L2 CHEMISTRY MIDTERM REVIEW
Name
KEY
ability to flow ability to compress
A. Properties of Matter (Text Chapter 2)
Properties – Define the following:
Forces
shape
volume
motion
1. solid-
v. strong
def.
def
vibrate
no
no
2.
strong
indef.
Def
slide
yes
no
weak
indef
indef.
yes
yes
liquid-
3. gas-
Rapid random straight
4. physical properties- property of matter that can be observed or measured without changing composition
examples: color, odor, melting & Boiling points, density
5.
chemical properties- of matter describe its “potential” to undergo some chemical change or reaction
examples: ability to burn, ability to rust
6.
physical changes- change appearance of matter but not its composition
examples: cutting, grinding, tearing, ALL PHASE CHANGES!! CHANGES IN STATE OF MATTER (gas to liquid etc.)
7. chemical changes- changes the composition of matter (HAVE SOMETHING NEW)
indicators of chemical change
CHANGE IN COLOR/ODOR,; ENERGY CHANGE; FORMATION OF GAS /PRECIPITATE;
List 3 elements, 3 compounds and 3 mixtures:
8. Elements:
9. Compounds:
10. Mixtures:
11.
Draw a picture using symbols that could represent the molecular view of a sample of EACH of the following:
Mixture
Compound
Substance (one kind of matter)
Homogeneous mixture (solution)
Element
Heterogeneous mixture
B. Dimensional Analysis (Conversion Factors) and Significant Figures (Text Chapter 3)
12. How many significant figures are in the following:
a. 0.0340 g (3 )
b. 2340 L (3 )
c. 5.032 cm ( 4 )
d. 46.000 kg ( 5 )
e. 5000 mm ( 1 )
f. 5300. mL ( 2 )
RULES: When multiplying and dividing: Least number of sig. figs.
When adding and subtracting: Least place value.
13. Report the following to the correct number of significant figures and with correct units:
a. 2.7m3 / 5.27m
b. 4.5mm x 9.56mm
o.51m2
43mm2
e. 89.5kg - 2.0kg
f. (2.1 x 102m)(3.019 x 104m)
105
d. 4.00kg x 2.658m
7.6cm
10.6kgm
g. (3.20 x 10-3m) ÷ (2.00 x 104s)
6.3x106m2
87.5kg
megaM
106
c. 2.0cm + 5.62cm
kilok
103
104
hectoh
102
decada
101
Base
m,L,g
1
decid
10-1
1.60x10-7m/s
centic
10-2
millim
10-3
10-4
10-5
microµ
10-6
10-7
10-8
14. Write the metric equivalents for each of the following:
a. 1.00 km =
0.001 Mm
b. 4.26 cm =
42600µm
d. 25.3 km =
2,530,000 cm
e. 5.42 L =
542 cL
c. 1.23 g =
f. 4.30 g =
0.0043 kg
Use the following conversion factors to answer the questions below.
Write out all conversion factors necessary to solve each problem.
You may need to use some additional metric conversions to solve the problems.
453.6 g = 1 lb
1 mi = 1.609 km
1 gal = 3.785 L
5280 ft = 1 mi
1 in = 2.54 cm
15. How many kilometers are in 4.51 miles?
4.51miles X
1
1.609 km = 7.26 km
1 mile
16. How many gallons are in a 0.524-L bucket of water?
0.524 L
1
X
1 gallon = 0.138 gallons
3.785 L
123 cg
nanon
10-9
453.6 g = 1 lb
1 mi = 1.609 km
1 gal = 3.785 L
5280 ft = 1 mi
1 in = 2.54 cm
17. If an inchworm travels 3.456 x 106 inches, how many miles has he gone?
3.456x106 inches
1
X 1foot X 1 mile = 54.55miles
12 inches 5280 foot
18. You receive a 523 kg package in the mail. If it costs 5.00 cents per pound to ship, how much money was needed to
ship your package?
