Download 1st Semester Final Exam Review Guide

Chapter 1
Chapter 2
Chapter 3
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CHEMISTRY: Semester 1 Final Exam Review
What is the difference between a physical and chemical change? Give an example of each.
Elements in the same group on the periodic table have similar _______________.
A reaction is exothermic when it ______________ _______. It is endothermic when it ______________ __________.
Perform the following conversions:
35.0 ml = ______ L
5.36 L = ______ ml
.615 kg = ______ g
What is the mass of an object that measures 1.55 cm x 3.00 cm x 2.50 cm and has a density of 2.70 g/cm3?
Compare the mass of a proton, neutron, and electron. Describe the location of each within an atom.
Determine the number of atoms in each of these compounds: a) C6H22O11 b) Mg3(PO4)2
How many atoms are in 1 mole of zinc?
Define each: a) alpha particle
b) beta particle
c) fission
d) fusion
Given 238U. Determine the number of protons, neutrons, and electrons. Also determine its mass number.
92
6.
Complete the nuclear reaction:
208Po
84

204Pb
+ ______
82
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8.
Chapters 4 & 5
Chapters 6 & 7
How many grams are there in 2.00 mol of calcium?
Calculate the atomic mass of carbon based on the following mass numbers of the isotopes of carbon found in nature.
mass # 12 at 81.00%
mass # 13 at 10.00%
mass # 14 at 9.000%
9. Define “isotope”. Give an example of 2 isotopes of a particular element.
1. In Bohr’s model of the atom, electrons that gain energy become “excited”. When these electrons drop to a lower energy
level, they release energy in the form of ___________________.
2. Write the electron configuration for oxygen.
3. How many electrons does O-2 have?
4. Write the electron configurations for: a) N-3
b) Al+3
c) Sr+2
5. Draw the orbital diagram for the outer electron configuration of fluorine.
6. Draw the electron dot structure for a nitrogen atom.
7. Know the trend for atomic size on the periodic table. Which element is the largest?
8. Define and describe the trend of electronegativity on the periodic table. Which element is the most electronegative?
1. Describe covalent, ionic, and metallic bonding in terms of sharing or transferring electrons.
2. Ionic compounds are formed from a _____________________ and ______________________.
3. List the diatomic elements (by name not formula).
4. Name the following: a) LiNO3
b) Ca(OH)2
c) Fe2O3
d) N2O4
5. Write the formula: a) sodium oxide b) carbon tetrachloride c) copper (II) fluoride
d) potassium phosphate
6. Identify as either a molecule or a formula unit: a) NiBr2
b) Ca(OH)2
c) CuS
d) HCl
e) Al2S3
7. Draw the Lewis structure for both H2S and SiCl4.
8. In an H2S molecule, how many unshared electron pairs (lone pairs) are there?
9. What two things must be considered when determining whether or not a molecule is polar or non polar?
10. Is CCl4 polar? Is H2S polar? Explain.
11. Determine the shape of the following molecules:
a.
b.
c.
d.
e.
f.
Cl -- Cl
H—O—H
H—N—H
H—C—H
S
O=C=O
Pressure
Chapters 7,8, & 9 1. A certain compound contains 47.3 % copper and 52.7% chlorine. Calculate the empirical (simplest) formula.
2. A compound has the simplest formula CH3, with a molecular mass of 30.0 g/mol. Determine the molecular formula.
3. Given: _____N2(g) + _____ H2(g)  _____ NH3(g)
a. Balance the equation.
b. If .6100 mol of hydrogen were to react, how many moles of nitrogen would be needed?
c. If 25.70 grams of nitrogen react, how many grams of ammonia would be produced?
d. What type of reaction is this?
4. Determine the number of grams in .055 moles of oxygen gas.
5. Calculate the percent of sodium in table salt (sodium chloride).
C
D
#2
6. How many total atoms are in 1.50 mol of H2O?
Chapter 10
1. Identify the states of matter in areas #1, #2, and #3
#1
Label the line segments that indicate:
B
Boiling point _____
Melting point _____
#3
A
Freezing point _____
Temperature
2. Crystalline solids. Give examples of each type (ionic, metallic, covalent network, and covalent molecular)
3. What is the relationship between IMFs and the boiling point of a substance?
Miscellaneous information
1. To what decimal place do the balances in our lab measure?
2. Differentiate between chemical…element, formula, equation, mixture