Formulas to memorize:
Percent error = | literature value – measured value | x 100 =
D = m/v
Unit 1 Objectives:
•
Be able to name the correct unit for common chemistry measurements (i.e., length = meter, mass = gram, temp = K)
•
Distinguish between precision and accuracy
•
Be able to count significant figures and use them accurately in math problems
•
Be able to calculate percent error from lab data.
•
Be able to visually identify and correctly use lab equipment and safety equipment.
•
Be able to distinguish between qualitative and quantitative descriptions
Practice Questions:
1. The following are descriptions of our classroom. Choose the one that is a quantitative description:
A. 20 meters wide B. white walls C. large D. full of students ___2. Energy in a car
2. Is a measure of the average kinetic energy of the particles making up a substance:
A. velocity B. temperature C. mass D. density
3. When measuring in lab, your readings should have ___ estimated digits:
4. In lab, Cathy measures a container that has an actual mass of 25.23 g, she records masses of 32.25 and 31.98 g.
Her results are:
A. accurate B. precise
5. The symbol for lead is: Pb
6. The name for the element Na is: sodium
C. precise & accurate D. neither precise nor accurate
7. The balances in lab measure: mass
8. In lab you need to accurately measure out 5.2 ml of water, you would use a(n)_ graduated cylinder __ for the best results.
9. How many significant digits in the following numbers? (2 each)
A. 100 _____ 1 _____ B. 0.020010 ____ 5 ____ 4 ________
10. Perform the following calculations with your calculator. Put your answers in the correct number of significant figures and include the proper units.
A.
7.4 mm x 8.902 mm = 66
B.
8.765 g + 14.0 g = 22.8
C.
97 ml + 14.402 ml = 111
D.
9.84 g / 34 ml = 0.29
11. A student measures the density of cork and finds it to be 0.32 g/cm
3
. The actual density is 0.24 g/cm
3
. What is her percent error? 33%
12.
Convert 135 g to oz (1 oz = 28.35 g) 4.76
13.
Convert 4.2 gallons to L. (1 gal = 3.8 L)
14.
Convert 53,000,000 mm
2
to m
2
.
16
53
15.
Convert to 11.4
16.
Convert 45 miles/hr to cm/sec. 2.0x10
3
Unit 2 Objectives:
•
Be able to use the density formula to solve for any of the variables.
•
Distinguish between endothermic and exothermic reactions, physical and chemical changes, physical and chemical properties, and extensive and intensive properties
•
Name the three indicators that a chemical change has occurred
•
Distinguish between elements, compounds, heterogeneous and homogenous solutions
•
Recognize organic compounds
1. m = 42 g v = 30.0 cm
3
What is the density (in g/cm
3
2. D = 0.823 g/cm
3
)?
m = 4.2 kg What is the volume (in cm
3
)?
1.4 g/cm
3
4200g/.823 cm
3
= 5100 g/cm
3

For each of the following, indicate if the phrase refers to an exothermic or endothermic reaction.
4.
releases exo-
5.
absorbs endo-
6.
products have more potential energy than reactants endo-
7.
8.
products have less potential energy than reactants feels hot to the touch
9.
feels cool to the touch exoexoendo-
Classify the following as chemical or physical changes:
10.
burning coal C
11.
tearing paper P
12.
exploding TNT C
Classify the following as physical or chemical properties:
15.
density P
16.
melting point P
13.
14.
17.
18.
dissolving sugar digesting food length flammability
P
C
P
C
Classify the following as extrinsic or intrinsic properties:
19.
mass
20.
color
21.
ductility ex in in
22.
23.
length melting point
Ex
In
Color change, temp. change, appearance of a new substance
24. Name three indicators that a chemical change has occurred.
Classify the following as elements or compounds
25.
calcium element
26.
water
27.
sugar
28.
aluminum compound compound element
Classify the following as organic or inorganic compounds:
29. C
6
H
12
O org
6 2
O
inorg
31. HCl
inorg
32. C
2
H
5
O
2 org
N

Unit 3 Objectives:
•
Describe the 4 major historical atomic models.