523 kg X 1000 g X 1 pound X 5.00 cents = 5.76x103 cents
1
1 kg
453.6 g
1 pound
1 mole of atoms = 6.022 x 1023 atoms
1 mole of atoms = mass in grams on periodic table
19. How many moles are in 5.36 grams of sodium?
5.36 g Na X 1 mole = 0.233 moles Na
1
23g
20. How many moles of magnesium are in a sample containing 3.54 x 1023 magnesium atoms?
3.54x1023 atoms X _____1 mole_____ = 0.588 moles Mg
1
6.02x1023 atoms
21. How many grams will 1.25 x 1024 atoms of sulfur weigh?
1.25x1024 atoms X _____1 mole_____ X 32 grams = 66.4 g S
1
6.02x1023 atoms
1 mole
22. How many atoms are in 2.56 moles of iron?
2.56 moles X 6.02x1023 atoms = 1.54x1024 atoms Fe
1
1 mole
23. How many grams of carbon are in a 2.53 mole sample of carbon?
2.53 moles X 12 g = 30.4 g C
1
1mole
24. You and your lab partner collected the following data during a lab:
Mass of large graduated cylinder
Mass of large graduated cylinder w/ water
Volume of water
Calculate the density of the water:
D = mass/volume
Mass = 183.67 g
-118.92 8
64.75 g
118.92 g
183.67 g
65.1 mL
D = 64.75 g = 0.995 g/ml
65.1 ml
If water at that temperature has an actual density of 0.998 g/mL, calculate the percent error for your
measurements.
% error = (|Your Result - Accepted Value| / Accepted Value) x 100
% error = (|0.995g/ml – 0.998g/ml| ÷ 0.988g/ml ) X 100 = 0.3%
25. A class measured the density of an unknown liquid. Their measurements were as follows:
0.789 g/mL
0.792 g/mL
0.788 g/mL
The liquid was determined to be ethyl alcohol, which has an accepted density of 0.785 g/mL. Are the class data
accurate, precise, neither, or both? What was the percent error for the class data?
Find the average of the class data and use this value to determine if the class data is accurate, precise, both or neither.
Use the class average to determine percent error for the class as well.
0.789 g/mL
0.792 g/mL
+0.788 g/mL
2.369 ÷ 3 = 0.7897 g/ml
(|0.7897g/ml – 0.785g/ml| ÷ 0.785g/ml) X 100 = 0.6%
The class data is both precise (repeatable) and accurate (close to the
accepted value).
C. Atoms, Isotopes, Ions (Text Chapter 4,5,7)
Definitions- Define or fill in the following:
33. atomic number34. mass number-
number of protons in a nucleus
number of protons + number of neutrons in the nucleus of an atom
35. isotope-
atoms that contain the same number of protons but a different number of neutrons
36. ion-
charged atom or group of atoms formed by atom(s) gaining or losing electron(s) {unequal number of
protons and electrons}
37. protons have a
positive charge and are found
inside the nucleus .
38. neutrons have a
neutral charge and are found
inside the nucleus
39. electrons have a
negative
40. The protons
and
The
charge and are found surrounding the nucleus
neutrons
electrons
.
.
are about the same size and make up the mass of the atom.
are MUCH smaller and do not really affect the mass of the atom.
41. Two isotopes are the same in this way:
same number of protons, electrons and properties
. They are different in this way:
different number of neutrons, mass number and mass .
42. An ion is different from an atom because it has gained or lost
electrons . It therefore has a
whereas an atom is neutral.
43. Complete the following table:
Symbol
Name
O-2
Oxide ion
8
16
8
8
10
Ba
Barium atom
56
138
56
82
56
2
1𝐻
Hydrogen-two
1
2
1
1
1
carbon-fourteen
6
14
6
8
6
40
2+
20𝐶𝑎
Calcium ion
20
20
20
20
18
S-2
sulfide ion
16
32
16
16
18
14
C
Atomic # Mass # # p+ # n # e-
charge
44. Another name for a positive ion is cation .
45. Another name for a negative ion is anion.
46. Metals lose electrons to form
positive .
47. Nonmetals __gain__ electrons to form
anions .
48. Explain the difference between the mass number of an element and the average atomic mass of the element.
Mass number is the total number of protons + neutrons in the nucleus of an atom. (whole number)
Whereas Average atomic mass is the AVERAGE mass of all the isotopes of an element (usually decimal)
Problems
49. A sample of silver is 52.0% 107Ag and 48.0% 108Ag. Calculate its average atomic mass in amu.
O.52 X 107 = 55.64
0.48 X 108 = +51.84
107.48 amu
50. Isotopic data for lead is below. Use that to calculate the average atomic mass of lead in amu.
Isotope Percent Abundance
204
0.0137 X 204 =
2.795
Pb
1.37%
206
0.2626 X 206 = 54.9
Pb
26.26%
207
0.2028 X 207 = 43.10
Pb
20.82%
208
0.5155 X 208 = +107.2
Pb
51.55%
207.185 amu
51. Tell what each of the following scientists contributed to the atomic theory (give details of how they came up with
their theories from experiments). Additionally, draw a picture or describe the atomic model that resulted from each
person’s work:
John Dalton – All elements are composed of submicroscopic indivisible particles called atoms.
Atoms of the same element are identical. Atoms of different elements are different.
Atoms of different elements can physically mix together or chemically combine.
Chemical reactions occur when atoms are separated, joined, or rearranged.
