Chemistry Midterm Review Study Guide 2013
Chapter 1: Matter and Change
1. Classify each of the following as a pure substance or a mixture.
a. sugar PS b. iron filings PS
c. milk M
d. plastic PS e. cement M
2. Identify each example as a physical property or a chemical property.
a. sodium chloride is a solid physical
b. water’s boiling point is 100°C physical
c. ammonia is very soluble in cold water physical d. sodium reacts violently with water chemical
3. What is the difference between extensive and intensive properties? Give examples of each.
Extensive: depend on the amount of matter present ex. Mass and volume
Intensive: independent of the amount of matter present ex. Odor, color, state, texture
4. Consider the burning of gasoline and the evaporation of gasoline. Which process represents a
chemical change and which represents a physical change? Give a reason for your answer.
Burning of gasoline- chemical change because gas is being changed into different substances.
Evaporation of gasoline physical change because; gas is simply undergoing a change in state,
still gasoline.
5. Classify each of the following as either a physical change or a chemical change.
a. melting an ice cube PC
d.. sharpening a pencil PC
b. burning a piece of paper CC
e. decomposing mercury(II) oxide CC
c. slicing a loaf of bread PC
f. dissolving sugar in water PC
6. Explain the differences between solid, liquid, and gaseous states in terms of the arrangement
of the particles.
Solids: definite volume and shape, particle vibrate around a fixed position
Liquids: Definite volume but no definite shape, particles can slide past one another.
Gases: No definite volume or shape, mainly empty space, particles move rapidly past one
another.
7. Compare the characteristics of metals, non metals and metalloids. Give examples of each.
Metals: Ductile, malleable, conduct heat and electricity luster, Ex. Fe, Cu, Ni, Na
Non-metal: Brittle, lack luster, poor conductors of heat and electricity ex. O2, Cl2
Metalloids: Semi conductor’s, solid, have characteristics between metals/ nonmetal,
B,Si,Ge,As,Sb,Te
8. What do elements in the same group have in common? The same period?
Group or family- vertical, have similar chemical and physical properties, have the same number
of valence electrons.
Period: horizontal, chemical and physical properties vary.
9. What is the difference between a homogeneous and heterogeneous mixture? Give an example
of each.
Pure substance: can’t be physically separated. Elements (Na, Ca) and compounds CO2, NH3
Mixture: blend of two or more kinds of matter, each of which retains its own identity and
properties. Can be physically separated. Homogenous mixture: same consistency throughout ex
salt water. Heterogeneous mixtures: Not uniform throughout ex. Muddy water
Chemistry Midterm Review Study Guide 2013
Chapter 2: Measurements and Calculations
1. How many significant figures are in each of the following measurements?
a. 0.0450 kg 3
b. 45.9 m 3 c. 0.05306 L 4
d. 689.0 cm 4
e. 7.20 x 1023 g 3
2. Round the following to have 3 significant figures:
a. 25.666 mL 25.7 mL
b. 25,550 cm 25,600 cm
c. 50 g 50.0 g
3. Complete the following conversions: Show the work!
a. 54 mL=_______ L
b. 0.25 g = ________ mg
54 𝑚𝐿
1
𝑥
1𝐿
1000 𝑚𝐿
= 0.054 L
c. 400 cm3 =________ L
400 𝑐𝑚3
1
𝑥
1𝐿
1000 𝑐𝑚3
0.25 𝑔
1
𝑥
1000 𝑚𝑔
1𝑔
= 250 𝑚𝑔
d. 67 cm =_________ m
67 𝑐𝑚
= 0.4 𝐿
1
𝑥
1𝑚
100 𝑐𝑚
= 0.67 𝑚
4. Aluminum has a density of 2.70 g/cm3. What would be the mass of a sample whose volume is
10.0 cm3?
D = M/V
D = 2.70 g/cm3
3
M=DxV
V= 10.0 cm
M = 2.70 g/cm3 x 10.0 cm3 = 270 g
M=?
5. The mass of a 5.00 cm3 sample of clay is 11.0 g. What is the density of the clay?
D=?
D = M/V
M = 11.0 g
D = 11.0 g/5.00 cm3
3
V = 5.00 cm
D = 2.20 g/cm3
6. If the density of gold is 19.8 g/mL, what is the volume of a 75.0 g piece of gold?
D = 19.8 g/mL
D = M/V
M = 75.0 g
M=DxV
V=?
