Name____________________________
The Most SuperBad
Chemistry Study Guide
Mrs. Magat
Spring 2012
1
2
Table of Contents
Practice Q
Completed
Page
Standards
Topic
4-5
1a, 1e
6-7
8-9
1b, 1d,
1c
2e
Atomic Number, Atomic Mass and
The Atom (2)
* The Groups and Periodic Trends (4)
10-11
2a, 2b
* Chemical Bonding (3)
12-13
2c,2d
Salt Crystals and Intermolecular
Forces (2)
14-15
3a
Balancing Chemical Equations (2)
16-17
3b, 3c
The Mole (2)
18-19
3d
* Molar Conversions (3)
20-21
3e
* Equation Stoichiometry (3)
22-23
4a, 4b
Pressure and Diffusion (2)
24-25
* The Gas Laws and Celsius Scale (4)
26-27
4c, 4d,
4e
5a, 5d
28-29
5b, 5c
Types of Acids and Bases (2)
30-31
6a, 6b
32-33
6c, 6d
34-35
7a, 7b,
7c
7d
Solute, Solvent and the Dissolving
Process (2)
* Dissolving and Solution
Concentrations (3)
* Exothermic and Endothermic
Reactions (3)
Specific Heat Calculations (2)
36-37
38-39
Lewis Dot Structures (2)
* The pH Scale (3)
8a, 8b,
8c
9a, 9b
* Changing the Rate of a Reaction (4)
Organic Chemistry (2)
44-45
10a,
10b, 10c
11a- 11e
46-55
all
CST Released Questions
40-41
42-43
* Le Chatelier’s Principle (4)
Nuclear Chemistry (2)
3
Video
Completed
Date
Completed
Parent
Signature
Teacher
Signature
Atomic Number, Atomic Mass, and The Atom
Standard:
1a. Students know how to relate the position of an
element in the periodic table to its atomic number and atomic
mass. 1e. Students know the nucleus of the atom is much smaller
than the atom yet contains most of its mass. (2 questions)
Practice Test Q:
1, 2, 5, 10, 11, 12, 18, 19
Video:
http://www.khanacademy.org/science/chemistry/v/elements-and-atoms
http://www.khanacademy.org/science/chemistry/v/introduction-to-the-atom
Online Tutorial:
http://www.learner.org/interactives/periodic/index.html
What is an atom's atomic number?
The number of protons in the nucleus of an atom determines an element's atomic number. In other words,
each element has a unique number that identifies how many protons are in one atom of that element. For
example, all hydrogen atoms, and only hydrogen atoms, contain one proton and have an atomic number of 1.
All carbon atoms, and only carbon atoms, contain six protons and have an atomic number of 6. Oxygen
atoms contain 8 protons and have an atomic number of 8. The atomic number of an element never changes,
meaning that the number of protons in the nucleus of every atom in an element is always the same.
How do we determine an atom's mass number?
All atoms have a mass number which is derived as follows.
How big is an atom?
An atom is incredibly small. The diameter of an atom would have to be increased 200 million times to have
the diameter of a penny. If an apple were enlarged to the size of the Earth, the atoms in the apple would be
the size of cherries.
What is an atom made of? Subatomic particles: Neutron, proton, electrons. A central, dense nucleus
(neutrons & protons) surrounded by electrons. Electrons have a negative charge (-). Protons have a positive
charge (+). The atom is held together by the attraction of electrons and protons. Neutrons are neutral.
Nucleus is Small but Heavy Compared to the Rest of the Atom
Most of the mass of an atom is in the nucleus: the protons and neutrons. The size of the neutron relative to
the size of the atoms is like a penny in the middle of a baseball field. The number of protons and electrons
need to be equal so that the atom has no charge.
4
Periodic Table and the Atom Practice Problems
Complete the table. There is enough information given for each element to determine all missing numbers.
Symbol
23
Atomic
Number
Mass
Number
Number
of
Protons
Number
of
Electrons
Number
of
Neutrons
Na
K
40
19
38
38
F
52
10
20
41
50
18
50
72
131
I
26
Mg
1. An element's or isotope's atomic number tells you _________________________________.
2. An element's or isotope's mass number tells you ___________________________________.
3. The heaviest part of an atom is the ___________________, which contains both _______________
and _________________. The __________________ are found in a cloud surrounding the nucleus.
4. If an atom was a penny inside a football stadium, the penny would represent the
_________________________ and the football stadium would represent _____________________.
5. Which has a higher atomic number?
Helium or Hydrogen ____________________
Magnesium or Manganese __________________
6. Which has a lower atomic mass?
Carbon or Calcium ________________________
Xenon or Radon ___________________________.
7. Generally speaking, how does atomic mass change throughout the periodic table? ________________
____________________________________________________________________________________.
5
The Groups and Periodic Trends
Standard:
1b. Students know how to use the periodic table to identify metals, semimetals, non-metals,
and halogens.1d. Students know how to use the periodic table to determine the number of
electrons available for bonding. 1c. Students know how to use the periodic table to identify
alkali metals, alkaline earth metals and transition metals, trends in ionization energy,
electronegativity, and the relative sizes of ions and atoms. (4 questions)
Practice Test Q:
3, 4, 13, 14, 15, 16, 17
Video:
http://www.khanacademy.org/science/chemistry/v/orbitals
http://www.wiziq.com/tutorial/97462-Periodic-Table-Groups
Online Tutorial:
http://www.chemtutor.com/perich.htm
Metals are malleable, ductile, and have luster; most of the elements on the periodic table are metals. They
oxidize (rust and tarnish) readily and form positive ions (cations). They are excellent conductors of both heat
and electricity. The metals can be broken down
into several groups. Transition metals (also
called the transition elements) are known for
their ability to refract light as a result of their
unpaired electrons. They also have several
possible oxidation states. Ionic solutions of these
metals are usually colored, so these metals are
often used in pigments. Uranium is the
last naturally occurring element; the rest
are man-made. Nonmetals lie to the
right of the staircase and do not
conduct electricity well because they do not have a sea of electrons. All the elemental gases are included in
the nonmetals. Notice that hydrogen is placed with the metals because it has only one valence electron, but it
is a nonmetal. Here are some specific families you should know about, within the three main groups (metals,
nonmetals, and metalloids):Alkali metals (1A)—The most reactive metal family, these must be stored under
oil because they react violently with water! They dissolve and create an alkaline, or basic, solution, hence
their name. Alkaline earth metals (2A)—These also are reactive metals, but
they don’t explode in water; pastes of these are used in
batteries. Halogens (7A)—Known as the “salt formers,”
they are used in modern lighting and always
exist as diatomic molecules in their
elemental form. Noble gases (8A)—Known
for their extremely slow reactivity, these
were once thought to never react; neon, one
of the noble gases, is used to make bright
signs. The number of electrons
available for bonding depends on how
many unpaired valence electrons- Group 1
has 1, Group 2 has 2, Group 13 has 3. Group 14 has 4,
Group 15 has 3 (2 are paired), Group 16 has 2 (4 are paired), Group 17 has 1 and Group 18 has none.
Here we have a summary of the periodic trends. You will need to be able to identify differences in atomic
radius (size of the atom), ionization energy (the energy required to remove an electron from the atom in the
gas phase), and electronegativity (a measure of the attraction an atom has for electrons when it is involved
in a chemical bond).
6
The Groups and Periodic Trends Practice Problems
Identify the following elements as metal, nonmetals, or metalloid.
1) Boron
2) Carbon
3) Gold
4) Lead
5) Hydrogen
Identify the following elements by which group they belong to on the periodic table.
6) Flourine
7) Argon
8) Calcium
9) Potassium
10) Carbon
11) Which of these elements has the largest atomic radius?
a) aluminum
b) calcium
c) fluorine
d) potassium
e) sulfur
12) Which of these elements has the smallest atomic radius?
a) potassium
b) iron
c) arsenic
d) bromine
e) krypton
13) Which of these elements has the highest first ionization energy?
a) oxygen
b) oxygen
c) fluorine
d) carbon
e) boron
14) Which of these elements has the highest electronegativity?
a) lithium
b) nitrogen
c) potassium
d) arsenic
e) beryllium
7
Lewis Dot Structures
Standard: 2e. Students know how to draw Lewis dot structures. (2 questions)
Practice Test Q:
32
Video:
www.youtube.com/watch?v=ulyopnxjAZ8
Online Tutorial:
http://www.chem.ucla.edu/harding/lewisdots.html
The simplest way to represent and describe molecules is to use a Lewis structure. The Lewis structure model
generally follows the octet rule (atoms prefer 8 electrons in their outer shell) and provides a framework to
understand covalent bonding. Lewis structures represent valence electrons as dots and bonding electrons as
lines. Lewis structures do not represent inner electrons; only valence electrons are shown.
Here we give a step-by-step procedure for writing valid Lewis structures for any given molecular formula:
1. Count the total number of valence electrons by summing the group numbers of all the atoms. If there
is a net positive charge, subtract that number from the total electron count. If there is a net negative
charge, add that number to the total electron count.
2. Draw single bonds to form the desired connectivity.
3. Add lone pairs and multiple bonds, keeping the octet rule in mind.
4. Add formal charges as needed.
An example:
Some Common Bonding Motifs in Organic
Molecules
You have seen that carbon tends to form four
bonds, nitrogen three, oxygen two, and
hydrogen/halogens one (remember also: as the
number of bonds of an atom decreases, the
number of its lone pairs increases). The number
of bonds that a neutral atom forms is called its
valence. Hence carbon is tetravalent, nitrogen is
trivalent, oxygen is divalent, and so on.
However, a carbon atom, for example, can be
tetravalent in a number of different ways. The
following chart shows a number of common
bonding motifs for C, N, O and H
The majority of motifs in the table above obey
the octet rule.
While atoms can occasionally be short of a full octet,
elements in the first two rows of the periodic table
can never exceed the octet. Students often make
the dreaded mistake of drawing pentavalent
carbons. Never do this!
8
Lewis Structure Practice Problems
Draw the Lewis structures for the following compounds:
1)
PBr3
2)
N2H2
3)
CH3OH
4)
NO2-1
5)
C2H4
6)
BSF
7)
HBr
8)
C2H5OH (ethanol)
9)
N2F4
10)
SF6
9
Chemical Bonding
Standard: 2a. Students know atoms combine to form molecules by sharing electrons to form covalent or
metallic bonds or by exchanging electrons to form ionic bonds. 2b. Students know chemical bonds
between atoms in molecules such as H2, CH4, NH3, H2CCH2, N2, Cl2 and many large biological molecules
are covalent. (3 questions)
Practice Test Q:
24, 25, 26, 27, 28
Video:
http://www.khanacademy.org/science/chemistry/v/ionic--covalent--and-metallic-bonds
Online Tutorial:
http://www.syvum.com/cgi/online/serve.cgi/squizzes/chem/bonds1.tdf?0
Atoms are the building blocks of all substances. But what is it that keeps atoms connected together? They are
held together by CHEMICAL BONDS, strong attractive forces between atoms. Without these ties that bind,
the universe would be nothing more than a mass chaos of individual atoms.
So what constitutes a chemical bond? A bond is formed when electrons from two atoms interact with each
other and their atoms become joined. The electrons that interact with each other are VALENCE
ELECTRONS, the ones that reside in the outermost electron shell of an atom.
There are two main types of bonding discussed here. A COVALENT BOND results when two atoms
"share" valence electrons between them. An IONIC BOND occurs when one atom gains a valence electron
from a different atom, forming a negative ion (ANION) and a positive ion (CATION), respectively. These
oppositely charged ions are attracted to each other, forming an ionic bond.
Why are chemical bonds important? The type of chemical bond that occurs in a molecule or
substance in part defines its properties. For example, consider sodium chloride (NaCl) and
hydrogen chloride(HCl). Both substances contain chlorine, but NaCl is the white solid
crystalline substance sprinkled on French fries, and HCl is a foul smelling gas.(note: when
this gas is dissolved in water, it forms a solution known as hydrochloric acid. This is the
acid that your stomach uses to digest food.)
Table salt
How can this be if both materials have chlorine in them? The chemical bonding that takes
(NaCl)
place in NaCl is different than that in HCl. This gives NaCl and HCl very different
structures, appearances, and properties.What other differences are noticeable among
molecules that result from different types of chemical bonding? Think about what happens when a bunch of
sodium and chlorine ions join together to form rock salt. If we hit this with a hammer, it shatters into tiny
pieces. It does this because the bonds between the atoms in rock salt are ionic. The particles are arranged in
such a way that they line up along rows of positive and negative charge. Under enough stress, the salt
crystals break along those lines into much smaller pieces.
The bonds that hold the carbon and hydrogen atoms in
rubber together, on the other hand, are not ionic but covalent.
