Download Worksheet Significant Figures

MOUNT VERNON HIGH SCHOOL
IB CHEMISTRY 1: 2011-2012
IB CHEMISTRY 1
WORKBOOK
1
2
Course Content
IB Chemistry 1
Mrs. McManus and Mrs. Townley
Measurement and Data Processing
Atomic Theory (SL)
Periodicity (SL)
Bonding (SL)
Stoichiometry (SL)
Energetics (SL)
Kinetics (SL) and Equilibrium (SL)
Acids and Bases (SL)
Oxidation and Reduction (SL)
Topic 11
Topic 2
Topic 3
Topic 4
Topic 1
Topic 5
Topics 6 and 7
Topic 8
Topics 9
The topics above will be covered during the course of the year. The numbered topics are taken
directly from the IB Program syllabus. This workbook contains objectives, worksheets, labs
and study guides for the topics shown.
3
FCPS Student Laboratory Safety Rules
You and your classmates will be participating in many hands-on laboratory activities this year. Some
of these activities may require the use of materials or pieces of equipment that are potentially harmful if
not handled in a safe manner. Carefully review the following rules for student conduct. After you have
reviewed these rules with your instructor, sign the statement indicating you understand and agree to
follow these rules. Review these rules with your parent(s)/guardian(s) and have them sign the form,
indicating that they understand the risks and rules you will follow and support your adherence to these
safety rules. Failure to abide by these rules will result in removal from the lab and forfeiture of
the grade.
1. No unauthorized experiments are permitted. Do not modify or change the design of your
experiment without instructor’s written approval.
2. Avoid loose fitting clothing, open-toed shoes or sandals, dangling jewelry, and tie back long hair.
Ponytails that could fall over the shoulders must be folded over and tied. Loose clothing or hair can
catch on fire or knock over equipment causing accidents.
3. Wear appropriate eye protection at all times (safety goggles). Chemicals and sharp objects can
damage unprotected eyes. If you splash chemicals in your eyes, notify your instructor immediately
and begin flushing eyes for 15-30 minutes in the eye wash.
4. Absolutely no horseplay of any kind is permitted in lab. This can cause accidents.
5. No visiting by friends is allowed during lab sessions. Keep your attention on the laboratory at hand
– more attention equals fewer accidents.
6. Eating, drinking, applying cosmetics, chewing on gum, pencils or fingernails is never permitted in
the labs. Assume everything you handle in lab is contaminated. Eating or drinking in the lab could
allow you to ingest poisons. Wash your hands before leaving the lab. This prevents your taking
trace chemicals to other locations.
7. Do not deliberately smell or taste chemicals unless instructed to do so. When testing odors, use a
wafting motion of your hand to direct the odors to your nose.
8. Fire is to be used with caution in the laboratory. Some organic solvents are highly flammable. No
flammable solvents should be around an open flame. Follow your teacher’s instructions when
heating liquids so as to avoid fires.
9. Burners and hot plates require special care in usage. Never leave a heat source unattended when in
use. Remember that these as well as heated apparatus remain hot long after heating, so handle with
care. Be aware of cords and heating surfaces. If a fire alarm occurs, turn off Bunsen burners and
hot plates.
10. Never look directly into a test tube or flask. Do not point the opening of these at anyone especially
when heating. Liquids can quickly become boiling hot which can cause their contents to eject
forcefully from a tube. Do not use fingers as rubber stoppers.
11. Read bottle labels carefully and thoroughly. Always verify that you have selected the correct
chemical. Report unlabelled containers to the instructor. Mixing the incorrect chemicals can have
harmful results.
12. Never contaminate reagents by pouring unused portions into stock bottles.
4
13. Follow waste disposal procedures specified by your teacher. Do not dump chemicals into trashcans
or put into the sink unless instructed to do so. Never put any solid material into the sink!!
14. Clean up broken glass immediately.
15. If you spill anything during lab, notify your instructor immediately (acids and bases must be
neutralized before cleanup).
16. Wash your skin if contact is made with any chemical immediately and notify your instructor.
Chemical burns must be washed for 15-30 minutes.
17. Immediately report any unusual odors, broken equipment or unsafe situations to the instructor.
18. Immediately report any spill, cut, burn or other injury to the instructor.
19. No chemicals or supplies may be taken from storerooms or labs for home use unless approved by
teacher for independent (science fair) projects.
20. Know the location and use of the following laboratory safety equipment: eye wash station, fire
extinguisher, safety shower, fire blanket, fume hood.
21. Know the basic procedures to be followed in the event of a fire at your lab table, fire on a person,
cut or chemical burn, or chemical in the eyes.
5
Internal Assessment Form
Student:____________________________
Teacher:_______________________
Date:_________
Lab Title:___________________________
Topic/Option:_____________________ Hours:________
Students receive 2 marks for completely fulfilling an aspect, 1 mark for
partially fulfilling it and 0 marks for failing to fulfill the aspect.
Criterion
Design
Data
Collection/Processing
Conclusion and
Evaluation
Aspect 1
Aspect 2
Aspect 3
Formulates a
focused
problem/research
question and
identifies relevant
variables.
Designs a
method for
effective
control of the
variables.
Develops a
method that
allows for the
collection of
sufficient
relevant data.
2 1 0
Records
appropriate
quantitative and
associated
qualitative raw
data, including
units and
uncertainties
where relevant.
2 1 0
Processes
the
quantitative
raw data
correctly.
2 1 0
Presents
processes
data
appropriately
and, where
relevant,
includes errors
and
uncertainties.
2 1 0
States a
conclusion, with
justification,
based on a
reasonable
interpretation of
the data.
2 1 0
Evaluates
weaknesses
and
limitations.
2 1 0
Suggests
realistic
improvements
in respect of
identified
weaknesses
and
limitations.
2
1
0
2
6
1
0
2
1
0
Total
Marks
7
8
Topic 11: Measurement and Data Processing
11.1 Uncertainty and Error in Measurement





11.1.1 Describe and give examples of random uncertainties and systematic errors.
11.1.2 Distinguish between precision and accuracy.
11.1.3 Describe how the effects of random uncertainties may be reduced.
11.1.4 State random uncertainty as an uncertainly range(±).
11.1.5 State the results of calculations to the appropriate number of significant figures.
11.2 Uncertainties in Calculated Results.


11.2.1 State uncertainties as absolute and percentage uncertainties.
11.2.2 Determine the uncertainties in results.
11.3 Graphical Techniques




11.3.1 Sketch graphs to represent dependences and interpret graph behavior.
11.3.2 Construct graphs from experimental data.
11.3.3 Draw best-fit lines through data points on a graph.
11.3.4 Determine the values of physical quantities from graphs.
9
Topic 11 Notes
10
Measurement & Calculations
Metric System
We use several types of measurements:
length - meters (km, cm and mm)
mass - grams (kg and mg)
amount of substance – moles (mol)
time – seconds (s)
volume - Liters (mL) or cubic centimeters (cm3)
We commonly use the prefixes:
deci – 1 / 10 th
centi - 1 / 100 th
milli - 1 / 1000 th
kilo - 1000
Occasionally you will encounter micro (µ), nano, pico, mega, and giga.
Significant Digits: What do they mean?
In a measurement or a calculation, it is important to know which digits of the reported number are significant. That
means…if the same measurement were repeated again and again, some numbers would be consistent and some might
vary.
All of the digits that you are absolutely certain of plus one more that is a judgment are significant.
Consider: 16.82394 cm. If all the digits are significant, everyone who measures the object will agree that it is 16.8239
cm, but everyone will not agree about the final digit; some would say …94 cm while others might say …95 cm.
Rules for Recognizing Significant Digits




All non-zero digits are significant. 523 g (3); 972,366 sec (6)
0’s in the MIDDLE between two non-zero digits are ALWAYS significant. 5082 m (4); 0.002008 L (4)
0’s in the FRONT of a number are NEVER significant. 0.0032 kg (2); 0.00000751 m (3)
0’s at the END of a number with a decimal point present are ALWAYS significant. 2.000 L (4); 0.00500 g (3)
Significant Digits in Calculations
When you perform a calculation using measurements, often the calculator gives you an incorrect number of significant
digits. Here are the rules to follow to report your answers:
x and ÷ : The answer has the same # of sig. digits as the measurement in the problem with the least number of sig. digits.
Example: 3.71 cm x 8.1 cm = 30.051? 30. cm2 (2 sig. digits)
+ and – : The answer has the same # of decimal places as the measurement in the problem with the least number of
decimal places. example: 3.70 cm + 8.1 cm = 11.8 cm (1 decimal place)
Scientific Notation
Scientific notation uses a number between 1 and 9.99… x 10 n where n is an integer.
Know how to put numbers into scientific notation:
5392 = 5.392 x 103
0.000328 = 3.28 x 10-4
1.03 = 1.03
5500 = 5.5 x 103
Some 0’s in numbers are placeholders and are not a significant part of the measurement so they disappear when written in
scientific notation. Ex: 0.000328 above. In scientific notation, only the three sig. digits (3.28) are written.
Accuracy vs. Precision
Accuracy refers to how close a measurement is to some accepted or true value (a standard).
Ex: an experimental value of the density of Al is 2.69 g/mL. The accepted value is 2.70 g/mL. Your value is
accurate to within 0.4%. Percent error is used to express accuracy.
Precision refers to the reliability, repeatability, or consistency of a measurement. ± uncertainty and sig. digits are used
to express precision.
11
Worksheet Significant Figures
State the number of significant figures in each of the following measurements:
1. 8.523 cm
5. 0.00560 m
5
2. 9.50 x 10 s
6. 27 students
3. 0.0040040 kg
7. 950 g
4. 8.0000 m
8. 0.0000001 cm
For each of the following measurements (a) state whether or not the zero(s) are significant and (b)
state the rule that applies:
9.
9502 m
10.
52.0 cm
11.
0.0045 m
12.
370 s
Calculate the answer for the following problems using the proper number of significant figures and
appropriate units.
13.
15.2 cm x 0.0021 cm
14.
25. 6 m + 2365.2 m + 152.123 m
15.
28.7 cm2 / 4.00 cm
16.
(3.52 x 105 m) x (1.522 x 107 m)
17.
(9.0 x 1025 m) x (6.00 x 1010 m)
18.
10.000 cm – 8.99 cm
19.
0.00043256 m x 1.0 m
20.
(1.0 x 103 m) x (6 x 10-8 m)
Write the SI prefix that corresponds to the following:
21.
10 times smaller (101)
22.
1000 times smaller (103)
23.
1000 times larger (103)
24.
100 times smaller (102)
25.
1 000 000 times smaller (106)
Read the following measurement devices to the proper number of significant digits:
26.
cm
2
3
4
20
30 40
5
6
27.
cm
50 60
28.
cm
4
29.
5
6
5
4
mL
12
Calculations Using Significant Figures


When multiplying and dividing, limit and round your answer to the least number of
significant figures in any of the factors. You are only as good as your least precise
measurement.
Ex. 23.0 cm  432 cm  19 cm = 188.784 cm3
The answer is expressed as 1.9  102 cm3 since 19 cm has only two significant
figures.
When adding and subtracting, limit and round your answer to the least amount of decimal
places in any of the measurements.
Ex. 123.25 mL + 46.0 mL + 86.257 mL = 255.507 mL
The answer is expressed as 255.5 mL since 46.0 mL has only one decimal
place.
Perform the following operations expressing the answer in the correct number of significant figures.
1.
1.35 m  2.467 m =
____________
_____________
(Calculator answer)
(Answer w/ sig. Figs)
2.
1,305 m2  42 m =
____________
_____________
3.
12.01 mL + 35.2 mL + 6 mL =
____________
_____________
4.
55.46 g – 28.9 g =
____________
_____________
5.
0.21 cm  3.2 cm  100.1 cm =
____________
_____________
6.
0.15 cm + 1.15 cm + 1.051 cm = ____________
_____________
3
7.
150 L  4 L =
____________
_____________
3
2
8.
1.278  10 m  1.4267  10 m = ____________
_____________
Percentage Error
Percentage error is a way to express how far off an experimentally determined value is from the
accepted or true value.
% error = (accepted – experimental)  accepted value  100
Determine the percentage error in the following problems. Remember the rules for significant
figures.
1.
Experimental value = 1.24 g, Accepted value = 1.30 g
2.
Experimental value = 0.124 g, Accepted value = 0.0998 g
3.
Experimental value = 252 mL, Accepted value = 225 mL
4.
Experimental value = 22.2 L, Accepted value = 22.4 L
5.
Experimental value = 125.2 mg, Accepted value = 124.8 mg
13
Dimensional Analysis
We can use equalities to change from one unit to another. For example:
1 minute = 60 seconds
Therefore:
1 min
60 sec
=1=
60 sec
1 min
To convert one unit to another use the following steps:
1.
Write the given number and unit.
2.
Set up a conversion factor (like the fraction shown above) comparing the given unit to
another.
a.
Place the given unit in the denominator of the conversion factor
b.
Place desired unit as numerator
c.
Place a “1” in front of the larger unit
d.
Determine the number of small units in “1” larger unit
3.
Cancel units. Solve the problem
Example: 25 km = _____ cm
25 km
1000 m
1 km
100 cm
1m
= 2.5 x 106 cm
Problems. Convert the following measurements. Show all work and conversion factors for full
credit.
1.
2.5 hours = ______ sec
2.
15 cm = _____ m
3.
4.3 x 105 mm = _____ m
4.
8.00 x 10-3 km = _____ dm
5.
0.0075 m = _____ cm
6.
3.5 days = _____ sec
7.
2.5 x 103 mL = _____ L
8.
8,943 mg = _____ kg
14
Experimental Design
When designing an experiment “from scratch”, there are several things to consider. The experiment
should be designed so that there is a hypothesis, independent and dependent variables, constants, controls,
measurable data, repeated trials and graphs. What are these components?
HYPOTHESIS- an educated guess as to the relationship between the variables that can be tested. It can be
put into the form of: “If I change (the independent variable), then (the dependent variable) will do
something.
INDEPENDENT VARIABLE- (manipulated variable) - the variable that is purposely changed by the
experimenter.
DEPENDENT VARIABLE- (responding variable) - the variable that responds to the manipulated changes. It
is usually measured with items such as stopwatches, measuring devices, balances, etc. Try to eliminate
subjectivity. Do a ranking of change if there is no other way to measure.
CONSTANTS- all factors that remain the same and have a fixed value throughout the repeated trial of the
experiment. Constantly check to make sure they do stay the same.
CONTROL- the standard for comparing experimental effects. With living organisms, it will generally be the
ones grown at the optimal conditions for that organism.
REPEATED TRIALS- the number of experimental repetitions, object or organisms tested at each level of the
independent variable. They help to reduce the effects of errors. How many should you have? It depends
on the nature of the experiment- living things require more trials since they have more variation.
GRAPH- the independent variable is always on the X axis while the dependent variable is on the Y axis. Bar
graphs are used when the data is qualitative (descriptive, based on observations or categories of data).
Line graphs are used when the data is quantitative (more precise, measured with tools).
**VERY IMPORTANT** When designing an experiment, you should have only one independent and one
dependent variable. You can have several experiments with different independent variables, but they must
NOT be done at the same time!
Answer the following questions about experiment design.
A chemical experiment is performed by mixing substance A with substance B. A color change indicates the
completion of the reaction. The time it takes for the color change to occur is recorded below. Each
experiment is done in the same beaker.
Experiment
#
1
2
3
Volume of A
(mL)
5.0
5.0
5.0
Volume of B
(mL)
5.0
10.0
15.0
Temperature
(oC)
25
26
24
Time of
reaction (sec)
45
61
78
1.
Identify the independent variable.
2.
Identify the dependent variable.
3.
Identify the constants.
4.
Name two things that can be done to improve this experiment.
5.
If a graph were to be made, what measurement (remember units) would be placed on the x-axis?
6.
If a graph were to be made, what measurement (remember units) would be placed on the y-axis?
15
Experimental Design (cont.)
When designing your own experiment you must choose appropriate variables and constants, create
appropriate tables to collect measurable data and present the findings with charts and graphs.
Let’s design an experiment! (You won’t actually have to perform the experiment, just design it!)
We want to study plant growth (something you know about from biology).
1. Make a list of at least 5 variables affecting plant growth.
a.
b.
c.
d.
e.
2. From the list you created, choose an independent variable, the variable you will manipulate in
your experiment. ____________________________
3. Choose a dependent variable, the variable you will measure the changes in based on your
manipulations of the independent variable. ______________________
4. List your constants in the experiment.
5. Write an appropriate title for your experiment:
6. Create a data table for your experiment (you don’t have actual numbers but you can make a table
labeled with appropriate units.)
7. To make a graph of your data, label the x and y axes below with the appropriate units.
16
Uncertainty in Measurement
Look at the following measurement device.
5
10
cm
1. What place value does each mark on the device represent? (tens, ones, tenths etc.)
2. Read the device to the proper place value by guessing one place smaller than the smallest place
value on the device.
3. In what place value does uncertainty exist? (tens, ones, tenths, etc.)
4. Record the measure using the proper uncertainty. (± notation in the uncertain place value)
Density Calculations using uncertainty
Density is calculated with two measurements, mass ÷ volume. Since both measurements have
uncertainty, any calculated density will have uncertainty as well. There are two rules to keep in mind
When adding or subtracting measurements, add the uncertainties.
When multiplying or dividing measurements, add the percent uncertainties.
 Example1 (add/sub): 12.34 ± .01 g
- 10.11 ± .01 g
──────────
2.23 ± .02 g

Example 2 (mult/div):
1.23 ± .01 g ÷ 2.0 ± .1 ml
First determine the % uncertainty:
.01 / 1.23 × 100 = 0.81 %
.1 / 2.0 × 100 = 5 %
Now divide the values and add the uncertainties:
1.23 g ÷ 2.0 ml = 0.61 g/ml
0.81 % + 5 % = 5.81%
Write the final answer by converting the % back to a numerical value:
5.81% × 0.61 = 0.04
Final answer: 0.61 ± 0.04 g/ml
17
Uncertainty Worksheet
Addition and Subtraction
1. 5.67 ± .01 g + 8.28 ± .01 g
2. 89.6 ± .2 cm - 25.7 ± .2 cm
3. 28 ± 1 g + 165 ± 2 g
Multiplication and Division
Remember, you must first find the percent uncertainties!
4.
8.75 ± .02 m × 5.67 ± .02 m
5.
21.5 ± .2 g ÷
6.
125 ± 2 g × 3.7 ± .2 m
5.78 ± .02 m
Putting it all together! Given the following measurements, calculate the density of the unknown
liquid showing the uncertainty.
Mass of empty graduated cylinder: 25.00 ± .01 g
Mass of cylinder plus liquid:
28.54 ± .01 g
Volume of liquid in cylinder:
2.0 ± .1 ml
(Show each step of the calculations)
Mass of liquid:
Density of liquid:
18
Lab- Density Determination
Purpose: To practice measurement and calculation techniques with uncertainty.
Procedure
Part 1
1. Determine the mass of the regularly shaped object.
2. Measure the length, width, and height of the object. Remember the precision of the
measuring device you use! Note the smallest marking on the measurement device and record
your measurements with the proper uncertainty.
3. Using the same object, determine the volume by water displacement. Again, note the smallest
marking on the measurement device and record your measurement with the proper
uncertainty.
Part 2
You will determine the density of water by devising a procedure to determine the relationship
that exists between volume and mass of water. Collect at least 5 data points and graphically represent
your findings. Be sure to write your procedure and data in your lab notebook. Identify your
independent and dependent variable.
Data
Record all data in a neatly labeled format
Calculations
Show all work with units and use uncertainty notation.
1. Calculate the volume of the object using its length, width and height.
2. Calculate the density of the object using the volume found in #1.
3. Calculate the volume of the object using the water displacement data.
4. Calculate the density of the object using the volume found in #3. This answer may be
slightly different than the answer in #2.
5. Make a graph of the mass and volume data collected in part 2 of the procedure. Title the
graph and label the axes with units. Determine the slope of the line (remember units!)
Questions
Answer in complete sentences
1. Compare the two methods you used in part 1 to determine volume. Which measurement was less
uncertainty? Justify your answer.
2. Define the terms intensive and extensive property. Is density an intensive or extensive property?
3. The accepted value for the density of water is 1.00 g/mL. What is your percent error for the
density of water?
Concluding Statements- answer the following questions
1. What was determined about the two density values found in Part 1? Explain your findings.
2. What type of relationship is shown on the density graph?
3. Why do scientists use uncertainties and significant digits?
19
Topic 11 Notes
20
Topic 2: Atomic Structure
2.1 The Atom