•
Explain the experiments or discoveries made by certain scientists
•
Relate the experimental results to the conclusions made by those scientists
•
Write electron configurations for elements
•
Use the periodic table to predict values for the 4 quantum numbers
•
Relate frequency, wavelength, and energy mathematically
•
Calculate average atomic mass for a sample of isotopes
•
Relate atomic structure to atomic/ionic properties (charge, mass)
Match the name of the person with the accomplishment. You may use each name once, more than once, or not at all.
1. Discovered electrons B
2. First to proposed the idea that all matter is made of atoms. E
3. Discovered the neutron.
G
4. Said that the nucleus is solid, but the atom is mostly space
5.
A
Found the exact charge of an electron. C
6.
Found the mass to charge ratio of an electron B
7.
Introduced the modern atomic theory D
8.
Proposed the planetary (solar system) model of the atom F
9.
Proposed the plum pudding/chocolate chip
A.
B.
Rutherford
J.J. Thompson
C.
Millikan
D.
Dalton
E.
Democritus
F.
Bohr
G.
Chadwick model of an atom B
10.
proposed the billiard ball model of an atom E/D
True or false?
11.
Electrons and neutrons have similar masses. F
12.
Atoms (by definition) are always neutral. T
13.
An ion is a charged atom. T
14.
The number of neutrons in an atom is equal to the number of protons.
F
15.
Rutherford used a radioactive substance as his source of particles for his gold foil experiment.
T
16.
In Rutherford’s experiment, the alpha particles were negatively charged. F
17.
If wave A has a larger wavelength than wave B, it should also have a greater frequency.
F
18.
The set of wavelengths absorbed or emitted by an element is its spectrum. T
19.
Black light is composed of all the colors of the rainbow. F (black is the absence of colors)
20.
Democritus’ beliefs about the nature of matter were well accepted at the time.
F
21.
If an ion contains 2 more protons than electrons, it will have a +2 charge. T
Complete the following chart. You will need to use the periodic table.
Element (use Atomic mass Atomic number Protons isotopic notation)
13
C
+4
13 6 6
Electrons Neutrons
2 7
23
Na 23 11 11
35
Cl
-2
35 17 17
108
Ag
+2
108 47 47
11 12
19 18
45 61
22.
When an atom absorbs energy, what happens to the electrons (where do they move)? Away from the nucleus
23.
A. Describe Rutherford’s gold foil experiment. What evidence led him to conclude: a.
the atom is mostly empty space.
Most alphas passed through undeflected

b.
The atom has a positively charged nucleus. Some alphas deflected and some severely
24.
List two ways an electron is different from a proton. Mass, spin, charge, composition, electrons not subject to the strong nuclear force.
Calculations: Show all work and circle your final answer. Don’t forget to put units on your final answer.
25.
Find the average atomic mass of a fictitious element, X, if 89.6% of its atoms have an atomic mass of 124.988 u and 10.4% have an atomic mass of 123.958 u. 124.881 u
26.
What is the frequency of a wave that has a wavelength of 4.24 x 10
-9
m? 7.08 x 10
-16
Hz
27.
How much energy is contained in a form of electromagnetic energy with a frequency of 3.62 x 10
-5
Hz?
2.39x10
-38
J a. When an atom jumps down from the 3 nd
What is the frequency of light emitted?
energy level back to the 1 st f = 6.6 x 10
14
Hz
, it releases 4.4 x 10
-19
joules of energy. b.What is the wavelength associated with this frequency? 4.5 x 10
-7
m c. What color is the light (refer to the chart in your course packet)?
Visible light (these are approximate)
Violet – 4.0 x 10
Indigo/Blue – 4.2 x 10
-7
Green – 4.6 x 10
-7
– 4.2 x 10
-7 m
– 4.6 x 10
-7
– 5.0 x 10
-7 m
-7 m
Yellow – 5.0 x 10
-7
Orange – 5.8 x 10
-7
Red – 6.3 x 10
-7
– 5.8 x 10
– 6.3 x 10
– 7.0 x 10
-7 m
-7
-7 m m
28.