Each carbon atom shares four of its outermost electrons with its
immediate neighbors. Under stress, the bonds stretch, then snap
back as each atom pulls on the shared electrons. And that's the
way the ball bounces. There is a third type of bonding, called
METALLIC BONDING. As the name implies, metallic bonding
usually occurs in metals, such as copper. A piece of copper metal These silver atoms are joined by metallic
bonds.
has a certain arrangement of copper atoms. The valence electrons
of these atoms are free to move about the piece of metal and are
attracted to the positive cores of copper, thus holding the atoms together.Essential to understanding all types
of chemical bonding is realizing that all bonds use electron "glue." Every substance is made up of atoms, and
all atoms are surrounded by the charged particles called electrons. In large part, the difference between
materials as diverse as diamonds and pencils is how they're glued together.
10
Chemical Bonding Practice Problems
1) How are ionic bonds and covalent bonds different?
2) Describe the relationship between the length of a bond and the strength of that bond.
3) Identify the type(s) of bond(s) found in the following molecules:
a. CCl4
___________________________
b. Li2O
__________________________
c. NF3
___________________________
d. CaSO4
___________________________
e. SO2
___________________________
f.
Mg(OH)2 ___________________________
4) Determine if the bond between atoms in each example below is nonpolar covalent, polar covalent, or
ionic.
a. H2
_______________
b. PCl
________________
c. F2
_______________
d. NaBr
________________
e. NF
________________
f. MgO
______________
g. CH
_______________
h. HCl
________________
5) Proteins are large biological molecules. What type of bonds do they form?
6) Carbohydrates are large biological molecules. What type of bonds do they form?
7) Lipids are large biological molecules. What type of bonds do they form?
8) Sugars are large biological molecules. What type of bonds do they form?
11
Salt Crystals and Intermolecular Forces
Standard: 2c. Students know salt crystals, such as NaCl, are repeating patterns of positive and negative ions
held together by electrostatic attraction. 2d. Students know the atoms and molecules in liquids move in a
random pattern relative to one another because the intermolecular forces are too weak to hold the atoms or
molecules in a solid form. (2 questions)
Practice Test Q:
29, 30, 31
Video: http://www.khanacademy.org/science/chemistry/v/covalent-networks--metallic--and-ionic-crystals
http://www.khanacademy.org/science/chemistry/v/van-der-waals-forces
http://www.khanacademy.org/science/chemistry/v/vapor-pressure
Online Tutorial:
http://misterguch.brinkster.net/intermolecularforces.html
Relative strength of Intermolecular Forces:
 Intermolecular forces (dispersion forces, dipole-dipole interactions and hydrogen
bonds) are much weaker than intramolecular forces (covalent bonds, ionic bonds
or metallic bonds)
 Dispersion forces are the weakest intermolecular force (one hundredth-one
thousandth the strength of a covalent bond), hydrogen bonds are the strongest
intermolecular force (about one-tenth the strength of a covalent bond).
 dispersion forces < dipole-dipole interactions < hydrogen bonds
Dispersion Forces (London Forces, Weak Intermolecular Forces, van der Waal's
Forces)
 are very weak forces of attraction between molecules resulting from:
1. momentary dipoles occurring due to uneven electron distributions in neighboring molecules
as they approach one another
2. the weak residual attraction of the nuclei in one molecule for the electrons in a neighboring
molecule.
 The more electrons that are present in the molecule, the stronger the dispersion forces will be.
 Dispersion forces are the only type of intermolecular force operating between non-polar molecules,
for example, dispersion forces operate between hydrogen (H2) molecules, chlorine (Cl2) molecules,
carbon dioxide (CO2) molecules, dinitrogen tetroxide (N2O4) molecules and methane (CH4)
molecules.
Dipole-dipole Interactions
 are stronger intermolecular forces than Dispersion forces
 occur between molecules that have permanent net dipoles (polar
molecules), for example, dipole-dipole interactions occur between SCl2 molecules, PCl3 molecules and
CH3Cl molecules.
If the permanent net dipole within the polar molecules results from a covalent bond between a hydrogen
atom and either fluorine, oxygen or nitrogen, the resulting intermolecular force is referred to as a hydrogen
bond (see below).
 The partial positive charge on one molecule is electrostatically attracted to the partial negative charge on a
neighboring molecule.
Hydrogen bonds
 are a stronger intermolecular force than either Dispersion forces or
dipole-dipole interactions since the hydrogen nucleus is extremely small and positively charged and
fluorine, oxygen and nitrogen being very electronegative so that the electron on the hydrogen atom is
strongly attracted to the fluorine, oxygen or nitrogen atom, leaving a highly localized positive charge
on the hydrogen atom and highly negative localized charge on the fluorine, oxygen or nitrogen atom.
12
Types of Intermolecular Forces Practice Problems
What is the strongest intermolecular force present for each of the following compounds?
1)
water _____________________________________
2)
carbon tetrachloride _____________________________________
3)
ammonia _____________________________________
4)
carbon dioxide _____________________________________
5)
phosphorus trichloride _____________________________________
6)
nitrogen _____________________________________
7)
ethane (C2H6) _____________________________________
8)
acetone (CH2O) _____________________________________
9)
methanol (CH3OH) _____________________________________
10)
borane (BH3) _____________________________________
For each of the following compounds, determine the main intermolecular force. You may find it useful to
draw Lewis structures for some of these molecules:
11)
nitrogen ________________________________
12)
carbon tetrachloride ________________________________
13)
H2S ________________________________
14)
sulfur monoxide ________________________________
15)
N2H2 ________________________________
16)
boron trihydride ________________________________
17)
CH4O ________________________________
18)
SiH2O ________________________________
13
Balancing Chemical Equations
Standard:
3a. Students know how to describe chemical reactions by writing balanced
equations. (2 questions)
78, 79, 80, 81
http://www.khanacademy.org/science/chemistry/v/balancing-chemical-equations
Practice Test Q:
Video:
Online Tutorial:
http://education.jlab.org/elementbalancing/index.html
You may remember that the law of conservation of mass says that
matter is neither created nor destroyed during a chemical reaction.
This means that all chemical reactions must be balanced—the number
of atoms, moles, and ultimately the total mass must be conserved
during a chemical process. Here are the rules to follow when balancing
equations:
1) Determine the correct formulas for all the reactants and products
in the reaction.
2) Begin balancing with the most complicated-looking group. A polyatomic ion that appears unchanged on
both sides of the equation can be counted as a single unit.
3) Save the elemental (single elements) reactant and products for last, especially if it is hydrogen or oxygen.
Keep your eye out for diatomic molecules such as oxygen, hydrogen, and the halogens.
4) If you get stuck, double the most complicated-looking group and try again.
5) Finally, make sure that all coefficients are in the lowest-possible ratio.
6) Know when to quit! None of the reactions you will encounter will be that difficult. If the coefficients are
getting wild, double-check what you’ve done since you may have a simple mistake.
When balancing reactions, keep your hands off the subscripts! Use only coefficients to balance chemical equations.
Now let’s try an example. When you solve it yourself, make sure to follow the steps!
Example #1
Write the balanced equation for the reaction between chlorine and sodium bromide, which produces bromine and sodium
chloride.
Explanation
First write the chemical formulas, be on the lookout for the diatomic elements (such as Cl 2):
Next, find the reagent with the scariest subscripts. In this case, start with Cl2. You need a coefficient of 2 in front of
NaCl, which then requires a coefficient of 2 in front of NaBr. The balanced equation becomes
Cl2 + 2NaBr
Br2 + 2NaCl
Finally, count up everything to make sure you balanced the equation correctly. You have 2 chlorine atoms, 2 sodium,
and 2 bromine on the reactant side and 2 bromine, 2 sodium, and 2 chlorine on the product side.
Example #2
Write the balanced equation for the reaction between aluminum sulfate and calcium chloride, which produces aluminum
chloride and calcium sulfate.
Explanation
Write the chemical formulas on their correct sides:
Al2(SO4)3 + CaCl2
AlCl3 + CaSO4
In this reaction, the aluminum sulfate looks the most complicated, so start there. Look at what happens with sulfate—since
it remains sulfate on the right side of the reaction, treat it as a unit. You have three on the left side and only one on the right
side, so place a coefficient of 3 in front of calcium sulfate. Now deal with the aluminum. You have three on the left and
one on the right, so place a coefficient of 2 in front of aluminum chloride. Last, you must place a coefficient of 3 in front
of calcium chloride.
Al2(SO4)3 + 3CaCl2
2AlCl3 + 3CaSO4
Count the atoms on both sides of the reaction and you’ll see that
14you’re done.
Balancing Chemical Equations
Balance the equations below:
1)
____ N2 + ____ H2  ____ NH3
2)
____ KClO3  ____ KCl + ____ O2
3)
____ NaCl + ____ F2  ____ NaF + ____ Cl2
4)
____ H2 + ____ O2  ____ H2O
5)
____ Pb(OH)2 + ____ HCl  ____ H2O + ____ PbCl2
6)
____ AlBr3 + ____ K2SO4  ____ KBr + ____ Al2(SO4)3
7)
____ CH4 + ____ O2  ____ CO2 + ____ H2O
8)
____ C3H8 + ____ O2  ____ CO2 + ____ H2O
9)
____ C8H18 + ____ O2  ____ CO2 + ____ H2O
10)
____ FeCl3 + ____ NaOH  ____ Fe(OH)3 + ____NaCl
11)
____ P + ____O2  ____P2O5
12)
____ Na + ____ H2O  ____ NaOH + ____H2
13)
____ Ag2O  ____ Ag + ____O2
14)
____ S8 + ____O2  ____ SO3
15)
____ CO2 + ____ H2O  ____ C6H12O6 + ____O2
16)
____ K + ____ MgBr  ____ KBr + ____ Mg
17)
____ HCl + ____ CaCO3  ____ CaCl2 + ____H2O + ____ CO2
18)
____ HNO3 + ____ NaHCO3  ____ NaNO3 + ____ H2O + ____ CO2
19)
____ H2O + ____ O2  ____ H2O2
20)
____ NaBr + ____ CaF2  ____ NaF + ____ CaBr2
21)
____ H2SO4 + ____ NaNO2  ____ HNO2 + ____ Na2SO4
15
The Mole
Standard:
Practice Test Q:
Video:
Online Tutorial:
3b. Students know the quantity one mole is set by defining one mole of carbon 12
atoms to have a mass of exactly 12 grams. 3c. Students know one mole equals 6.02 x
1023 particles (atoms or molecules). (2 questions)
82, 83, 84
http://www.khanacademy.org/science/chemistry/v/the-mole-and-avogadro-s-number
http://askthenerd.com/phone/PhoneE1.html
You have already reviewed the process of
balancing equations and based the rules for
balancing equations on the principle that matter is
neither created nor destroyed in the course of a
chemical reaction. With this idea still in mind,
let’s begin our discussion of moles and formula
weights.
When you look at the periodic table, you see that
one of the pieces of data given for each element is
its atomic weight. But what exactly is the atomic
weight of a substance? It is the mass of one mole
of a substance. In turn,
1 mole substance = 6.02 x 1023 atoms or molecules
of the substance (depending on what it is), and
finally, the number 6.02 x 1023 is known as
Avogadro’s number. For example, carbon’s atomic weight is roughly 12 amu; this means that 6.02 x 1023
carbon atoms, in a pile, weigh 12 grams.
1 mole substance = atomic mass of the substance.
In order to find the formula weight of a substance, you simply add up the atomic masses of all of the atoms
in the molecular formula of a compound. But don’t forget to multiply the atomic mass of each element by the
subscript behind that element. Formula weights have the units amu, or atomic mass units; for example, the
formula weight of water, H2O, is about 18 amu. (O = 16 amu + 2 times H = 1 amu = 18 amu.) Similarly, the
molar mass of a molecule is the mass (in grams) of 1 mol of a substance; so the molar mass of H2O is also
roughly 18.
Example #1
.
Now
try calculating some molar masses and formula weights on your own by filling in the following
chart.
Substance
Molar Mass
Number of Moles Mass in Grams
Number of particles
Carbon dioxide, CO2
3.0
Oxygen, O2
64.0
Methane, CH4
0.279
Explanation
Molar Mass of CO2 = 1 C (12 amu) + 2 O (2 x 16 = 32 amu) = 44 amu = 44 g/mol CO2. Use this below.
3.0 moles CO2 (44.01 g CO2) = 132 grams CO2
(1 mole CO2)
Here’s the table, filled in.
Substance
Molar Mass
Carbon dioxide, CO2 44.01
Oxygen, O2
32.00
Methane, CH4
16.05
3.0 moles CO2 (6.02 x 1023 molecules) = 1.81 x 1024 particles
(1 mole CO2)
Number of Moles
3.0
2.00
0.279
16
Mass in Grams
132
64.0
4.48
Number of particles
1.81 x 1024
1.20 1024
1.68 1023
Mole Problems
Question 1
How many moles of copper are in 6,000,000 atoms of copper?