2.1.1 State the position of protons, neutrons and electrons in the atom.
2.1.2 State the relative masses and relative charges of protons, neutrons and electrons.
2.1.3 Define the terms mass number (A), atomic number (Z) and isotopes of an element.
2.1.4 Deduce the symbol for an isotope given its mass number and atomic number.
2.1.5 Calculate the number of protons, neutrons and electrons in atoms and ions from the
mass number, atomic number and charge.
2.1.6 Compare the properties of the isotopes of an element.
2.1.7 Discuss the uses of radioisotopes.
2.2 The Mass Spectrometer



2.2.1 Describe and explain the operation of a mass spectrometer.
2.2.2 Describe how the mass spectrometer may be used to determine relative atomic mass
using the 12C scale.
2.2.3 Calculate non-integer relative atomic masses and abundance of isotopes from given
data.
2.3 Electron Movement




2.3.1 Describe the electromagnetic spectrum.
2.3.2 Distinguish between a continuous spectrum and a line spectrum.
2.3.3 Explain how the lines in the emission spectrum of hydrogen are related to electron
energy levels.
2.3.4 Deduce the electron arrangement for atoms and ions up to Z=20.
12.1.3 Electron Configuration




12.1.3 State the relative energies of s, p, d, and f orbitals in a single energy level.
12.1.4 State the maximum number of orbitals in a given energy level.
12.1.5 Draw the shape of an s orbital and the shapes of the px, py, and pz orbitals.
12.1.6 Apply the Aufbau principle, Hund’s rule and the Pauli Exclusion principle to write
electron configurations for atoms up to Z = 54.
21
Topic 2 Notes
22
Subatomic Particles
Atoms are composed of even smaller particles known as subatomic particles. Fill in the following
table summarizing the properties of these particles.
Particle
Mass
(amu and g)
Charge
Location
Proton
Neutron
Electron
Define the following terms:
Atomic number (Z):
Mass number (A):
Isotope:
Ion:
Atomic mass:
Isotopic Notation (Shorthand!)
In the boxes below, identify what each letter represents.
A
X
Z
n+/-
How many electrons in 23Na+?
How many neutrons in 23Na+?
Write the shorthand notation for the negative 2 charged ion of sulfur with a mass number of 33.
23
Isotopes Worksheet
1. Fill in the table for each isotope listed
Symbol
20
Z
A
Number of
protons
Number of
electrons
Number of
Neutrons
Number of
electrons
Number of
Neutrons
Ne
201
66
Hg
Zn
27
Al
2. Complete the following table for the following atoms or ions.
Symbol
Z
A
36
84
Number of
protons
36
35
53
127
35
54
27
27
Zn
32
36
Cd2+
112
38
X2X3+
X3-
45
103
36
50
54
75
42
33
42
24
Isotopes and Average Atomic Mass
1. Elements contain a variety of isotopes. The atomic mass is the weighted average of all the
isotopes of an element.
 Example: A sample of cesium is 75% 133Cs, 20% 132Cs, and 5% 134Cs. What is the
average atomic mass?
0.75  133 = 99.75
0.20  132 = 26.4
0.05  134 = 6.7
total
= 132.85 amu = average atomic mass
Determine the average atomic mass of the following mixtures of isotopes.
a. 80% 127I, 17% 126I, 3% 128I
b. 50% 197Au, 50% 198Au
c. 15% 55Fe, 85% 56Fe
d. 99% 1H, 0.8% 2H, 0.2% 3H
2. The atomic mass of boron is given in the periodic table as 10.81, yet no single atom of boron has a
mass of 10.81 amu. Explain.
3. The element Europium exists in nature as two isotopes: Eu-151 and Eu-153. Their masses are
150.9196 amu and 152.9209 respectively. The average atomic mass of europium is 151.96 amu.
Calculate the relative abundance of the two europium isotopes.
25
Mass Spectrometry
Mass spectrometry is used to determine the relative abundance of isotopes of elements.
It works by vaporizing a substance then ionizing it (turning it into an ion by hitting it with electrons) so that it
can be accelerated through a tube and deflected by a magnet. The magnetic field required to deflect it gives a
measure of the mass to charge ratio. The more massive the ion, the stronger the magnetic field required to
deflect it so that it hits the detector. The more of a particular kind of ion, the bigger the peak on the spectra.
2. Substance is IONIZED
1. Substance is
VAPORIZED
4. Ions are
DEFLECTED
Spectrometer is under vacuum to
keep as few molecules in the
machine as possible to reduce the
number of particles that collide
with each other in the machine.
3. Ions are
ACCELERATED
5. Ions are DETECTED.
Sample spectra of the isotopes of a made-up element, E.
100
Relative
number of
Ions
50
0
24
25
m/z
(m = mass, z = charge)
Clearly the most abundant isotope is E-24. The other isotope, E-25, is only about 10% as common as E-24.
You can figure out the average atomic mass by:
1. Abundance of E-24: 100
Abundance of E-25: 10
2. Total amount: 100 + 10 = 110
3. % abundance E-24 = (100/110)*100 = 90.91 %
4. % abundance E-25 = (10/110)*100 = 9.09 %
5. Average atomic mass = (24 * 0.9091) + (25* 0.0909) = 24.1
(It’s important to know the percentages to several significant figures because the whole point of this is to get
the average atomic mass of an element to several significant figures.)
Mass spectrometry is also useful for determining the structure of organic compounds. See organic notes for
info on that.
26
Mass Spectrometer
The existence of isotopes can be proven with the use of a mass spectrometer. This device has 5
stages of operation: vaporization, ionization, acceleration, deflection and detection. (VIADD)
 Draw a simple diagram of a mass spectrometer showing where all five stages take place.

Describe what happens in each of the five stages.
Vaporization—
Ionization—
Acceleration—
Deflection—
Detection—

What factors determine the how much a particle is deflected?
The following results were obtained for a pure element (X) placed in a mass spec.
6
Relative
Abundance
2
85

86
87
88 (mass/charge)
Calculate the relative (average) atomic mass of element X from the above spectrum.
27
Bohr and the Atom
1. According to Bohr, electrons orbit the nucleus in what he called ________________.
2. The lowest energy level, represented by the letter n, is n = ______ .
3. Elements currently discovered have electrons occupying up to how many energy levels? ______
4. An electron going from n =1 to n = 2 must absorb/release energy. (circle one)
5. An electron going from n = 5 to n = 4 must absorb/release energy.
6. Which is more stable, an electron in n = 4 or in n = 3?
7. Define electromagnetic radiation.
8. Define quantum.
9. Name six types of electromagnetic radiation.
10. What properties are different between infrared and ultraviolet radiation? What property is the
same?
11. Wavelength and frequency are directly/inversely related.
12. Frequency and energy are directly/inversely related.
13. Write the equations showing (1) the relationship between wavelength and frequency, (2) and the
relationship between frequency and energy.
14. A wave with a frequency of 1.0 x 1020 Hz is what type of radiation?
15. A wave with a frequency of 1.0 x 109 Hz is what type of radiation?
16. What is the frequency of a light with a wavelength of 555 nm?
17. What is the energy of light with a frequency of 1.5 x 1014 Hz?
18. What is the energy of a quantum of energy with a wavelength of 175 nm?
28
Electron Configurations (Schrodinger)
1. Energy levels are further classified as having sublevels. How many sublevels are contained in the
following energy levels?
Energy Level
#of sublevels
2
4
n
2. List four types of sublevels in order of increasing energy.
3. The actual region in space having the highest probability of containing the electron is known as
an ___________.
4. How many orbitals are found in each of the following types of sublevels?
sublevel
# of orbitals
s
p
d
f
5. What is the maximum number of electrons that can occupy any orbital?
6. State the Pauli exclusion principle.
7. State Hund's rule.
8. How many electrons can occupy each of the following sublevels?
sublevel
# of electrons
s
p
d
f
9. Fill in the orbital filling diagrams below for the electrons in N, Mg, S, and Fe.
10. Write the electron configurations for N, Mg, S and Fe based on your diagrams above.
29
Electrons, Configurations and Orbital Diagrams
Fill in the following chart:
Element
# of
Valence
Electrons
Si
4
Electron configuration
Core
Outer electrons
config
Orbital filling diagram for
outer electrons
3s
3p

3s23p2
[Ne]
Fe
S2-
Cr
K+
Fill in the corresponding outer electron configurations as shown.
1s1 1s2
2p3
3d3
*
#
4f7
*
#
30
.
2.8.4
N
6s1
Electron
Arrangement
Spectroscopy: Element Identification and Emission Spectra
The energy levels in atoms and ions are the key to the production and detection of light. Energy
levels or "shells" exist for electrons in atoms and molecules. The colors of dyes and other
compounds results from electron jumps between these shells or levels. The colors of fireworks
result from jumps of electrons from one shell to another. An observation of light emitted by the
elements is also evidence for the existence of shells, sublevels and energy levels. The kinds of
light that interact with atoms indicate the energy differences between shells and energy levels in
the quantum theory model of the atom. Typically the valence electrons are the ones involved in
these jumps.
The "quantum" theory was proposed more than 90 years ago, and has been confirmed by
thousands of experiments. Science and education has failed to clearly describe the energy level
concept to almost four generations of citizens. This experiment is an exercise aimed at throwing a
little more light on the subject. (Don't laugh too hard at the joke.)
Atoms have two kinds of states; a ground state and an excited state. The ground state is the state
in which the electrons in the atom are in their lowest energy levels possible (atoms naturally are in
the ground state). This means the electrons have the lowest possible values for "n" the principal
quantum number.
Specific quantized amounts of energy are needed to excite an electron in an atom and produce an
excited state. An excited hydrogen atom with an electron in the n = 3 shell can release energy. If
the electron in hydrogen only drops to the n = 2 shell the energy matches a pulse of red light.
Energy can be added to atoms many different ways. It can be in the form of light, an electric
discharge or heat. This added or extra energy is emitted when the excited electrons in the atoms
give off light and fall back to lower shells. The light emitted has wavelengths and colors that
depend on the amount of energy originally absorbed by the atoms. Usually each individual excited
atom will emit one type of light. Since we have billions and billions of atoms we get billions of
excitations and emissions.
Not all atoms in a sample will absorb or be excited exactly the same. For example in hydrogen the
ground state has the electron in the n= 1 shell or level. The electron in some hydrogen atoms may
be excited into the n = 2 level. Other hydrogen atoms can have the electron excited into the n = 4
level.
Different elements emit different emission spectra when they are excited because each type of
element has a unique energy shell or energy level system. Each element has a different set of
emission colors because they have different energy level spacings. We will see the emission
spectra or pattern of wavelengths (atomic spectra) emitted by several different elements in this
lab. We will then identify an unknown element by comparing the color of the unknown with the
flame color of our knowns.
You need to know that white light is the combination of all colors of the spectrum.
31
Each color has a characteristic wavelength. The wavelength is the distance between the beginning
and end of a complete cycle of the light wave. All colors of light travel at the same speed, 3.0 x 108
meters/ second. The animation shows how a prism separates photons of red light from photons of
blue light. The photons of different colors fall in different positions on the color spectrum. The
position is determined by the wavelength.
Red light has longer wavelength and is lower in energy than blue light. The wavelength of red light
corresponds to the range of 700 to 600 nanometers, (7000 Ångstrom or 0.0000007 meters).
Blue light has shorter wavelength in the range of 400 nm (4000 Ångstrom or 0.00000004 meter, 1
Å = 1 x 10 -10 m = 0.0000000001 meter = 1 x 10-1 nanometer).
Spectroscopy is the analysis of light spectra and the way in which light interacts with matter. When
light is analyzed it is commonly separated into its component colors. The light source is directed
on a slit and the "beam" of light is separated using a prism or grating.
The reason that the images are lines is that the light from the lamp is focused on a narrow slit. The
illustration shows the separation of a light beam into its component colors.
This produces an image of the slit which has the shape of a line. The resulting beam of light can
be broken into the color spectrum or into its components of the spectrum emitted by the atom. You
can see the specific colors emitted by the light source. A white light source will give a spectrum
like the one shown above.
32
Experimental Procedure
Part 1 Flame tests and identification of an unknown metal.
Observe and record the color of the flame for each metal ion. Remember the metal ions are paired
with a nonmetal ion in an ionic formula unit. The electrical charges have to add to zero. The metal
ions are converted to atoms in the flame and then excited by the heat from the Bunsen burner flame.
The nonmetal ions, anions, do not get converted to atoms and do not and emit visible light like the
metals do.
Repeat the procedure for each known. Record the color observed for the unknown and use the color
to identify the cation in the unknown.
Metal ion
Observed Flame color
barium
_________________________
calcium
_________________________
sodium
_________________________
copper (II)
_________________________
potassium
_________________________
lithium
_________________________
strontium
_________________________
Part 1 Flame tests for unknown elements
Unknowns
Flame color
Identity of metal ion
based on flame test
Unknown 1
____________
__________
Unknown 2
____________
__________
Part 2 Observing line spectra with the spectroscope
In the second part of the experiment you will observe the color of light emitted by excited gases of
elements in sealed glass tubes called "spectrum" tubes. Direct current, DC, high voltage electrons
are used to excite the atoms in the spectrum tube. High voltage means 1000 to 2000 volts. This is
more than 10 times normal household voltage which is 120 volts AC.
The excited atoms release the energy they gained. Some of this energy is in the form of heat and
some is in the form of light. The billions of excited atoms release energy. Each excited atom
releases a single pulse of light energy as it returns to the "ground" state or low energy state. There
are so many pulses emitted the light appears to be continuous.
The excited atoms do not all emit the same energy light because the amount of energy that excited
them may differ, but there are limitations on the colors they do emit. The kind of light depends on the
size of the gaps between the "shells" or energy levels in the atom. The electrons are changing "n"
values in the atom. Remember "n" can have only positive whole number values like 1, 2, 3, ... up to
infinity.
33
The kind of light energy that can be emitted by excited atoms is unique for an element. The pattern
of "lines” or colors emitted can be used to identify an element. A powerful extension of this is the
ability to measure amounts of an element by measuring the brightness of the emitted light.
A spectroscope can separate the light produced by an emission tube. The color seen by the naked
eye is a combination of a number of colors of light. These are separated by a prism or a diffraction
grating which acts like a prism. The emission lines can be seen when you look through the
spectroscope at the light source. You will be able to observe the "line" spectrum for the elements
and record the spectral lines.
Element
Part 2 Emission line spectra for selected elements
Color with naked eye
Emission spectrum
Hydrogen
Mercury
Neon
Questions and observations
How do these emission spectra compare in terms of colors and numbers of
emission line positions?
Are the spectra identical?
What if anything is similar?
What is different?
FILL IN THE FOLLOWING TABLE WITH YOUR ANSWERS
Element with greatest number of visible emission lines
________________
_
Color of light for the longest wavelength
What produces the colors seen in the flame tests and the emission
spectra?
34
Spectroscopy lab (cont.)
Questions.
1.
Explain what causes the color that you see in the flame test portion of the lab. Be very
detailed in your description of the electron’s behavior. Write at least four COMPLETE
sentences in proper English to describe each aspect of the process.
2.
What is the significance of line spectrum? What did it prove the existence of? Why does it
verify that? Write at least four COMPLETE sentences in proper English.
3.
Discuss the color of light as it relates to wavelength and energy. Is red light more energetic?
What is the relationship between wavelength and energy? Feel free to bring in mathematical
relationships to clarify your answer. Write at least four COMPLETE sentences in proper
English.
35
Emission Spectrum of Hydrogen
http://www.avogadro.co.uk/light/bohr/spectra.htm
The diagram above shows several energy transitions as the electron falls from one energy level to
another. Each energy transition represents a very specific quantity of energy as mathematically
defined by the wave equation and Planck’s constant. Let’s look specifically at the emission spectrum
of hydrogen.
On the spectrum below color the lines representing the VISIBLE line spectrum of hydrogen. This
spectrum occurs only when electrons transition to n=2 as their final energy state. These transitions
create what is known as the Balmer Series. Label the lines with initial and final energy levels.
656 nm
n=3 to n=2
700 nm
1.
2.
3.
4.
350 nm
Which end of the spectrum represents more energy, 700 nm or 350 nm?
What color is at the 700 nm end of the spectrum?
Which transition represents the GREATEST amount of energy, n=3 to n=2 or n=4 to n=2?
Notice that the lines get closer and closer to one another as you approach the violet end of the
visible spectrum. Why do they converge?
5. An entire new series of lines was discovered when the electrons transitioned to n=1 (Lyman Series
in the UV). Would this represent smaller or larger energy differences? Would these lines
converge at a maximum energy just like the visible light series?
36
Models of the Atom
(SOL Content)
Match the following scientists with their model/experiment:
____ 1. Dalton
a. “Father of the Periodic Table”
____ 2. Thomson
b. quantum theory of energy
____ 3. Heisenberg
c. first to say all matter is made of atoms
____ 4. Democritus
d. oil drop experiment proved charge of
electron
____ 5. Bohr
e. electrons act as both waves and particles
____ 6. Rutherford
f. uncertainty principle
____ 7. deBroglie
g. electrons occupy an energy level
____ 8. Millikan
h. atom is mostly empty space, nucleus at
the center
____ 9. Mendeleev
i. plum pudding model
____ 10. Planck
j. atom is solid sphere
Answer each of the following questions.
11.
Why did Dalton rely so heavily on the results of others?
12.
What device did Thomson use in his experiments? What particle was found using this
device?
13.
Describe the Rutherford gold foil experiment. (Draw a diagram) What did it show that had
not been known previously?
14.
Explain how Bohr’s model of the atom corresponds with what you observed in the flame
test/emission spectra lab.
37
Introduction to Nuclear Chemistry
(SOL Content)
1.
What can be said about the nucleus of a radioactive isotope?
2.
Name two types particles a nucleus can emit in an attempt to become more stable.
3.
Fill in the following table regarding types of radioactive emissions.
Type of
emission
Alpha ()
Beta ()
Gamma ()
Mass
Charge
Composition
Shielding required to stop
particle
4. A beta particle is a high speed electron emitted from the nucleus. This is not to be confused
with electrons in the electron cloud. (Electrons in the cloud are ignored in nuclear reactions.) If
the nucleus consists of only protons and neutrons, how is it possible for an electron to be emitted
from it?
5. When balancing nuclear equations, what two properties must be conserved?
6. In isotopic notation of atoms, i.e. 146C, the number in the upper left refers to the
_____________, and the number in the lower left refers to the _______________.
7. Define half-life.
8. Complete the following table regarding half-lives of a radioactive isotope.
Number of half-lives
Amount remaining
Amount decayed
0
100% or 1
0% or 0
1
50% or ½
50% or ½
2
25% or
75% or
3
4
5
38
Nuclear Equations and Half-Life
(SOL Content)
Nuclear equations are balanced in a different manner than ordinary chemical reactions. Atoms are
NOT conserved! Instead the nuclear properties of mass and charge are conserved. Balance the
following nuclear equations. Remember in isotopic notation the superscript number is the mass
number (protons + neutrons), and the subscript, if shown, is the atomic number (number of protons).
Every element has a unique and constant atomic number that can be found on the periodic table.
H + 3H → _________
1.
1
2.
238
3.
42
K →
4.
27
Al + 4He →
5.
9
U →
4
2 He
0
-1e
+ _________
+ ________
30
P + _________
Be + 1H → _______ +
6. _______ + 01n →
142
Ba +
4
2He
91
Kr + 3 01n
Half-life is the time it takes for one half of a radioactive sample to decay. Answer the following
questions pertaining to half-life.
1. The half-life of iodine-131 is 8 days. What mass of I-131 remains from a 4.00 g sample after 32
days?
2. A sample of radioactive isotope with an original mass of 10.0 grams is observed for 24 days.
After that time, 1.25 grams of the isotope remains. What is its half-life?
3. After a period of four half-lives of a radioactive isotope has passed, what fraction of its original
mass will remain? What fraction has decayed?
4. The half-life of K-42 is 12.4 hours. How much of a 240.0 g sample will remain after 62.0 hours?
5. A 50.0 g sample of N-16 decays to 12.5 g in 14.4 seconds. What is its half-life?
39
TOPIC 2- NOTES
40
Study Guide: Topic 2 Atomic Theory
1.
Which statement is correct about the isotopes of an element? (2.1)
A.
They have the same mass number
B.
They have the same electron arrangement
C.
They have more protons than neutrons
D.
They have the same numbers of protons and neutrons
(Total 1 mark)
2.
How many electrons are there in one
A.
10
B.
12
C.
14
D.
22
24
+2
12 Mg
ion? (2.1)
(Total 1 mark)
3.
What is the symbol for a species that contains 15 protons, 16 neutrons and 18 electrons? (2.1)
A.
31
16
B.
31
16
S 3
C.
31
15
P
D.
31
15
P 3
S
(Total 1 mark)
4.
What is the difference between two neutral atoms represented by the symbols
(2.1)
A.
B.
C.
D.
59
27
Co and
59
28
Ni?
The number of neutrons only.
The number of protons and electrons only.
The number of protons and neutrons only.
The number of protons, neutrons and electrons.
(Total 1 mark)
5.
What is the correct number of each particle in a fluoride ion, 19F–? (2.1)
A.
B.
C.
D.
protons
9
9
9
9
neutrons
10
10
10
19
electrons
8
9
10
10
(Total 1 mark)
6.
A certain sample of element Z contains 60% of 69Z and 40% of 71Z. What is the relative atomic
mass of element Z in this sample? (2.2)
A.
69.2
B.
69.8
C.
70.0
D.
70.2
(Total 1 mark)
41
7.
Which ion would undergo the greatest deflection in a mass spectrometer? (2.2)
16 +
A.
O
16 2+
B.
O
18 2+
C.
O
16 18 +
D.
( O O)
(Total 1 mark)
8.
Which statement is correct about a line emission spectrum? (2.3)
A.
Electrons absorb energy as they move from low to high energy levels.
B.
Electrons absorb energy as they move from high to low energy levels.
C.
Electrons release energy as they move from low to high energy levels.
D.
Electrons release energy as they move from high to low energy levels.
(Total 1 mark)
9.
Which statement is correct for the emission spectrum of the hydrogen atom? (2.3)
A.
The lines converge at lower energies.
B.
The lines are produced when electrons move from lower to higher energy levels.
C.
The lines in the visible region involve electron transitions into the energy level closest to
the nucleus.
D.
The line corresponding to the greatest emission of energy is in the ultraviolet region.
(Total 1 mark)
10.
What is the electron arrangement of an Al3+ ion? (2.3)
A.
2, 8
B.
2, 3
C.
2, 8, 3
D.
2, 8, 8
(Total 1 mark)
11.
The electron arrangement of sodium is 2.8.1. How many occupied main electron energy levels are
there in an atom of sodium? (2.3)
A.
1
B.
3
C.
10
D.
11
(Total 1 mark)
12.
The element vanadium has two isotopes,
(2.1)
(a)
50
23
V and
51
23V,
and a relative atomic mass of 50.94.
Define the term isotope.
……………………………………………………………………………………….
……………………………………………………………………………………….
(1)
(b)
State the number of protons, electrons and neutrons in
50
23V.
……………………………………………………………………………………….
……………………………………………………………………………………….
42
(2)
(c)
State and explain which is the more abundant isotope.
……………………………………………………………………………………….
……………………………………………………………………………………….
(1)
(d)
State the name and the mass number of the isotope relative to which all atomic masses are
measured.
……………………………………………………………………………………….
(1)
(Total 5 marks)
13.
(a)
Define the term isotope. (2.2)
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
(b)
A sample of argon exists as a mixture of three isotopes.
mass number 36, relative abundance 0.337%
mass number 38, relative abundance 0.0630%
mass number 40, relative abundance 99.6%
Calculate the relative atomic mass of argon.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
(c)
State the number of electrons, protons and neutrons in the ion 56Fe3+.
electrons: ............................. protons: ............................. neutrons: ...........................
(2)
(Total 6 marks)
14.
Identify the numbers of protons, neutrons and electrons in the species 33S2–. (2.1)
..............................................................................................................................................
..............................................................................................................................................
(Total 1 mark)
43
15.
(a)
Describe the following stages in the operation of the mass spectrometer. (2.2)
(i)
ionization
(2)
(ii)
deflection
(2)
(iii)
acceleration
(1)
(b)
(i)
State the meaning of the term isotopes of an element.
(1)
(ii)
Calculate the percentage abundance of the two isotopes of rubidium 85Rb and 87Rb.
(2)
(iii)
State two physical properties that would differ for each of the rubidium isotopes.
(1)
(iv)
Determine the full electron configuration of an atom of Si, an Fe3+ ion and a P3– ion.
(3)
(Total 12 marks)
16.
(a)
Evidence for the existence of energy levels in atoms is provided by line spectra.
State how a line spectrum differs from a continuous spectrum. (2.3)
.....................................................................................................................................
.....................................................................................................................................
(1)
(b)
On the diagram below draw four lines in the visible line spectrum of hydrogen.
(1)
Low energy
(c)
High energy
Explain how the formation of lines indicates the presence of energy levels.
.....................................................................................................................................
.....................................................................................................................................
(1)
(Total 3 marks)
17.
State the electron arrangement, electron configuration and orbital filling diagrams for the
following particles: aluminium3+, nitrogen and fluorine1-. (2.3)
(Total 6 marks)
44
Topic 3: Periodicity
3.1 The Periodic Table




