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Period: _____
Testing Room: ________
Chemistry Level 2- Midterm Review
Chapters 1-8
TEST DATE: February 2, 2016
The midterm will consist of 90 multiple-choice questions and 5
dimensional analysis questions (90 minute period). You will be given a
periodic table whose front and back resemble the one you currently
possess. This will be the last sheet of your test. During the test, you may
remove this page if you like and write anything on it. You are
responsible for bringing your own calculator to the test along with a # 2
pencil. Arrive Prepared!
The midterm average will be
equally averaged with the
result of your final exam to
make up 1/5 of your final
course average.
Chapter 1 (5 questions)
1. What is chemistry?
2. What is matter?
3. Be able to describe each step of the scientific method. Provide a couple examples where you used the
scientific method to solve a problem in your life and label each part.
4. Theory v. Scientific Law --- What is the difference? Think of examples for both and identify the
limitations. Can they be proved/disproved supported/falsified?
5. What are variables? What step of the scientific method involves manipulating variables?
6. In large, why has science advanced as much as it has over the years?
--------------------------------------------------------------------------------------------------------------------------------------Chapter 2 (10 questions)
1. Name some examples of physical properties.
2. Define the particle arrangement of a solid, liquid and gas. What can you say about their shapes and the
volume they occupy?
3. Identify physical properties from chemical properties. For example zinc metal is hard, silver in color, is
easily shaped and reacts vigorously with hydrochloric acid. What are some physical and chemical
properties of zinc?
4. What is the difference between a physical and chemical change? How do you know when one
occurred? Think of examples.
5. What are symbols? What are formulas? When should you capitalize letters? What are subscripts?
6. An object has a volume of 12mL and a mass of 3.5g. What is the object’s density?
7. An object has a mass of 357g and volume of 75mL. What is the object’s density?
8. An object has a mass of 4.2g and a volume of 19.3mL. What is the object’s density?
--------------------------------------------------------------------------------------------------------------------------------------Chapter 3 (15 questions)
1. What is an advantage of the metric system compared to the English system?
2. What is the SI unit of mass? What is the SI unit for temperature?
3. Define weight. Does weight change based on location?
4. Which temperature scale has no negative numbers?
5.
Convert the following:
a. 50oC = __________K
b. -25oC = _________K
Count the number of significant figures in these problems
1. 50,400 mg = _____________
2. 0.00046 kg = _____________
3. 923.110 mg = ____________
4. 1.0045 L = _______________
Round problems 1-4 to two significant figures and write the answer in scientific notation.
1. 50,400 mg = ____________________
2. 0.00046 kg = ____________________
3. 923.110 mg = ____________________
4. 1,045 L = ______________________
Review rules regarding adding/subtracting and multiplying/dividing
1. What is the result of adding
2.01g +4.5g = ____________
2. What is result of subtracting
4.356m – 3.6m = ___________
3. What is the result of multiplying
5.2 X 102 by 1.367 X 10-4m = _____________
4. What is the result of dividing
4.3 X102 by 2.0 X 104g = ____________________
Density
1. Define density.
2. volume = 90.0 g
3. mass = 55 g
4. mass = 7.0 g
density = 0.70g/mL
density = 12.3 g/mL
volume = 9.1 cm3
mass = _________
volume = _________
density = ___________
5. What is the density of an object that has a mass of 5.6g and a volume of 25cm3?
Dimensional Analysis problems
Go back and review the problems that we worked on in our book for section 3.3. Make sure you know how to
do those problems.
1. An atom of gold has a mass of 3.271 X 10 -22g. How many atoms of gold are in a 5,050mg sample of
gold?
2. The distance from the Sun to Earth is roughly 1.50 X 108 km. Express this in units of centimeters.
3. Copper has a density of 8.3 g/cm3. What is the density of copper expressed in kg/m3?
4. Convert 23ft/s to mi/hr. {5,280ft = 1mi
60s = 1min
5. Convert 9.80 m/s to km/min {1000m = 1km
60min = 1hr}
60s = 1min}
--------------------------------------------------------------------------------------------------------------------------------------Chapter 4 (10 questions)
1. Fill in the table below:
Electrical Charge
Relative Atomic Mass
Location
Proton
Neutron
Electron
2. What does the atomic number of an element indicate?
3. What does the mass number of an element indicate?
4. If an atom has an atomic number of 36 and a mass number of 84 how many protons, neutrons and
electrons are present in this atom?