Question 2
How many atoms are in 5 moles of silver?
Question 3
How many atoms of gold are in 1 gram of gold?
Question 4
How many moles of sulfur are in 53.7 grams of sulfur?
Question 5
How many grams is a sample containing 2.71 x 1024 atoms of iron?
Question 6
How many moles of lithium (Li) are in 1 mole of lithium hydride (LiH)?
Question 7
How many moles of oxygen (O) are in 1 mole of calcium carbonate (CaCO3)?
Question 8
How many atoms of hydrogen are in 1 mole of water (H20)?
Question 9
How many atoms of oxygen are in 2 moles of O2?
Question 10
How many moles of oxygen are in 2.71 x 1025 molecules of carbon dioxide (CO2)
Question 11
Predict the mass of a mole of magnesium atoms.
Question 12
Calculate the molecular weights of carbon dioxide (CO2) and sugar (C12H22O11) and the mass of a mole of
each compound.
Question 13
Describe the difference between the mass of a mole of oxygen atoms and a mole of O2 molecules.
Question 14
Calculate the mass in grams of a single 12-C atom.
17
Molar Conversions
Standard:
3d. Students know how to determine the molar mass of a molecule from its
chemical formula and a table of atomic masses and how to convert the mass of
a molecular substance to moles, number of particles, or volume of gas at standard
temperature and pressure. (3 questions)
Practice Test Q:
85, 86, 87
Video:
http://www.youtube.com/watch?v=Kg-zaG0ckVg
Tutorial:
http://www.wiley.com/college/chem/spencer053872/tutorial/gramsmoles/gramsmoles1.html
The following chart shows the conversions between mass, particles and volume (gfm
means gram formula mass or molar mass).
.
Example #1
The density of CCl4 (l) is 1.59 g/mL. How many molecules of CCl4 are there in 2.59 L of CCl4?
Explanation
2.59 L CCl4 x (1.59 g CCl4) x (1000 mL) x (1 mol CCl4) x (6.02 x 1023 molecules CCl4) = 1.55 x 1025
(1 mL)
(1 L)
(154.0 g CCl4)
(1.0 mol CCl4)
Example 1 Given: Density of CCl4 = 1.59 grams
2.59 L CCl4
mole
Solution: grams x density x avogadro’s number
molar mass
Conclusion:
2.59 L CCl4 x (1.59 g CCl4) x (1000 mL) x (1 mol CCl4) x (6.02 x 1023 molecules CCl4) =
(1 mL)
(1 L)
(154.0 g CCl4)
(1.0 mol CCl4)
Answer: 1.55 x 1025 molecules
Example #2
What is the mass of 250.0 mL of C3H8?
Explanation
Mass = 0.2500 L C3H8 x (1 mol C3H8) x (42.0 g C3H8) = 0.469 g C3H8
(22.4 L)
1 mol C3H8
Example #3
How many moles are there in 6.00 L of NO3F (g) at STP?
Explanation
# of moles= 6.00 L NO3F x (1 mol NO3F) = 0.268 mol NO3F
(22.418
L)
Moles, Molecules, and Grams Worksheet
1)
How many molecules are there in 24 grams of FeF3?
2)
How many molecules are there in 450 grams of Na2SO4?
3)
How many grams are there in 2.3 x 1024 atoms of silver?
4)
How many grams are there in 7.4 x 1023 molecules of AgNO3?
5)
How many grams are there in 7.5 x 1023 molecules of H2SO4?
6)
How many molecules are there in 122 grams of Cu(NO3)2?
7)
How many grams are there in 9.4 x 1025 molecules of H2?
8)
How many molecules are there in 230 grams of CoCl2?
9)
How many molecules are there in 2.3 grams of NH4SO2?
10)
How many grams are there in 3.3 x 1023 molecules of N2I6?
11)
How many molecules are there in 200 grams of CCl4?
12)
How many grams are there in 1 x 1024 molecules of BCl3?
13)
How many grams are there in 4.5 x 1022 molecules of Ba(NO2)2?
14)
How many moles are in 15 grams of lithium?
15)
How many grams are in 2.4 moles of sulfur?
16)
How many moles are in 22 grams of argon?
17)
How many grams are in 88.1 moles of magnesium?
18)
How many moles are in 2.3 grams of phosphorus?
19)
How many grams are in 11.9 moles of chromium?
20)
How many moles are in 9.8 grams of calcium?
21)
How many grams are in 238 moles of arsenic?
19
Equation Stoichiometry
Standard:
3e. Students know how to calculate the masses of reactants and
products in a chemical reaction from the mass of one of the
reactants or products and the relevant atomic masses. (3 questions)
Practice Test Q: 88, 89, 90
Video:
http://www.khanacademy.org/science/chemistry/v/stoichiometry
Tutorial:
http://bertrand.home.mindspring.com/chem/st1frame.htm
One thing to realize when doing any calculations is that moles and coefficients are interchangeable. Both
mean number of particles, or multiples thereof. So in 2A + 3B ---> 4C, it either means 2 particles of A
combine with 3 particles of B to make 4 particles of C, or 2 moles reacts with 3 moles to make 4 moles.
Example #1
The formation of water from hydrogen and oxygen gas is: 2H2 (g) + O2 (g) ---> 2H2O (l). What mass of water will
form from 12.0 grams of hydrogen and excess oxygen (assuming the reaction goes to completion?)
Explanation
First what you must do in any of these problems is get all given masses into moles. You are given 12.0 grams of
hydrogen, let's see how many moles that is:
12.0 g H2 x (1 mol H2) = 5.94 mol H2
(2.02 g H2)
So how much water is formed? According to the equation, for every 2 moles of hydrogen, 2 moles of water are
produced, or in other words, a 1 to 1 ratio. So 5.94 moles of water will be formed. The question asks for what mass, so
we're not quite done yet.
5.94 mol H2O x (18.0 g H2O) = 107. grams of H2O.
(1 mol H2O)
Example #2
The newly discovered element Takalahium (symbol Tak; molecular mass = 411 g/mol) combines with oxygen to form
Takalahium Oxide. The unbalanced equation is:
Tak + O2 ---> Tak2O3
How many grams of Tak Oxide are formed when burning 8.00 kilograms of Tak?
Explanation
First and foremost, the balanced equation
is needed. That would be:
.
4Tak + 3O2 ---> 2Tak2O3
Then convert all given masses to moles:
8.00 kg Tak = 8000 g x (1 mol Tak) = 19.5 moles Tak.
(411 g Tak)
Since there are 2 Tak's for every 1 Tak Oxide, there must be half as many moles of Tak Oxide, or 9.50 moles. You can
also use Dimensional Analysis to do the same mole ratios.
Before you can get grams, you must first find the molar mass of Tak Oxide, which is no problem:
Mass = 2 x 411 g + 3 x 16.0 g = 870 g/mol.
Then you find the mass:
9.50 moles Tak Oxide x (870 g Tak Oxide) = 8260 grams Tak Oxide = 8.26 kilograms Tak Oxide
(1 mol Tak Oxide)
20
Stoichiometry Practice Problems
Solve the following stoichiometry grams-grams problems:
1)
Using the following equation:
2 NaOH + H2SO4  2 H2O + Na2SO4
How many grams of sodium sulfate will be formed if you start with 200 grams of sodium hydroxide
and you have an excess of sulfuric acid?
2)
Using the following equation:
Pb(SO4)2 + 4 LiNO3  Pb(NO3)4 + 2 Li2SO4
How many grams of lithium nitrate will be needed to make 250 grams of lithium sulfate, assuming
that you have an adequate amount of lead (IV) sulfate to do the reaction?
3)
Write the balanced equation for the reaction of acetic acid with aluminum hydroxide to form water
and aluminum acetate:
4)
Using the equation from problem #1, determine the mass of aluminum acetate that can be made if I
do this reaction with 125 grams of acetic acid and 275 grams of aluminum hydroxide.
5)
What is the limiting reagent in problem #2?
6)
How much of the excess reagent will be left over after the reaction is complete?
21
Pressure and Diffusion
Standard: 4a. Students know the random motion of molecules and their collisions with a surface create the
observable pressure on that surface. 4b. Students know the random motion of molecules explains the
diffusion of gases (2 questions).
Practice Test Q:
37, 38, 39
Video: http://www.chemthink.com/samples/sampleprob.htm
Online Tutorial: http://highered.mcgrawhill.com/sites/0072495855/student_view0/chapter2/animation__how_diffusion_works.html
All of the gas laws rely on some basic assumptions that are made about gases, and together they constitute
what it means for a gas to be in an ideal state. In an ideal state:
1. All gas particles are in constant, random motion.
2. All collisions between gas particles are perfectly elastic (meaning that the kinetic energy of the
system is conserved).
3. The volume of the gas molecules in a gas is negligible.
4. Gases have no intermolecular attractive or repulsive forces.
5. The average kinetic energy of the gas is directly proportional to its Kelvin temperature and is the
same for all gases at a specified temperature.
Only four measurable properties are used to describe a gas: its quantity, temperature, volume, and pressure.
The quantity (amount) of the gas is usually expressed in moles (n). The temperature, T, of gases must always
be converted to the Kelvin temperature scale (the absolute temperature scale). The volume, V, of a gas is
usually given in liters. Finally, the pressure, P, of a gas is usually expressed in atmospheres. Gases are often
discussed in terms of standard temperature and pressure (STP), which means 273K (or 0ºC) and 1 atm.
Example
Which of the following statements is not true of ideal gases?
1. The volume occupied by gas particles is only significant at very low pressures.
2. Gas molecules occupy an insignificant volume compared to the volume of the container that
holds them.
3. The particles of a gas move in random straight line paths until a collision occurs.
4. The collisions that occur between gas particles are considered elastic.
5. At a given temperature, all gas molecules within a sample possess the same average kinetic
energy.
Explanation
In this example, choice 1 is incorrect. Choices 2, 3, 4, and 5 all describe an ideal gas. Choice 1 makes
an incorrect assumption: it begins with a true statement about volume not being very significant but
then turns around and gives the incorrect scenario—if the pressure is low, then gas particles undergo
very few collisions, so the volume is insignificant. The volume only becomes significant if gas
particles collide often, increasing the chances that intermolecular forces will hold them together.
Pressure is the application of force to a surface, and the concentration of that force in a given area. A finger
can be pressed against a wall without making any lasting impression; however, the same finger pushing a
thumbtack can easily damage the wall, even though the force applied is the same, because
the point concentrates that force into a smaller area.
More formally, pressure (symbol: p or P) is the measure of the force that acts on a unit area,
Diffusion Movement of a fluid from an area of higher concentration to an area of lower concentration.
Diffusion is a result of the kinetic properties of particles of matter. The particles will mix until they are
evenly distributed.
Example: H2S(g) in a test tube will slowly diffuse into the air of a lab until equilibrium is reached.
22
Pressure and Diffusion Practice Problems
1) Simple diffusion is defined as
a) molecules from areas of higher concentration to areas of lower concentration.
b) molecules from areas of lower concentration to areas of higher concentration.
c) water molecules across a membrane.
d) gas molecules across a membrane.
e) water or gas molecules across a membrane.
2) When sugar is mixed with water, equilibrium is reached when
a) molecules of sugar stop moving.
b) water and sugar molecules are moving at the same speed.
c) the dissolved sugar molecules are evenly distributed throughout the solution.
d) there are the same number of water molecules as dissolved sugar molecules.
e) two tablespoons of coffee are added.
3) The rate of diffusion is affected by which of the following?
a) temperature
b) size of molecules
c) steepness of the concentration gradient
d) A and B
e) A, B, and C
4) The molecules in a solid lump of sugar do not move.
a) True
b) False
5) Which of the following is the best explanation of why a decrease in volume causes an increase in
pressure?
a) At a smaller volume the atoms will move faster and hit the sides more often.
b) At a smaller volume the atoms will slow down and so they will have more contact with the walls
of the container.
c) At a smaller volume the atoms will have less room to move around, so they will collide with the
sides more often.
d) The initial statement is false. Gas pressures do not increase when the volume is decreased.
6) What are the five assumptions we make about an ideal gas?
1)
2)
3)
4)
5)
23
The Gas Laws and Celsius Scale
Standard: 4c. Students know how to apply the gas laws to relations between the
pressure, temperature, and volume of any amount of an ideal gas or any mixture of
ideal gases. 4d. Students know the values and meanings of standard temperature
and pressure (STP). 4e. Students know how to convert between the Celsius and
Kelvin temperature scales. 4f. Students know there is no temperature lower than 0
Kelvin. (4 questions)
Practice Test Q:
40, 41, 42, 43, 44, 45, 46, 47
Video:
http://www.khanacademy.org/science/chemistry/v/ideal-gas-equation--pv-nrt
Online Tutorial:
http://www.ausetute.com.au/tempconv.html
Boyle’s law simply states that the volume of a confined gas at a fixed temperature is inversely proportional
to the pressure exerted on the gas. This can also be expressed as PV = a constant. This makes sense if you
think of a balloon. When the pressure around a balloon increases, the volume of the balloon decreases, and
likewise, when you decrease the pressure around a balloon, its volume will increase.