Atomic Model: Atom is a solid ball.
J.J. Thomson – English physicist discovered electrons in 1897. Experimented with flow of electron current
through gases. Electrodes connected to high energy electricity source. Positive side – anode
Negative side- cathode. Cathode ray – glowing beam, which travels from the cathode to the anode.
Thomson proposed that the cathode ray had a stream of small negatively charged particles.
Atomic Model: Plum Pudding Theory – Electrons are imbedded in a positive sphere.
Ernest Rutherford –
1911 – Gold Foil Experiment Proposed that almost all of the mass and all the positive charge are concentrated in
a small region at the center of the atom called the nucleus. Nucleus- the central core of the atom, composed of
protons and neutrons.
Atomic Model: Positive dense tiny nucleus is surrounded by electrons in empty space.
Niels Bohr –
1913, Danish physicist. Electrons have a fixed path.
Atomic Model: Planetary Model.
52. How are wavelength and frequency related (directly or inversely)? Inversely { Frequency ˅ wavelength ˄ }
53. How are wavelength and energy related (directly or inversely)?
Inversely ( greater energy = smaller wavelength)
54. Does ROYGBIV go from high to low or low to high energy? Low energy to high energy
55. What is an atomic emission spectrum?
Photon is just the name for a quantum of light
Electron Transition – when an electron moves from one level to another
a. When an electron transitions to a higher energy level, a photon is absorbed.
b. When an electron transitions to a lower energy level, a photon is emitted.
The emission spectrum of a chemical element or chemical compound is the spectrum of frequencies of
electromagnetic radiation emitted by the element's atoms or the compound's molecules when they are returned to a
lower energy state.
a. You should know how a line on the atomic emission spectrum relates to the ground and excited states
of the atom. Draw a picture of the energy levels in an atom and show two electron transitions that
would each emit a photon. Compare the energies of the two photons that would be emitted from these
transitions.
++
D. Electrons In Atoms (Text Chapter 5)
56. Fill in the following charts:
Sub# of
# of
Level
Orbitals
Electrons
s
1
2
p
3
6
d
5
10
f
7
14
Major Energy Level
The sublevels that exist
in this level
1
s
2
s&p
3
s,p &d
4
s, p, d & f
57. Draw a picture of the shape of an orbital in each of the sublevels: s,p and d
S
P
d
58. Give the electron configurations for the following atoms (both regular and box diagram). You may use the shortcut
for any atom in period 3 or larger:
Electron configuration
Orbital Box Diagram
1s2 2s1
a. Lithium (3)
1s
2s
Lewis Dot Diagram: Li ●
# of valence electrons 1
How many energy levels are occupied? 2
How many unpaired e- are in this atom? 1
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 5d1 4f8
b. Terbium (65)
[Xe]
6s
Lewis Dot Diagram:
5d
●
Tb ●
4f
valence electron configuration: 6s2
What is the highest occupied energy level in this atom? 6
c. Iodine (53)
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5px2y2z1
[Kr]
5s
Lewis Dot Diagram:
4d
●●
●● I ●●
●●
5px
py
# of valence electrons 7
Name one element with a similar valence configuration.
d. Rhodium (45)
pz
Any Halogen (group 17; 7A)
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d7
[Kr]
5s
Lewis Dot Diagram:
4d
●
Rh ●
# of valence electrons 2
How many unpaired electrons are in this atom?
E. Periodic Trends (Text Chapter 6)
3
59. For the following pairs of elements, list the ones that have the higher electronegativity:
a. Oxygen and Fluorine - ______F______
(bigger love of electrons will be smaller in size)
b. Lithium and Cesium - _______Li______
across table increases
c. Silicon and Sulfur - _______Si________
down table decreases
d. Boron and Aluminum - _____B_______
60. For the following pairs of elements, list the ones with the smaller atomic radius:
a. Magnesium and Barium - ______Mg______
stronger nuclear charge when going across the table
b. Arsenic and Bromine - _________Br___
less shells as you go down the table
c. Potassium and Zinc - _______Zn_______
d. Neon and Xenon - _______Ne_________
61. Define ionization energy: Energy needed to remove an electron from the outer shell.
a. What are the general trends in ionization energy across a period and down a group?
Across a period it increases because the smaller the atom the harder it is to remove an
electron and you are moving towards the nonmetals which want to gain electrons.
Down a family it decreases because the atoms get larger and the nucleus can not hold onto
the valence electrons as tightly due to shielding effect.
b. Use the above answer to describe why alkali metals are more reactive than alkaline earth metals.
Make sure to state what happens to the electron configuration of metals when they form a
compound.