5. What does the number 13 represent in carbon-13?
6. How do you calculate the number of neutrons in an atom?
7. Who was Democritus?
8. What is the unit that is used to measure a weighted average atomic mass?
9. What is an isotope?
10. What is an ion?
11. How many protons, neutrons and electrons are present in this neutral atom?
39 K
19
--------------------------------------------------------------------------------------------------------------------------------------Chapter 5 (20 questions)
1. Describe the quantum mechanical model of an atom?
2. Which scientist developed the quantum mechanical model of an atom?
3. What are the shapes of an s and p orbitals?
4. What is a principal energy level, sublevel and atomic orbital?
5. What is the maximum number in each s, p, d and f orbitals?
6. What types of atomic orbitals are in the 1st, 2nd and 3rd principal energy levels?
7. If the spin of one electron is clockwise in an orbital the spin on the second electron must be
________________________?
8. Using the Aufbau diagram what orbital would come after 3p?
9. What is the number of electrons in the outermost energy level of sulfur?
10. What happens when an electron moves from a lower to higher energy level? (absorb a quantum of
energy or lose a quantum of energy)
11. When an electron moves from a higher to lower energy level what happens?
12. When does emission of light occur?
13. Which of the following quantum leaps would be associated with the greatest energy emitted? Why?
a. n = 3 to n=2
b. n = 6 to n= 2
14. How are frequency and wavelength of light related?
15. Which variable is directly proportional to energy?
16. Define a photon.
17. What is the wavelength of a radiation with a frequency of 5 X 1018 Hz (s-1)?
(c = 3.00 X 108 m/s)
18. What is the frequency of light having a wavelength of 4.2 X 10-8m?
19. What is the energy of a photon with a frequency of 2.5 X 1015 Hz?
[E = h v] (h = 6.63 X 10-34 J s)
20. What are the electron configurations for the following elements:
a. Na
b. Ar
c. Zn
d. As
e. Br
f. Sr
--------------------------------------------------------------------------------------------------------------------------------------Chapter 6 (10 questions)
1. Another name for the representative elements is ________________. Where are these elements located
on the periodic table?
2. Who was the first scientist to arrange the elements according to similar chemical and physical properties
in order of increasing atomic mass?
3. What is characteristic of the electron configurations of noble gases?
4. Which subatomic particle plays the greatest role in determining the properties of an element?
5. What is the periodic trend for:
a. Atomic radius
b. Ionic radius
c. Ionization energy
d. Electronegativity
6. What is the energy required to remove an electron from an atom in the gaseous state called?
7.
Which element would have the lowest first ionization energy?
8.
Which element has the largest atomic radius?
Na, Mg, Al, Si, S, or Cl?
a. H, Li, Na, K, Rb, Cs or Fr
b. Li, Be, B, C, N, O, or F
9.
Which element has the highest electronegativity value on the periodic table?
10. Define ion, cation, and anion.
11. What are the ionic charges for the representative elements?
12. Write the charges for each one of these ions as a superscript
a. Li
e. Ca
b. O
f. Al
c. S
g. Br
d. Na
h. Rb
--------------------------------------------------------------------------------------------------------------------------------------Chapter 7
(5 questions)
1. How many valence electrons would these atoms have?
a. Li
b. N
c. O
d. Cl
2. What is an ionic compound?
3. What are some characteristics of ionic compounds?
4. What is the electron configurations of the following ions:
a. N3b. Al3+
c. Cl1d. Ca2+
5. What is a metallic bond?
6. What are some properties of metals?
7. Why do metals have these properties?
--------------------------------------------------------------------------------------------------------------------------------------Chapter 8 (5 questions)
1. In a single covalent bond, how would two atoms still achieve a stable noble-gas electron configuration?
2. How many electrons are required to create a single, double and triple covalent bond?
3. What is an unshared pair of electrons?
4. According to VSEPR theory, why do molecules adjust their shapes?
5. What is a polar covalent bond? What is a non-polar covalent bond? How do you determine which type
of bond would form?
6. Which type of solid has the strongest bond structure and therefore the highest melting point?
a. Ionic Solids
b. Metallic Solids
c. Network Solids
7. How are ionic compounds and molecular compounds different?
Ionic Compounds
Molecular Compounds
8. Draw the electron (or Lewis) dot structure for the following ionic compounds.
a. Aluminum chloride
b. Potassium nitride
c. Magnesium sulfide
9.
Draw the electron (or Lewis) dot structure for the following molecular compounds
a. Nitrogen dioxide
b. Carbon monoxide
c. Oxygen
d. Carbon tetrachloride