P1V1 = P2V2
Example #1
Sulfur dioxide (SO2) gas is a component of car exhaust and power plant discharge, and it plays a major role in the
formation of acid rain. Consider a 3.0 L sample of gaseous SO2 at a pressure of 1.0 atm. If the pressure is changed to
1.5 atm at a constant temperature, what will be the new volume of the gas?
Explanation
If P1V1 = P2V2, then (1.0 atm) (3.0 L) = (1.5 atm) (V2), so V2 = 2.0 L. This answer makes sense according to Boyle’s
law—as the pressure of the system increases, the volume should decrease.
Charles’s law states that if a given quantity of gas is held at a constant pressure, its volume is
directly proportional to the absolute temperature. Think of it this way. As the temperature of the
gas increases, the gas molecules will begin to move around more quickly and hit the walls of their
container with more force—thus the volume will increase. Keep in mind that you must use only the Kelvin
temperature scale when working with temperature in all gas law formulas!
Example #2
A sample of gas at 15ºC and 1 atm has a volume of 2.50 L. What volume will the gas occupy at 30ºC and 1 atm?
Explanation
The pressure remains the same, while the volume and temperature change—this is the hallmark of a Charles’s law
question.
So,
, then 2.50 L/288K = V2/303K, and V2 = 2.63 L
This makes sense—the temperature is increasing slightly, so the volume should increase slightly. Be careful of questions
like this—it’s tempting to just use the Celsius temperature, but you must first convert to Kelvin temperature (by adding
273) to get the correct relationships!
The ideal gas law is the most important gas law for you to know: it combines all of the laws you learned about
in this chapter thus far, under a set of standard conditions. The four conditions used to describe a gas—pressure,
volume, temperature, and number of moles (quantity)—are all related, along with R, the universal gas law
constant, in the following formula:
PV = nRT
where P = pressure (atm), V = volume (L), n = moles (mol), R = 0.08206 L · atm/mol · K, and T = temp (K).
Example #3
A 16.0 g sample of methane gas, CH4, the gas used in chemistry lab, has a volume of 5.0 L at 27ºC. Calculate the pressure.
Explanation
Looking at all the information given, you have a mass, a volume, and a temperature, and you need to find the pressure of
the system. As always, start by checking your units. You must first convert 16.0 g of CH4 into moles: 16.0 g CH4
1 mol
CH4/16.0 g CH4 = 1 mol of methane. The volume is in the correct units, but you must convert the temperature into
Kelvins: 27 + 273 = 300K. Now you’re ready to plug these numbers
24 into the ideal gas law equation:
PV = nRT
(P) (5.0 L) = (1.0 mol) (0.0821 L
atm/mol
K) (300K), so P = 4.9 atm
Gas Laws Practice Problems
Use Boyles’ Law to answer the following questions:
1)
1.00 L of a gas at standard temperature and pressure is compressed to 473 mL. What is the new
pressure of the gas?
2)
In a thermonuclear device, the pressure of 0.050 liters of gas within the bomb casing reaches 4.0 x
106 atm. When the bomb casing is destroyed by the explosion, the gas is released into the
atmosphere where it reaches a pressure of 1.00 atm. What is the volume of the gas after the
explosion?
Use Charles’ Law to answer the following questions:
3)
The temperature inside my refrigerator is about 40 Celsius. If I place a balloon in my fridge that
initially has a temperature of 220 C and a volume of 0.5 liters, what will be the volume of the balloon
when it is fully cooled by my refrigerator?
4)
A man heats a balloon in the oven. If the balloon initially has a volume of 0.4 liters and a
temperature of 20 0C, what will the volume of the balloon be after he heats it to a temperature of 250
0
C?
Use the ideal gas law to solve the following problems:
5)
If I have 4 moles of a gas at a pressure of 5.6 atm and a volume of 12 liters, what is the temperature?
6)
If I have an unknown quantity of gas at a pressure of 1.2 atm, a volume of 31 liters, and a temperature
of 87 0C, how many moles of gas do I have?
Convert the following temperatures into the unit required.
7) 100 K into C
8) 323 K into C
9) 100 C into K
10) 25 C into K
25
The pH Scale
Standard: 5a. Students know the observable properties of acids, bases, and salt solutions.
5d. Students know how to use the pH scale to characterize acid and base solutions (3 questions).
Practice Test Q:
60, 61, 62, 65
Video:
http://www.khanacademy.org/science/chemistry/v/acid-base-introduction.
Online Tutorial:
http://www.elmhurst.edu/~chm/vchembook/184ph.html
Properties of Acids and Bases
ACIDS
Taste sour
Reach with certain metals (Zn, Fe, etc.) to produce
hydrogen gas
cause certain organic dyes to change color
react with limestone (CaCO3) to produce carbon
dioxide
React with bases to form salts and water
BASES
Taste bitter
feel slippery or soapy
react with oils and grease
cause certain organic dyes to change color
react with acids to form salts and water
Define:Acid - a substance that produces protons, H+
Base - a substance that produces hydroxide
ions, OHReaction of acids and bases with water:
Acids and bases form ions in solution:
HCl(aq)  H+(aq) + Cl-(aq)
H3O+ - hydronium ion H+ and H3O+ are equivalent in aq. solution
When we look at the reactions of acids - can be generalized using hydrogen ion
Reaction with zinc yields hydrogen gas
Reaction with limestone - produce CO2(g)
Acids react with bases to produce a salt
Similarly for bases, produce hydroxide ions
Introduction and Definitions:
Acidic and basic are two extremes that describe a chemical property about chemicals. Mixing acids and
bases can cancel out or neutralize their extreme effects. A substance that is neither acidic nor basic is neutral.
The pH scale measures how acidic or basic a substance is. The pH scale ranges from 0 to 14. A pH of 7 is
neutral. A pH less than 7 is acidic. A pH greater than 7 is basic.
The pH scale is logarithmic and as a result, each whole pH value below 7 is ten times more acidic than the
next higher value. For example, pH 4 is ten times more acidic than pH 5 and 100 times (10 times 10) more
acidic than pH 6. The same holds true for pH values above 7, each of which is ten times more alkaline
(another way to say basic) than the next lower whole value. For example, pH 10 is ten times more alkaline
than pH 9 and 100 times (10 times 10) more alkaline than pH 8.
Pure water is neutral. But when chemicals are mixed with water, the mixture can become either acidic or
basic. Examples of acidic substances are vinegar and lemon juice. Lye, milk of magnesia, and ammonia are
examples of basic substances.
26
pH Scale Practice Problems
1) Five solutions A, B, C, D, E when tested with universal indicator showed a pH of 4, 1, 11, 7, and 9
respectively.
a) Which solution is (i) neutral, (ii) strongly alkaline, (iii) strongly acidic, (iv) weakly acidic, and (v) weakly
alkaline?
b) Arrange the pH in increasing order of hydrogen ion concentration.
2) Define the term "pH"; what does" pH" stand for?
3) What is 'pH' scale? Explain briefly.
4) What is the 'pH' of pure water and that of rain water? Explain the difference.
5) What is the pH of solution 'A' which liberates CO2 gas with a carbonate salt? Give the reason?
6) What is the pH of solution 'B' which liberates NH3 gas with an ammonium salt? Give reason?
7) How do you increase or decrease the pH of pure water?
8) What are indicators?
27
Types of Acids and Bases
Standard: 5b. Students know acids are hydrogen-ion-donating and bases are
hydrogen-ion-accepting substances. 5c. Students know strong acids and bases
fully dissociate and weak acids and bases partially dissociate. (2 questions)
Practice Test Q:
63, 64
Video:
http://www.khanacademy.org/science/chemistry/v/ph-of-a-weak-acid
http://www.khanacademy.org/science/chemistry/v/ph--poh-of-strong-acids-and-bases
Online Tutorial:
http://www.msdiehl.com/resources/notes2.pdf
1. Strong Acids: completely dissociate in water, forming H+ and an anion.
example: HN03 dissociates completely in water to form H+ and N031-. The reaction is
HNO3(aq) → H+(aq) + N031-(aq)
A 0.01 M solution of nitric acid contains 0.01 M of H+ and 0.01 M N03- ions and almost no HN03
molecules. The pH of the solution would be 2.0.
 There are only 6 strong acids: You must learn them. The remainding acids are weak acids.
HCl
H2SO4
HNO3
HClO4
HBr
HI
Note: when a strong acid dissociates only one H+ ion is removed. H2S04 dissociates giving H+ and
HS04- ions.
H2SO4 → H+ + HSO41A 0.01 M solution of sulfuric acid would contain 0.01 M H+ and 0.01 M HSO41- (hydrogen sulfate ion).

2. Weak acids: a weak acid only partially dissociates in water to give H+ and the anion
for example, HF dissociates in water to give H+ and F-. It is a weak acid. with a dissociation equation that is
HF(aq) ↔ H+(aq) + F-(aq)

Which are the weak acids? Anything that dissociates in water to produce H+ and is not one of
the 6 strong acids.
1.
Molecules containing an ionizable proton. (If the formula starts with H then it is a prime candidate
for being an acid.) Also: organic acids have at least one carboxyl group, -COOH, with the H being ionizable.
2.
Anions that contain an ionizable proton. ( HSO41- → H+ + SO42- )
3.
Cations: (transition metal cations and heavy metal cations with high charge)
also NH4+ dissociates into NH3 + H+
3. Strong Bases: They dissociate 100% into the cation and OH- (hydroxide ion).
example: NaOH(aq) → Na+(aq) + OH-(aq)
a. 0.010 M NaOH solution has 0.010 M OH- ions (as well as 0.010 M Na+ ions) and have a pH of 12.

Which are the strong bases?
The hydroxides of Groups I and II.
 Note: the hydroxides of Group II metals produce 2 mol of OH- ions for every mole of base that
dissociates. These hydroxides are not very soluble, but what amount that does completely dissociates.
example: Ba(OH)2(aq) → Ba2+(aq) + 2OH-(aq)
a. 0.000100 M Ba(OH)2 solution will be 0.000200 M in OH- ions and will have a pH of 10.3.
4. Weak Bases: What compounds are considered to be weak bases?
1. Most weak bases are anions of weak acids.
2. Weak bases do not furnish OH- ions by dissociation. They react with water to form OH- ions.
3. When a weak base reacts with water the OH- comes from the water and the remaining H+ attaches
itself to the weak base, giving a weak acid as one of the products.
General reaction: weak base(aq) + H2O(aq) → weak acid(aq) + OH-(aq)
28
Types of Acids and Bases Practice Problems
1) What is the pH of the solution with a hydronium concentration [H3O+] 1.47 x 10-4?
What is the pOH of this solution?
2) What is the pOH of the solution with a hydroxyl concentration [OH-] 2.98 x 10-2?
What is the hydronium concentration [H3O+] of this solution?
3) What is the hydronium concentration [H3O+] of a solution with a pH of 7.84?
What is the hydroxyl concentration [OH-] of this solution?
4) What is the hydroxyl concentration [OH-] of a solution with a pH of 3.76?
5) What is the hydronium concentration [H3O+] of a solution with a pOH of 2.47?
Identify the following compounds as a strong acid, weak acid, strong base, or a weak base.
NaOH
HCl
H2CO3
KOH
H2SO4
LiOH
NH3
29
Solute, Solvent and the Dissolving Process
Standard: 6a. Students know the definitions of solute and solvent. 6b. Students know how to describe the
dissolving process at the molecular level by using the concept of random molecular motion. (2 questions)
Practice Test Q:
48, 49, 50, 51
Video:
http://www.khanacademy.org/science/chemistry/v/states-of-matter
Online Tutorial:
http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch3/solution.html
A solution is a homogeneous mixture of a solute and a solvent. Solutions can be formed in any state of
matter; that is they may be solid, liquid, or gas. A solution is prepared by dissolving a solute into the solvent.
Solute is either the smaller component of a mixture or, when liquid solutions are considered, the gaseous or
solid substance added to the solution. Solutions could be composed of either complete molecules molecular solution, or ions - ionic solution. The latter usually is referred to aqueous solutions of salts.
Fluids that mix or dissolve in each other in all proportions are call miscible fluids, lacking that property
fluids are called immiscible. So gases are always miscible.
The number of grams of solute that can just be dissolved in 100 ml of solvent at 20°C is defined as the
solubility (36g for NaCl). At the maximum solubility the solution is saturated and in dynamic equilibrium
with the insoluble part of solute. Such a solution is
called saturated. Solution with less concentration is call
unsaturated.