3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic
number.
3.1.2 Distinguish between the terms group and period.
3.1.3 Apply the relationship between the electron arrangement of elements and their position
in the periodic table up to Z=54.
3.1.4 Apply the relationship between the number of electrons in the highest occupied energy
level for an element and its position in the periodic table.
3.2 Physical Properties




3.2.1 Define the terms first ionization energy and electronegativity.
3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization, energies,
electronegativity and melting points for the alkali metals (Li Cs) and the halogens (F  I).
3.2.3 Describe and explain the trends in atomic radii, ionic radii, first ionization, energies,
electronegativity for elements across period 3.
3.2.4 Compare the relative electronegativity values of two or more elements based on their
positions in the periodic table.
3.3 Chemical Properties


3.3.1 Discuss the similarities and differences in the chemical properties of elements in the
same group.
3.3.2 Discuss the changes in nature, from ionic to covalent and from basic to acidic, of the
oxides across period 3.
45
Topic 3 NOTES
46
Periodic Table Exercise
Use the periodic table on the following page to complete the following activities.
1. Label the columns 1-18 across the top.
2. Label the rows 1-7 along the left edge.
3. Fill in the atomic numbers for the elements.
4. Fill in the symbols for the elements (be sure to write in small print).
5. Fill in the atomic masses of the elements.
6. Color H red (Color lightly enough to see the information underneath).
7. Color the remaining column one elements orange.
8. Color column two elements yellow.
9. Color columns 3-12 blue.
10. Color column 17 green.
11. Color column 18 purple.
12. Put an asterisk (*) in the top right corner of elements #80 and #35.
13. Put a dot () in the top left hand corner of elements #1, 2, 7-10, 17, 18, 36, 54, and 86.
Answer the following questions about the periodic table.
1. What phase of matter are most elements at room temperature and pressure?
2. At room temperature and pressure, what phase of matter are the elements you marked with an
asterisk (*)?
3. At room temperature and pressure, what phase of matter are the elements you marked with a
dot?
4. If a symbol for an element has two letters in it, is the second letter upper or lower case?
5. Why do some elements have symbols that seem unrelated to their common names, for
example sodium has the symbol Na?
6. Seven elements on the table occur in nature as diatomic molecules. Name those seven
elements and write their symbols.
7. As new elements are created/discovered, how are they initially named and how many letters
are in their symbols?
47
48
Periodicity
What two physical properties determine the strength of the force of attraction between two oppositely
charged particles?
(1)
(2)
Atomic Radii
Define atomic radius:
In general, atomic radii (increase/decrease) as you move down the periodic table.
Why?
In general, atomic radii (increase/decrease) as you move across the periodic table.
Why?
Define ionic radius:
When an atom gains electrons it forms a (cation/anion) with a (negative/positive) charge. The
newly formed ion is (larger/smaller) that the atom from which it was formed.
When an atom loses electrons it forms a (cation/anion) with a (negative/positive) charge. The
newly formed ion is (larger/smaller) that the atom from which it was formed.
For each list of particles, place them in order of increasing size.
a) K, Na, Rb
b) Fe2+, Fe, Fe3+
c) N-, N, N+
d) F, N, O
e) Cl-, S2-, Ar, K+1
Electronegativity
Define electronegativity:
In general, electronegativity (increases/decreases) as you move down the periodic table. Why?
In general, electronegativity (increases/decreases) as you move across the periodic table.
Why?
In a bond formed between Na and F, which atom more strongly attracts the electrons in the
bond?
Metals tend to have a (high/low) electronegativity and tend to (gain/lose) electrons when
bonded. They tend to form (positive/negative) ions.
Nonmetals tend to have a (high/low) electronegativity and tend to (gain/lose) electrons when
bonded. They tend to form (positive/negative) ions.
49
Ionization Energy (Ionisation energy)
Define first ionization energy:
In general, ionization energy (increases/decreases) as you move down the periodic table.
Why?
In general, ionization energy (increases/decreases) as you move across the periodic table.
Why?
The more tightly an electron is held to the nucleus the (higher/lower) the ionization energy.
From each pair listed below pick the species with the specified property:
a)
largest radius
Rb
or
Cs
b)
largest electronegativity
Rb
or
Cs
c)
largest ionization energy
Rb
or
Cs
d)
largest radius
Cl
or
I
e)
largest elctronegativity
Cl
or
I
f)
largest ionization energy
Cl
or
I
g)
largest radius
Na
or
Mg
h)
largest electronegativity
Na
or
Mg
i)
largest ionization energy
Na
or
Mg
j)
largest radius
K
or
K+
k)
largest radius
Br
or
Br-
50
Halogens
1.
Define the term diatomic.
2.
List the seven diatomic elements.
3.
How many valence electrons does a halogen atom have?
4.
What is the valence shell electron configuration for all halogen atoms?
5.
What charge ion do halogens tend to make?
6.
Write the appropriate symbols for the following species:
a.
chlorine
b.
chloride
7.
Explain the difference between a halogen and a halide.
8.
What is the periodic trend of the melting point of the halogens?
9.
All the halogens have similar chemical properties. Why?
10.
If aluminum forms a compound with chlorine with the formula of AlCl3, what would you
expect to be the formula of the compound formed between aluminum and bromine?
11.
The compound potassium fluoride has the formula KF. What would you expect to be the
formula of potassium iodide?
51
Periodic Trends in the Alkali Metals and Halogens
1. What is the trend in melting point as you go down the group of alkali metals?
2. What is the trend in melting point as you go down the group of halogens?
3. pH is a measure of the acidity of a solution. A pH of less than 7 means the solution is
acidic/basic. A pH greater than 7 means the solution is acidic/basic. Acidic solutions are
associated with high concentrations of the hydrogen (H+)/hydroxide ion(OH-). Basic solutions
are associated with high concentrations of the hydrogen (H+)/hydroxide ion(OH-).
4. Describe the reaction of alkali metals in water:
a. List visual observations
b. What gas is produced?
c. Is heat produced or absorbed?
d. Complete the general form of the equation of an alkali metal reacting with water.
Alkali metal + water → _______ + _______ + _______
e. Is the resulting solution acidic or basic? Justify.
f. What is the trend of reactivity as you go down the group of alkali metals?
g. What charge ion do alkali metals make?
5. Answer the following questions about reactions with halogens:
a. What charge ion would you expect the halogens to make?
b. What is the general trend in reactivity as you go down the group?
c. Is fluorine capable of replacing bromide?
d. Is iodine capable of replacing chloride?
e. Complete the following equation: chlorine + bromide → _____ + _____
f. Write the symbols for the equation in “e” above.
6. Answer the following questions about oxide compounds:
a. What is the formula for sodium oxide?
b. What is the formula for aluminum oxide?
c. When a metallic oxide is place in water a basic/acidic solution is formed with a pH
above/below 7.
d. When a nonmetallic oxide is place in water a basic/acidic solution is formed with a
pH above/below 7.
e. Define amphoteric.
f. An amphoteric substance will behave like a base/acid when in the presence of a strong
acid.
g. List an example of an amphoteric substance.
h. What is the trend of acidic/basic nature of the period 3 oxides?
52
Study Guide: Topic 3 Periodicity
1.
For which element are the group number and the period number the same?
A.
Li
B.
Be
C.
B
D.
Mg
(Total 1 mark)
2.
What is the total number of p orbitals containing one or more electrons in germanium (atomic
number 32)?
A.
2
B.
3
C.
5
D.
8
(Total 1 mark)
3.
Which is correct about the element tin (Sn) (Z = 50)?
Number of main energy levels
Number of electrons in
containing electrons
highest main energy level
A.
4
4
B.
4
14
C.
5
4
D.
5
14
(Total 1 mark)
4.
Which equation represents the first ionization energy of fluorine?
A.
F(g) + e–  F–(g)
B.
F–(g)  F(g) + e–
C.
F+(g)  F(g) + e–
D.
F(g)  F+(g) + e–
(Total 1 mark)
5.
Which factors lead to an element having a low value of first ionization energy?
I.
large atomic radius
II.
high number of occupied energy levels
III. high nuclear charge
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
6.
Which of the following properties of the halogens increase from F to I?
I.
Atomic radius
II.
Melting point
III. Electronegativity
A.
I only
B.
I and II only
C.
I and III only
D.
I, II and III
(Total 1 mark)
(Total 1 mark)
53
7.
Which statement is correct for a periodic trend?
A.
Ionization energy increases from Li to Cs.
B.
Melting point increases from Li to Cs.
C.
Ionization energy increases from F to I.
D.
Melting point increases from F to I.
(Total 1 mark)
8.
Which series is arranged in order of increasing radius?
A.
Ca2+ < Cl– < K+
B.
K+ < Ca2+ < Cl–
C.
Ca2+ < K+ < Cl–
D.
Cl– < K+ < Ca2+
(Total 1 mark)
9.
Which pair of elements reacts most readily?
A.
Li + Br2
B.
Li + Cl2
C.
K + Br2
D.
K + Cl2
(Total 1 mark)
10.
Rubidium is an element in the same group of the periodic table as lithium and sodium.
It is likely to be a metal which has a
A.
B.
C.
D.
high melting point and reacts slowly with water.
high melting point and reacts vigorously with water.
low melting point and reacts vigorously with water.
low melting point and reacts slowly with water.
(Total 1 mark)
11.
Explain why
(i)
the first ionization energy of magnesium is lower than that of fluorine.
(2)
(ii)
magnesium has a higher melting point than sodium.
(3)
(Total 5 marks)
54
12.
(a)
(i)
State the meaning of the term electronegativity and explain why the noble gases are
not assigned electronegativity values.
(2)
(ii)
State and explain the trend in electronegativity across period 3 from Na to Cl.
(2)
(iii)
Explain why Cl2 rather than Br2 would react more vigorously with a solution of I–.
(2)
(b)
State the acid-base properties of the following period 3 oxides.
MgO
Al2O3
P4O6
Write equations to demonstrate the acid-base properties of each compound.
(7)
(Total 13 marks)
13.
Describe the acid-base character of the oxides of the period 3 elements Na to Ar. For sodium
oxide and sulfur trioxide, write balanced equations to illustrate their acid-base character.
(Total 3 marks)
55
Topic 3 NOTES
56
Topic 4: Bonding
4.1 Ionic Bonding








4.1.1 Describe the ionic bond as the electrostatic attraction between oppositely charged ions.
4.1.2 Describe how ions can be formed as a result of electron transfer.
4.1.3 Deduce which ions will be formed when elements in groups 1, 2, and 3 lose electrons.
4.1.4 Deduce which ions will be formed when elements in groups 5, 6, and 7 gain electrons.
4.1.5 State that transition elements can form more than one ion.
4.1.6 Predict whether a compound of two elements would be ionic from the position of the
elements in the periodic table or from their electronegativity values.
4.1.7 State the formula of common polyatomic ions formed by nonmetals in periods 2 and 3.
4.1.8 Describe the lattice structure of ionic compounds.
4.2 Covalent Bonding










4.2.1 Describe the covalent bond as the electrostatic attraction between a pair of electrons and
positively charged nuclei.
4.2.2 Describe how the covalent bond is formed as a result of electron sharing.
4.2.3 Deduce the Lewis (electron dot) structures of molecules and ions for up to four electron
pairs on each atom.
4.2.4 State and explain the relationship between the number of bonds, bond length and bond
strength.
4.2.5 Predict whether a compound of two elements would be covalent from the position of the
elements in the periodic able or from their electronegativity values.
4.2.6 Predict the relative polarity of bonds from electronegativity values.
4.2.7 Predict the shape and bond angles for species with four, three and two negative charge
centres on the central atom using the valence shell electron pair repulsion theory (VSEPR).
4.2.8 Predict whether or not a molecule is polar from its molecular shape and bond polarities.
4.2.9 Describe and compare the structure and bonding in the three allotropes of carbon
(diamond, graphite and C60 fullerene).
4.2.10 Describe the structure of and bonding in silicon and silicon dioxide.
4.3 Intermolecular Forces


4.3.1 Describe the types of intermolecular forces (attractions between molecules that have
temporary dipoles, permanent dipoles or hydrogen bonding) and explain how they arise from
the structural features of molecules.
4.3.2 Describe and explain how intermolecular forces affect the boiling points of substances.
4.4 Metallic Bonding


