cm
Atomic Structure
11. Identify the scientists who made the following discoveries.
a. Atoms contain negative particles called electrons.
b. The mass of an electron is 9.11  10-28 g.
c. Atoms contain neutral particles called neutrons.
d. Atoms contain a dense, positive nucleus.
e. Atoms are indivisible and resemble BB’s.
12. Describe the evolution of the atomic model from the BB (solid Sphere) model to the electron cloud
model.
13. Magnesium combines with oxygen to form magnesium oxide. If 10.57 grams of magnesium reacts
completely with 6.96 g of oxygen, what is the percent by mass of oxygen in magnesium oxide?
14. When mercury oxide is heated, it decomposes into mercury and oxygen. If 28.4 g of mercury oxide
decomposes, producing 2.0 g of oxygen, what is the percent by mass of mercury in mercury oxide?
15. Write the isotope symbol, including atomic number & mass number, for the following isotopes.
a. carbon-14
c. nickel-63
b. chromium-53
d. zirconium-92
16. Complete the table for the following isotopes.
Symbol
Zn
Atomic #
20
Mass #
65
74
# of protons
34
# of neutrons
21
# of electrons
40
18
17. Calculate the average atomic mass of copper if 69.17% of the copper atoms occurring in nature are
63
Cu and 30.83% are 65Cu.
18. Determine the half-life of a radioactive isotope if 100 grams decayed into 6.25 grams in two hours.
19. If the half-life of iodine-131 is 8.10 days, how long will it take a 50.00 g sample to decay to 6.25 g?
20. If the half-life of an isotope is 525 years, how long will it take for 64 g to become 1.0g?
21. Carbon-14 has a half-life of 5730 years. How much of a 144 g sample of carbon-14 will remain after
1.719 × 104 years?
22. Uranium -238 decays by emitting an alpha particle. Finish, Balance, and classify the reaction
23. Carbon – 14 emits a beta particle. Finish, Balance, and classify the reaction
24. The isotope C-11 decays by an electron capture. Finish, Balance, and classify the reaction
25. The isotope O-15 decay by position emission. Finish, Balance, and classify the reaction
Isotopes of elements found in nature are all located within the gray area on the graph below called
the band of stability. Those elements found in the middle of the “band” have a very stable nucleus,
while those elements on the outer edges of the band have an unstable nucleus and are said to be
“radioactive”.
1. Locate & plot where the following atoms would be on the graph. Label each atom after it has
been plotted
Mg-24
Nd-142
Ir-195
Br-81
U-238
Ir-191
2. Did any of your atoms land outside the gray area? Explain why or why not.
3. How can there be two different atoms of iridium? How are they different
4. Would a small atom (less than 40 protons) be found in nature if it has the same number of
protons & neutrons (1:1 ratio)? Explain.
5. Would a large atom (more than 40 protons) be found in nature if it has the same number of
protons & neutrons (1:1 ratio)? Explain.
6. Two of the atoms you plotted are naturally radioactive, that is, their nuclei fall apart over time.
Which two do you think they are? What is your reasoning?
7. Imagine a chemist was trying to create an atom with 60 protons and a mass number of 155.
Would this be possible? Why or why not? (SHOW where it would fall on the graph).
8. If an element had 90 protons, how many neutrons would be a good number for it to have in order
to be considered a stable element? What element would this be? (SHOW where it would fall on
the graph).
Use the table below to answer questions A and B below.
Percent by Mass of Carbon, Hydrogen, and Oxygen in Selected Compounds
Compound
%H
%C
%O
3.2
19.4
77.4
Carbonic Acid (H 2 CO 3 )
Acetic acid (CH3COOH)
6.7
40.0
53.3
Methanol (CH 3 OH)
12.5
37.5
40.0
Methanal (H 2 CO)
6.7
40.0
53.3
Isopropanol (C 3 H 8 O)
A. You have a 125-g sample of one of these substances. You determine that it is made of 16.7 g H,
75.0 g C, and 33.3 g O. Which compound it it?
B. In another experiment, you determine that a sample of acetic acid consists of 56.8% oxygen. What
is your percent error?
Ch 5 review
To solve these problems, remember the following:
8

All electromagnetic waves move at the speed of light: 3.00 x 10 m/s.


Energy is measured in Joules and can be found using: E = hf, where h is Planck’s constant and equals 6.626 x 10
J  s.