NaCl(s) <==> Na+(aq) + Cl (aq)
If there is more solute dissolved than saturation allows,
the solution is said to be supersaturated.
2.1a The changes that occur in the dissolving process
The most electronegative part of the highly polar water
molecule (>Oδ-) will be attracted to the positive ion and,
since the oxygen atom has two lone pairs of electrons, it
is also the source of the dative covalent bond by donation of one of these pairs of electrons into a vacant
metal ion orbital.
In the case of anions, the positive ends of the water molecules (-Hδ+) will orientate themselves towards the
negative anion and the water molecules become weakly associated with anion, but no covalent bonds are
formed.
2.1b Diagram illustrating the dissolving-solvation-hydration process for sodium chloride crystals
forming salt solution
 For diagram simplicity, each ion is surrounded by four water molecules, though in reality, much more
than this - see later.
 The strong crystal lattice (giant ionic structure) is broken down by the solvation process so the salt
dissolves and the ions are free to move around in the water solvent.
30
Solute, Solvent/ Dissolving Process Practice Problems
Identify the solute and solvent in the following solutions. .
nail polish in acetone
acetone dissolving glue
eggshells in vinegar
iodine in hexane
chromium in hydrochloric acid
Kool Aid in water
Describe the difference between unsaturated, saturated, and supersaturated solutions.
Draw a picture showing a salt crystal.
Now draw how that salt crustal dissolves in a water solution.
31
Dissolving and Solution Concentration
Standard: 6c. Students know temperature, pressure, and surface area affect the dissolving process. 6d.
Students know how to calculate the concentration of a solute in terms of grams per liter, molarity, parts
per million, and percent composition. (3 questions)
Practice Test Q:
52, 53, 54, 55
Video:
http://www.khanacademy.org/science/chemistry/v/solubility
Online Tutorial:
http://www.nauticus.org/chemistry/chemconcunits.html
Factors affecting solubility include intermolecular forces, viscosity, and entropy.
Solubility varies dramatically. General 'rule' - "like dissolves like". The solubility of a
solute in a solvent (that is, the extent of the mixing of the solute and solvent species)
depends on a balance between the natural tendency for the solute and solvent species to
mix and the tendency for a system to have the lowest energy possible.
.The solubility of a solute in both molecular solutions and ionic solutions is dependent on temperature
and pressure.
Temperature. Most gases are less soluble at high T but solid compounds usually (but not always) are more
soluble with T. Heat of solution could be either positive or negative. Depending on that - different
applications. Dissolving of ammonium nitrate in water is the basis for instant cold packs used in hospitals
and elsewhere (NH4NO3 crystals inside a bag of water). When the inner bag is broken, NH4NO3 dissolves in
the water. Heat is absorbed, so the bag feels cold. Similarly, hot packs are available, containing either CaCl2,
MgSO4 or sodium acetate, which dissolve in water with the evolution of heat.
Pressure. In general, pressure change has little effect on the solubility of a liquid or solid in water, except for
gases.
Concentration
The concentration of a solute is the amount of solute dissolved in a given quantity of solvent or solution.
The quantity of solvent or solution can be expressed in terms of volume or in terms of mass or molar amount.
The common expressions for concentration are: Molarity (M), Mass Percent, Mole Fraction (XA) and
Molality (m)
Molarity (M)
Molarity =
moles of solute
liters of solution
Mole Fraction (XA)
moles of substance A
XA =
total moles of solution
Mass Percent Composition
Mass percentage mass of solute
x100 %
of solute =
mass of solution
Parts Per Million (ppm)
mg of solute
ppm =
L of solution
Perhaps the most important property of a solution is its concentration. A dilute acetic acid solution, also
called vinegar, is used in cooking while a concentrated solution of acetic acid would kill you if ingested. The
only difference between such solutions is the concentration of the solute. In order to quantify the
concentrations of solutions, chemists have devised many different units of concentration each of which is
useful for different purposes.
Molarity, the number of moles of solute per liter of solution, has the units moles / L which are abbreviated
M. This unit is the most commonly used measure of concentration. It is useful when you would like to know
the number of moles of solute when you know both the molarity and the volume of a solution.
32
Dissolving and Solution Concentration Practice Problems
1)
How many grams of beryllium chloride are needed to make 125 mL of a 0.050 M solution?
2)
How many grams of beryllium chloride would you need to add to 125 mL of water to make a 0.050
molal solution?
3)
The density of ethanol is 0.789 g/mL. How many grams of ethanol should be mixed with 225 mL of
water to make a 4.5% (v/v) mixture?
4)
Explain how to make at least one liter of a 1.25 molal ammonium hydroxide solution.
5)
What is the molarity of a solution in which 0.45 grams of sodium nitrate are dissolved in 265 mL of
solution.
6)
What is the mole fraction of sulfuric acid in a solution made by adding 3.4 grams of sulfuric acid to
3,500 mL of water?
7)
What will the volume of a 0.50 M solution be if it contains 25 grams of calcium hydroxide?
8)
How many grams of ammonia are present in 5.0 L of a 0.050 M solution?
33
Endothermic and Exothermic Reactions
Standard: 7a. Students know how to describe temperature and heat flow in terms of the motion of
molecules (or atoms). 7b. Students know chemical processes can either release (exothermic) or
absorb (endothermic) thermal energy. 7c. Students know energy is released when a material
condenses or freezes and is absorbed when a material evaporates or melts. (3 questions)
Practice Test Q:
56, 57, 58
Video:
http://www.youtube.com/watch?v=5RJLvQXce4A
Online Tutorial:
http://www.mikeshultzfiction.com/School/endoquiz.html
Heat Changes in Chemical Reactions
When chemical reactions occur, as well as the formation of the products - the chemical change, there is also
a heat energy change which can often be detected as a temperature change. This means the products have
different energy content than the original reactants (see the reaction profile diagrams below).
If the products contain less energy than the reactants, heat is released or given out to the surroundings and
the change is called exothermic. The temperature of the system will be observed to rise in an exothermic
change. (examples: combustion, reaction of metal with acid, acid neutralization)
If the products contain more energy than the reactants, heat is taken in or absorbed from the surroundings
and the change is called endothermic. If the change can take place spontaneously, the temperature of the
reacting system will fall but, as is more likely, the reactants must be heated to speed up the reaction and
provide the absorbed heat. (examples: thermal decomposition of limestone, cracking of oil fractions)
The difference between the energy levels of the reactants and products gives the overall energy change (delta
H, ΔH) for the reaction (the activation energies are NOT shown on the diagrams below).
ΔH is negative (-ve) for exothermic reactions i.e. heat energy is given out and lost from the system to the
surroundings which warm up.ΔH is positive (+ve) for endothermic reactions i.e. heat energy is gained by
the system and taken in from the surroundings which cool down OR, as is more likely, the system is heated
to provide the energy needed to effect the change.
1b. Heat changes in physical
changes of state
 Changes of physical
state i.e. gas <==>
liquid <==> solid
are also
accompanied by
energy changes.
 To melt a solid, or boil/evaporate a liquid, heat energy must be absorbed or taken in from the
surroundings, so these are endothermic energy changes (ΔH +ve). The system is heated to effect
these changes.
 To condense a gas, or freeze a solid, heat energy must be removed or given out to the surroundings,
so these are exothermic energy changes (ΔH -ve). The system is cooled to effect these changes. The
energy required to boil or evaporate a substance is usually much more than that required to melt the
solid.
o This is because in a liquid the particles are still quite close together with attractive forces
holding together the liquid, but in a gas the particles of the structure must be completely
separated with virtually no attraction between them.
 The stronger the forces between the individual molecules, the more energy are needed to melt or boil.
o As this is shown by the varying energy requirements to melt or boil a substance.
 For simple covalent molecules, the energy absorbed by the material is relatively small to melt or
vaporize the substance and the bigger the molecule the greater the inter-molecular forces.
34
Endothermic and Exothermic Practice Problems
1) What type of chemical reaction absorbs energy and requires energy for the reaction to occur?
a. endothermic
b. exothermic
c. synthesis
d. both A and B
2) What type of reaction releases energy and does not require initial energy to occur?
a. endothermic
b. exothermic
c. decomposition
d. both A and B
3) Which of the following are examples of an exothermic chemical reaction? Check all that apply.
a. photosynthesis
b. burning a piece of wood
c. freezing ice into water
d. none of the above
4) Which type of reactions cannot occur spontaneously?
a. endothermic
b. exothermic
c. neither
5) Any type of reaction that involves burning (combustion) can be classified as which of the following types
of reactions?
a. synthesis
b. endothermic
c. exothermic
d. all of the above
6) Burning sugar is an exothermic process.
a. True
b. False
7) What is enthalpy?
a. heat content
b. absolute amount of energy is a chemical system
c. the reactants of the chemical reaction
d. none of the above
8) A 100 g sample of water at 25 oC and 1 atm. of pressure _?_ than 100 g of water that has recently been
heated to 100 oC from 0 oC and then cooled to 25 oC at 1 atm. of pressure.
a. has more internal energy than
b. has less internal energy than
c. has the same internal energy as
9) Water has a specific heat of 4.184 J/g deg while glass (Pyrex) has a specific heat of 0.780 J/g deg. If 10.0 J
of heat is added to 1.00 g of each of these, which will experience the larger increase of temperature?
a. glass
b. water
c. They both will experience the same change in temperature since only the amount of a substance
relates to the increase in temperature.
35
Specific Heat Calculations
Standard: 7d. Students know how to solve problems involving heat flow and temperature changes, using
known values of specific heat and latent heat of phase change (2 questions).
Practice Test Q:
59
Video: http://www.khanacademy.org/science/chemistry/v/specific-heat--heat-of-fusion-and-vaporization
Online Tutorial:
http://quizlet.com/1554348/science-heat-transfer-specific-heat-notes-flash-cards/
Specific heat is another physical property of matter. All matter has a temperature associated with it. The
temperature of matter is a direct measure of the motion of the molecules: The greater the motion the higher
the temperature:
Motion requires energy: The more energy matter has
the higher temperature it will also have. Typically this
energy is supplied by heat. Heat loss or gain by matter
is equivalent energy loss or gain.
With the observation above understood we can now
ask the following question: by how much will the
temperature of an object increase or decrease by the
gain or loss of heat energy? The answer is given by the
specific heat (S) of the object. The specific heat of an
object is defined in the following way: Take an object
of mass m, put in x amount of heat and carefully note the temperature rise, then S is given by
In this definition mass is usually in either grams or kilograms and temperature is either in kelvin or degrees
Celsius. Note that the specific heat is "per unit mass". Thus, the specific heat of a gallon of milk is equal to
the specific heat of a quart of milk. A related quantity is called the heat capacity (C). of an object. The
relation between S and C is C = (mass of object) x (specific heat of object).
Consider the specific heat of copper , 0.385 J/g 0C. What this means is that it takes 0.385 Joules of heat to
raise 1 gram of copper 1 degree Celsius. Thus, if we take 1 gram of copper at 25 0C and add 1 Joule of heat
to it, we will find that the temperature of the copper will have risen to 26 0C. We can then ask: How much
heat wil it take to raise by 1 0C 2 g of copper?. Clearly the answer is 0.385 J for each gram or 2 x 0.385 J =
0.770 J. What about a pound of copper? A simple way of dealing with different masses of matter is to
dtermine the heat capacity C as defined above. Note that C depends upon the size of the object as opposed to
S that does not.
Example 1: How much energy does it take to
raise the temperature of 50 g of copper by 10 0C?
Example 2: If we add 30 J of heat to 10 g of
aluminum, by how much will its temperature
increase?
36
Specific Heat Practice Problems
1) What is the specific heat of a substance that absorbs 2.5 x 103 joules of heat when a sample
of 1.0 x 104 g of the substance increases in temperature from 10.0C to 70.0C?
2) A 1.0 kg sample of metal with a specific heat of 0.50 KJ/KgC is heated to 100.0C and then
placed in a 50.0 g sample of water at 20.0C. What is the final temperature of the metal and
the water?
3) A 2.8 kg sample of a metal with a specific heat of 0.43KJ/KgC is heated to 100.0C then
placed in a 50.0 g sample of water at 30.0C. What is the final temperature of the metal and
the water?