4.4.1 Describe the metallic bond as the electrostatic attraction between a lattice of positive
ions and delocalized electrons.
4.4.2 Explain the electrical conductivity and malleability of metals.
4.5 Physical Properties

4.5.1 Compare and explain the properties of substances resulting from different types of
bonding.
57
Topic 4 NOTES
58
Ionic Bonding Basics
Define the following terms:
Ion:
Cation:
Anion:
Stable Octet:
Polyatomic ion:
1. When an atom loses electrons, what type of ion is formed?
2. When an atom gains electrons, what type of ion is formed?
3. Why do atoms form ions?
4. The most common charges of ions can be determined by placement in the periodic table. Give the
most common charge of an ion formed from the following families of atoms:
a. alkali metals
b. alkaline earth metals
c. transition metals
d. column 15 or 5A
e. column 16 or 6A
f. halogens
g. noble gases
5. When metals form ions they tend to (lose/gain) electrons.
6. When nonmetals form ions they tend to ______ electrons.
7. When ions come together to form an ionic compound, what is the overall charge of the
compound?
Fill in the following chart for polyatomic ions
Name
Formula Charge
ammonium
carbonate
hydrogen carbonate
acetate
nitrate
nitrite
phosphate
sulfate
hydrogen sulfate
sulfite
chlorate
hydroxide
59
Naming and Writing Formulas for Ionic Compounds
Write the names for the following:
1. BaCl2
11.
K2O
2. Zn(NO3)2
12.
Li2SO4
3. Ca(OH)2
13.
Al2S3
4. Sr(ClO3)2
14.
(NH4)3PO4
5. Mg3N2
15.
SrF2
6. NaC2H3O2
16.
NaHCO3
7. AgI
17.
KNO2
8. Mg(HSO4)2
18.
NH4C2H3O2
9. K2SO3
19.
BaSO4
10. AlP
20.
Na2O
21. calcium carbonate
31.
sodium sulfate
22. sodium sulfide
32.
aluminum nitrate
23. ammonium chloride
33.
barium acetate
24. strontium bromide
34.
magnesium phosphate
25. aluminum hydroxide
35.
lithium sulfite
26. barium nitrate
36.
potassium oxide
27. calcium acetate
37.
ammonium sulfide
28. magnesium iodide
38.
sodium nitride
29. potassium phosphide
39.
calcium hydrogen sulfate
30. lithium carbonate
40.
aluminum oxide
Write the formulas for the following:
60
Stock System
Most transition metals and the elements tin and lead have multiple oxidation states (they can have more than
one ionic charge.) To differentiate between the different charges a Roman numeral is used to express the
CHARGE of the ion. Name the following compounds. All require the use of the stock system.
1. CuCl
6. CuSO4
2. PbI2
7. Fe(NO3)3
3. SnCl4
8. Pb(OH)4
4. Fe2O3
9. SnSO4
5. CrF3
10. HgO
Write the formulas for the following compounds.
11. chromium (III) sulfate
16. tin (IV) oxide
12. iron (III) bromide
17. copper (II) phosphate
13. lead (II) acetate
18. iron (II) sulfide
14. copper (I) oxide
19. tin (II) fluoride
15. mercury (II) hydroxide
20. lead (IV) carbonate
Naming Simple Organic Compounds (Alkanes)
Hydrocarbons with the formula of CnH2n+2 are known as alkanes. They have their own naming system using
root words for the number of carbons in the compound. List the root words for the following number of
carbon atoms in the compound:
13572468Name the following compounds:
1. C3H8
4.
C8H18
2. CH4
5.
C2H6
3. C5H12
6.
C4H10
Write formulas for the following compounds:
7. hexane
9.
ethane
8. methane
10.
propane
61
Types of Chemical Bonds
Classify the following compounds as containing either ionic bonds (metal cation bonded to a
nonmetal anion), covalent bonds (nonmetal atom + nonmetal atom), or both (compound containing a
polyatomic ion substituted for one of the ions in an ionic compound.
1.
2.
3.
4.
5.
11.
CO2
Li2O
H2O
NaNO3
NO2
6.
7.
8.
9.
10.
KOH
N2O5
CaSO4
SrBr2
FeCl3
When two atoms have greatly different electronegativities the bond formed between them is
said to be which type, ionic or covalent?
Naming Molecular Compounds
Molecular compounds contain covalently bonded nonmetal atoms. A prefix system is used to
indicate the number of atoms of each element. The second element ends in the suffix
“–ide.” Following is a list of prefixes:
one: mono-*
four: tetratwo: difive: pentathree: trisix: hexa*Note: mono- is not used on the first element, only the second.
Name the following molecular compounds.
1.
2.
3.
4.
5.
CO2
CO
N2O
N2O5
SF6
6. N2O4
7. ClF3
8. PCl5
9. SiO2
10. SO3
Write formulas for the following compounds.
1. sulfur dioxide
2. dinitrogen trioxide
3. nitrogen monoxide
4. sulfur tetrafluoride
5. carbon disulfide
6. oxygen difluoride
62
Lewis Dot Diagrams and Review of Covalent Bonding
1. What is a single covalent bond?
2. What is a double covalent bond?
3. What is a triple covalent bond?
4. What is a dative (coordinate covalent) bond?
Electron dot diagrams are ways to represent the 2-dimensional structures of two types of
structures formed by covalent bonding, molecules (such as water or carbon dioxide) and polyatomic
ions (such as nitrate or ammonium). The symbol for each element represents the nucleus and core or
kernel electrons. The dots represent the valence electrons.
Steps for constructing Lewis dot diagrams
There are many methods for drawing dot diagrams. Your teacher will demonstrate the steps for you.
Use the space below to summarize the steps or rules.
1.
2.
3.
4.
5.
Examples:
HCl
C2H6
CH2Cl2
NH3
63
Practice Exercises
1. CH4
2. Cl2
3. H2
4. PH3
5. CHI3
6. SO2
7. CH3OH
8. N2
9. H2Te
10. H2CO
11. OF2
12. HCN
13. BF3
14. PCl3
15. SiO2
16. CO2
17. C2H4
18. C2H2
64
Polyatomic ions: The only difference between these and the regular covalent compounds is you
have to account for the charge when you add up your valence electrons. A positive charge means you
have less electrons, a negative charge means you have more electrons.
19. ClO3-
20. NO3-
21. ClO2-
22. NO2-
23. ClO-
24. NH4+
25. SO4-2
26. NH2-
27. H3O+
28. OH-
When you have completed your structures, circle the number of any structure that contains a
dative (coordinate covalent) bond.
65
Bond Length and Strength
Draw Lewis dot diagrams for the following compounds:
1a. C2H6
b. C2H4
c. C2H2
2. Identify the number of bonds between the two carbon atoms in each of the structures above.
a.
b.
c.
3. Which compound has the strongest C-C bond?
4. Which compound has the longest C-C bond?
5. How are bond lengths and bond strength related?
6. Given the following compounds: CO2, CO and CH3OCH3. (dot diagrams are essential!)
a. Which has the strongest C-O bond?
b. Which has the longest C-O bond?
7. Given N2 and N2H4.
a. Which has the strongest N-N bond?
b. Which has the longest N-N bond?
66
Electronegativity Related to Bonding
1. Define electronegativity.
2. What is the trend in electronegativity as you move across the periodic table from left to right?
3. What is the trend in electronegativity as you move down the periodic table from top to bottom?
4. From the following pairs, choose the atom with the highest electronegativity.
a. B or C
b. Se or Te
c. Ca or Ba
5. If the difference in electronegativity between two atoms in a bond is greater than 1.7, the bond is
said to be more than 50% ionic/covalent in character.
6. If the difference in electronegativity between two atoms in a bond is less than 1.7, the bond is said
to be more than 50% ionic/covalent in character.
7. Is there any bond that is 100% ionic? (Complete transfer of electrons) If so, give an example.
8. Is there any bond that is 100% covalent? (evenly shared electrons) If so, give an example.
9. Define "polar covalent bond."
10. State whether or not the following bonds are ionic, polar covalent or nonpolar covalent:
a. C -- F
b. O -- O
c. N -- O
d. Na -- F
11. Draw electron dot diagrams for the following molecular compounds (all covalent bonds)
a. H2O
b. O2
c. NF3
d CO2
e. C2H2
12. In #11, how many bonds are in each compound? Are they polar, nonpolar or a mixture of both?
Number of Bonds
polarity
a
b
c
d
e
67
Bonding and Molecular Geometry
Lewis dot diagrams show how valence electrons are involved in bonding. They do not, however,
show the geometric arrangements of the atoms in a molecule. We will use VSEPR theory for this
purpose. To understand the geometry, we must be able to draw the Lewis diagram and look at the
electrons around the central atom. Once we know the geometry, we are better able to predict
properties of molecules.
1. What does VSEPR stand for?
2. What is the difference between a bonded pair of electrons and a nonbonded (or unshared) pair of
electrons?
3. How many bonded and nonbonded pairs of electrons are on the central atom in the following
molecules? (Note: you need to draw the electron dot diagram to be able to answer this)
a. H2O
b. CO2
c. NH3
d. SO2
4. Why do electron pairs repel one another?
5. If two electron pairs are surrounding a central atom, what is the maximum angle of separation
they can achieve?
6. If three electron pairs are surrounding a central atom, what is the maximum angle of separation
they can achieve?
7. If four electron pairs are surrounding a central atom, what is the maximum angle of separation
they can achieve?
8. Which takes up more room, a bonded pair or nonbonded (unshared) pair of electrons?
9. Fill in the following table:
Number of electron
pairs on central
atom
Angle of separation
between pairs
2
3
4
68
Name of
geometrical shape
made by the
electron pairs
Molecular Geometry and Polarity
1. How does a polar bond differ from a nonpolar bond?
2. How does a polar bond differ from an ionic bond?
3. How is electronegativity difference used to predict bond type? What value separates ionic from
polar covalent bonds?
4. What is a dipole (polar molecule)?
5. What two criteria must be met for a molecule to be polar?
6. How can a molecule such as CO2 contain polar bonds yet still be a nonpolar molecule?
7. Fill in the following table:
Number of
Number of bonds
Number of
bond angle
molecular
bonds and
(remember
unshared (lone,
shape
unshared pairs
multiple bonds
unbonded) pairs
on central atom count as one bond)
2
2
0
3
3
0
3
2
1
4
4
0
4
3
1
4
2
2
8. Of the shapes you listed in the table above, which can be symmetrical if all bonds are alike?
9. Complete the following table:
molecule electron dot diagram
number of
bonds and
unshared
pairs
geometric
shape
BF3
NBr3
SO2
CI4
69
polar
bonds?
polarity of
molecule
Molecular Geometry/Polarity Review
Fill in the following table.
Molecule
Electron Dot
Diagram
Molecular
Geometry
H2S
PF3
SiCl4
CO2
SO3
70
Bond
Angle
Molecular
Polarity
Allotropes of Carbon
1. Define allotrope:
2. List the names of the three common allotropes of carbon:
Navigate to the following website: (The links are posted on Blackboard 24/7)
http://www.edinformatics.com/interactive_molecules/diamond.htm
The website poses the question, “Why is graphite soft and diamond hard if both are made of pure carbon?”
Look at the macromolecule of graphite. Hold down the left mouse button over the structure and move it
around. Study the structure of graphite. If you have a scrolling mouse, scroll in and out over the graphic.
Enlarge the graphic enough so that the individual atoms are clearly visible.
 Double click on the center of one atom and move the mouse to the center of the next atom
 Click once on the second atom.
 Move the mouse to a third atom and double click.
 A bond angle should be displayed.
Answer the following questions about the structure of graphite:
3. Each carbon atom is bonded to how many others?
4. What is the bond angle between all the C-C-C bonds?
5. What holds carbon atoms in one layer together?
6. What holds the layers to other layers?
7. What are the physical properties of graphite?
Now look at the macromolecule of diamond which is further down the web page. Follow the instructions
above to determine the bond angle.
Answer the following questions about diamond:
8. Each carbon atom is bonded to how many others?
9. What is the bond angle between all the C-C-C bonds?
10. What are the physical properties of diamond?
11. Why is graphite soft and slippery while diamond is so hard when both are pure carbon?
Now study the structure of fullerene at the following website:
http://www.edinformatics.com/interactive_molecules/fullerene.htm
12. Describe the bonding in fullerene, C60.
Silicon is in the same family as Carbon. Study the information at the website below:
http://web1.caryacademy.org/chemistry/rushin/studentprojects/elementwebsites/silicon/Structure.htm
13. What is the structure of pure silicon crystals?
14. What are the properties of pure silicon?
15. Research the structure and properties of the common compound, SiO2. Are the properties and structure
more like graphite or diamond?
71
Intermolecular Forces
1.
Covalent bonds hold atoms/molecules to one another.
Intermolecular forces (weak forces) hold atoms/molecules to one another.
2.
Intermolecular forces are found in what type of solid?
3.
List the three intermolecular forces in increasing order of strength.
a.
b.
c.
4.
Describe hydrogen bonding.
5.
A molecule must contain a hydrogen atom covalently bonded to what other element(s) in
order for a hydrogen bond to form between two molecules?
6.
What are the properties of the elements listed in number 5 that give rise to hydrogen bonding?
7.
Label the covalent bonds and the hydrogen bonds on the following diagram of HF molecules.
H
F
---- H
F
--- H
F
---- H
F
8.
What properties are associated with molecular compounds that contain hydrogen bonds versus
those that do not?
9.
Describe dipole-dipole attractions. (permanent dipoles)
10.
What types of molecules exhibit this type of attraction?
11.
What is the only intermolecular force that attracts nonpolar molecules to one another in the
solid phase?
12.
Describe van der Waals forces (London dispersion forces, or temporary dipoles).
13.
What factors determine the magnitude of van der Waals forces?
72
14.
What properties are associated with molecules attracted solely by van der Waals forces?
15.
For each substance listed below, state the type of intermolecular force that will hold the
substance together as a solid.
a.
H2O
b.
HCl
c.
C2H6
d.
CH3OH
e.
CO2
f.
H2S
g.
HF
16.
For each pair listed below, choose the one that you would expect to have the highest melting
point and give an explanation for your choice.
a.
H2O or H2S
b.
C2H6 or C5H12
73
Types of Solids
1.
Complete the following table about the four types of solids.
Type of Solid
Properties
Type of Particles
Force
attracting
particles
metallic
ionic
covalent
network
molecular
(including
group 18)
2.
Describe metallic bonding. How is it different from ionic and covalent bonding?
3.
A solid has a high melting point, does not dissolve in water and does not conduct electricity.
What type of solid is it most likely to be?
4.
A solid has a moderate melting point, does not dissolve in water but conducts electricity.
What type of solid is it?
5.
A solid has a high melting point, dissolves in water and does not conduct electricity. What
type of solid is it?
6.
State which type of solid would be formed by each substance listed below.
a.
Fe
g.
graphite
b.
NaCl
h.
water
c.
carbon dioxide
i.
brass
d.
diamond
j.
potassium sulfate
e.
sodium
k.
sugar, C12H22O11
f.
copper (II) chloride
l.
helium
74
Unit 4 Notes
75
Bonds, Polarity and Solubility
Purpose: To relate solubilities of various combinations of substances to their molecular
polarities.
Pre-lab assignment: Fill in the following table for the substances used in this lab
Substance
bond type (ionic
molecular
polarity
or covalent)
I2
CuCl2
C2H5OH
(ethanol)
Mineral oil
H2O
Procedure
Part A.
1. Place three clean, dry test tubes in a test tube rack.
2. Put a small amount of water in the first test tube. (Two-finger depth is fine.) Place an
equal amount of ethanol in the second test tube, and an equal amount of mineral oil in
the third. (***Caution: Avoid breathing the vapors and do not let the liquids come in contact
with your skin!!)
3. To each test tube add 1-2 crystals of iodine. (Just the tiny little crystals!!) Stopper the
tubes and shake. Record your observations. (Write "soluble" if the crystals dissolved, or
"insoluble" if they did not.) Dispose of the liquids in the waste beaker.
4. Repeat steps 1-2.
5. This time add 1-2 crystals of copper (II) chloride to each test tube. Record your
observations as before.
Part B.
1. You will be using the three solvents that you used in Part A, in addition to hexane, a
nonpolar solvent. If two liquids dissolve in one another in all proportions they are said to
be miscible. If the two liquids do not dissolve you will see two distinct layers and they
are said to be immiscible, and the least dense liquid will float on top of the more dense
liquid. For each combination of liquids (see data table B for combinations) pour about 1
ml of each into a clean test tube. Record your observations. (Write either miscible or
immiscible.)
76
Data Table Part A
Solvent
water
I2
CuCl2
ethanol
mineral oil
Data Table Part B
Liquid
water
ethanol
mineral oil
ethanol
mineral oil
hexane
Questions
1. Why did iodine dissolve in mineral oil and not water?
2. Why did copper (II) chloride dissolve in water but not in mineral oil?
3. Did all substances dissolve in ethanol? For any that did not dissolve, describe the
bonding.
4. Alcohols (such as ethanol) are said to be intermediates. Why do you think this is so?
5. Fill in the following blanks:
a. Polar solvents tend to dissolve __________ solutes.
b. Nonpolar solvents tend to dissolve __________ solutes.
c. Polar liquids are miscible with __________ liquids.
d. Nonpolar liquids are miscible with __________ liquids.
e. Alcohols are miscible with ____________ liquids.
f. Ionic compounds only dissolve in _____________ solvents.
77
Evaporation and
Intermolecular Attractions
In this experiment, Temperature Probes are placed in various liquids. Evaporation occurs when the
probe is removed from the liquid’s container. This evaporation is an endothermic process that results
in a temperature decrease. The magnitude of a temperature decrease is, like viscosity and boiling
temperature, related to the strength of intermolecular forces of attraction. In this experiment, you will
study temperature changes caused by the evaporation of several liquids and relate the temperature
changes to the strength of intermolecular forces of attraction. You will use the results to predict, and
then measure, the temperature change for several other liquids.
You will encounter two types of organic compounds in this experiment—alkanes and alcohols. The
two alkanes are n-pentane, C5H12, and n-hexane, C6H14. In addition to carbon and hydrogen atoms,
alcohols also contain the -OH functional group. Methanol, CH3OH, and ethanol, C2H5OH, are two of
the alcohols that we will use in this experiment. You will examine the molecular structure of alkanes
and alcohols for the presence and relative strength of two intermolecular forces—hydrogen bonding
and dispersion forces.
Figure 1
MATERIALS
Power Macintosh or Windows PC
Vernier computer interface
Logger Pro
two Temperature Probes
6 pieces of filter paper (2.5 cm X 2.5 cm)
2 small rubber bands (orthodontic bands)
masking tape
methanol (methyl alcohol)
ethanol (ethyl alcohol)
1-propanol
1-butanol
n-pentane
n-hexane
78
PROCEDURE
1. Obtain and wear goggles!
2. Prepare the computer for data collection. Make sure that the vertical axis has temperature is
scaled from -10 to 30C and the horizontal axis is scaled from 0 to 250 seconds.
3. Wrap Probe 1 and Probe 2 with square pieces of filter paper secured by small rubber bands as
shown in Figure 1 (rubber bands from the orthodontist work best). Roll the filter paper around the
probe tip in the shape of a cylinder. The paper should be even with the probe end.
4. Stand Probe 1 in the 1-propanol container and Probe 2 in the 1-butanol container. Make sure the
containers do not tip over.
5. After the probes have been in the liquids for at least 45 seconds, begin data collection by clicking
Collect . Monitor the temperature for 15 seconds to establish the initial temperature of each liquid.
Then simultaneously remove the probes from the liquids and hold them so the probe tips extend 5
cm over the edge of the table top as shown in Figure 1.
6. When both temperatures have reached minimums and have begun to increase, click Stop to end
data collection. Click the Statistics button, , then click OK to display a box for both probes.
Record the maximum (t1) and minimum (t2) values for 1-propanol and 1-butanol.
7. For each liquid, subtract the minimum temperature from the maximum temperature to determine
t, the temperature change during evaporation.
8. Clear data by selecting “Data” and “Clear All Data”.
9. Repeat Steps 2-8 using methanol (Probe 1) and ethanol (Probe 2).
10.
Repeat Steps 2-8 using pentane (Probe 1) and hexane (Probe 2).
10. Clean up.
79
PRE-LAB EXERCISE
Prior to doing the experiment, complete the Pre-Lab table. The name and formula are given for each
compound. Draw a structural formula for a molecule of each compound. Then determine the
molecular mass of each of the molecules. Dispersion forces exist between any two molecules, and
generally increase as the molecular mass of the molecule increases. Next, examine each molecule for
the presence of hydrogen bonding. Before hydrogen bonding can occur, a hydrogen atom must be
bonded directly to an N, O, or F atom within the molecule. Tell whether or not each molecule has
hydrogen-bonding capability.
Substance
Formula
ethanol
C2H5OH
1-propanol
C3H7OH
1-butanol
C4H9OH
n-pentane
C5H12
methanol
CH3OH
n-hexane
C6H14
Structural Formulas
80
Molecular Mass
Intermolecular
Force
DATA TABLE
Substance
t1
(°C)
t2
(°C)
t (t1–t2)
(°C)
ethanol
1-propanol
1-butanol
n-pentane
methanol
n-hexane
Data Analysis and Questions
We would like to determine the relationship between the strength of the intermolecular forces and the
molar mass. From the data you collected, construct a graph that shows this relationship.
81
EXPLAIN (Analysis Questions):
1. Two of the liquids, n-pentane and 1-butanol, had nearly the same molecular weights, but
significantly different t values. Explain the difference in t values of these substances, based
on their intermolecular forces.
2. Which of the alcohols studied has the strongest intermolecular forces of attraction? The
weakest intermolecular forces? Explain using the results of this experiment.
3. Which of the alkanes studied has the stronger intermolecular forces of attraction? The weaker
intermolecular forces? Explain using the results of this experiment.
4. Compare the molar masses of butane (C4H10) and pentane. Can you make a prediction of Δt
based on these masses? Which of the two will have stronger intermolecular forces? Explain.
5. How is the length of the carbon chain related to the molar mass? How does the length of the
carbon chain relate to the strength of the intermolecular forces?
82
EVALUATE
1. Which of the following compounds has the strongest intermolecular forces? (Circle your
answer and give an explanation as to why you chose that molecule.
Ethylene Glycol
isobutane
OH OH