V = M/D
V = 75.0 g / 19.8 g/mL
V = 3.79 mL
7. Rewrite the following values in scientific notation.
a. 530000 L =
5.3 x 105 L
b. 0.00053 L =
5.3 x 10-4 L
8. Complete the following calculations in scientific notation.
a. 5.0 x 104 m + 2.4 x 104 m = 7.4 x 104 m
b. 9.3 x 1012 g ÷ 3.0 x 108 mL = 3.1 x 104 g/mL
9. A length measurement is 1.40 cm. The correct value is 1.36 cm. Calculate % error
|1.36−1.40|
𝑡ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙−𝑒𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙
% 𝑒𝑟𝑟𝑜𝑟 =
𝑥 100 =
𝑥 100 = 2.94%
𝑡ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙
1.36
Chemistry Midterm Review Study Guide 2013
10. If x is directly proportional to y, when x increases, y increases Sketch what this graph would look like below:
24
Y
0
0
12
X
11. If x is inversely proportional to y, when x increases, y decreases Sketch what this graph would look like below:
12
Y
0
0
12
X
Chapter 3:Atoms the Building Blocks of Matter
1.
What does the atomic number tell you about an element?
number of protons in the nucleus of an atom
2.
List the number of protons, neutrons and electrons found in Zn-66:
Protons 30
Neutrons 36
Electrons 30
3. Complete the following table
Isotope Name
Isotope Symbol Protons
Neutrons
Mass Number
Copper-65
29
36
65
65
29
Potassium-42
42
19
Calcium-42
K
42
𝐶𝑎
20
Sulfur-34
4.
Cu
34
𝑆
16
19
23
42
20
22
42
16
18
34
Complete the following table:
Ion Symbol
Number of Protons
Number of electrons
Ca+2
20
18
8
10
53
54
O
-2
I -1
Chemistry Midterm Review Study Guide 2013
5. Consider a hypothetical element, D, which has three isotopes. Calculate the weightedaverage atomic mass of the element from the following data.
Isotope
Abundance
Atomic Mass
D-179
7.00%
.07 x 178.91 amu= 12.523
D-177
42.00%
.42 x 176.84 amu= 74.272
D-176
51.00%
.51 x 175.92 amu =89.719
176.51 amu
6. Describe the contributions of the following scientists:
Mendeleev : Credited with the 1st version of the periodic table. Grouped elements according to
atomic mass. Predicted the existence and properties of new elements.
Rutherford: Gold foil experiment. Discovered atoms have a very tiny dense nucleus with positive
change.Most alpha particles passed through atoms, but 1/8000 were deflected by nucleus
Dalton: Developed atomic theory- 1) all matter is made of atoms 2) atoms of the same element
are identical in their properties 3) atoms can’t be created nor destroyed. 4) atoms combine in
whole number rations to create compounds 5) atoms are combined, separated or rearranged in
chemical reactions.
Thompson: Cathode ray experiment proved the existence of negatively charged particles later
called electrons.
Bohr: : Electrons can circle the nucleus only in allowed paths , orbits and have fixed energy
Chapter 4:Arrangements of Electrons in Atoms
1.
2.
Arrange the colors of light (rainbow) in order of:
a.
Increasing wavelength (VIBGYOR)
b.
Increasing frequency (ROYGBIV)
c.
Increasing energy (ROYGBIV)
What is the relationship between:
a.
Wavelength and frequency inversely proportional
b.
Frequency and energy directly proportional
3.
Draw a picture of Bohr’s Model for a hydrogen atom. Using the term ground state and
excited state, explain how an atom of hydrogen can emit a red photon of light and a violet
photon of light.
Each atom contains a certain number of electrons which move around the nucleus
in fixed energy levels. The e- are on orbits. When an atom is in the ground state,
the e- s are in the lowest possible energy level. When heat or electricity is applied
the e- jumps to higher energy levels that are further from the nucleus. This is
called a quantum leap and the atom is now in the excited state. When it falls back
down it releases specific amounts of energy in the form of light.
