Name: ___________________________
Block: ________ Date: __________
Chemistry Midterm Exam Review
A. Measurements and Calculations
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accuracy
precision
scientific notation
dimensional analysis
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percent error
qualitative
quantitative
SI Units
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significant figures (counting,
multiplying, dividing, adding,
subtracting, averaging)
uncertainty in measurement
1. Determine the number of significant figures in the following numbers.
a. 0.00007690 g ______
c. 90070 km _______
b. 0.00978 mm
______
d. 9.10 cg
_______
2. Record your answer with the correct number of significant figures and units.
a. 41.6 g + 3.259 g =
__________ c. (26.2 cm)(1.234 cm) = __________
b. 0.00134 mL - 0.000233 mL = ________
d. (32.20 kg) / (4.0 kg) = _________
c. Convert the following numbers from scientific notation to ordinary notation.
a. 3.02 x 10-3 g = _________________ b. 5.791 x 105 m = ______________
d. Convert the following numbers from ordinary notation to scientific notation.
a. 4560 cm = ______________
b. 0.0076 g = ______________
3. How many centigrams are in 234 mg?
4. How many meters are in 0.000325 km?
5. What is the percent error if you measure the density of a substance to be 3.679g/ml but the
known true density of that material is 4.115g/mL
6. What is the difference between accuracy and precision?
B. Density and C. Matter
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matter
phases
physical property
chemical property
density
volume
mass
displacement
element
molecule
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monatomic/diatomic
compound
mixture
homogenous mixture
heterogeneous mixture
pure substance
solution
melting point
boiling point
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conductivity
luster
malleability
ductility
magnetism
chromatography
filtration
distillation
chemical/physical
change
7. The density of Mercury is 13.0 g/mL. If you have 24.3 mL of Mercury, what is its mass?
8. A cube of wood that weighs 23.5 g measures 224.21 cm by 1.45 cm by 7.34 cm. What is
the density of the wood?
9. You measure 23.6 mL of water into a graduate cylinder. You place a 56.56 g chunk of
metal into the cylinder and the volume increases to 29.3 mL. What is the density of the
chunk of metal?
10. Classify the following as a chemical or physical property.
a. Color
b. Flammability
c. Solubility
11. Classify the following as a chemical or physical change.
a. Tearing paper
b. Burning wood
c. Boiling water
12. Classify the following as an element or a compound.
a. Phosphorus
b. Carbon dioxide
c. Water
13. Classify the following as a mixture or pure substance.
a. A multivitamin tablet
b. distilled water
c. tap water
14. Classify the following as a homogeneous or heterogeneous mixture.
a. chunky peanut butter
b. a solution of copper (II) sulfate
c. a bag of trail mix
15. What would you use to separate two liquid solvents from each other?
a. Chromatography
c. Distillation
b. Filtration
d. Nuclear separation
D. Chemical Foundations
atom
atomic mass
atomic number
proton
neutron
electron
ion
cation/anion
average atomic mass
isotope
half-life
α (alpha) radiation
β (beta) radiation
γ (gamma) radiation
16. Give the symbols for the following elements
a. Iron
b. Fluorine
e. Helium
f. Hydrogen
Democritus
Dalton
Thomson
Millikan
Rutherford
Bohr
c. Beryllium
g. Silicon
d. Boron
h. Vanadium
17. Write the formula for the compound containing
a. a two to three ratio of Manganese to sulfur
b. six carbon, twelve hydrogen and six oxygen
18. Determine the number of protons, neutrons, and electrons in the following.
p
n
e
p
n
e
a.
42
Ca0
d.
62
b.
6
Li+1
e.
30 -3
c.
52
f.
124 -1
Cr
Co+3
P
I
19. If element X consists of 78.7% of atoms with a mass of 24.0 amu, 10.1% of atoms with a
mass of 25.0 amu, and 11.2% of the atoms with a mass of 26.0 amu, what is the atomic
mass of element X?
20. The average atomic mass of rubidium is 85.47 amu. There are two naturally occurring
isotopes of rubidium: 85Rb, mass of 84.91 amu and 87Rb, mass of 86.92 amu. What is the
percent abundance of each of the isotopes?
21. What is the half-life of a sample if 5.00 kg decays to 0.63kg in 27 days?
22. If the half-life of a chemical is 25 years and you have 1.000 gram, how much will be left in
100 years?
E. Modern Atomic Theory and the Periodic Table
electron configuration
excited state
ground state
orbital
Principle energy level
sublevel
Aufbau Order
Pauli Exclusion Principle
Hund’s Rule
Bohr
Planck
Heisenberg
De Broglie
element
group/family
period
metal
nonmetal
metalloid
alkali metal
alkaline earth metal
transition metal
halogen
noble gas
Mendeleev
Moseley
atomic radius
electronegativity
ionization energy
shielding effect
ion
oxidation state
cation
anion
23. Define the following terms:
a. principle energy level
b. orbitals
c. sublevel
24. Write the complete electron configuration and the orbital diagram for
Aluminum
Chlorine
Argon
Copper
25. Using your periodic table, write the noble gas configurations for
Fluorine
Silicon
Cesium
Lead
Iodine
26. Explain the following rules:
Aufbau Order
Pauli Exclusion Principle
Hund’s Rule
F. Chemical Bonding
ionic bond
covalent bond
polar covalent bond
electronegativity
polyatomic ion
molecule
27.
chemical formula
molecular formula
structural formula
Lewis structure
monatomic
diatomic
intramolecular force
VSEPR
linear
bent
tetrahedral
trigonal planar
trigonal pyramidal
For each of the following pairs of bonds, choose the bond that will be more polar.
a. H-P or H-C
b. O-F or O-I
c. N-O or S-O
d. N-H or Si-H
28. Draw the Lewis structure for each of the following
a. CO2
g. CF4
b. N2
h. NO+
c. NH3
i. NO3-
d. NH4+
j. H2S
e. SO4-2
k. ClO3-
f. NF3
l. BeF2
29. Determine the molecular shape for each of the previous drawings: Tetrahedral, Trigonal
pyramid, bent, trigonal planar, linear.