H C  C  H
 
H H
propane
H CH3 H



H C  C  C  H
 

H H
H
H H
H

 
HC  C  C  H

 
H
H H
2. Which of the following compounds would have the strongest intermolecular forces? Circle
your answer and explain your choice.
Isobutane
(See above for structure)
Benzene
(C6H6 carbon ring)
Propane
(See above for structure)
3. Without using any resources, rank the following chemicals in order of increasing boiling
point, 7 is the substances with the highest boiling point. Next to the rank, indicate what
intermolecular forces are present in each substance.
Compound
Boiling Point Rank
Intermolecular Forces Present
NaCl
CO2
H2
CH3OH
C6H12O6
H2O
CH4
4. Based on polarity, which of the following substance would you expect to dissolve in water?
CH4
N2
NH3
CaCl2
83
H2S
CO2
Study Guide: Topic 4 Bonding
1.
Which statement is true for most ionic compounds?
A.
They contain elements of similar electronegativity.
B.
They conduct electricity in the solid state.
C.
They are coloured.
D.
They have high melting and boiling points.
2.
Which fluoride is the most ionic?
A.
NaF
B.
CsF
C.
MgF2
D.
BaF2
3.
Which statement is a correct description of electron loss in this reaction?
2Al + 3S  Al2S3
A.
Each aluminium atom loses two electrons.
B.
Each aluminium atom loses three electrons.
C.
Each sulfur atom loses two electrons.
D.
Each sulfur atom loses three electrons.
(Total 1 mark)
(Total 1 mark)
(Total 1 mark)
4.
Element X is in group 2, and element Y in group 7, of the periodic table. Which ions will be
present in the compound formed when X and Y react together?
A.
X+ and Y–
B.
X 2+ and Y–
C.
X+ and Y2–
D.
X2– and Y+
(Total 1 mark)
5.
What is the formula for the compound formed by calcium and nitrogen?
A.
CaN
B.
Ca2N
C.
Ca2N3
D.
Ca3N2
(Total 1 mark)
6.
When the species BF2+, BF3 and BF4– are arranged in order of increasing F−B−F bond angle,
what is the correct order?
A.
BF3, BF4–, BF2+
B.
BF4–, BF3, BF2+
C.
BF2+, BF4–, BF3
D.
BF2+, BF3, BF4–
7.
Which statement is true for compounds containing only covalent bonds?
A.
They are held together by electrostatic forces of attraction between oppositely charged ions.
B.
They are made up of metal elements only.
C.
They are made up of a metal from the far left of the periodic table and a non-metal from the
far right of the periodic table.
D.
They are made up of non-metal elements only.
(Total 1 mark)
(Total 1 mark)
84
8.
Which molecule is non-polar?
A.
H2CO
B.
SO3
C.
NF3
D.
CHCl3
9.
How many electrons are used in the carbon-carbon bond in C2H2?
A.
4
B.
6
C.
10
D.
12
(Total 1 mark)
(Total 1 mark)
10.
What is the valence shell electron pair repulsion (VSEPR) theory used to predict?
A.
The energy levels in an atom
B.
The shapes of molecules and ions
C.
The electronegativities of elements
D.
The type of bonding in compounds
(Total 1 mark)
11.
According to VSEPR theory, repulsion between electron pairs in a valence shell decreases in
the order
A.
lone pair-lone pair > lone pair-bond pair > bond pair-bond pair.
B.
bond pair-bond pair > lone pair-bond pair > lone pair-lone pair.
C.
lone pair-lone pair > bond pair-bond pair > bond pair-lone pair.
D.
bond pair-bond pair > lone pair-lone pair > lone pair-bond pair.
(Total 1 mark)
12.
When C2H4, C2H2 and C2H6 are arranged in order of increasing C–C bond length, what is the
correct order?
A.
C2H6, C2H2, C2H4
B.
C2H4, C2H2, C2H6
C.
C2H2, C2H4, C2H6
D.
C2H4, C2H6, C2H2
(Total 1 mark)
13.
In the molecules N2H4, N2H2, and N2, the nitrogen atoms are linked by single, double and triple
bonds, respectively. When these molecules are arranged in increasing order of the lengths of their
nitrogen to nitrogen bonds (shortest bond first) which order is correct?
A.
N2H4, N2, N2H2
B.
N2H4, N2H2, N2
C.
N2H2, N2, N2H4
D.
N2, N2H2, N2H4
(Total 1 mark)
14.
Which species has a trigonal planar shape?
A.
CO32–
B.
SO32–
C.
NF3
D.
PCl3
(Total 1 mark)
85
15.
In which substance is hydrogen bonding present?
A.
CH4
B.
CH2F2
C.
CH3CHO
D.
CH3OH
16.
Which compound has the highest boiling point?
A.
CH3CH2CH3
B.
CH3CH2OH
C.
CH3OCH3
D.
CH3CHO
(Total 1 mark)
(Total 1 mark)
17.
Which substance is most similar in shape to NH3?
A.
GaI3
B.
BF3
C.
FeCl3
D.
PBr3
(Total 1 mark)
18.
What are responsible for the high electrical conductivity of metals?
A.
Delocalized positive ions
B.
Delocalized valence electrons
C.
Delocalized atoms
D.
Delocalized negative ions
(Total 1 mark)
19.
What intermolecular forces are present in gaseous hydrogen?
A.
Hydrogen bonds
B.
Covalent bonds
C.
Dipole-dipole attractions
D.
Van der Waals’ forces
(Total 1 mark)
20.
Which statement best describes the attraction present in metallic bonding?
A.
the attraction between nuclei and electrons
B.
the attraction between positive ions and electrons
C.
the attraction between positive ions and negative ions
D.
the attraction between protons and electrons
(Total 1 mark)
21.
Which substance has the lowest electrical conductivity?
A.
Cu(s)
B.
Hg(l)
C.
H2(g)
D.
LiOH(aq)
(Total 1 mark)
22.
Which substance is most soluble in water (in mol dm–3) at 298 K?
A.
CH3CH3
B.
CH3OCH3
C.
CH3CH2OH
D.
CH3CH2CH2CH2OH
(Total 1 mark)
86
23.
Arrange the following in decreasing order of bond angle (largest one first), and explain your
reasoning.
NH2–, NH3, NH4+
(Total 5 marks)
24.
(a)
An important compound of nitrogen is ammonia, NH3. The chemistry of ammonia is
influenced by its polarity and its ability to form hydrogen bonds. Polarity can be explained
in terms of electronegativity.
(i)
Explain the term electronegativity.
……………………………………………………………………………………
……………………………………………………………………………………
(2)
(ii)
Draw a diagram to show hydrogen bonding between two molecules of NH3.
The diagram should include any dipoles and/or lone pairs of electrons
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
(3)
(iii)
State the H–N–H bond angle in an ammonia molecule.
………………………………………………………………………………………
(1)
(iv)
Explain why the ammonia molecule is polar.
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
(1)
87
(b)
Ammonia reacts with hydrogen ions forming ammonium ions, NH4+.
(i)
State the H–N–H bond angle in an ammonium ion.
……………………………………………………………………………………
(1)
(ii)
Explain why the H–N–H bond angle of NH3 is different from the H–N–H bond angle
of NH4+; referring to both species in your answer.
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
(3)
(Total 11 marks)
25.
(i)
Use the VSEPR theory to predict and explain the shape and the bond angle of each of the
molecules SCl2 and C2Cl2
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(6)
88
(ii)
Deduce whether or not each molecule is polar, giving a reason for your answer.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(3)
(Total 9 marks)
26.
Three scientists shared the Chemistry Nobel Prize in 1996 for the discovery of fullerenes.
Fullerenes, like diamond and graphite, are allotropes of the element carbon.
(i)
State the structures of and the bonding in diamond and graphite.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(2)
(ii)
Compare and explain the hardness and electrical conductivity of diamond and graphite.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(4)
(iii) Predict and explain how the hardness and electrical conductivity of C60 fullerene would
compare with that of diamond and graphite.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(4)
(Total 10 marks)
89
27.
(i)
List the following substances in order of increasing boiling point (lowest first).
CH3CHO
C2H6
CH3COOH
C2H5OH
……………………………………………………………………………………………
(2)
(ii)
State whether each compound is polar or non-polar, and explain the order of boiling points
in (i).
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(8)
(Total 10 marks)
28.
State two physical properties associated with metals and explain them at the atomic level.
(Total 4 marks)
90
Topic 1: Quantitative Chemistry
1.1 The Mole Concept and Avogadro’s Constant
1.1.1 Apply the mole concept to substances.
1.1.2 Determine the number of particles and the amount of substance (in moles).
1.2 Formulas
1.2.1 Define the terms relative atomic mass (Ar) and relative molecular mass (Mr).
1.2.2 Calculate the mass of one mole of a species from its formula.
1.2.3 Solve problems involving the relationship between the amount of substance in moles,
mass and molar mass.
1.2.4 Distinguish between the terms empirical formula and molecular formula.
1.2.5 Determine the empirical formula from the percentage composition or from other
experimental data.
1.2.6 Determine the molecular formula when given both the empirical formula and
experimental data.
1.3 Chemical Equations
1.3.1 Deduce chemical equations when all reactants and products are given.
1.3.2 Identify the mole ratio of any two species in a chemical equation.
1.3.3 Apply the state symbols (s), (l), (g) and (aq).
1.4 Mass and Gaseous Volume Relationships in Chemical Reactions
1.4.1 Calculate theoretical yields from chemical equations.
1.4.2 Determine the limiting reactant and the reactant in excess when quantities of reacting
substances are given.
1.4.3 Solve problems involving theoretical, experimental and percentage yield.
1.4.4 Apply Avogadro’s law to calculate reacting volumes of gases.
1.4.5 Apply the concept of molar volume at standard temperature and pressure in calculations.
1.4.6 Solve problems involving the relationship between temperature, pressure and volume for a
fixed mass of an ideal gas.
1.4.7 Solve problems using the ideal gas equation, PV=nRT.
1.4.8 Analyze graphs relating to the ideal gas equation.
1.5 Solutions
1.5.1 Distinguish between the terms solute, solvent, solution and concentration (g dm-3 and
mole dm-3).
1.5.2 Solve problems involving concentration, amount of solute and volume of solution.
91
Topic 1 NOTES
92
Introduction to the Mole
1 dozen indicates 12 of something. For instance 1 dozen eggs = 12 eggs.
Similarly 1 dozen atoms = 12 atoms
Unfortunately, atoms are so small that we need large numbers of them before we can start
experimenting with them. Chemists use a unit known as the “mole”. Like a dozen, it simply
indicates a specific number of items. Instead of 12 (like the dozen), 1 mole is equal to …..
602 000 000 000 000 000 000 000
or
6.02 x 1023
This number is extremely important in chemistry and is known as Avogadro’s Number or NA.
Complete the following examples:
1.75 dozen atoms = _______ atoms
1.75 moles of atoms = ________ atoms
Complete the following problems, showing your calculations.
1. 0.250 moles of atoms = ______ atoms
2. 4.5 x 1025 atoms = _______ moles of atoms
3. 5.67 x 1019 atoms = _______ moles of atoms
4. 0.0056 moles of molecules = _______ molecules
5. 3.5 moles of H2O molecules = _______ molecules
6. 3.5 moles of H2O molecules = _______ atoms
93
Percentage Composition
The law of definite proportions states the ratio or proportion by mass of the elements in a compound
is always the same. This means that the ratio of hydrogen to oxygen in a sample of water is the same
no matter what the sample size. Percentage composition by mass is an easy way to apply this
principle.
 Example:
H2O
Element present Atomic mass Number of atoms Total mass
H
1.01