Chemistry Midterm Review Study Guide 2013
4. Write the full electron configuration of the following atoms:
a. carbon 1s22s22p2
b. potassium 1s2 2s2 2p6 3s2 3p6 4s1
5. Write the full electron configuration of the following ions:
a. sodium ion 1s22s22p6
(+1 loses 1 electron to form ion)
b. oxygen ion 1s22s22p6
(-2 gains 2 electrons to form ion)
6. Write the noble-gas notation for the following element in the space provided.
a. silicon [Ne]3s23p2
b. barium [Xe]6s2
7. Write the orbital notation for the following element.
carbon
8. Which of the following electron configurations represent a(n) noble gas, transition metal,
alkali metal, halogen and alkaline earth metal. Identify each element
a. 1s22s22p63s2 Magnesium – Alkaline Earth Metal
b. 1s22s22p63s23p64s23d104p5 Bromine - Halogen
c. 1s22s22p63s23p64s23d104p6 Krypton – Noble Gas
d 1s22s22p63s23p64s23d3 Vanadium – Transition Metal
e. 1s22s22p63s23p64s1 Potassium – Alkali Metal
9. On the line at the left of each expression in the first column, write the letter of the expression
in the second column that is most closely related.
D i. An electron occupies the lowest
energy orbital that can receive it.
A ii. Orbitals of equal energy are
occupied by one electron before
any orbital is occupied by a second
electron.
B iii. No two electrons will have the
same spin in the same orbital.
C iv. The single electron of hydrogen
orbits the nucleus only in allowed
orbits, each with a fixed energy.
a. Hund’s rule
b. Pauli exclusion principle
c. Bohr model of the atom
d. Aufbau principle
Chapter 5: The Periodic Law
1. a. Metalloids are found in which block, s, p, d, or f ? p
b. The hardest, densest metals are found in which block, s, p, d, or f ? d
Chemistry Midterm Review Study Guide 2013
2. a. Name the most chemically active halogen. Fluorine
b. Write its electron configuration. 1s22s22p5
c. Write the configuration of the most-stable ion this element makes. 1s22s22p6
3. Define the following terms and describe the general trends as you go across a period AND
down a group:
a. atomic radius
b. electronegativity
c. ionization energy
size of atom
attraction for electrons
energy required to remove eincreases down group
decreases down group
decreases down group
decreases across row
increases across row (halogens highest)
increases across row
4. a. Which has the larger radius, Al or In? In
b. Which has the larger radius, Se or Ca? Ca
c. Which has a larger radius, Ca or Ca+2 Ca (would get smaller if lost 2 e-)
d. Which has greater ionization energies as a class, metals or nonmetals? nonmetals
e. Which has the greater ionization energy, As or Cl? Cl
f. In general, which has a stronger electron attraction, large atoms or small atoms? Small atoms
g. Which has greater electronegativity, O or F? F
h. In the covalent bond between O and F, to which atom is the electron pair more closely drawn?
Fluorine, greater electronegativity, smaller atom with greater nuclear charge (protons)
5. a. Elements whose atoms contain partially filled d sublevels are called Transition Metals
b. The electrons available to be gained, lost, or shared in the formation of chemical compounds
are called Valence electrons.
c. The measure of the ability of an atom in a chemical compound to attract electrons is called
electronegativity
d. The energy required to remove an electron from an atom is called its ionization energy
e. An atom or group of atoms that has a positive or negative charge is called a(n) cation
6. Explain which of the elements in each pair has the larger ionization energy and WHY.
a.Mg or Cl they have the same number of energy levels of electrons, Cl has a greater nuclear
charge with more protons in its nucleus with a greater attraction for valence electrons.
b. Ca or Ba they both have two valence electrons, but barium has 6 energy levels of electrons
while calcium has 4. The valence electrons in Ca are closer to nucleus than barium and are
better attracted to the nucleus.
Chapter 6: Chemical Bonding
1. a. A covalent bond in which the bonded atoms have an unequal attraction for the shared electrons
is called a(n) polar covalent bond.
b. The degree to which bonding between atoms of two elements is ionic or covalent can be
determined from the differences in the electronegativity of the elements.
c. A charged group of covalently bonded atoms is called a(n)polyatomic ion
d. A chemical bond that results from the attraction between metal atoms and the surrounding sea
of electrons is called a(n) metallic bond.
Chemistry Midterm Review Study Guide 2013
e. If the electronegativity difference between two atoms is between 0-.3 the bond is nonpolar
covalent, if the electronegativity difference is between .3 and 1.7 the bond is polar covalent and
if the electronegativity difference is greater than 1.7 then the bond is ionic.