30. Determine if the previous drawings (#26) were Polar or Nonpolar
G. Chemical Nomenclature
polyatomic ions
cation
anion
common charges
acids
molecules
ionic compounds
binary covalent molecules
metals/non-metals
“-ate/-ic acids”
“-ite/-ous acids”
“-hydro_____-ic acids”
31. Identify the following elements as a metal, a nonmetal or a semi-metal
a. Strontium
c. Antimony
e. Selenium
b. Cadmium
d. Silicon
f. Manganese
32. Write the name for the following compounds.
a. AgBr
b. Ni(CN)3
c. Sr3(PO4)2
d. HF
e. NH3
f. Au2SO4
g. Ca(OH)2
h. PF5
33. Write the formula for the following compounds.
a. Barium chloride
e. tetraphosphorus decoxide
b. Cesium sulfite
f. Manganese (II) acetate
c. Sodium chlorite
g. Nitric acid
d. tricarbon hexanitride
h. hydronitric acid
H. Chemical Composition and the Mole
Atomic mass unit (amu)
Mole
Avagadro’s Number
Molar Mass
Molecular Weight
Formula Unit
Percent composition
Empirical formula
Molecular formula
34. Calculate the molar mass of the following compounds? (these will be used in #35-38)
a. AgBr
e. Ni(CN)3
b. HNO2
f. Sr3(PO4)2
c. HF
g. CO2
35. How many moles are in 7.23 grams of the following compounds?
a. Ni(CN)3
b. HNO2
c. Sr3(PO4)2
36. How many moles are in 3.02 x 1023 formula units of the following?
a. AgBr
b. Sr3(PO4)2
c. HF
37. How many grams are in 7.2 x 1026 formula units or molecules of the following? (two steps)
a. Ni(CN)3
b. HNO2
c. CO2
38. How many particles are in 3.45 grams of the following? (two steps)
a. AgBr
b. HNO2
c. HF
d. Ni(CN)3
e. Sr3(PO4)2
f. CO2
39. A compound used as an additive in gasoline is 71.65% Cl, 24.27% C, and 4.07% H. The
molar mass is 98.96 g. Determine the empirical formula and the molecular formula.
40. What is the percent composition of sodium, phosphorous, and oxygen in sodium
phosphate?
I. Chemical Reactions – An Introduction
Chemical Equation
Reactants
Products
Coefficients
Subscripts
Solid (s)
Liquid (l)
Gas (g)
Aqueous (aq)
Diatomic elements
41. Balance the following equations.
a. ___ C2H6 + ___ O2  ___ CO2 + ___ H2O
b. ___ CaCl2 + ___ H2O  ___ Ca(OH)2 + ___ HCl
c. ___ Cl2 + ___ NaI  ___ NaCl + ___ I2
d. ___ NH3 + ___ Cl2  ___ NH4Cl + ___ NCl3
e. ___ AuCl2 + ___ K2SO4  ___ AuSO4 + ___ KCl
42. Write and balance the following equations and include state symbols.
a. Solid iron (III) oxide is heated strongly in carbon monoxide gas; it produces
elemental iron and carbon dioxide gas.
b. Acetylene gas (C2H2) is burned in oxygen to produce carbon dioxide gas and
water vapor.
c. Calcium metal is added to water to produce hydrogen gas and the solid calcium
hydroxide.
J: Chemical Reaction Types
Precipitate
Soluble/Insoluble
Dissociation
Solubility Rules
Activity Series
Double Replacement
Single Replacement
Decomposition
Synthesis
Combustion (Complete
and Incomplete)
Acid
Base
Neutralization Reaction (same as
double displacement acid/base
reaction)
Write Out the Generic Reaction Formula for the Following Reaction Types
- Synthesis
- decomposition
- single replacement (metal cation exchange)
- double replacement
- neutralization (a type of double replacement)
- combustion - complete
- combustion - incomplete
43. Determine whether the following compounds are soluble or insoluble.
a. sodium acetate
c. silver hydroxide
b. lithium sulfide
d. colbalt (II) sulfate
44. Balance the following equations. Include physical state symbols. Classify the reactions in as
many ways as possible.
a. 2 C4H10 (g) + 13 O2 (g)  ___ CO2 (
)
+ ____ H2O (
b. ___ Ca (s) + ____ H2O (l)  ____ Ca(OH)2
(
)
)
+ ____
H2 (
)
Will this reation go? Why or why not?
c.
____ AgC2H3O2 (aq) + _____ KBr (aq)  ___ KC2H3O2 (
)
+ ___ AgBr(
Will this reation go? Why or why not?
d. ____ SO2 (g) + ___ H2O (l)  ____ H2SO3(aq)
e. ____ KClO3 (s) 
_____ KCl (s) + _____ O2 (
f. ____ NaOH(aq) + ____ H2SO4(aq) 
)
____ Na2SO4(aq) + ____ HOH(l)
)