2
= 2.02
O
16.0

1
= 16.0
18.0 amu total mass
Percent Hydrogen in water
(part/whole)  100
(2.02 / 18.0 )  100 = 11.2 %
Percent Oxygen in water
(part/whole)  100
(16.0 / 18.0 )  100 = 88.8 %
Therefore any sample of water is composed of 11.2 % hydrogen by mass. A 100 g sample contains
11.2 g of hydrogen. A 500 g sample of water contains (500 g )  (11.2 %) = 56 g of hydrogen.
Determine the percentage composition of each of the following compounds below.
1. KMnO4
2. HCl
3. Mg(NO3)2
4. (NH4)3PO4
5. Al2(SO4)3
Using your answers from the problems above, answer the following questions:
6. How many grams of oxygen are contained in a 75 g sample of KMnO4?
7. How much magnesium is contained in a 125 g sample of Mg(NO3)2?
8. What mass of aluminum is found in a 25 g sample of Al2(SO4)3?
94
Mass and the Mole
Fill in the following table.
Substance
Symbol
Molar Mass (Ar or Mr)
g·mol-1
Aluminum
Sn
Oxygen
Iron (II)chloride
MgSO4
Ammonium phosphate
1 mole of atoms equals 6.02 x 1023 atoms
1 mole of molecules equals 6.02 x 1023 molecules
Avogadro’s number expresses the number of particles but not the mass. How do we find the mass of
a mole? The answer is molar mass. Use the relative atomic masses, Ar, of the elements in a
compound to find the relative molecular mass of a molecule, Mr.
Answer the following questions about atoms, moles and mass.
1.
What is the mass of 1.65 moles of gold atoms?
2.
How many atoms are found in 2.25 moles of copper?
3.
1.75 moles of Na2O has what mass?
4.
How many moles are contained in 175 g of Cu(NO3)2?
5.
How many atoms are contained in 2.45 grams of helium?
6.
What is the mass of 1.25 × 1022 molecules of CO2?
7.
How many molecules are contained in 25.0 g of H2O?
8.
How many atoms of hydrogen are contained in 10.0 g of H2O?
95
Empirical and Molecular Formulas
An empirical formula is the simplest ratio of atoms in a compound. A molecular formula is the actual
number of atoms in a compound. The molecular formula is always a whole number multiple of the
empirical formula.
1. What is the empirical formula of the following compounds?
a. H2O2
b. C8H18
c. H2SO4
d. N2O4
2. A compound contains 64.3 g of carbon and 10.7 g of hydrogen. What is its empirical formula?
3. A compound has the empirical formula of CH. If its molar mass is 104 g mol-1, what is its
molecular formula?
4. A compound is 22.1% aluminum, 25.4% phosphorus, and 52.5% oxygen by mass. Determine its
empirical formula.
5. A compound is 69.9% iron and 30.1% oxygen by mass. Determine its empirical formula.
6. A compound is 54.5% carbon, 9.1% hydrogen, and 36.4% oxygen by mass. Its molar mass is 88
g mol-1. What is its empirical and molecular formula?
7. A compound is 43.7% phosphorus and 56.3% oxygen by mass. Its molar mass is 284 g mol-1.
Determine its empirical and molecular formula.
8. A hydrated ionic compound is 25.3% copper, 12.9% sulfur, 25.7% oxygen, and 36.1% water.
Determine the empirical formula of this hydrate.
96
Percentage Water in a Hydrated Salt
Name ______________
Purpose
To determine the percent water in a hydrated salt. To determine the empirical formula of the hydrate.
Background
When water molecules are chemically attached to a compound it is called a hydrate. Water can be
separated from the compound by heating it. As water is added back to the compound, the water
molecules can reattach.
Equipment
Evaporating dish
Spatula
Hot plate
Crucible tongs
Balance
Goggles and apron
Materials
Copper (II) sulfate hydrate, CuSO4  xH2O
Procedure
1.
2.
3.
4.
5.
6.
7.
Find the mass of a clean, dry evaporating dish. Record this mass.
With the dish still on the balance, press the TARE button to rezero the balance. Add the
copper (II) sulfate hydrate until you have exactly 2.00 g.
Place the evaporating dish on the hot plate and turn it to a low heat at first. Gradually
increase the heat as the hydrate begins to change color. Avoid any popping and
spattering. (Immediately remove the evaporating dish if it starts to pop with the crucible
tongs and turn down the heat before returning the dish to the hot plate.)
After the blue color has disappeared, leave the dish on medium-high heat for an additional
three minutes to assure that all of the water has been removed. You may use a spatula to
break up any lumps or “caked” portions of the hydrate. If the edges of the solid appear to
be turning brown, remove the dish from the heat momentarily and resume heating at a
gentler rate.
Allow the dish to cool on an insulating square for about a minute. Immediately find the
mass of the dish + anhydrous salt, and record the data below.
If you would like, you can attempt to reattach the water by adding some drops of water
back to the anhydrous salt.
Dispose of the copper sulfate in the appropriate container and wash and dry the
evaporating dish when you are finished.
Data
a.
b.
c.
Mass of empty evaporating dish
Mass of copper (II) sulfate hydrate
Mass of evaporating dish + anhydrous salt
anhydrous means without water
97
________ ±______ g
________  _____ g
________ ± _____ g
Calculations-You must show all of your work. (Remember uncertainty!)
1. Find the mass of the anhydrous salt in grams.
2. Find the mass of the water lost in grams.
3. Find the percentage of water in the hydrate:
Questions
1.
The true value for the percent of water in this hydrate is 36.0%. What is your
experimental error? Show your work.
2.
Why must you measure the mass of the anhydrous salt immediately upon cooling?
3.
Calculate the moles of water. Show your work.
4.
Calculate the number of moles of anhydrous salt (CuSO4). Show your work.
5.
How many times more moles of water are there than anhydrous salt? (Look at your
answers to questions 3 and 4.) Divide your moles of water by your moles of CuSO4 to get
a ratio. Based on this answer, what is the true formula for copper (II) sulfate hydrate?
(CuSO4  xH2O, where x is the number of moles of water for every one mole of CuSO4)
6.
If you didn’t heat your hydrate long enough, all of the water may not have been removed.
If this happened, would your moles of water (x, in question #5 above) be too high or too
low. Explain your answer.
98
Balancing Chemical Equations
1. ___ Cl2 + ___ NaBr  ___ NaCl + ___ Br2
2. ___ Ca(OH)2 + ___ HNO3  ___ Ca(NO3)2 + ___ HOH
3. ___ C2H4 + ___ O2  ___ CO2 + ___ H2O
4. ___ Fe(OH)3  ___ Fe2O3 + ___ HOH
5. ___ P2O5 + ___ H2O  ___ H3PO4
6. ___ Al(NO3)3 + ___ NaOH  ___ Al(OH)3 + ___ NaNO3
7. ___ KClO3  ___ KCl + ___ O2
8. ___ C3H8 + ___ O2  ___ CO2 + ___ H2O
9. ___ H3PO4 + ___ Mg(OH)2  ___ Mg3(PO4)2 + ___ HOH
10. ___ NH3 + ___ O2  ___ NO + ___ H2O
11. ___ Na2SO4 + ___ Ba(NO3)2  ___ BaSO4 + ___ NaNO3
12. ___ CaO + ___ P2O5  ___ Ca3(PO4)2
13. ___ Al + ___ CuCl2  ___ AlCl3 + ___Cu
14. ___ Ca(OH)2 + ___ H3PO4  ___ Ca3(PO4)2 + ___ HOH
15. ___ NaHCO3  ___ Na2CO3 + ___ H2O + ___ CO2
16. ___ C2H5OH + ___ O2  ___ CO2 + ___ H2O
17. ___ Fe + ___ HCl  ___ FeCl3 + ___ H2
18. ___ Co(OH)3 + ___ HNO3  ___ Co(NO3)3 + ___ HOH
19. ___ Mg + ___ O2  ___ MgO
20. ___ Na + ___ Sn(NO3)2  ___ Sn + ___ NaNO3
99
Writing and Balancing Equations
Write a balanced equation for each reaction. Classify as:
Synthesis (S) – the combination of two or more reactants into one product
Decomposition (D) – one reactant forming two or more simpler products
Single Displacement (SD) – One element takes the place of another in a compound
Double Displacement (DD) – Two compounds exchange their ions
Combustion (C) – A compound reacts with oxygen to form oxygen-containing compounds
____ 1. sodium bromide + fluorine  sodium fluoride + bromine
____ 2. calcium nitrate + copper (II) sulfate  calcium sulfate + copper (II) nitrate
____ 3. potassium chlorate  potassium chloride + oxygen gas
____ 4. propane + oxygen  carbon dioxide + water
____ 5. calcium bromide + chromium (III) nitrate  calcium nitrate + chromium (III) bromide
____ 6. iron + copper (I) nitrate  iron (II) nitrate + copper
____ 7. nitrogen + hydrogen  ammonia (NH3)
____ 8. hydrogen sulfate  water + oxygen + sulfur dioxide
____ 9. zinc sulfide + oxygen  zinc oxide + sulfur dioxide
____ 10. ammonium phosphate + lithium hydroxide  ammonium hydroxide + lithium phosphate
100
Lab: Types of Chemical Reactions
Pre-Lab Discussion
There are many kinds of chemical reactions and several ways classify them. One useful method
classifies reactions into four major types. These are:
1). Direct combination, or synthesis
2). Decomposition, or analysis
3). Single replacement, or single displacement
4). Double replacement (displacement) or exchange of ions.
Not all reactions can be put into one of these categories. Many, however, can.
In a synthesis reaction, two or more substances (elements, or compounds) combine to form a more
complex substance. Equations for synthesis reactions have the general form A + B  AB. For
example, the formation of water from hydrogen and oxygenic written, 2H2 + O2  2H2O.
A decomposition reaction is the opposite of a synthesis reaction. In decomposition, a compound
breaks down into two or more simpler substances (elements or compounds). Equations for
decomposition reactions have the form of AB  A + B. The breakdown of water into its elements is
an example of such a reaction: 2H20  2H2 + O2.
In a single replacement reaction, one element in a compound is replaced by another, more active,
element. Equations for single replacement reactions have tow general forms. In reactions in which
one metal replaces another metal, the general equation is M1 + M2X ---> M1X + M2. In those cases
where a nonmetal in a compound is replaced by another nonmetal the general equation is MX + Y 
MY + X. The following equations illustrate these types of reactions:
Zinc metal replaces copper(II) ion:
Zn(cr) + CuSO(aq)  ZnSO4(aq) + Cu(cr)
Chlorine (a nonmetal) replaces bromide ions:
KBr(aq) + Cl2(g)  KCl(aq) + Br2(l)
In a double replacement reaction, the metal ions of two different ionic compounds can be thought of
a “replacing one another.” Equations for the type of reaction have the general form AB + CD 
AD + CB. Most replacement reactions, both single and double, take place in aqueous solutions
which contain free ions. In a double replacement reaction, one of the products must be either a
precipitate, and insoluble gas, or additional water molecules. An example is the reactions between
silver nitrate and sodium chloride in which the precipitate silver chloride is formed:
AgNO3(aq) + NaCl(aq)  AgCl(cr) + NaNO(aq).
The production of a gas is an indication of chemical reaction. A burning splint placed at the mouth of
test tube in which hydrogen is being generated will cause a high pitched “pop”. If a burning splint is
extinguished when it is placed in a test tube, it indicates the production of carbon dioxide.
101
Purpose
In this lab you will observe examples of chemical reactions and will classify each of them into one of
the four types of reactions described above. In reactions in which you use a burning splint to test for
a gas, the gas produced will be identified.
Equipment
goggles and apron
burner and striker
crucible tongs
spatula
wood splints
test tubes
test tube holder
fine sandpaper
evaporating dish
Materials
zinc, mossy (Zn)
copper wire, 10 cm (Cu)
magnesium ribbon (Mg)
copper (II) carbonate (CuCO3)
6 M hydrochloric acid (HCl(aq))
1 M copper(II) sulfate (CuSO4)
0.1 M zinc acetate [Zn(C2H3O2)]
0.1 M sodium phosphate (Na3PO4)
1 M sodium carbonate (Na2CO3)
Safety