Specific Heat Capacities of Various Materials
SUBSTANCE SPECIFIC HEAT
SUBSTANCE
CAPACITY (J/KG ºC)
Aluminum
9.0 x 102
Alcohol (ethyl)
Brass
3.8 x 102
Alcohol (methyl)
2
Copper
3.9 x 10
Glycerine
2
Glass (crown)
6.7 x 10
Mercury
Glass (pyrex)
7.8 x 102
Nitrogen (liquid)
2
Gold
1.3 x 10
Water (liquid)
2
Iron
4.5 x 10
Water (ice)
Lead
1.3 x 102
Water (steam)
2
Sand
8.0 x 10
air
2
Silver
2.3 x 10
SPECIFIC HEAT
CAPACITY (J/KG ºC)
2.3 x 102
2.5 x 102
2.4 x 102
1.4 x 102
1.1 x 102
4.2 x 103
2.1 x 103
2.0 x 103
1.0 x 103
1. When 3.0 kg of water is cooled from 80.0C to 10.0C, how much heat energy is lost?
2. How much heat is needed to raise a 0.30 kg piece of aluminum from 30.C to 150C?
3. Calculate the temperature change when:
a) 10.0 kg of water loses 232 kJ of heat.
b) 1.96 kJ of heat are added to 500. g of copper.
37
Changing the Rate of a Reaction
Standard: 8a. Students know the rate of reaction is the decrease in concentration of reactants or the increase
in concentration of products with time. 8b. Students know how reaction rates depend on such factors as
concentration, temperature, and pressure. 8c. Students know the role a catalyst plays in increasing the
reaction rate. (4 questions)
Practice Test Q:
66, 67, 68, 69, 70, 71
Video:
http://www.khanacademy.org/science/chemistry/v/reactions-in-equilibrium
Online Tutorial:
http://www.nauticus.org/chemistry/chemkinetics.html
Rate of reaction is concerned with how quickly a reaction reaches a certain point. It can be defined as the
decrease in concentration of the reactants per unit time or the increase in concentration of the products per
unit time. A graph may be plotted of concentration against time, with time on the x-axis and some measure of
how far the reaction has gone (ie concentration, volume, mass loss etc) on the y-axis. This will produce a
curve and the rate at any given point is the gradient of the tangent to this curve.
Collision theory -- reactions take place as a result of particles (atoms or molecules) colliding and then
undergoing a reaction. Not all collisions cause reaction, however, even in a system where the reaction is
spontaneous. The particles must have sufficient kinetic energy, and the correct orientation with respect to
each other for them to react.
Activation energy
This is the minimum energy that
particles colliding must have in
order to produce successful reaction.
It is given the symbol Ea (Energy of
Activation). The energy of particles
is expressed by their speed.
Changing the conditions
Increasing the temperature of a
substance increases the average
speed (Energy) of the particles and consequently the number of particles colliding with sufficient energy (Ea)
to react. At higher temperatures there are more successful collisions and therefore a faster reaction. At higher
concentrations there are more collisions and consequently a faster reaction. Catalysts lower the activation
energy by providing an alternative mechanism for the reaction/ greater probability of proper orientation. This
results in a faster reaction. In heterogeneous reactions (where the reactants are in different states) the size of
the particles of a solid may change reaction rate, since the surface is where the reaction takes place, and the
surface area is increased when the particles are more finely divided (therefore smaller solid particles in a
heterogeneous reaction tend to produce a faster reaction).
Condition
Effect on rate
Explanation
Temperature
increasing the temperature increases the
rate of a reaction
Two reasons:
1. There are more particles with sufficient energy to react
(most important) - more successful collisions
2. There are more collisions
Concentration
Increasing the concentration of a reactant
increases the rate of the reaction (usually)
There are more collisions as there are more particles in
closer proximity
Particle size
The smaller the particles the faster the
reaction. (note: the solute particles in
solutions have the smallest particle size
possible. and so solutions react fastest)
Collisions occur at the surface of particles. The larger the
particle size the smaller the surface area and the fewer
collisions can occur.
Catalysts
The presence of a catalyst increases the
rate of a reaction
Catalysts provide an alternative mechanism with a lower
activation energy
38
Changing the Rate of a Reaction Practice Problems
Use the following diagram to answer questions 1-3.
1) Which letter corresponds to the activation energy of the reaction?
a. A
b. B
c. C
d. Y
E. X
2) Which letter corresponds to the change in energy for the overall reaction?
a. A
b. B
c. C
d. Y
E. X
True or False:
3) The reaction shown above is exothermic BECAUSE energy difference B is greater than difference A.
______________________
4) A system is at equilibrium when the rate of the forward reaction is equal to the rate of the reverse
reactions BECAUSE At equilibrium, the concentration of the products is equal to that of the reactants.
________________________
5) Which of the following statements best describes the condition(s) needed for a successful formation of a
product in a chemical reaction?
a. The collision must involve a sufficient amount of energy, provided from the motion of the
particles, to overcome the activation energy.
b. The relative orientation of the particles has little or no effect on the formation of the product.
c. The relative orientation of the particles has an effect only if the kinetic energy of the particles is
below some minimum value.
d. The relative orientation of the particles must allow for formation of the new bonds in the product.
e. The energy of the incoming particles must be above a certain minimum value and the relative
orientation of the particles must allow for formation of new bonds in the product.
6) The catalyzed pathway in a reaction mechanism has a _____ activation energy and thus causes a _____
reaction rate.
a. higher, lower
b. higher, higher
c. lower, higher
d. lower, steady
39
Le Chatelier’s Principle
Standard: 9a. Students know how to use LeChatelier’s principle to predict the effect of changes in
concentration, temperature, and pressure. 9b. Students know equilibrium is established when forward and
reverse reaction rates are equal. (4 questions)
Practice Test Q:
72, 73, 74, 75, 76, 77
Video:
http://www.khanacademy.org/science/chemistry/v/le-chatelier-s-principle
Online Tutorial:
http://www.sciencegeek.net/Chemistry/taters/LeChatelier.htm
Using Le Chatelier's Principle- If a dynamic equilibrium (when the forward and reverse reaction rates are equal) is
disturbed by changing the conditions, the position of equilibrium moves to counteract the change.
Using Le Chatelier's Principle with a change of concentration
Suppose you have an equilibrium established between four substances A, B, C and D.
What would happen if you changed the conditions by increasing the concentration of A?
According to Le Chatelier, the position of equilibrium will move in such a
way
as to counteract the change. That means that the position of equilibrium will
move so that the concentration of A decreases again - by reacting it with B and
turning it into C + D. The position of equilibrium moves to the right.
What would happen if you changed the conditions by decreasing the concentration of A?
According to Le Chatelier, the position of equilibrium will move so that the
concentration of A increases again. That means that more C and D will react
to replace the A that has been removed. The position of equilibrium moves to the
left. This is essentially what happens if you remove one of the products of the
reaction as soon as it is formed. If, for example, you removed C as soon as it
was formed, the position of equilibrium would move to the right to replace it. If you kept on removing it, the
equilibrium position would keep on moving rightwards - turning this into a one-way reaction.
This
Using Le Chatelier's Principle with a change of pressure
only applies to reactions involving gases:
What would happen if you changed the conditions by increasing the pressure?
According to Le Chatelier, the position of equilibrium will move in such a way as
to counteract the change. That means that the position of equilibrium will move so that the pressure is reduced again.
Pressure is caused by gas molecules hitting the sides of their container. The more molecules you have in the container,
the higher the pressure will be. The system can reduce the pressure by reacting in such a way as to produce fewer
molecules. In this case, there are 3 molecules on the left-hand side of the equation, but only 2 on the right. By forming
more C and D, the system causes the pressure to reduce. Increasing the pressure on a gas reaction shifts the position of
equilibrium towards the side with fewer molecules.What happens if there are the same number of molecules on both
sides of the equilibrium reaction? In this case, increasing the pressure has no effect whatsoever on the position of the
equilibrium. Because you have the same numbers of molecules on both sides, the equilibrium can't move in any way
that will reduce the pressure again.
Using Le Chatelier's Principle with a change of temperature
For this, you need to know whether heat is given out or absorbed during the reaction. Assume that our forward reaction
is exothermic (heat is evolved):
This shows that 250 kJ is evolved
(hence the negative sign) when 1 mole of A reacts completely with 2 moles of B. For reversible reactions, the value is
always given as if the reaction was one-way in the forward direction.
 Increasing the temperature of a system in dynamic equilibrium favors the endothermic reaction. The system
counteracts the change you have made by absorbing the extra heat.
 Decreasing the temperature of a system in dynamic equilibrium favors the exothermic reaction. The system
counteracts the change you have made by producing more heat.
40
Le Chatelier’s Principle Practice Problems
1) Consider the system below at equilibrium. Which of the following changes will shift the equilibrium to
the right?
N2 (g) + 3H2 (g) ↔ 2NH3 (g) + 92.94 kJ
I. Increasing the temperature
II. Decreasing the temperature
III. Increasing the pressure on the system
(A) I only (B) II only (C) III only (D) I and III (E) II and III
2) Suggest four ways to increase the concentration of SO3 in the following equilibrium reaction.
Be specific.
2 SO2(g) + O2(g)
2 SO3(g) + 192.3 kJ
1. _____________________________________________________________________________________
2. _____________________________________________________________________________________
3. _____________________________________________________________________________________
4. _____________________________________________________________________________________
3) Use Le Chatelier's Principle to predict how the changes listed will affect the following equilibrium
reaction:
2 HI (g) + 9.4 kJ
H2 (g) + I2 (g)
a) Will the concentration of HI increase, decrease, or remain the same if more H2 is added?
b) What is the effect on the concentration of HI if the pressure of the system is increased?
c) What is the effect on the concentration of HI if the temperature of the system is increased?
d) What is the effect on the concentration of HI if a catalyst is added to the system?
e) Write the equilibrium constant expression for this reaction.
41
Organic Chemistry
Standard: 10a. Students know large molecules (polymers), such as proteins, nucleic acids, and starch, are
formed by repetitive combinations of simple subunits. 10b. Students know the bonding
characteristics of carbon that result in the formation of a large variety of structures ranging
from simple hydrocarbons to complex polymers and biological molecules. 10c. Students
know amino acids are the building blocks of proteins. (2 questions)
Practice Test Q:
33, 34, 35, 36
Video:
www.youtube.com/watch?v=MGNRna3JnCk
Online Tutorial:
http://www.execulink.com/~ekimmel/onclick0.htm
Alcohols
Organic molecules with a hydroxyl group (-OH). Methanol [CH3OH] and ethanol (beverage alcohol)[CH3CH2OH]
are common examples. Sugars are also alcohols.
Carboxylic Acids
Contain one or more carboxyl groups [-COOH], also the
intermediates in food breakdown by cellular respiration.
Aldehydes
Contain a carbon atom to which is attached one hydrogen atom and — by a
double bond — one oxygen atom. Formaldehyde [HCHO] is a powerful
disinfectant and preservative (it denatures proteins). Acetaldehyde is
produced during the conversion of pyruvic acid to ethanol when yeast
ferment sugars. The converse is also true - acetaldehyde is produced in the
liver as it metabolizes ingested ethanol (and may be the prime culprit in a
"hangover").
Ethers
Formed when two carbon atoms are linked by an oxygen atom.
Diethyl ether is a commonly-used anesthetic.
Esters
The removal of a molecule of water between the -OH group of an alcohol
and the -OH group of a
 carboxylic acid (-COOH) [shown in the diagram] or
 phosphoric acid
 Fats are triesters of three fatty acids and glycerol (the alcohol).
Ketones
Organic molecules with a carbonyl group (-C=O) between two hydrocarbon
portions. Ketones are synthesized in the liver, usually from fatty acids. When
glucose metabolism is suppressed, during starvation or in diabetics, fatty
acids are used as a source of energy. But instead of entering the citric acid
cycle, the acetyl-CoA produced from them is converted into the ketone
acetoacetate. Some of this is then converted into acetone (which can be
smelled on the breath of patients whose diabetes is out of control).
Amines
Organic molecules with an amino group, -NH2. Some examples:
 all the amino acids (lysine has two of them).
 the thyroid hormones thyroxine (T4) and triiodothyronine (T3)
 Many neurotransmitters:
o adrenaline and noradrenaline, dopamine, serotonin (5hydroxytryptamine), histamine
Amides
Amides are organic molecules containing a carbonyl group (-C=O) attached to a nitrogen atom. The peptide
bond between the amino acids linked in a polypeptide is also called an amide bond.
42
Organic Chemistry Practice Problems
For the following six molecular structures identify all the functional groups present and determine the
overall properties of the molecule (e.g. acid, base, neutral).
1.
2.
3.
4.
5.
6.
43
Nuclear Chemistry
Standard: 11a. Students know protons and neutrons in the nucleus are held together by nuclear forces that
overcome the electromagnetic repulsion between the protons. 11b. Students know the energy release per
gram of material is much larger in nuclear fusion or fission reactions than in chemical reactions. The change
in mass (calculated by E=mc2) is small but significant in nuclear reactions. 11c. Students know some
naturally occurring isotopes of elements are radioactive, as are isotopes formed in nuclear reactions. 11d.
