Download Chemistry B1A - Bakersfield College

Chemistry B1A
CRN:________
Name:_______________________
Date:________________________
I. Choose the correct answer(s) for the following questions:
1.
The following are examples of physical and chemical changes. Circle the
physical changes.
a. sugar dissolving in water
b. iron rusting
c. melting of sodium chloride
d. wood decaying
2.
Following are properties of substances. Circle the properties that are chemical
properties.
a. Salt substitute , potassium chloride, dissolving in water
b. Pure water freezing at 0oC.
c. Steel wool burning in air
d. Limestone fizzing when it is dissolved in vinegar
e. Bread cooking in the oven
3.
Following are some pure substances and some mixtures. Which are pure
substances?
a. a sugar solution with crystals of sugar at the bottom
b. ink containing a liquid solution with fine carbon particles
c. liquid water and steam at 100oC
d. a sand containing quartz and calcite
4.
Which of the following temperatures might be a good day to go to the beach?
a. 0 K
b. 100 K
c. 100 oF
d. 100 oC
e. 35 oC
5.
Which of the following are metals?
a. Si
b. Ar
c. K
d. Ba
e. Br
The correct formula mass of CaCO3 is:
a. 40 g
b. 100 g
c. 100 amu
d.56 amu
e. none of these
The oxidation number of Cl in HClO3 is:
a. –1
b. +1
c. +3
d. +5
e. +7
6.
7.
8.
Which of the following compounds will evolve a gas when reacted with a strong
acid?
a. CaCO3
b. Na2S
c. CaSO4
d. Na3PO4
e. Na2SO3
9.
Oxidation and reduction are said to;
a. never occur
b. involve neutrons
c. always occur together
d. always involves oxygen
10.
Page 1 was:
a. easy
b. not so easy
c. impossible
d. torturous
II.
Solve discuss or answer with explanation the following questions. Show all
calculations and use appropriate significant figures.
10.
A mineral has a mass of 31.5 g and a volume of 7.9 cm3. The mineral is either
sphalerite (density = 4.0 g/cm3), casserite (density = 6.99 g/cm3) or cinnabar
(density 8.10 g/cm3).. Which is it?
11.
How many significant figures are there in each of the following numbers?
a. 73,000 g_______
12.
b. 73.0000 g________ c. 0.0120 g _________
Consider the following compounds and their densities.
Substance
Density (g/mL)
Isopropyl alcohol
0.785
n-butyl alcohol
0.810
Substance
Toluene
Ethylene glycol
Density (g/mL)
0.866
1.114
You create a column of the liquids in a glass
cylinder. Draw a sketch and indicate which
liquid is at which level in the column. Then
explain what would happen if you did the
following:
a. First you drop a plastic bead that has a density
of 0.24 g/cm3 into the column.
b. You drop a bead in that makes it all the way
to the bottom. What can you say about the
density of this bead?
c. You drop a bead with a volume of 0.043 mL
and a mass of 3.92 x 10-2 g into the column.
What happens?
13.
Some iron wire weighing 5.6 g is placed in a beaker and covered with 15.1 g of
hydrochloric acid. The acid reacts with the metal and gives off hydrogen gas,
which escapes into the surrounding air. After reaction the contents of the beaker
weighs 20.4 g. What is the mass of hydrogen produced? Write the reaction. What
law do you use to solve this problem?
14.
The label on a soda can indicates that the can contains 12 fluid ounces and 355
mL. How could you determine the conversion from ounces to mL from this
information?
15.
Which of the following models represent a(n)
a. Element__________________
b. Compound_______________
c. Mixture_________________
d. Ionic Solid_______________
e. Mixture of elements_______
f. Solid element_____________
g. Solid___________________
h. Liquid__________________
i. Gas containing an element and a compound____________
16.
Dalton explained elements atoms compounds and reactions with his atomic
theory. What were the four postulates of the atomic theory and how does his
theory explain chemical reactions?
17.
An element has two naturally occurring isotopes. The following data was
determined for the element:
Isotopic Mass (amu)
63.390
64.928
Fractional Abundance
0.6909
0.3091
What is the elements atomic weight and what is the element?
18.
In the following equation (unbalanced), (a) give the oxidation state of each
element in the equation, (b) tell what is oxidized and what is reduced, and (c) give
the name of each substance in the equation.
MnO2 +
19.
20.
HCl
 MnCl2 + Cl2 + H2O
Give the systematic names of the following binary compounds:
a. N2O ______________________
b. CCl4 ______________________
c. Na2O______________________
d. CaS_______________________
Write the formula from the following names:
a. dichlorine pentoxide _____________
b. oxygen difluoride_____________
c. copper (I) hydroxide_____________
d. ammonium bromide___________
21.
Write the nuclide symbol for and element containing 12 protons, 12 neutrons and
12 electrons. What would be the nuclide symbol if the element lost two
electrons?
22.
You have in your hand 35.6 g of BaCr2O7. Calculate the following for this
quantity of substance.
a. Moles of BaCr2O7
b. Formula Units of BaCr2O7
c. Total number of atoms present
23.
What six evidences might be observed when two substances are combined?
24.
Potassium superoxide, K2O, is used in rebreathing gas masks to generate oxygen.
If a reaction vessel contains 0.20 mol K2O and 0.25 moles of H2O, what is the
limiting reactant and how much oxygen can be produced?
4 K2O (s) + 2 H2O (l) 
4 KOH (s) + 3 H2O (g)
25.
Using the half reaction method, balance the following redox equation.
CuCl2 (aq) + Al (s)  AlCl3 (aq) + Cu (s)
26.
Using your solubility chart, determine if a reaction occurs when the following set
of substances are combined. If there is a reaction complete it and if none occurs
write NR. Write the total ionic equation and the net ionic equation for each
reaction.
AgNO3 + NaI 
Na2SO4 + Ba(NO3) 
27.
Explain the difference between a weak acid and a strong acid and describe how
you could identify each.
28.
A moth repellent, p-dichlorobenzene, has the composition 49.1% C, 2.7% H, and
48.2 % Cl. Its molecular weight is 112 amu. What is the molecular formula?
29.
Explain the mole concept in chemistry? Why is it important?
30.
Describe the development of the periodic table, how it was originally arranged,
how it is currently arranged, what standard is used to determine the atomic weight
and what information can be determine from the table.
III.
Match the following Definitions
a. ____ Precise
1. Equal protons but unequal neutrons
b. ____ Mixture
2. Mixture that is the same throughout
c. ____ Matter
3. Smallest whole number ratio
d. ____ Metal
4. Sum of atomic masses in grams
e. ____Isotope
5. Exact quantity calculated from an equation
f. ____ Atomic number
6. Substance having luster and usually conducts heat
g. ____Theoretical yield
7. Combination of two or more substances
h. ____Empirical Formula
8. All the substances around us
i. ____ Homegenous mixture
9. Mixture in which components can be identified
j. ____ Heterogeneous mixture
10. Product that should be recovered
k. ____Stoichiometric quantity
11. Data that is reproducible
l. ____ Molar mass
12. Data that is near theoretical number
m. ___ Accurate
13. Number of protons in the nucleus
n. ____ Carbon-12
14. Simplest form of matter
o. ____ Atom
15. Mass standard for the periodic chart