Wear safety goggles and apron at all times.
You will be working with open flames, heating chemicals, handling acids, and generating gaseous
products.
Do not handle equipment that may be hot, test it first with the back of your hand.
Burning magnesium produces a very bright, hot flame. Make sure you hold the burning metal
away from yourself and other students. Do not look directly at it as it burns!
Neutralize any acid spill with the sodium bicarbonate solution before you wipe it up.
Procedure
Reaction 1. Use fine sandpaper to clean a piece of copper wire until the wire is shiny. Note the
appearance of the wire in your data table. Using crucible tongs, hold the wire in the hottest part
of the burner flame for 1-2 minutes. Examine the wire and note any changes caused by the
heating.
Reaction 2. Place an evaporating dish near the base of the burner. Examine a piece of magnesium
ribbon. Using crucible tongs, hold the magnesium in the flame until it starts to burn. Do not
look directly at the flame. Hold the burning magnesium away from you and directly over the
evaporating dish. When the ribbon stops burning, put the remains in the evaporating dish.
Examine this product carefully.
Reaction 3. Place a heaping spatulaful of copper (II) carbonate into a large, clean dry test tube. Note
its appearance. Using a test tube holder, heat the CuCO3 strongly for about 3 minutes. Turn off
the burner and insert a burning splint into the top part of the test tube. Write your observations
of what happens, and the appearances of the test tube and the product formed.
102
Reaction 4. Stand a clean dry small test tube in the test tube rack. Add about 5 ml of 6 M
hydrochloric acid to the test tube. Note its appearance. CAUTION. Handle acids with care,
they can cause painful burns. Do not inhale any HCl fumes. Obtain and observe a small piece
of zinc metal. Carefully place it into the acid in the test tube. Observe and record what happens.
Using a test tube holder, invert a second test tube over the mouth of the test tube in which the
reaction is taking place. After about 30 seconds, remove the top tube, keeping it inverted, and
immediately insert a burning wood splint into its mouth. Record your observations of what
happens, and the appearances of the test tube’s interior.
Reaction 5. Obtain another small clean dry test tube and add about 5 ml of 1M copper (II) sulfate
solution to it. Place a small piece of zinc metal into the solution. Note the appearance of the
solution and the zinc before and after the reaction.
Reaction 6. To a clean dry test tube, add about 2 ml of 0.1M zinc acetate. Next add about 2 ml of
0.1M sodium phosphate solution to the test tube. Note observations of the solutions before they
were combined, and any changes once they were combined, in your data table.
Reaction 7. Into another small, clean, dry test tube add about 1 ml of 1M sodium carbonate solution.
To this solution (cautiously) add about 10 drops of 6 M HCl. After the reaction has stopped,
place a burning wood splint into the test tube, but not so far in that it touches the solution Note
any changes.
Carefully wash all equipment, being certain no solids go down the sink. Be sure your sink is clean!
Return all equipment it to its proper place. Carefully wash your hands before you leave the lab.
Post Lab
For all seven reactions, classify the reaction and then write a complete balanced equation with phase
notation.
103
Ions and Precipitates
1. Write dissociation reactions for the following aqueous solutions.
a. LiNO3
b. K2SO4
c. AlBr3
d. FeCl3
2. Predict products and write balanced molecular equations for the following double replacement
reactions. Include phase notation to show which product is a precipitate.
a. silver nitrate + magnesium chloride
b. barium nitrate + aluminum sulfate
c. sodium phosphate + calcium bromide
d. potassium hydroxide + iron (III) chloride
3.
Write ionic and net ionic equations for each of the reactions in number 2.
4. Silver nitrate reacts with sodium chloride to form a white precipitate. Complete the visual
representation of the reaction below. (Remember the law of conservation of matter!)
= Ag+
= NO3-
= Na+
= Cl-
104
Lab- Precipitates and Solubility Rules
Purpose: To observe the formation of precipitates, and to write balanced double displacement and
net ionic equations.
Procedure.
You are given a set of 5 solutions. React every combination of two solutions together in the multiwell plate. React only 1 drop of one solution with 1 drop of another solution. This gives ten
combinations. Fill in the following data table with “PPT” if a precipitate formed, or “NR” if no
reaction was observed.
Solution
AgNO3
KOH
CoCl2
Al2(SO4)3
FeCl3
1)
2)
4)
7)
AgNO3
3)
5)
8)
KOH
6)
9)
CoCl2
10)
Post Lab
1. For each reaction that produced a precipitate, PPT, identify the precipitate. Use page 920 in
your text to help. (Use the numbers in the data table to identify the reaction.)
Ex: 1)
AgCl(s)
2. For each reaction that produced a precipitate, write a balanced double displacement reaction
using phase notation, (s), for the precipitate.
Ex: 1)
3AgNO3 + FeCl3 → 3AgCl(s) + Fe(NO3)3
3. For each reation that produced a precipitate, write a net ionic equation.
Ex: 1)
Ag+ + Cl- → AgCl(s)
105
Mole-Map Worksheet #1
Mass (g)
A
molar mass
A
Moles
A
coefficient
ratio
6.02
x1023
Particles
A
Moles
B
molar mass
B
Mass (g)
B
6.02
x1023
coefficient
ratio
Particles
B
2H2(g) + O2(g)  2H2O(g)
1. If 5.0 moles of hydrogen gas are consumed in this reaction, how many moles of water are
formed?
2. If 4.0 moles of hydrogen gas are consumed in this reaction, how many moles of oxygen are
required?
3. How many molecules of water can be produced from the consumption of 2.0 x 1022 molecules of
hydrogen?
4. How many molecules of hydrogen are required to react completely with 4.5 x 1024 molecules of
oxygen
5. If 2.5 moles of hydrogen gas are consumed how many molecules of water are produced?
6. How many moles of oxygen gas are required to use up 5.5 x 1023 molecules of hydrogen gas?
7. How many moles of water can be produced from the consumption of 25.0 grams of oxygen?
8. How many grams of water can be produced from the consumption of 10.0 grams of oxygen?
9. How many grams of hydrogen are required to completely consume 5.0 grams of oxygen?
10. How many molecules of water can be produced by the consumption of 2.5 grams of oxygen?
106
Mole-Map Worksheet #2
1.
KClO3  KCl + O2
a. Balance the equation.
b. If 2.5 grams of potassium chlorate are decomposed, what theoretical mass of oxygen can be
formed?
c. If 2.5 x 1023 molecules of oxygen are formed, what theoretical mass of potassium chloride can
be produced?
d. When 45.0 grams of potassium chlorate is decomposed, how many molecules of oxygen are
formed?
2. Potassium bromide reacts with chlorine to form potassium chloride and bromine.
a. Write a balanced equation for the reaction.
b. If 25.0 grams of potassium bromide are reacted, what is the theoretical mass of bromine
formed?
c. If 15.2 g of bromine are actually produced, what is the percent yield?
3. Sodium phosphate reacts with barium nitrate to form two products.
a. Write a balanced equation for the reaction.
b. If 10.0 grams of sodium phosphate are reacted, what is the theoretical mass of mass of barium
phosphate that could be produced?
c. If the actual yield of barium phosphate is 17.8 g, what is the percent yield?
107
Limiting Reactants
Let’s say that you want to make a cheese sandwich. It takes 2 slices of bread and 1 slice of cheese to
accomplish this task. It can be written as an equation:
2 Br + Ch → Br2Ch
(2 bread + 1 cheese make 1 sandwich)
How many sandwiches can be made with the following starting ingredients?
a. 10 slices of bread and 8 slices of cheese?
b. 18 slices of bread and 7 slices of cheese?
In “a” and “b” above, which reactant was limiting?
a.
b.
In “a” and “b” above, which reactant was in excess and by how much?
a.
b.
Now let’s try it with a real chemical system:
2 H2 + O2 → 2 H2O
For each starting amount given below, state the limiting reactant and how many moles of water can
be formed.
Starting moles Starting moles
Moles of
of hydrogen
of oxygen
water formed
1.0
1.0
3.0
2.0
10.0
4.0
Now let’s add the relationship of mass to the equation. Remember coefficient ratios apply to moles,
not mass. Therefore all masses must be converted to moles to clearly see the relationships of
amounts of reactants present. Use the same equation above to answer the following questions:
2 H2 + O2 → 2 H2O
1. If 10.0 g of hydrogen and 10.0 g of oxygen are reacted, which reactant is limiting? What mass of
water can be produced?
2. If 1.50 g of hydrogen and 50.0 g of oxygen are reacted, which reactant is limiting? What mass of
water can be produced?
108
More Limiting Reactant Problems
1. 50.0 g of methane and 50.0 g of oxygen are combusted in a reaction vessel.
a. Write a balanced equation to represent this process.
b. Determine the moles of each reactant present.
c. Determine which reactant is the limiting reactant.
d. What mass of carbon dioxide can be formed?
2. The synthesis of ammonia is represented by the following equation:
3 H2(g) + N2(g) → 2 NH3(g)
If 15.0 g of hydrogen and 75.0 g of nitrogen are reacted, what is theoretical mass of ammonia
can be produced?
3. 125 g of iron and 125 g of oxygen are reacted according to the reaction shown below.
4 Fe(s) + 3 O2(g) → 2 Fe2O3(s)
What mass of iron (III) oxide will be formed?
For the reactant in excess, what mass will be left unreacted? (You must calculate the amount
consumed to determine this mass!!!)
109
An Introduction to the Gas Laws (internet reqd.)
Boyle’s Law
We are going to take a peek into properties of gases and the relationships between these properties.
For an introduction, navigate to the following site and scroll through the animation.
http://preparatorychemistry.com/Bishop_KMT_frames.htm
Once you have read through the introductory material in the animation, look at the menu list on the
left side of the page. (There is a scroll bar as well.) Scroll down to Chapter 13, and click on the
“Boyle’s Law animation”. As you read through the pages in the animation answer the following
questions:
1. What are the four inter-related properties of a gas?
a.
b.
c.
d.
2. What creates pressure in a gas?
3. In Boyle’s Law, what two properties are we changing?
a.
b.
4. In Boyle’s Law, what two properties are we keeping constant?
a.
b.
5. When the volume is cut in half, what happens to the pressure?
6. As the volume decreases, the pressure ___________. Is this a direct or inverse relationship?
Now that we have the concepts, let’s collect some data. Navigate to the following website:
http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/gaslaw/boyles_law_graph.html
You will see a device for measuring the pressure and volume of a gas. We will use the mouse to drag
the plunger to change the volume of the gas. Hold down the mouse and drag it to the desired volume.
When you release the mouse it will automatically record a data point. Collect a total of 6 data points
from a variety of different volumes. Start with a small volume and gradually increase the volume in
your experiment. Record your data points in the table below:
Volume (mL)
Pressure (psi)
110
Once you have a variety of points recorded, press the “Graph” button on the animation. Sketch this
graph on the axes below:
This graph shows
the relationship
P×V = constant
Pressure
Volume (mL)
Boyle’s Law Summary:
If temperature and number of particles of gas are kept constant, the pressure and volume of a
gas are (directly/inversely) related. This means if you increase the volume, the pressure will
_____________.
Let’s look at our next gas law: Charles’ Law
Navigate back to http://preparatorychemistry.com/Bishop_KMT_frames.htm
Scroll down to Chapter 13, Charles’ Law and click through the animation. Answer the following
questions as you go along.
1. What two properties are you keeping constant in Charles’ Law?
a.
b.
2. What two properties are you changing?
a.
b.
3. When you increase the temperature of a gas, what is happening to the velocity of the
particles?
4. To keep the pressure (number of collisions) constant, when the temperature increases what
must happen to the volume?
5. As temperature increases, volume _______________.
6. Is this a direct or inverse relationship?
7. Would a graph of temperature vs. volume be a curve like Boyle’s law or a straight line?
Sketch the graph below:
This graph represents the relationship
T
V ÷ T = constant
V
8. If the temperature drops to zero Kelvin, what happens to the volume of the gas?
111
9. What temperature is absolute zero in Kelvins? In ºC?
10. What happens to molecular motion at absolute zero?
Avogadro’s Principle
Navigate back to http://preparatorychemistry.com/Bishop_KMT_frames.htm
Scroll down to Chapter 13, “Volume/moles animation” and click through the animation. Answer the
following questions as you go along.
You recognize the name from the number of particles in a mole, but Avogadro also did a lot of
experimenting with gases.
1. What two properties of a gas are changing ?
a.
b.
2. What two properties must be kept constant?
a.
b.
3. As the number of moles of gas increases what happens to the volume?
4. Is this an inverse or direct relationship?
5. Would a graph of temperature vs. volume be a curve like Boyle’s law or a straight line?
Sketch the graph below:
This graph represents the relationship
n
V÷ n = constant
V
6. Equal volumes of gas at the same temperature and pressure contain ________ moles of gas.
7. Which statement is true about 1 L of hydrogen gas at the same temperature and pressure as 1
L of oxygen gas? Circle the best answer.
a. 1 L of hydrogen contains the most moles of gas.
b. 1 L of oxygen contains the most moles of gas.
c. Both gases contain the same number of moles of gas.
8. Which gas in number 7 above has the highest mass?
Summarizing Graphical Relationships
Navigate to the following website and play with the variables. Freeze two of the variables (click on
them in the upper left hand corner) and determine the relationship between the remaining two
variables.
http://www.grc.nasa.gov/WWW/K-12/airplane/Animation/frglab2.html
A straight line on a graph represents what type of relationship?
A descending curved line on a graph represents what type of relationship?
112
Dalton’s Law
Dalton’s law is a little different than the ones we learned above. Unlike Boyle’s, Charles’ and
Avogadro’s Laws, this one deals with a mixture of gases. Navigate to the website below and click
through the 8 pages and answer the practice questions.
http://www.wwnorton.com/college/chemistry/gilbert2/tutorials/interface.asp?chapter=chapter_
06&folder=daltons_law
1. What is the equation for Dalton’s Law?
2. In the space below show your work and answers for the practice questions on section 4 of the
web site.
Practice question 1:
Practice question 2:
Practice question 3:
Practice question 4:
3. If a mixture of oxygen and nitrogen has a total pressure of 2.50 atm and the partial pressure of
the oxygen is 1.73 atm, what is the partial pressure of the nitrogen?
4. A mixture of gases is 80% helium and 20% neon (by moles). If the total pressure is 760 mm
Hg, what is the partial pressure of the helium?
5. 2.00 moles of gas A, 3.00 moles of gas B and 5.00 moles of gas C have a total pressure of 800
mm Hg. What is the partial pressure of each individual gas?
Extra information about pressure and its units!
Pressure is measured in many units. The official SI unit is the kilopascal (kPa). We also use two
other units as well, the atmosphere (atm) and millimeter of mercury (mm Hg). Standard pressure
(close to the air pressure in the room on any given day) is defined as the following values:
101.3 kPa = 1.00 atm = 760 mm Hg
Using dimensional analysis, convert the following pressure values into the desired unit.
1. 75 kPa = _____ atm
3. 655 mm Hg = _____ atm
2. 1.55 atm = _____ kPa
4. 128 kPa = _____ mm Hg
113
Ideal Gas Law
The ideal gas equation relates all four measurable properties of a gas. If you know three of the
properties, this equation allows you to calculate the fourth. It has the form:
PV = nRT
List what each of the letters represents and the units for each:
P=
V=
n=
T=
R is defined as the gas law constant. Its value does not change.
State the value and units for R:
Solve the following problems using the ideal gas equation.
1. What volume does 1.25 moles of gas occupy at 35ºC and 145 kPa?
2. What is the pressure of 2.38 moles of gas that has a volume of 38 dm3 at a temperature of -55ºC?
3. What volume does 10.0 g of carbon dioxide occupy at 75ºC and 85 kPa?
4. How many grams of sulfur dioxide are contained in a 2.50 dm3 container at a pressure of 135 kPa
and a temperature of 28ºC?
5. What is the molar mass of a 2.50 g sample of gas that occupies 375 cm3 at a temperature of 15ºC
and a pressure of 175 kPa?
114
Gas Laws -- Calculations
1.
A gas occupies 1.00 cm3 at STP. What volume does it occupy at 710 mm Hg and 55C?
2.
What volume does a gas occupy at 1.50 atm, if at standard pressure it has a volume of
L?
3.
A gas occupies 48.0 mL at 75C. What volume does it occupy at 15C?
4.
A gas occupies 0.75 dm3 at 125C and 95.5 kPa. What volume does it occupy at STP?
5.
Two gases are mixed in a 1.00 L container. The total pressure of the combined gases is 2.25
atm. If the pressure of one gas is 0.75 atm, what is the pressure of the second gas?
6.
A mixture of three gases (A,B and C) has a pressure of 800 mm Hg. There are 1.00 mole of
gas A, 0.75 mole of gas B, and 0.25 mole of gas C. What are the partial pressures of each
gas?
7
What is the volume of 16.0 g of oxygen gas at 25 C and 135 kPa?
8
A fixed amount of gas has its temperature doubled and its pressure doubled. What happens to
its volume?
9
A fixed amount of gas has its temperature tripled and its pressure cut in half. What happens
to its volume?
115
25.0
Molar Volume of a Gas
1.
What does STP stand for?
2.
What is the volume of 1.0 mole of ANY gas at STP? (Remember units!!)
3.
State Avogadro's hypothesis.
4.
Solve the following problems using your roadmap and the following balanced equation
C2H4(g) + 3 O2(g)  2 CO2(g) + 2 H2O(g)
a. What volume does 8.0 g of carbon dioxide gas occupy at STP?
b. What is the volume of 1.50 x 1022 molecules of oxygen gas at STP?
c. What volume of carbon dioxide is produced when 10.0 dm3 of oxygen is consumed?
d. What mass of water is produced when 10.0 dm3 of ethene (C2H4) at STP undergoes combustion?
e. What volume of carbon dioxide gas at STP is produced when 24.0 g of oxygen gas is reacted?
f. How many molecules of water are produced by the consumption of 10.0 L of ethene gas at STP?
116
Molarity
Define the following terms:
SoluteSolventSolution
ConcentrationMolarityDilution-
Solve the following problems.
1. What is the molarity of 1.75 dm3 of solution that contains 2.45 moles of NaCl?
2. What is the molarity of 0.755 dm3 of solution that contains 25 g of NaCl?
3. 1.5 dm3 of 0.35 mol dm-3 solution contains what mass of KNO3?
4. What volume of 1.45 mol dm-3 solution contains 84.0 g of MgBr2?
5. What is the molarity of 282 cm3 of solution that contains 18.7 g of potassium oxide?
6. What mass of lithium nitrate is contained in 583 mL of 2.50 mol dm-3 solution?
117
7. What is the concentration of Cl- in a 2.50 mol dm-3 solution of AlCl3?
8. What is the concentration of K+ ions in a solution prepared with 10.0 g of potassium oxide in 255
cm3 of solution?
9. 15 mL of 12 mol dm-3 HCl solution is diluted to a total volume of 225 mL. What is the new
concentration?
10. What volume of 6.0 M HCl solution must be added to water to make 365 mL of 1.75 M solution?
11.
2 NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2 H2O(l)
Using the reaction above, what mass of H2O can be formed from the reaction of 100.0 cm3 of
1.25 mol dm-3 H2SO4 solution with excess NaOH solution?
12.
2 NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2 H2O(l)
Using the reaction above, what mass of H2O can be formed from the reaction of 150.0 cm3 of
1.00 mol dm-3 H2SO4 solution with 200.0 cm3 of 2.00 mol dm-3 NaOH solution? (This is a
limiting reactant problem!)
118
Stoichiometry Review
Use the following equation to solve problems 1-3.
N2(g) + 3 H2(g) → 2 NH3(g)
1. How many molecules of NH3 can be produced from the reaction of 10.0 dm3 of N2 at STP with
excess hydrogen?
2. If 2.0 L of N2 gas and 5.0 L of H2 gas are reacted together, what is the theoretical volume of NH3
gas that can be produced? Assume constant temperature and volume.
3. In problem #2, which reactant is in excess and what volume will be left unreacted?
Use the following equation to solve problems 4-5
Ba(OH)2(aq) + 2 HCl(aq) → BaCl2(aq) + 2 H2O(l)
4. What mass of water will be produced from the reaction of 100.0 cm3 of 0.100 mol dm-3 Ba(OH)2
solution and 100.0 cm3 of 0.100 mol dm-3 HCl?
5. What volume of 0.250 mol dm-3 HCl would be required to completely neutralize 1.25 dm3 of
0.500 mol dm-3 Ba(OH)2?
119
Topic 1 NOTES
120
Study Guide: Topic 1 Quantitative Chemistry
1.
Which of the following quantities has units?
A.
Relative atomic mass
B.
Relative molecular mass
C.
Molar mass
D.
Mass number
(Total 1 mark)
2.
What is the total number of atoms in 0.20 mol of propanone, CH3COCH3?
A.
1.2×1022
B.
6.0×1023
C.
1.2×1024
D.
6.0×1024
(Total 1 mark)
3.
Which contains the same number of ions as the value of Avogadro’s constant?
A.
0.5 mol NaCl
B.
0.5 mol MgCl2
C.
1.0 mol Na2O
D.
1.0 mol MgO
(Total 1 mark)
4.
Which is a correct definition of the term empirical formula?
A.
formula showing the numbers of atoms present in a compound
B.
formula showing the numbers of elements present in a compound
C.
formula showing the actual numbers of atoms of each element in a compound
D.
formula showing the simplest ratio of numbers of atoms of each element in a compound
(Total 1 mark)
5.
The complete oxidation of propane produces carbon dioxide and water as shown below.
C3H8 + __O2 __CO2 + __H2O
What is the total of the coefficients for the products in the balanced equation for 1 mole of
propane?
A.
6
B.
7
C.
12
D.
13
(Total 1 mark)
6.
When the equation below is balanced for 1 mol of C3H4, what is the coefficient for O2?
C3H4 + O2  CO2 + H2O
A.
2
B.
3
C.
4
D.
5
(Total 1 mark)
121
7.
The equation for a reaction occurring in the synthesis of methanol is
CO2 + 3H2  CH3OH + H2O
What is the maximum amount of methanol that can be formed from 2 mol of carbon dioxide and
3 mol of hydrogen?
A.
1 mol
B.
2 mol
C.
3 mol
D.
5 mol
(Total 1 mark)
8.
Lithium hydroxide reacts with carbon dioxide as follows.
2LiOH + CO2 → Li2 CO3 + H2O
What mass (in grams) of lithium hydroxide is needed to react with 11 g of carbon dioxide?
A.
6
B.
12
C.
24
D.
48
(Total 1 mark)
9.
3.0 dm3 of sulfur dioxide is reacted with 2.0 dm3 of oxygen according to the equation below.
2SO2(g) + O2(g) → 2SO3(g)
What volume of sulfur trioxide (in dm3) is formed? (Assume the reaction goes to completion and
all gases are measured at the same temperature and pressure.)
A.
5.0
B.
4.0
C.
3.0
D.
2.0
(Total 1 mark)
10.
Which change in conditions would increase the volume of a fixed mass of gas?
A.
B.
C.
D.
Pressure /kPa
Doubled
Halved
Doubled
Halved
Temperature /K
Doubled
Halved
Halved
Doubled
(Total 1 mark)
11.
The temperature in Kelvin of 2.0 dm3 of an ideal gas is doubled and its pressure is increased by a
factor of four. What is the final volume of the gas?
A.
1.0 dm3
B.
2.0 dm3
C.
3.0 dm3
D.
4.0 dm3
(Total 1 mark)
12.
What volume (in dm3) of 0.30 mol dm–3 NaCl solution can be prepared from 0.060 mol of solute?
A.
0.018
B.
0.20
C.
0.50
D.
5.0
(Total 1 mark)
122
13.
Which solution contains 0.1 mol of sodium hydroxide?
A.
1 cm3 of 0.1 mol dm–3 NaOH
B.
10 cm3 of 0.1 mol dm–3 NaOH
C.
100 cm3 of 1.0 mol dm–3 NaOH
D.
1000 cm3 of 1.0 mol dm–3 NaOH
(Total 1 mark)
14.
Assuming complete reaction, what volume of 0.200 mol dm–3 HCl(aq) is required to neutralize
25.0 cm3 of 0.200 mol dm–3 Ba(OH)2(aq)?
A.
12.5 cm3
B.
25.0 cm3
C.
50.0 cm3
D.
75.0 cm3
(Total 1 mark)
15.
An organic compound A contains 62.0 by mass of carbon, 24.1 by mass of nitrogen, the
remainder being hydrogen.
(i)
Determine the percentage by mass of hydrogen and the empirical formula of A.
...................................................................................................................................
...................................................................................................................................
...................................................................................................................................
...................................................................................................................................
...................................................................................................................................
...................................................................................................................................
(3)
(ii)
Define the term relative molecular mass.
...................................................................................................................................
...................................................................................................................................
...................................................................................................................................