Frequency and wavelength of anything in the electromagnetic spectrum is related by c = f , where  is the wavelength in
meters, f = frequency in Hertz and c is the speed of light.
34
Keep the proper significance and correct units on all problems. Show your work.
SOLVE THE FOLLOWING LIGHT & ENERGY PROBLEMS:
1. Find the energy of an infrared photon whose frequency is 2.0  1012 Hz.
2. What is the wavelength of a red light wave whose frequency is 4.3  1014 Hz?
3. Calculate the frequency of an X-ray wave that has a wavelength of 9.2 nm.
4. What is the energy of a cosmic ray photon whose wavelength is 2.0 pm?
5. Find the frequency of a microwave that carries 3.4  10-25 J of energy.
NAME THE TERM DESCRIBED BY THE FOLLOWING DEFINITIONS:
6. A packet of light energy that carries a quantum of energy.
7. The state when all electrons of an atom are in the lowest possible energy levels.
8. When an electron jumps up to a higher energy level, the atom is in its ___.
9. The scientist who applied Einstein’s particle-wave theory to electrons.
10. The theory that it is impossible to know both the position and speed of an electron
simultaneously.
11. The theory that no two electrons in an atom can share the same 4 quantum numbers.
12. The theory that electrons fill the lowest energy orbitals first.
13. The theory that, within a sublevel, electrons prefer to occupy their own orbital.
14. A term describing the outermost electrons in an atom.
15. The most stable type of electron configuration.
16. A three-dimensional region in space where an electron is likely to exist.
1. Using the element Nickel, answer the following questions.
a. Give the long hand summary notation.
b. Give the orbital notation using arrows and boxes
c. Give the short hand (noble gas) notation.
d. Give the electron dot structure.
2. Using the element Bromine, answer the following questions.
a. Give the long hand summary notation.
b. Give the orbital notation using arrows and boxes
c. Give the short hand (noble gas) notation.
d. Give the electron dot structure.
3. Using the element Barium, answer the following questions.
a. Give the long hand summary notation.
b. Give the orbital notation using arrows and boxes
c. Give the short hand (noble gas) notation.
d. Give the electron dot structure.
4. If n=3, how many sublevels ____________, orbitals _________, and electrons _________
5. How many orbitals are present in an (p) sublevel_______________.
6. How many electrons can a (f) sublevel hold_____________.
7. T or F there are 5 orbitals present in a (p) sublevel.
8. T or F n=3 can hold up to 18 electrons.
9. T or F (f) sublevels can hold 14 electrons
10. Give the electron dot notation for the following elements:
a. Se
b. Sr
c. Rb
d. Br
c. Fe
d. Sb
11. Give the noble gas configuration for the following:
a. Os
b. Si
12. Determine what rule(s) is being violated in each of the following configurations: (Hund’s Rule,
Pauli Exclusion principle, Aufbau principle)
a.
1s
2s
2p
_______________________
b.
1s
2s
2p
_______________________
c.
1s
2s
2p
_______________________
d.
1s
2s
2p
Chapter 7 Writing and Naming Ionic Compounds
Name each of the following ionic binary compounds.
1. NaCl
____________________
2. LiBr
____________________
3. FeCl2
____________________
4. Cr2S3
____________________
5. CaO
____________________
6. K3P
____________________
7. PbF4
____________________
8. Sr3As2
____________________
9. NaBr
____________________
10. Mn2O3
____________________
11. Mg(NO3)2
____________________
12. K2SO4
____________________
13. NH4Cl
____________________
14. Fe(NO2)3
____________________
15. ZnSO3
____________________
16. K3PO4
____________________
17. Li2CO3
____________________
18. Cu(OH)2
____________________
19. Pb(SO4)2
____________________
20. NaHCO3
____________________
________________________
Give the formula for each of the following ionic compounds.
1.
Sodium Chloride
2.
Calcium Nitride
3.
Aluminum Bromide
4.
Iron (II) Oxide
5.
Zinc Sulfide
6.
Barium Fluoride
7.
Chromium (III) Nitride
8.
Manganese (IV) Phosphide
9.
Magnesium Oxide
10. Lithium Iodide
11. Potassium Permanganate
12. Lead (II) Nitrate
13. Sodium Sulfate
14. Cesium Acetate
15. Ammonium Chloride
16. Cobalt (III) Carbonate
17. Silver Nitrite
18. Iron (III) Hydroxide
19. Mercury (II) Chromate
20. Hydrogen Peroxide