2. Compare the characteristics of ionic and covalent compounds. Give an example of each.
Ionic: transfer of electrons to form ions, forms between a metal/nonmetal, strong attraction
between charged ions, solids, high melting points, solubility in water varies, high conductivity in
solution
Covalent: Sharing of electrons to form molecule, forms between a nonmetal/nonmetal, polar
(unequal sharing) and nonpolar (sharing) bonds, weak attraction between neutral molecules, low
melting points, solid, liquid or gas, solubility varies, does not conduct electricity
3. Draw a Lewis Dot diagrams for the following atoms
a. Ca
b. C
c. O
4. Draw a Lewis structure for each of the following formulas. Use the VSEPR theory to predict the
molecular shape of the molecules and then determine whether the molecule is polar or nonpolar. Give the
name of the type of intermolecular force its molecules will exhibit.
a. H2S Bent, Polar, Dipole-Dipole
b. N2
c. PCl3 trigonal pyramidal, polar, Dipole-Dipole
d. CH2O trigonal planar, polar, Dipole-Dipole
e.
CO2 linear, nonpolar, London Dispersion
g.
ICN Linear, Polar, Dipole-Dipole
Linear, nonpolar, London Dispersion
f. CH4 tetrahedral, nonpolar, London Dispersion
h. BH3 Trigonal Planar, nonpolar, London Dispersion
I--C≡N:
5.Draw the Lewis Dot diagram for the following polyatomic ion:
a. BrO3-1
b. SO4-2
Chemistry Midterm Review Study Guide 2013
Chapter 7: Nomenclature/The Mole
1. Write the correct formula for the following ionic compounds:
a. copper (II) oxide CuO
b. tin (II) sulfide SnS
c. barium nitride Ba3N2
d. ammonium phosphate (NH4)3PO4
e. potassium carbonate K2CO3
f. calcium hydroxide Ca(OH)2
2. Give the correct name for each of the following ionic compounds:
a.
Al(CN)3 aluminum cyanide
b.
BaSO4 barium sulfate
c.
FeCl3 iron (III) chloride or ferric chloride
3. Write the correct formula for the following binary, molecular compounds:
a. iodine heptafluoride IF7
b. diphosphorus trisulfide P2S3
c. selenium dioxide SeO2
4. Write the name or the formula for the following acids:
a. HCl hydrochloric acid
b. H2SO4 sulfuric acid
c. Phosphoric acid H3PO4
d. Nitrous acid HNO2
e. Hydrobromic acid HBr
_____ 5. Which completion is false? One mole of carbon dioxide, CO2 ….
a. contains the same number of molecules as one mole of water, H2O
b. has the same mass as one mole of water, H2O
is 6.022 x 1023 molecules of CO2
c.
_____ 6. How many atoms of iron are in 100 g of iron?
a. 1.09 x1026 atoms
b. 1.08 x1024 atoms
c. 2.69 x1023 atoms
_____7. A compound whose molecular mass is 90g contains 40.0% carbon, 6.67% hydrogen, and
53.33% oxygen, What is the true formula of the compound?
a. C2H2O4
b. CH2O4
c. C3H6O
d. C3HO3
e. C3H6O3
8. What is the molar mass of calcium chloride?
CaCl2 40.1 + (2x35.5) = 111.1 g/mol
Chemistry Midterm Review Study Guide 2013
9. What is the mass in grams of 0.25 mol of sodium hydroxide?
0.25 𝑚𝑜𝑙
1
40.0 𝑔
𝑥
= 10. g NaOH
1 𝑚𝑜𝑙
10. How many grams are in 6.3 moles of NH3?
6.3 𝑚𝑜𝑙𝑒
1
𝑥
17.0 𝑔
1 𝑚𝑜𝑙𝑒
= 110 g NH3
11. How many molecules are in 4.25 moles of carbon dioxide?
4.25 𝑚𝑜𝑙𝑒
1
𝑥
6.02 𝑥 1023
1 𝑚𝑜𝑙𝑒
= 2.56 𝑥1024 CO2 molecules
12. How many moles of copper are 5.34 x 1024 atoms of copper?
5.34 x 1024 atoms
1
𝑥
1 𝑚𝑜𝑙𝑒
6.02 𝑥 1023 𝑎𝑡𝑜𝑚𝑠
= 8.87 moles Cu
13. Find the percentage due to mass of hydrogen, sulfur and oxygen in H2SO4
2x 1.0= 2.0
H= 2.0/98.1 x 100 = 2.0%
1x32.1= 32.1
S= 32.1/98.1 x 100= 32.7%
4 x 16.0= 64
O= 64/ 98.1 X 100 = 65.2 %
98.1 g H2SO4
14. What is the percentage composition of ethane gas, C2H6?
2 x 12.0= 24.0
C= 24/30 x 100 = 80.0%
6 x 1.0 = + 6.0
H = 6/30 x 100 = 20.0%
30g C2H6?