Alkali metals are more reactive because they only need to lose one electron to obtain a stable
octet. (isoelectronic to a noble gas ie. Noble gas electron configuration)
62. List three properties of metals.
Shiny; malleable; ductile; good conductors
63. List three properties of nonmetals.
Dull; brittle; poor conductors
64. List a possible set of three properties for an element that is a metalloid.
Have properties of both the metals and nonmetals therefore a metalloid could be a semiconductor; shiny and
brittle.
65. Periodic Table Terms: Make sure you are familiar with, and can properly label the terms below on the blank
periodic chart.
Groups/Families
Periods 
Noble Gases
Halogens
Alkali Metals
Lanthanide Series
Actinide Series
Alkaline Earth Met.
Periodic table
group and family numbers
period
representative elements(groups 1,2; 13-18)GroupA
transition elements (groups 3-12) Group B Middle
inner transition elements (
) pulled out
metals (everything left of staircase)
nonmetals (small upper right triangle of staircase)
Metalloids (either side of staircase)
alkali metals
alkaline earth metals
halogens
noble gases
lanthanide series
actinide series
F. Ionic Compounds (Text Chapter 7)
66. When metal atoms react to form a compound, do they form positive ions or negative ions? (+ ion)
a. What happens to the structure of the metal atom in order to become an ion?
Loses outer shell (valence) electrons, shell underneath contains a stable octet; it becomes
ioelectonic to a noble gas.
67. When nonmetal atoms react to form a compound, do they form positive ions or negative ions? (- ion)
b. What happens to the structure of the nonmetal atom in order to become an ion?
Gains electrons to complete its outer shell (stable octet) and become isoelectronic to a noble gas.
68. Draw electron dot structures for any three metal atoms.
Na ●
Mg ●
●Al●
●
●
69. Draw electron dot structures for any three nonmetal atoms.
●●
●●
●●
● N●
●O●●
● F ●●
●
●
●●
70. What is the octet rule?
Having 8 valence electrons (outer shell) most stable and lowest energy.
71. Why do metals form cations, but nonmetals form anions? (Hint: see Periodic Trends, ch. 6)
Metals lose electrons to form cations to get a stable octet. (shell underneath is complete)
Nonmetals gain electrons to complete their outer shell and obtain a stable octet.
Both achieve the electron configuration of a noble gas. Isoelectronic.
72. Use dot structures to show how aluminum forms a compound with sulfur. Draw arrows to show the
transfer of electrons from one atom to another. Write the formula unit that results from this transfer.
x
●●
Al x
●S●●
x
●
●●
S-2
+3
●S●●
Al s-2
Al2S3
+3
-2
x
●●
●
Al s
Al x
●S●●
x
●
73. What is an ionic bond?
Strong electrostatic attraction between oppositely charged particles.
74. Why do ionic compounds have high melting points?
The ionic bond is strong because the crystal structure has no discerning units there is just one kind of
bond the strong attraction between oppositely charged particles.
75. Why do ionic compounds conduct electricity when dissolved in water, but molecular compounds do not?
When ionic compounds dissolve they form ions. When molecular compounds dissolve they form
molecules (no Ions).
G. Names and Formulas (Text Chapter 9)
Writing Formulas- Identify each as ‘I’ for ionic or ‘M’ for molecular, then write the formula.
86. dinitrogen hexabromide
_M__ ____N2Br6____ 91. diphosphorus pentachloride
M P2Cl5___
87. ammonium carbonate
_I_ ___(NH4)2CO3___
92. iron (III) permanganate _I_ __Fe(MnO4)3_____
88. potassium phosphide
_I_ ___K3P______
93. calcium hydroxide _I_ ___Ca(OH)2____
89. mercury (II) acetate
_I_ ___Hg(C2H3O2)2_
94. carbon tetrachloride_M_ CCl4____
90. sulfur trioxide
_M_ ___SO3____
95. barium sulfide
_M_ ___BaS___
Naming Compounds- Identify each as ‘I’ for ionic or ‘M’ for molecular, then write the name.
86. Prefixes are used when naming
Molecular
87. Roman numerals are used to designate the
naming
Ionic
compounds.
88. CaO
_I_
Calcium oxide
89. Mg(OH)2
_I_
Magnesium hydroxide
90. CuC2H3O2 _I_
Copper I acetate
91. HgF2
_I_
Mercury II fluoride
92. K3N
_I_
Potassium nitride
93. Ba3(PO4)2 _I_
Barium phosphate
94. NO3
_M_
Nitrogen trioxide
95. Cl2O
_M_
Dichlorine monoxide
96. (NH4)2S
_I_
Ammonium sulfide
97. Ni2SO4
_I_
Nickle I sulfate
compounds.
Charge
of most transition metals when