Students know the three most common forms of radioactive decay (alpha, beta, and gamma) and know how
the nucleus changes in each type of decay. 11e. Students know alpha, beta, and gamma radiation produce
different amounts and kinds of damage in matter and have different penetrations. (2 questions)
Practice Test Q:
20, 21, 22, 23
Video:
http://www.khanacademy.org/science/chemistry/v/types-of-decay
Online Tutorial:
http://www.avon-chemistry.com/nuclear_practice.htm
What Makes an Element Radioactive?
To understand radioactivity, we need to explore the structure of an atomic
nucleus. The protons in the nucleus, all being positively charged, repel each
other! So if all the protons repel each other, how does the nucleus stay glued
together and remain stable? It is because of the 'Nuclear Force'.
This force is stronger than the electromagnetic force, but the range of this
force is only limited to size of the nucleus, unlike electromagnetic force whose range is infinite. This nuclear
force acts between the protons and neutrons, irrespective of the charge and it's always strongly attractive.
However, it has limitations of range. So, in the nucleus, there is a constant tussle between the repelling
electromagnetic coulomb force of protons and the attractive strong nuclear force.
In a nucleus like Uranium, which has almost 92 protons, coulomb repulsive force becomes too much for the
nuclear force to contain. Subsequently, the nucleus is very unstable and radioactive decay occurs and
Uranium decays into a more stable element. Such an unstable nucleus like Uranium, when gently tapped by a
neutron, splits up into two other nuclei through nuclear fission, releasing tremendous amount of energy in the
process! This is the principle on which nuclear energy and nuclear weapons are based.
The reason alpha decay occurs is because the nucleus has too many protons
which cause excessive repulsion. In an attempt to reduce the repulsion, a
Helium nucleus is emitted. The way it works is that the Helium nuclei are in
constant collision with the walls of the nucleus and because of its energy
and mass, there exists a nonzero probability of transmission. That is, an
alpha particle (Helium nucleus) will tunnel out of the nucleus. Here is an
example of alpha emission with americium-241:
Beta decay occurs when the neutron to proton ratio is too great in the
nucleus and causes instability. In basic
beta decay, a neutron is turned into a
proton and an electron. The electron is
then emitted.
The final type of beta decay is known as electron capture and
also occurs when the neutron to proton ratio in the nucleus is too
small. The nucleus captures an electron which basically turns a proton into a neutron. Here's a
diagram of electron capture with beryllium-7:
Gamma decay occurs because the nucleus is at too high an energy. The nucleus falls down to a lower energy
state and, in the process, emits a high energy photon known as a gamma particle.
44
Nuclear Chemistry Practice Problems
1) Select the correct equation when: Americium-244, Am, undergoes decay to form Curium-244, Cm..
2) Which of the following statements is true?
a. The atomic number is always greater than the atomic mass.
b. The mass number is the same for all atoms of the same element.
c. The atomic number is the sum of the number of particles in the nucleus
d. The difference between the atomic mass and the atomic number is the number of neutrons
e. All the above are correct
3) Neon-21 undergoes beta particle emission. Indicate the correct equation
Directions: Identify the following as alpha, beta, gamma, or neutron.
1.
1
0
n
2.
0
1
e
3.
4
2
He
4.
0
0
γ
7. Least penetrating nuclear decay
8. Most damaging nuclear decay to the human body
9. Nuclear decay that can be stopped by skin or paper.
10. Nuclear decay that can be stopped by aluminum.
Complete the following nuclear equations.
11.
42
19
13.
9
4
K →
Be
→
0
-1
e +
9
4
12.
__________
Be + __________
14.
45
239
Pu
94
235
92 U
→
4
He
2
+ __________
→ _________ +
231
90
Th
CST Released Questions Practice Test
1) A weather balloon with a 2-meter diameter at
ambient temperature holds 525 grams of helium. What
type of electronic probe could be used to determine the
pressure inside the balloon?
6) In order to advance to the level of a theory, a
hypothesis should be
a.
b.
c.
d.
a. spectrophotometric
b. calorimetric
c. thermometric
d. barometric
obviously accepted by most people.
repeatedly confirmed by experimentation.
in alignment with past theories.
a fully functional experiment.
7) Matter is made of atoms that have positive centers
of neutrons and protons surrounded by a cloud of
negatively charged electrons. This statement is
2) Which would be most appropriate for collecting data
during a neutralization reaction?
a.
b.
c.
d.
a. a statistics program
b. a thermometer
c. a graphing program
d. a pH probe
an inference.
a hypothesis.
a theory.
an observation.
8)
3) A scientist observed changes in the gas pressure of
one mole of a gas in a sealed chamber with a fixed
volume. To identify the source of the changes, the
scientist should check for variations in the
a.
b.
c.
d.
temperature of the chamber.
air pressure outside the chamber.
isotopes of the gas.
molecular formula of the gas.
4) Electrical fires cannot be safely put out by dousing
them with water. However, fire extinguishers that
spray solid carbon dioxide on the fire work very
effectively. This method works because carbon dioxide
a.
b.
c.
d.
The model of ideal gases shown above is useful
because it
forms water vapor.
renders the fire’s fuel non-flammable.
displaces the oxygen.
blows the fire out with strong wind
currents.
a.
predicts the behavior of other phases of
matter.
b. shows a linear relation between gas pressure
and volume.
c. gives precise explanations for nonideal gas
behavior.
d. accurately approximates the properties of
most gas molecules.
5) In the cubic crystal shown, if each edge is 2.0 A in
length, what is the diagonal distance, d, between atoms
1 and 3? (Assume that the Pythagorean theorem can be
used to solve.)
a.
b.
c.
d.
9) When a metal is heated in a flame, the flame has a
distinctive color. This information was eventually
extended to the study of stars because
2.5 Å
a.
the color spectra of stars indicate which
elements are present.
b. star color indicates absolute distance.
c. a red shift in star color indicates stars are
moving away.
d. it allows the observer to determine the size of
stars.
Å
Å
Å
46
10)
14) Which of the following elements is classified
as a metal?
a.
b.
c.
d.
sulfur
bromine
lithium
helium
15)
Which of the following ordered pairs of
elements shows an increase in atomic number
but a decrease in average atomic mass?
a. Co to Ni
b. Ag to Pd
c. Cr to Mo
d. Ge to Sn
11) Why is cobalt (Co) placed before nickel (Ni)
on the periodic table of the elements even
though it has a higher average atomic mass
than nickel?
a.
b.
c.
d.
Nickel has fewer electrons.
Nickel has one more proton.
Cobalt was discovered first.
Cobalt has a lower density.
12) Generally, how do atomic masses vary
throughout the periodic table of the elements?
a.
They increase from right to left and
bottom to top.
b. They increase from left to right and
bottom to top.
c. They increase from right to left and
top to bottom.
d. They increase from left to right and
top to bottom.
The chart above shows the relationship
between the first ionization energy and the
increase in atomic number. The letter on the
chart for the alkali family of elements is
a.
b.
c.
d.
X
Y
W
Z
16) Which of the following atoms has the largest
atomic radius?
a.
b.
c.
d.
chlorine (Cl)
magnesium (Mg)
barium (Ba)
iodine (I)
17) Which of the following atoms has six valence
electrons?
13)
a.
b.
c.
d.
Iodine would have chemical properties most
like
a. tellurium (Te).
b. xenon (Xe).
c. chlorine (Cl).
d. manganese (Mn).
magnesium (Mg)
silicon (Si)
argon (Ar)
sulfur (S)
18) Which statement best describes the density of
an atom’s nucleus?
a.
The nucleus occupies little of the
atom’s volume but contains most of the
mass.
The nucleus occupies most of the atom’s
volume and contains most of its mass.
The nucleus occupies very little of the
atom’s volume and contains little of its
mass.
The nucleus occupies most of the
atom’s volume but contains little of its
mass.
b.
c.
d.
21) The most abundant isotope of lead contains 82
protons and 124 neutrons packed closely
together in the nucleus. Why do the protons
stay together in the nucleus rather than fly
apart?
a. Electrons in neighboring atoms neutralize
repulsive forces between protons.
b. Nuclear forces overcome repulsive forces
between protons in the nucleus.
c.
Electrostatic forces between neutrons
and protons hold the nucleus together.
d.
Neutrons effectively block the protons
and keep them far apart to prevent
repulsion.
19)
22) Which equation correctly represents the alpha
decay of polonium-214?
a.
What information do the experimental results
above reveal about the nucleus of the gold
atom?
The nucleus is small and is the
densest part of the atom.
b. The nucleus contains small positive
and negative particles.
c. The nucleus is large and occupies
most of the atom’s space.
d. The nucleus contains less than half
the mass of the atom.
b.
a.
20) Why are enormous amounts of energy required
to separate a nucleus into its component
protons and neutrons even though the protons
in the nucleus repel each other?
a. The forces holding the nucleus together are
much stronger than the repulsion between
the protons.
b. The interactions between neutrons and
electrons neutralize the repulsive forces
between the protons.
c. The force of the protons repelling each
other is small compared to the attraction
of the neutrons to each other.
d. The electrostatic forces acting between
other atoms lowers the force of repulsion
of the protons.
c.
d.
23) A 2-cm-thick piece of cardboard placed over a
radiation source would be most effective in
protecting against which type of radiation?
a.
b.
c.
d.
alpha
x-ray
beta
gamma
24) Which of the following is a monatomic gas at
STP?
a. helium
b. nitrogen
c. chlorine
d. fluorine
25) When cations and anions join, they form what
kind of chemical bond?
a. covalent
b. hydrogen
c. ionic
d. metallic
26) Which of the following correctly shows how
carbon and hydrogen bond to form a
compound?
29) The reason salt crystals, such as KCl, hold
together so well is because the cations are
strongly attracted to
a. neighboring cations.
b. free electrons in the crystals.
c. neighboring anions.
d. the protons in the neighboring
nucleus.
30) What type of force holds ions together in salts
such as CaF2?
a. magnetic
b. nuclear
c. gravitational
d. electrostatic
31) Under the same conditions of pressure and
temperature, a liquid differs from a gas
because the molecules of the liquid
a.
a. take the shape of the container they are in.
b. have stronger forces of attraction between
them.
c. have no regular arrangement.
d. are in constant motion.
b.
32)
c.
d.
27) Some of the molecules found in the human
body are NH2CH2COOH (glycine), C6H12O6
(glucose), and CH3(CH2)16COOH (stearic
acid). The bonds they form are
a. metallic
b. covalent
c. ionic
d. nuclear
Which of the following elements has the same
Lewis dot structure as silicon?
a. gallium (Ga)
b. aluminum (Al)
c. germanium (Ge)
d. arsenic (As)
33) Which substance is made up of many
monomers joined together in long chains?
28) What type of bond do all of the molecules in
the table above have in common?
a.
b.
c.
d.
polar
covalent
metallic
ionic
a.
b.
c.
d.
ethanol
protein
salt
propane
34) For the polymer, polyvinyl chloride, (PVC), ~
CH2CH(Cl)CH2CH(Cl)CH2CH(Cl) ~
the repeating subunit is
a. CH(Cl)
b. CH(Cl)CHCH2.
c. CH2CH.
d. CH2CH(Cl)
39) Methane (CH4) gas diffuses through air
because the molecules are
35) Which element is capable of forming stable,
extended chains of atoms through single,
double, or triple bonds with itself?
40) The volume of 400 mL of chlorine gas at 400
mm Hg is decreased to 200 mL at constant
temperature. What is the new gas pressure?
a.
b.
c.
d.
nitrogen
oxygen
carbon
hydrogen
36) Proteins are large macromolecules composed
of thousands of subunits. The structure of the
protein depends on the sequence of
a.
b.
c.
d.
a.
b.
c.
d.
monosaccharides.
lipids.
amino acids.
nucleosides.
37) When a cold tire is inflated to a certain
pressure and then is warmed up due to friction
with the road, the pressure increases. This
happens because the
a. Tire rubber reacts with oxygen in the
atmosphere.
b. air molecules speed up and collide with the tire
walls more often.
c. air molecules hit walls of the tire less frequently.
d. air molecules diffuse rapidly through the walls of
the tire.
38) When someone standing at one end of a large
room opens a bottle of vinegar, it may take
several minutes for a person at the other end to
smell it. Gas molecules at room temperature
move at very high velocities, so what is
responsible for the delay in detection of the
vinegar?
a. the chemical reaction with nerves, which is
slower than other sensory processes
b. random collisions between the air and vinegar
molecules
c. the increase in the airspace occupied by vinegar
molecules
d. attractive force between air and vinegar molecules
a.
b.
c.
d.
expanding steadily
traveling slowly.
moving randomly.
dissolving quickly.
650 mm Hg
300 mm Hg
800 mm Hg
400 mm Hg
41) Under what circumstance might a gas decrease
in volume when heated?
a.
b.
c.
d.
The gas is placed under increasing pressure.
The gas undergoes a decrease in pressure.