15. What is the empirical formula of a compound composed of 85.7% C and 14.3% H?
Empirical Formula= CH2
C= 85.7 g C X 1mole C= 7.141mole C/7.141= 1
12g C
H= 14.3 gH x 1 mole H= 14.3mole H/ 7.141 = 2
1gH
Chapter 8: Chemical Reactions
1. What are the five indicators that a chemical reaction is occurring?
Production of a gas, formation of a precipitate, color change, production of heat or light
2. Balance the following equations and state what type of reaction they are.
synthesis a. H2 +
Cl2  2 HCl
decomposition b.
2HgO  2Hg + O2
single replacement c.
combustion
synthesis e.
d.
Ca + 2H2O 
CH4 + 2 O2 
N2 + 3H2 
decomposition f. 4KMnO4 
Ca(OH)2 + H2
CO2 + 2 H20
2 NH3
2 K2O + 4 MnO + 5 O2
double replacement g. Mg(NO3)2 + 2KOH  Mg(OH)2 + 2KNO3
Chemistry Midterm Review Study Guide 2013
3. Convert the following word equations into balanced chemical equations:
a. Aluminum metal and aqueous copper (II) fluoride yield aqueous aluminum fluoride and solid
copper.
2Al (s) + 3CuF2 (aq)  2AlF3(aq) + 3Cu (s)
b. Aqueous solutions of sodium chloride and silver nitrate produce an aqueous solution of sodium
nitrate and solid silver chloride.
NaCl (aq) + AgNO3 (aq)  NaNO3(aq) + AgCl (s)
c. Aluminum is above copper in the activity series. Will aluminum metal react with copper (II)
nitrate? If so, write the balanced equation for the reaction.
Yes, if Al is above the Cu ions it is more reactive and will replace it within solution.
2Al (s)+ 3Cu(NO3)2 (aq) - 3Cu (s) + 2Al(NO3)3 (aq)
4. Complete and balance the following equations:
a.
AgNO3(aq) + NaCl(aq) → NaNO3 (aq) + AgCl (s)
b.
3LiOH(aq) + Fe(NO3)3(aq) → 3LiNO3 (aq) + Fe(OH)3(s)
c.
Na +
d.
H2 + Cl2 →
H2O → NaOH + H2
2HCl
Chapter 9: Stoichiometry
1. Given the following balanced reaction:
2Fe (s) + 3CuCl2 (aq)  3Cu (s) + 2FeCl3 (aq)
a. How many moles of copper will be produced by the complete reaction of 5.0 moles of
iron?
5.0 𝑚𝑜𝑙𝑒 𝐹𝑒 3 𝑚𝑜𝑙𝑒 𝐶𝑢
𝑥
= 7.5 𝑚𝑜𝑙𝑒 𝐶𝑢
1
2 𝑚𝑜𝑙𝑒 𝐹𝑒
b. How many moles of iron chloride will be produced by the complete reaction of 12.0
grams of iron?
12.0 𝑔 𝐹𝑒 1 𝑚𝑜𝑙𝑒 𝐹𝑒 2 𝑚𝑜𝑙𝑒 𝐹𝑒𝐶𝑙3
𝑥
𝑥
= 0.215 𝑚𝑜𝑙𝑒 𝐹𝑒𝐶𝑙3
1
55.8 𝑔
2 𝑚𝑜𝑙𝑒 𝐹𝑒
c. How many grams of copper will be produced by the complete reaction of 0.25 moles of
copper (II) chloride?
0.25 𝑚𝑜𝑙𝑒 𝐶𝑢𝐶𝑙2
3 𝑚𝑜𝑙𝑒 𝐶𝑢
63.5 𝑔
𝑥
𝑥
= 16 𝑔 𝐶𝑢
1
3 𝑚𝑜𝑙𝑒 𝐶𝑢𝐶𝑙2 1 𝑚𝑜𝑙𝑒 𝐶𝑢
d. How many grams of copper will be produced by the complete reaction of 27.5 grams of
iron?
27.5 𝑔
1 𝑚𝑜𝑙𝑒 𝐹𝑒
3 𝑚𝑜𝑙𝑒 𝐶𝑢
63.5 𝑔
𝑥
𝑥
𝑥
= 46.9 g Cu
1
55.8 𝑔
2 𝑚𝑜𝑙𝑒 𝐹𝑒
1 𝑚𝑜𝑙𝑒