The photoelectric effect describes the emission of electrons from certain metals when they are radiated by a light source. The minimum energy needed to remove an electron from the metal is called the work function and is a characteristic of the metal. For example, chromium (Cr) metal will emit electrons when the wavelength of the radiation is 284 nm or less. Calculate the work function for chromium. (284 nm = 2.84 x 10
-7 m). Hint:
This is a 2-part question! W = 7.00 x 10
-19
J
29.
What does each of the four quantum numbers designate (n, l, m, and s)? n = energy level, l = sublevel, m = orbital, s = spin up/down
30.
State Heisenberg’s uncertainty principle. You cannot know exactly position and momentum of a particle simultaneously.
31.
State Pauli’s exclusion principle. No two electrons in an atom can have the same 4 quantum number values.
32.
Two electrons in the same orbital with different spins are said to be _ degenerate (equal in energy) _.
33.
What was de Broglie’s contribution to chemistry?
Proposed that matter can have wave-like properties.
34.
Who developed the wave-mechanical model of the atom?
Erwin Schrodinger
35.
Draw the orbital diagrams for the following elements ( try this online utility!
): a.
Rh ( skip ) b. Br c. Li d. O
36.
Give the symbols for the following elements: a.
1s
2
2s
2
2p
6
3s
2
3p
6
4s
2
3d
3 b.
1s c.
1s
2 d.
1s
2
2
2s
2s
2s
2
2
2
2p
2p
4
6
3s
2
3p
6
4s
2
3d
10
4p
6
5s
2
4d
10
5p
6
6s
2
4f
14
5d
8
V
Be
O
Pt

37.
How many electrons can be held in a single orbital? 2
38.
How many orbitals are in an s sublevel? 1 p sublevel? 3 d sublevel? 5 f sublevel? 7
39.
What is the oxidation state (charge) for the following elements? e.
S -2 -1 c. Ge +4 d. Be +2 e. Rb +1
40.
Look at the periodic table. Name one metal. _______________ Name one non-metal. _____________. Name one metalloid. ___________________. Name one transition metal. ______________________. Name one noble gas. __________________. (For the test, be able to identify where each is located on the periodic table).
Unit 4 Objectives:
•
List the seven diatomic molecules.
•
Know the formulas and charges of the 9 polyatomic ions on your homework sheet.
•
Know the formulas and names of the 7 acids on you homework sheet.
•
Know the difference between a cation and an anion.
•
Determine the number of valence electrons in an element
•
Draw Lewis structures for elements, molecules (covalent compounds), and ionic compounds
•
Use electronegativity differences to predict whether a bond will be ionic or covalent.
•
Describe how a spectrophotometer works o Describe how to prepare and use a standard curve o Give common error sources for spectrophotometric measurements
1.
2.
a.
b.
c.
d.
e.
Write the formulas for the 7 diatomic molecules. H
2
, N
2
, O
2
, F
2
, Cl
2
, I
2
, Br
2
Name the following:
NH
4
+1
HCO
NO
SO
3
-1
4
-2
__________________
3
-1
_________________
__________________
__________________
HCl ______________________ f.
g.
h.
i.
HI _______________________
HNO
3
______________________
CH
3
HSO
COOH ____________________
4
-
____________________
Write the formulas for the following: 3.
a.
b.
c.
d.
e.
f.
g.
Acetate ________________________
Hydroxide ______________________
Carbonate _______________________
Phosphate _________________________
Hydrobromic acid ______________
Sulfuric acid __________________
Perchloric acid ________________
4. Propose Lewis Structures for the following compounds:
O
N a b c d
O
-1
H H
O
S : H P :
H B
O O
5. Answer these questions about spectrophotometry
H a. Fingerprints or smudges can scatter light, causing the absorbance to be read incorrectly.
H b. The purpose of the blank is to remove the absorbances of other compounds you aren’t interested in from the absorbance spectrum for the solution, leaving only your species of interest. In other words, it sets the absorbance to 0 for a solution that doesn’t contain your unknown. c. About .55 (draw over from the Y axis to the line, then drop down to the X).