The gas remains under uniform temperature.
The gas is held constant at STP.
42) A sample of carbon dioxide gas occupies a
volume of 20 L at standard temperature and
pressure (STP). What will be the volume of a
sample of argon gas that has the same number
of moles and pressure but twice the absolute
temperature?
a.
b.
c.
d.
20 L
80 L
40 L
10 L
43) Standard temperature and pressure (STP) are
defined as
a.
b.
c.
d.
0 °C and 273 mm Hg pressure.
0 °C and 1.0 atm pressure.
0 K and 760 mm Hg pressure.
0 K and 1.0 atm pressure.
44) Under which of the following sets of
conditions will a 0.50 mole sample of helium
occupy a volume of 11.2 liters?
a.
b.
c.
d.
298 K and 0.90 atm
273 K and 1.10 atm
273 K and 1.00 atm
373 K and 0.50 atm
45) What is the equivalent of 423 kelvin in degrees
Celsius?
a.
b.
c.
d.
50) If the attractive forces among solid particles
are less than the attractive forces between the
solid and a liquid, the solid will
150 °C
–223 °C
–23 °C
696 °C
46) Theoretically, when an ideal gas in a closed
container cools, the pressure will drop steadily
until the pressure inside is essentially that of a
vacuum. At what temperature should this
occur?
a. −460 °C
b. 0 K
c. 0 °C
d. −273 K
a.
be unaffected because attractive forces
within the crystal lattice are too strong for
the dissolution to occur.
b. dissolve as particles are pulled away from
the crystal lattice by the liquid molecules.
c. probably form a new precipitate as its
crystal lattice is broken and re-formed.
d. begin the process of melting to form a
liquid.
51) Water is a polar solvent, while hexane is a
nonpolar solvent.
47) The temperature at which all molecular motion
stops is
a.
b.
c.
d.
0 °C.
0 K.
−273 K.
−460 °C
Which of the examples above illustrates a
nonpolar solute in a polar solvent?
48)
a.
b.
c.
d.
Which of the substances in the table can act as
either the solute or the solvent when mixed
with 100 grams of water at 20 °C?
a.
b.
c.
d.
CH3CH2OH
C6H5COOH
NH3
MgCl2
49) A teaspoon of dry coffee crystals dissolves
when mixed in a cup of hot water. This process
produces a coffee solution. The original
crystals are classified as a
a.
b.
c.
d.
solvent.
reactant.
solute.
product.
C2H5OH in hexane
CO(NH2)2 in hexane
NH4Cl in water
C10H8 in water
52) A technician prepared a solution by heating
100 milliliters of distilled water while adding
KCl crystals until no more KCl would
dissolve. She then capped the clear solution
and set it aside on the lab bench. After several
hours she noticed the solution had become
cloudy and some solid had settled to the
bottom of the flask. Which statement best
describes what happened?
a.
Water molecules, trapped with the KCl
crystals, were released after heating.
b. At increased temp the solubility of KCl
increased and remained high after cooling.
c. At lower temp the solubility of KCl
decreased and recrystallization occurred.
d. As the solution cooled, evaporation of water
increased the KCl concentration beyond its
solubility.
53) If the solubility of NaCl at 25 °C is 36.2 g/100
g H2O, what mass of NaCl can be dissolved in
50.0 g of H2O?
a.
b.
c.
d.
86.2 g
36.2 g
18.1 g
72.4 g
54) How many moles of HNO3 are needed to
prepare 5.0 liters of a 2.0 M solution of HNO3?
a.
b.
c.
d.
5
20
2.5
10
55) The Dead Sea is the saltiest sea in the world. It
contains 332 grams of salt per 1000 grams of
water. What is the concentration in parts per
million (ppm)?
a.
b.
c.
d.
59) The specific heat of copper is about 0.4
joules/gram °C. How much heat is needed to
change the temperature of a 30-gram sample of
copper from 20.0 °C to 60.0 °C?
a.
b.
c.
d.
240 J
1000 J
720 J
480 J
60) Equal volumes of 1 molar hydrochloric acid
(HCl) and 1 molar sodium hydroxide base
(NaOH) are mixed. After mixing, the solution
will be
a.
b.
c.
d.
weakly basic.
strongly acidic.
weakly acidic.
nearly neutral.
61)
33,200 ppm
332,000 ppm
332 ppm
0.332 ppm
56) The random molecular motion of a substance
is greatest when the substance is
a.
b.
c.
d.
a gas.
frozen.
a liquid.
condensed.
57) Which of these is an example of an exothermic
chemical process?
a.
b.
c.
d.
photosynthesis of glucose
melting ice
evaporation of water
combustion of gasoline
58) The boiling point of liquid nitrogen (LN) is 77
K. It is observed that ice forms at the opening
of a container of liquid nitrogen. The best
explanation for this observation is
a. water at 0 ˚ C is colder than LN and freezes.
b. water trapped in the LN escapes and freezes.
c. the water vapor in the air over the opening of
the liquid nitrogen freezes out.
d. the nitrogen boils and then cools to form a
solid at the opening of the container.
The above picture shows a light bulb
connected to a battery with the circuit
interrupted by a solution. When dissolved in
the water to form a 1.0 molar solution, all of
the following substances will complete a
circuit allowing the bulb to light except
a.
b.
c.
d.
sucrose.
sodium nitrate.
ammonium sulfate.
hydrochloric acid.
62) Which of the following is an observable
property of many acids?
a. They become slippery with water.
b. They produce salts when mixed with acids.
c. They become more acidic when mixed with a
base.
d. They react with metals to release hydrogen
gas.
63) Copper (II) nitrate and sodium hydroxide
solutions react in a test tube as shown below.
Cu(NO3)2(aq) + 2NaOH(aq) → Cu(OH)2(s) +
2NaNO3(aq)
If nitric acid is added to the test tube, the amount
of solid precipitate decreases. The best explanation
for this is that the acid
a. reacts with the copper (II) nitrate,
pulling the equilibrium to the left.
b. will react with the copper (II)
hydroxide to form water and soluble
copper (II) nitrate.
c. dilutes the solution making the
precipitate dissolve.
d. will dissolve most solids, including
sodium nitrate.
64) Potassium hydroxide (KOH) is a strong base
because it
a. does not conduct an electric current.
b. easily releases hydroxide ions.
c. does not dissolve in water.
d. reacts to form salt crystals in water.
68) If the reaction below takes place inside a
sealed reaction chamber, then which of these
procedures will cause a decrease in the rate of
reaction?
2CO + O2 → 2CO2
a. adding more CO to the reaction chamber
b. increasing the volume inside the chamber
c. raising the temperature of the chamber
d. removing the CO2 as it is formed
69) A catalyst can speed up the rate of a given
chemical reaction by
a. increasing the pressure of reactants, favors
products.
b. lowering the activation energy required
c. increasing equilibrium constant favors products.
d. raising the temperature at which the reaction
occurs
70) Which reaction diagram shows the effect of
using the appropriate catalyst in a chemical
reaction?
65) Of four different laboratory solutions, the
solution with the highest acidity has a pH of
a.
b.
c.
d.
11.
3.
5.
7.
a.
66)
Which of these describes the rate of this
chemical reaction?
a.
an increase in the concentration of
HCl with time
b. an increase in H2 and Cl2 with time
c. an increase in the concentration of
HCl and H2 with time
d. a decrease in HCl and Cl2 with time
67) Which of the following changes will cause an
increase in the rate of the reaction below?
a.
b.
c.
d.
increasing the concentration of Br2
decreasing the concentration of C6H6
decreasing the temperature
increasing the concentration of HBr
b.
c.
d.
76)
71) H2O2, hydrogen peroxide, naturally breaks
down into H2O and O2 over time. MnO2,
manganese dioxide, can be used to lower the
energy of activation needed for this reaction to
take place and, thus, increase the rate of
reaction. What type of substance is MnO2?
a. an enhancer
b. a reactant
c. a catalyst
d. an inhibitor
72) When a reaction is at equilibrium and more
reactant is added, which of the following
changes is the immediate result?
a.
The reverse reaction rate remains the
same.
b. The forward reaction rate increases.
c. The forward reaction rate remains the
same.
d. The reverse reaction rate decreases.
73) In which of the following reactions involving
gases would the forward reaction be favored
by an increase in pressure?
a.
b.
c.
d.
2A + B ↔ C + 2D
A + B ↔ AB
A+B↔C+D
AC ↔ A + C
74)
Which action will drive the reaction above to
the right?
a.
b.
c.
d.
increasing the system’s pressure
heating the equilibrium mixture
adding water to the system
decreasing the oxygen concentration
75)
The reaction shown above occurs inside a
closed flask. What action will shift the reaction
to the left?
a.
b.
c.
d.
venting some CO2 gas from the flask
pumping CO gas into the closed flask
raising the total pressure inside the flask
increasing the NO concentration in the
flask
What kind of change will shift the reaction
above to the right to form more products?
a.
b.
c.
a decrease in total pressure
a decrease in temperature
an increase in the concentration of
HCl
d. an increase in the pressure of NH3
77) In a sealed bottle that is half full of water,
equilibrium will be attained when water
molecules
a.
are equal amount for both liquid and gas
phase.
b. cease to evaporate.
c. begin to condense.
d. evaporate and condense at equal rates.
78) C3H8 + O2 → CO2 + H2O
This chemical equation represents the
combustion of propane. When correctly
balanced, the coefficient for water is
a. 16.
b. 8.
c. 2.
d. 4.
79) Which of the following is a balanced equation
for the combustion of ethanol (CH3CH2OH)?
a. CH3CH2OH + O2 → 2CO2 + 2H2O
b. CH3CH2OH + 3O2 → 2CO2 + 3H2O
c. CH3CH2OH + 2O2 → 3CO2 + 2H2O
d. CH3CH2OH + 3O2 → CO2 + 2H2O
80) N2H4, and N2O4 react to form gaseous nitrogen
and water. Which of these represents a
properly balanced equation for this reaction?
a. 2N2H4 + N2O4 → 2N2 + 4H2O
b. N2H4 + N2O4 → N2 + H2O
c. 2N2H4 + 3N2O4 → 5N2 + 6H2O
d. 2N2H4 + N2O4 → 3N2 + 4H2O
81)
__NH3(g) + __O2(g) → __N2(g) + __H2O(g)
When the reaction above is completely
balanced, the coefficient for NH3 will be
a.
b.
c.
d.
3
6
2
4
82) How many moles of carbon-12 are contained
in exactly 6 grams of carbon-12?
a. 3.01 ×1023 moles
b. 2.0 moles
c. 0.5 mole
d. 6.02 ×1023 moles
83) How many atoms are in 97.6 g of platinum
(Pt)?
a. 5.16 × 1030
b. 1.10 × 1028
c. 1.20 × 1024
d. 3.01 × 1023
84) When methane (CH4) gas is burned in the
presence of oxygen, the following chemical
reaction occurs.
CH4 + 2O2 → CO2 + 2H2O
If 1 mole of methane reacts with 2 moles of
oxygen, then
a.
23
24
6.02 ×10 molecules of CO2 and 1.2 ×10
molecules of H2Oare produced.
b. 1.2 ×1024 molecules of CO2 and 1.2 ×1024
molecules of H2Oare produced.
c. 6.02 ×1023 molecules of CO2 and 6.02 ×1023
molecules of H2Oare produced.
d. 1.2 ×1024 molecules of CO2 and 6.02 ×1023
molecules of H2Oare produced.
85) How many moles of CH4 are contained in 96.0
grams of CH4?
a. 6.00 moles
b. 3.00 moles
c. 12.0 moles
d. 16.0 moles
86) How many atoms are in a chromium sample
with a mass of 13 grams?
a. 3.3 ×1023
b. 1.5 ×1023
c. 1.9 ×1026
d. 2.4 ×1024
87) How many moles of chlorine gas are contained
in 9.03 ×1023 molecules?
a.
b.
c.
d.
2.0 moles
6.02 moles
1.5 moles
9.03 moles
88) Fe2O3 + 3CO → 2Fe + 3CO2
In this reaction, how many grams of Fe2O3 are
required to completely react with 84 grams of
CO?
a. 64 g
b. 160 g
c. 1400 g
d. 80 g
89)
Mg3N2(s) + 6H2O(l) → 2NH3(aq) +
3Mg(OH)2(s)
If 54.0 grams of water are mixed with excess
magnesium nitride, then how many grams of
ammonia are produced?
a.
b.
c.
d.
153
1.00
17.0
51.0
90) A mass of 5.4 grams of aluminum (Al) reacts
with an excess of copper (II) chloride (CuCl2)
in solution, as shown below.
3CuCl2 + 2Al → 2AlCl3 + 3Cu
What mass of solid copper (Cu) is produced?
a.
b.
c.
d.
19 g
8.5 g
0.65 g
13 g