GT CHEMISTRY MIDTERM REVIEW PACKET
CHAPTER 1
1. Why is there uncertainty in any measurement? What are two sources of this uncertainty?
2. What is the difference between accuracy and precision?
3. What is the formula for calculating percent error?
4. If you measure the mass of a book at 125 grams and the accepted value is really 130 grams,
what is your percent error?
5. A piece of iron (mass = 47.5 grams) is placed into a graduated cylinder filled with 50.0 ml of
water. The new volume (with the metal) is 62.0 ml. Calculate the density of iron.
6. List the metric prefixes in order from largest (kilo) to smallest (micro).
7. Convert 125 cm to km.
8. Convert 78.5 cm to µm. Report your answer in scientific notation.
CHAPTER 2
1. Define energy.
2. Differentiate between kinetic and potential energy.
3. State the Law of Conservation of Matter and Energy.
4. Write a formula to convert Celsius to Kelvin.
5. Convert 625 K to °C.
6. What temperature is absolute zero in Celsius and Kelvin?
7. Matter is
8. List three properties of a solid, a liquid, and a gas.
9. What is a physical property? List three examples.
10. What is a chemical property? List three examples.
11. What is a physical change? List three examples.
12. What is a chemical change? List three examples.
13. Define element.
14. Define compound.
15. What two things are classified as pure substances?
16. What is a mixture? What is a solution?
17. What is the difference between a homogeneous and heterogeneous mixture?
18. If you had salt sand and water mixed together in a beaker, what steps would you take to
separate it into its three components?
CHAPTER 3
What contributions were made by each of the following scientists in the development of the current atomic
model?
1. Democritus
2. Aristotle
3. Dalton
4. Thomson
5. Rutherford
6. List Dalton’s Atomic Theory.
7. What part of Dalton’s atomic theory has been changed? Why?
8. Atoms are composed of __ _, ______, and _____.
9. What does “amu” stand for? What is it defined as?
10. How many protons, neutrons and electrons are in each of the following?
a. carbon-14
b. chlorine-37
c. Na+1
d. Br-1
11. The mass number of an element is the total number of …
12.
Define atom.
13. There are two isotopes of Lithium, Li-6 and Li-7, found in nature at 7.50% and 92.5%
respectively. Calculate the atomic mass of lithium.
13. C-12 and C-14 are isotopes. What does that mean?
14. A certain “particle” has 17 protons, 18 electrons, and 18 neutrons. Write it’s isotopic notation.
15. A certain “particle” has 12 protons, 10 electrons, and 12 neutrons. Write it’s isotopic notation.
16. Complete the following nuclear reactions:
239
U 🡪
+
235
Th
92
140
Ba 🡪
56
90
+
140
La
57
CHAPTER 4
1. All waves have what four properties?
2. How are wavelength and frequency of light related?
3. List the colors of the visible spectrum in order from largest to smallest wavelength.
5. What is a line spectrum and why is it called “the fingerprint for an element”?
6. When an electron jumps from the excited state to the ground state, it ______ energy.
7.. What is the difference between a ground state electron configuration and an excited state configuration?
Provide an example.
9. State the Aufbau Principle.
10. State the Pauli exclusion principle.
11. State Hund’s Rule.
12. Write electron configurations for…
a. sodium –
b. iodine c. iron -
CHAPTER 5
1. What was Mendeleev’s contribution to the development of the periodic table?
2. How did Moseley change Mendeleev’s periodic table?
3. State the Periodic Law.
4. Rows are called…
5. Columns are called either …
6. Name the following:
a. Group I –
b. Group II –
c. Group VII d. Group VIII –
e. the “D” block –
f. the “F” block –
g. elements on the staircase –
7. A metal _______ electrons.
8. Nonmetals normally __________ electrons.
9. Define ionization energy and state its trend on the periodic table.
10. Define electronegativity and state its trend on the periodic table.
11. Define atomic radius and state its trend on the periodic table.
12. When K turns into K+1, does it get larger or smaller? Why?
13. When O turns into O-2, does it get larger or smaller? Why?
CHAPTER 7 (Nomenclature)
1. Name the following:
a. P2O5 –
i. K2SO4 –
b. CaCO3 –
j. CCl4 –
c. NH4F –
k. CrF2 –
d. AlN -
l. CuO –
e. FeO-
m. Cu2O–
f. SnO2 –
n. Al(C 2H3O2)3 –
g. H2SO3 –
o. LiSO3-
h. H3P -
p. P2O5-
2. Write formulas for the following:
a. lead (IV) sulfide -
g. magnesium chloride –
b. hydrosulfuric acid -
h. sulfur dibromide –
c. dinitrogen trioxide -
i. ammonium oxide –
d. chromium (III) hydroxide –
j. dinitrogen tetraoxide –
e. chromic acid –
k. iron (II) nitrite –
f. sodium carbonate –
l. calcium nitride –
CHAPTER 7 (Bonding)
1. All elements follow the _______________ rule, except for ____________
and _____________ which follow the _________ rule.
2. If an element fulfills either of the two rules mentioned in question #1, then that
element has the same _____________ as one of the _______.
3. The type of bond that is formed is determined by the difference in ______.
between the two bonding atoms. If the difference is between 0 and 0.4, then the bond will be
_________; if the difference is greater than 0.4 and less than 1.7, then the bond will
be__; if the difference is greater than 1.7, then the bond will be ____.
4. An ionic bond involves a ________ of electrons to the _______
electronegative element. A non-polar bond involves an _______ __________
of electrons. A polar bond involves an _______________ ___________ of
electrons.
5. Lewis Dot Diagrams use dots to represent the ____________ electrons and the
chemical symbol to represent the ____________ electrons.
6. A ___ double bond____________ involves the sharing of two pairs of electrons and a
__triple bond___________ involves a sharing of three pairs of electrons.
7. What is the difference between an ionic bond and a covalent bond?
Ionic Bond – transfer of electrons from one atom to another to form positive and negative ions
Covalent Bond – sharing of electrons between two atoms
8. Draw Lewis dot diagrams for:
a. C
b. P
c. Br
d. Xe
10. Draw Lewis structures for these ionic compounds: (show how the ionic bond is formed)
a. NaF
b. CaI2
11. Draw Lewis Structures for these covalent compounds:
a. CF4
b. NCl3
c. C2H2
d. H2S
12. What is the difference between a molecular formula (CH4) and a structural formula?
CHAPTER 8
1. What does “VSEPR” stand for?
2. What is the definition of polarity? What determines if a bond is polar? What about a molecule?
3. Each C-F bond in CF4 is polar but the molecule itself is nonpolar. Explain.
4. Draw the water molecule.
a.
b.
c.
d.
e.
f.
How many bonds are in the molecule?
This means __ electrons are “shared.
How many unpaired electrons are in the molecule?
What is the shape of this molecule?
Are that any polar bonds?
If the molecule polar? \
5. Fill in the following chart:
Formula
NF3
H2O
CCl4
CO2
CH3F
Lewis Structure
Geometry
Bond
Polarity
Molecular
Polarity