...................................................................................................................................
(2)
(iii)
The relative molecular mass of A is 116. Determine the molecular formula of A.
...................................................................................................................................
...................................................................................................................................
(1)
(Total 6 marks)
123
16.
Copper metal may be produced by the reaction of copper(I) oxide and copper(I) sulfide according
to the below equation.
2Cu2O + Cu2S  6Cu + SO2
A mixture of 10.0 kg of copper(I) oxide and 5.00 kg of copper(I) sulfide was heated until no
further reaction occurred.
(a)
Determine the limiting reagent in this reaction, showing your working.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(3)
(b)
Calculate the maximum mass of copper that could be obtained from these masses of
reactants.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(2)
(Total 5 marks)
124
Topic 5: Energetics
5.1 Exothermic and Endothermic Reactions
5.1.1 Define the terms exothermic reaction, endothermic reaction and standard enthalpy change of
reaction (∆H°).
5.1.2 State that combustion and neutralization are exothermic processes.
5.1.3 Apply the relationship between temperature change, enthalpy change and the classification of
a reaction as endothermic or exothermic.
5.1.4 Deduce, from an enthalpy level diagram, the relative stabilities of reactants and products, and
the sign of the enthalpy change for the reaction.
5.2 Calculation of Enthalpy Changes
5.2.1 Calculate the heat energy change when the temperature of a pure substance is changed.
5.2.2 Design suitable experimental procedures for measuring the heat energy changes of reactions.
5.2.3 Calculate the enthalpy change for a reaction using experimental data on temperature changes,
quantities of reactants and mass of water.
5.2.4 Evaluate the results of experiments to determine enthalpy changes.
5.3 Hess’s Law
5.3.1 Determine the enthalpy change of a reaction that is the sum of two or three reactions with
known enthalpy changes.
5.4 Bond Enthalpies
5.4.1 Define the term average bond enthalpy.
5.4.2 Explain, in terms of average bond enthalpies, why some reactions are exothermic and others
are endothermic.
125
Topic 5 NOTES
126
Average Bond Enthalpy and Enthalpy of Reaction (ΔH)
1. Define bond enthalpy.
2. When a bond is formed energy is absorbed/released. (choose one)
3. When a bond is broken energy is absorbed/released.
4. In a chemical reaction bonds are both broken and formed. The difference in the energies of the
bonds is known as the heat of reaction, ΔH. If energy is released during a chemical reaction it is
said to be endothermic/exothermic.
5. If energy is absorbed the reaction is said to be endothermic/exothermic.
6. If the heat of reaction is negative the reaction is said to be endothermic/exothermic.
7. If the heat of reaction is positive the reaction is said to be endothermic/exothermic.
8. In an exothermic reaction the products/reactants have greater energy.
9. In an endothermic reaction the products/reactants have greater energy.
10. In an exothermic reaction the surrounding’s temperature goes up/down.
11. In an endothermic reaction the surrounding’s temperature goes up/down
12. In the formation of H2O from its elements, what bonds must be broken? (look at the visual
representation of the reaction)
2H2
+
O2
→
2H2O
13. What bonds are formed in the equation above?
14. The H-H bond energy is 432 kJ/mol. The O=O bond energy is 495 kJ/mol. How much energy is
required to break up the reactants in this reaction?
15. The O-H bond energy is 467 kJ/mol. How much energy will be realeased in the formation of the
products of this reaction?
16. Calculate the H for this reaction. ΔH = sum of bonds broken - sum of bonds formed.
17. Is this reaction endothermic or exothermic? How do you know?
18. Which are more stable, the reactants or products?
127
128
Enthalpy of Reaction Worksheet
1. a. Write a balanced equation for the combustion of ethyne (C2H2).
b. Draw dot diagrams for each substance in your balanced equation.
c. Determine the bond enthalpies of the bonds broken in the reaction. (use the table on the
preceding page)
d. Determine the bond enthalpies of the bonds formed.
e. Determine the enthalpy of reaction, ΔH, for the reaction.
f. All combustion reactions are (exothermic/endothermic).
g. Which have more energy, the reactants or products?
h. Which are more stable, the reactants or products?
2. Determine the enthalpy of reaction for the following reaction:
N2(g) + 3 H2(g) → 2 NH3(g)
129
Enthalpy Level Diagrams
Answer the questions by referring to the diagrams of the potential energy of a reaction.
Potential Energy Diagram #1
Potential Energy Diagram #2
75
D
XY
50
C
25
XY
50
X+
Y
A
D
A
E
B
B
X+Y
25
C
0
E
0
Reaction Coordinate (X + Y  XY)
Reaction Coordinate (XY  X + Y)
Identify the letter that describes the following for each:
Diagram #1
Letter
Diagram #2
Value
Value
Letter
(kJ)
(kJ)
Total Energy of the Reactants
Total Energy of the Products
Total Energy of the Activated Complex
Heat of Reaction
Indicate if this reaction is endothermic or exothermic
1. If a catalyst were added, would the activation energy increase, decrease, or remain the same?
2. If a catalyst were added, would the heat of reaction increase, decrease, or remain the same?
Use the axes below to draw a potential energy diagram for the following conditions. The potential
energy of the reactants is 100 kJ. The potential energy of the products is 50 kJ. The activation energy
is 25 kJ.
Potential Energy
Potential Energy (kJ)
75
100
50
Reaction Coordinate
1. Is this reaction exothermic or endothermic?
2. What is the value of the heat of reaction?
3. Draw a dotted line representing the effect of adding a catalyst.
130
Enthalpy of Reaction Worksheet
2Al + Fe2O3 → Al2O3 + 2Fe
1.
a.
b.
c.
d.
e.
ΔH = -850 kJ
Is this reaction exothermic or endothermic?
Is heat a reactant or product in this reaction?
What type of reaction is this?
How many moles of aluminum are required to produce 850 kJ?
How much heat energy is produced when 20.0 g of aluminum are reacted?
f. How much heat energy is produced when 20.0 g of aluminum oxide are produced?
g. What mass of iron is produced if 20.0 g of iron (III) oxide are consumed?
N2 + 2O2 → 2NO2
2.
a.
b.
c.
d.
e.
ΔH = 67.7 kJ
Is this reaction exothermic or endothermic?
Is heat a reactant or product in this reaction?
What type of reaction is this?
Which are more stable in this reaction, the reactants or products?
How much heat is involved in the production of 10.0 g of NO2?
f. How many molecules of O2 are required to produce 10.0 g of NO2?
131
Energy and Calorimetry
1. What is the SI unit of energy?
2. The calorie is a non-SI unit of energy commonly used. How are the calorie and the joule
related?
3. Convert the following units:
a. 425 J = _____ cal
b. 275 cal = _____ J
4. How does a calorie relate to a nutritional Calorie?
5. How many joules are contained in a slice of chocolate cake with 255 Calories?
6. Define specific heat.
7. What is the specific heat of water in joules and in calories? (Remember to use all appropriate
units)
8. In the equation q = m  CP  T, what do each of the letters represent? (include units)
q = _____________
m = _____________
CP = _____________
T = _____________
Solve the following problems using the specific heat data listed in the Appendix of your textbook.
9. 40.0 grams of water are heated from 10.0C to 30.0C. Determine the heat absorbed in
joules.
10. 25.0 grams of aluminum are heated from 15.0C to 35.0C. Determine the heat absorbed.
11. 1.5 kg of lead is cooled from 136.4C to 21.7C. Determine the heat released.
12. What mass of copper can be heated by 3.00 x 103 J, if the temperature increases by 35.0C?
13. What is the change in temperature of 200.0 g of mercury that absorbs 4.20 x 103 J?
14. What is the final temperature of the mixture when 25.0 g of 15.0C water is mixed with 74.0
g of 95.0C water?
15. What is the final temperature of the mixture when a 55.0 g iron bar at 115C is placed into
155 g of water at 25.0C?
132
Heat of Reaction Problem Set
1. A 0.250 mol dm-3 solution of HCl is added to 1.00 mol dm-3 NaOH. 100.0 ml of each solution is
used. A 13.0 degree temperature increase is noted. Calculate the energy absorbed or released in
kJ mol-1 of HCl.
2. A sample of 3.50 g of NaOH is added to 125.0 ml of 1.00 mol dm-3 HCl. The temperature
increase was observed to be 5.00 degrees. Calculate the energy absorbed or released in kJ mol-1 of
HCl.
3. A 10.0 g sample of NH4Cl is added to 100.0 ml of water. The temperature decrease was observed
to be 5.00 degrees. Calculate the energy absorbed or released in kJ mol-1 of NH4Cl.
4. A 1.50 g sample of zinc is added to 50.0 ml of 0.250 mol dm-3 copper (II) sulfate solution. The
temperature increase was observed to be 7.00 degrees. Calculate the energy absorbed or released
in kJ mol-1 of zinc.
5. Equal volumes of 6.00 mol dm-3 HNO3 (nitric acid) and 1.00 mol dm-3 KOH are mixed. The total
volume of mixture is 500.0 cm3. The temperature increase was observed to be 15.0 degrees.
Calculate the energy absorbed or released in kJ mol-1 of KOH.
133
6. A 1.25 mol dm-3 solution of HCl is added to 2.00 mol dm-3 NaOH. 200.0 ml of each solution is
used. A 17.0 degree temperature increase is noted. Calculate the energy absorbed or released in
kJ mol-1 of HCl.
7. A sample of 10.5 g of NaOH is added to 125.0 ml of 1.00 mol dm-3 HCl. The temperature
increase was observed to be 5.00 degrees. Calculate the energy absorbed or released in kJ mol-1 of
HCl.
8. A 20.0 g sample of NH4NO3 is added to 100.0 ml of water. The temperature decrease was
observed to be 8.00 degrees. Calculate the energy absorbed or released in kJ mol-1 of NH4NO3.
9. A 1.50 g sample of zinc is added to 150.0 ml of 0.250 mol dm-3 copper (II) sulfate solution. The
temperature increase was observed to be 10.0 degrees. Calculate the energy absorbed or released
in kJ mol-1 of copper (II) sulfate.
10. Equal volumes of 1.00 mol dm-3 HNO3 (nitric acid) and 1.00 mol dm-3 Ca(OH)2 are mixed. The
total volume of mixture is 1000.0 cm3. The temperature increase was observed to be 25.0 degrees.
Calculate the energy absorbed or released in kJ mol-1 of HNO3.
134
Hess’s Law
1. Calculate the heat of reaction for the combustion of nitrogen monoxide gas, NO, to form
nitrogen dioxide gas, NO2, as given in the following thermochemical equation.
2 NO (g) + O2 (g)  2 NO2 (g)
You are given the following heat of formation data:
N2 (g) + O2 (g)  2 NO (g)
H0 = 180.58 kJ
N2 (g) + 2 O2 (g)  2 NO2 (g)
H0 = 66.4 kJ
2. Calculate the H for the following reaction
2 N2 (g) + 5 O2 (g)  2 N2O5 (g)
Use the following date in your calculations:
2 H2 (g) + O2 (g)  2 H2O (l)
N2O5 (g) + H2O (l)  2 HNO3 (l)
N2 (g) + 3 O2 (g) + H2 (g)  2 HNO3 (l)
H0 = -571.6 kJ
H0 = -76.6 kJ
H0 = -348.2 kJ
3. Calculate the heat of formation for pentane, C5H12.
5 C (s) + 6 H2 (g)  C5H12 (g)
Use the following data for your calculations:
C (s) + O2 (g)  CO2 (g)
2 H2 (g) + O2 (g)  2 H2O (l)
C5H12 (g) + 8 O2 (g)  5 CO2 (g) + 6 H2O (l)
H0 = -393.5 kJ
H0 = -571.6 kJ
H0 = -3535.6 kJ
135
4. Calculate the heat of formation for sulfur dioxide, from its elements.
Write the balanced equation you are looking for:
Use the following data for your calculations:
2 S (s) + 3 O2 (g)  2 SO3 (g)
2 SO2 (g) + O2 (g)  2 SO3 (g)
5. Calculate the ΔH for the process
A  2C + E
Use the following numbered processes:
1.
A  2B
ΔH1
2.
B  C +D
ΔH2
3.
E  2D
ΔH3
Express your answers in terms of ΔH1, ΔH2, and ΔH3.
136
H0 = -790.2 kJ
H0 = -198.2 kJ
The Enthalpy of Decomposition of Hydrogen Peroxide
Purpose: To experimentally determine the enthalpy of reaction using a calorimeter. To use Hess’
Law to verify experimental data.
Equipment:
Graduated cylinder
PC interfaced thermometer
H2O2 solution (3.00% by mass)
stirring rod or magnetic stirrer
calorimeter (Styrofoam cup!)
0.500 M Fe(NO3)3 solution
There is no single instrument that can directly measure heat in the way a balance measures mass or a
thermometer measures temperature. However, it is possible to determine the heat change when a
chemical reaction occurs. The change in heat is calculated from the mass, temperature change and
specific heat of the substance which gains or loses heat.
The equation that is used to calculate heat gain or loss is:
q = (grams of substance ×(specific heat) × (ΔT)
Where q = the heat gained or lost and ΔT is the change in temperature. Since ΔT = (final
temperature minus initial temperature), an increase in temperature will result in a positive value for
both ΔT and q, and a loss of heat will give a negative value.
A calorimeter itself will participate in the transfer of heat. It will be assumed that this heat transfer
will be minimal and will be neglected in calculations.
ΔH is related to the amount of substance involved in a reaction. To determine ΔH from q, you will
need to divide the value of q by the moles of substance.
In this lab you will be determining the enthalpy of reaction (in kJ/mol) for the decomposition of
hydrogen peroxide which slowly decomposes into water and oxygen gas.
Procedure
Add 50.0 mL of 3.0% H2O2 solution into the calorimeter. Place a temperature probe into the
calorimeter and secure it with a clamp so it doesn’t overturn. Connect the probe to a computer and
set it up to record the temperature for 15 minutes. (Your teacher will help you with this.) Start the
data collection and stir continuously. After two minutes have elapsed, add 10.0 mL of 0.50 M
Fe(NO3)3 solution to the solution. Let the computer continue to collect data for the remaining time.
Continue to stir! Print a copy of the resulting temperature vs. time graph.
Calculations (show your work and uncertainty)
1. Determine the moles of H2O2 present in the 50.0 mL sample. (Assume the density of the
solution is the same as water, 1.00 g/mL.)
2. Determine the heat transferred to the solution in the calorimeter, qcalorimeter. (The specific heat
of the solution can be assumed to be the same as that of water.)
137
3. From the law of conservation of energy, qreaction
=
- qcalorimeter. Determine qreaction.
4. Determine the heat of reaction, ΔH, for the decomposition of hydrogen peroxide in joules per
mole.
5. Write the balanced equation for the decomposition of hydrogen peroxide. Use ΔH notation.
Post Lab
1. How does the graphical temperature analysis improve the accuracy of your data?
2. What is the purpose of the Fe(NO3)3 solution? Why is it not included in the chemical
equation for the decomposition of hydrogen peroxide?
3.
2H2 + O2 → 2H2O
ΔH = -572 kJ
Look at the equation written directly above. Look at the equation you wrote in calculation
number 5. Using these two equations and Hess’ Law, calculate the ΔH for the following
overall equation:
H2 + O2 → H2O2
ΔH = ?
Hint: rearrange the equations to make the third equation.
4. The true value of ΔH for reaction H2 + O2 → H2O2, is -187.5 kJ. Calculate your percent
error.
138
Energy Change in Phase Changes
1. Define "heat (enthalpy) of vaporization"-2. Define "heat (enthalpy) of fusion"-3. What are the symbols used for heat of vaporization and heat of fusion?
4. When a substance is strongly attracted to other molecules in the liquid state, it will have a
(high/low) heat of vaporization.
5. Which would you expect to have a higher heat of vaporization, H2O or Cl2? (Think about the type
of intermolecular force holding the molecules together) Explain.
6. What are the values of the heat of vaporization and the heat of fusion of water? Remember to use
the proper units.
7. How much energy does it take to vaporize 2.5 moles of water at 100ºC?
8. How much energy does it take to melt 35.0 g of water at 0ºC?
9. What happens to the temperature (kinetic energy) during a phase change?
10. List 3 phase changes that are exothermic.
10. Draw a heating diagram for heating 45.0 g of ice from a temperature of -25.0º C to steam at
110ºC.
temperature
(ºC)
time
11. From the information given in problem number 10, how much energy is involved in heating the
ice from -25.0º C to steam at 110ºC? Hint: this will take five steps!
139
Study Guide: Topic 5 Energetics
1.
According to the enthalpy level diagram below, what is the sign for H and what term is used to
refer to the reaction?
H
reactants
products
reaction progress
H
reaction
A.
positive
endothermic
B.
negative
exothermic
C.
positive
exothermic
D.
negative
endothermic
(Total 1 mark)
2.
Which statement is correct about the reaction shown?
2SO2(g) + O2(g)  2SO3(g)
H = –196 kJ
A.
B.
C.
D.
196 kJ of energy are released for every mole of SO2(g) reacted.
196 kJ of energy are absorbed for every mole of SO2(g) reacted.
98 kJ of energy are released for every mole of SO2(g) reacted.
98 kJ of energy are absorbed for every mole of SO2(g) reacted.
(Total 1 mark)
3.
Which statement is correct for an endothermic reaction?
A.
The products are more stable than the reactants and H is positive.
B.
The products are less stable than the reactants and H is negative.
C.
The reactants are more stable than the products and H is positive.
D.
The reactants are less stable than the products and H is negative.
(Total 1 mark)
4.
The following equation shows the formation of magnesium oxide from magnesium metal.
2Mg(s) + O2(g)2MgO(s)
HӨ = –1204kJ
Which statement is correct for this reaction?
A.
1204 kJ of energy are released for every mol of magnesium reacted.
B.
602 kJ of energy are absorbed for every mol of magnesium oxide formed.
C.
602 kJ of energy are released for every mol of oxygen gas reacted.
D.
1204 kJ of energy are released for every two mol of magnesium oxide formed.
(Total 1 mark)
140
5.
Which of the quantities in the enthalpy level diagram below is (are) affected by the use of a
catalyst?
Enthalpy
I
II
III
Time
A.
B.
C.
D.
I only
III only
I and II only
II and III only
(Total 1 mark)
6.
When the solids Ba(OH)2 and NH4SCN are mixed, a solution is produced and the
temperature drops.
Ba(OH)2(s) + 2NH4SCN(s) → Ba(SCN)2(aq) + 2NH3(g) + 2H2O(l)
Which statement about the energetics of this reaction is correct?
A.
B.
C.
D.
The reaction is endothermic and H is negative.
The reaction is endothermic and H is positive.
The reaction is exothermic and H is negative.
The reaction is exothermic and H is positive.
(Total 1 mark)
7.
Which statements about exothermic reactions are correct?
I.
They have negative H values.
II.
The products have a lower enthalpy than the reactants.
III. The products are more energetically stable than the reactants.
A.
B.
C.
D.
I and II only
I and III only
II and III only
I, II and III
(Total 1 mark)
8.
Consider the specific heat capacity of the following metals.
Specific heat capacity / J kg–1 K–1
Cu
385
Ag
234
Au
130
Pt
134
Which metal will show the greatest temperature increase if 50 J of heat is supplied to a 0.001 kg
sample of each metal at the same initial temperature?
A.
Cu
B.
Ag
C.
Au
D.
Pt
(Total 1 mark)
Metal
141
9.
When 40 joules of heat are added to a sample of solid H2O at –16.0°C the temperature increases
to –8.0°C. What is the mass of the solid H2O sample?
[Specific heat capacity of H2O(s) = 2.0 J g–1K–1]
A.
2.5 g
B.
5.0 g
C.
10 g
D.
160 g
(Total 1 mark)
10.
The mass m (in g) of a substance of specific heat capacity c (in J g–1 K–1 ) increases by t°C. What
is the heat change in J?
A.
mct
B.
mc(t + 273)
C.
D.
mct
1000
mc (t  273)
1000
(Total 1 mark)
11.
Calculate the enthalpy change, H4 for the reaction
C + 2H2 +
1
2
O2  CH3OH
H4
using Hess’s Law and the following information.
CH3OH + 1 12 O2  CO2 + 2H2O
H1 = 676 kJ mol1
C + O2  CO2
H2 = 394 kJ mol1
H2 +
1
2
H3 = 242 kJ mol1
O2  H2O
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
(Total 4 marks)
142
12.
The data below is from an experiment used to measure the enthalpy change for the combustion of
1 mole of sucrose (common table sugar), C12H22O11(s). The time-temperature data was taken
from a data-logging software program. The sugar was burned and used to heat water whose
temperature is shown below.
Mass of sample of sucrose, m = 0.100 g
Mass of water heated by the combustion of sucrose, m = 300.0 g
Heat capacity of water c= 4.18 J g–1 C–1
(a)
Calculate ΔT, for the water, surrounding the chamber in the calorimeter.
.....................................................................................................................................
.....................................................................................................................................
(1)
(b)
Determine the amount, in moles, of sucrose.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(1)
143
(c)
(i)
Calculate the enthalpy change for the combustion of 1 mole of sucrose.
...........................................................................................................................
...........................................................................................................................
(1)
(ii)
The true value of the enthalpy of combustion of sucrose is -5644 kJ mol-1. Calculate
the percentage experimental error based on the data used in this experiment.
...........................................................................................................................
...........................................................................................................................
(1)
(d)
A hypothesis is suggested that TNT, 2-methyl-1,3,5-trinitrobenzene, is a powerful
explosive because it has:
• a large enthalpy of combustion
• a high reaction rate
• a large volume of gas generated upon combustion
Use your answer in part (c)(i) and the following data to evaluate this hypothesis:
Equation for combustion
Relative
rate of
combustion
Sucrose
C12H22O11(s) + 12O2(g)  12CO2(g) + 11H2O(g)
Low
TNT
2C7H5N3O6(s)  7CO(g) + 7C(s) + 5H2O(g) + 3N2(g)
High
Enthalpy of
combustion
/ kJ mol–1
3406
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(3)
(Total 7 marks)
144
13.
The standard enthalpy change of formation values of two oxides of phosphorus are:
P4(s) + 3O2(g)  P4O6(s)
HӨf= –1600 kJ mol–1
P4(s) + 5O2(g)  P4O10(s)
HӨf= –3000 kJ mol–1
What is the enthalpy change, in kJ mol–1, for the reaction below?
P4O6(s) + 2O2(g)  P4O10(s)
A.
B.
C.
D.
+4600
+1400
–1400
–4600
(Total 1 mark)
14.
The equations and enthalpy changes for two reactions used in the manufacture of sulfuric acid
are:
S(s) O2(g)  SO2(g)
HӨ = –300 kJ
2SO2(g) + O2(g)  2SO3(g)
HӨ = –200 kJ
What is the enthalpy change, in kJ, for the reaction below?
2S(s) + 3O2(g)  2SO3(g)
A.
B.
C.
D.
–100
–400
–500
–800
(Total 1 mark)
15.
For the reaction
2H2(g) + O2(g)  2H2O(g)
the bond enthalpies (in kJ mol–1) are
H–H
x
O=O
y
O–H
z
Which calculation will give the value, in kJ mol–1, of HӨ for the reaction?
A.
2x + y –2z
B.
4z – 2x – y
C.
2x + y – 4z
D.
2z –2x – y
(Total 1 mark)
145
16.
But–1–ene gas, burns in oxygen to produce carbon dioxide and water vapour according to the
following equation.
C4H8 + 6O2  4CO2 + 4H2O
(a)
Use the data below to calculate the value of HӨ for the combustion of but-1-ene.
Bond
Average bond
enthalpy / kJ mol–
CC
C=C
CH
O=O
C=O
O–H
348
612
412
496
743
463
1
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(3)
(b)
State and explain whether the reaction above is endothermic or exothermic.
....................................................................................................................................
....................................................................................................................................
(1)
(Total 4 marks)
146
Topic 6: Kinetics
6.1 Rates of Reaction
6.1.1 Define the term rate of reaction.
6.1.2 Describe suitable experimental procedures for measuring rates of reactions.
6.1.3 Analyze data from rate experiments.
6.2 Collision Theory
6.2.1 Describe the kinetic theory in terms of the movement of particles whose average energy is
proportional to temperature in kelvins.
6.2.2 Define the term activation energy, Ea.
6.2.3 Describe the collision theory.
6.2.4 Predict and explain, using the collision theory, the qualitatively effects of particle size,
temperature, concentration and pressure on the rate of a reaction.
6.2.5 Sketch and explain qualitatively the Maxwell-Boltzmann energy distribution curve for a fixed
amount of gas at different temperatures and its consequences for changes in reaction rate.
6.2.6 Describe the effect of a catalyst on a chemical reaction.
6.2.7 Sketch and explain Maxwell-Boltzmann curves for reactions with and without catalysts.
Topic 7: Equilibrium
7.1 Dynamic Equilibrium
7.1.1 Outline the characteristics of chemical and physical systems in a state of equilibrium.
7.2 The position of Equilibrium
7.2.1 Deduce the equilibrium constant expression from the equation for a homogeneous reaction.
7.2.2 Deduce the extent of a reaction from the magnitude of the equilibrium constant.
7.2.3 Apply Le Chatelier’s principle to predict the qualitative effects of changes of temperature,
pressure, and concentration on the position of equilibrium and on the value of the equilibrium
constant.
7.2.4 State and explain the effect of a catalyst on an equilibrium reaction.
7.2.5 Apply the concepts of kinetics and equilibrium to industrial processes.
147
Topic 6 & 7 NOTES
148
Rates of Reaction
1. Define “rate of reaction”
2. All rates are measured as a change in something over a change in time. What are the units of rate of
reaction?
3. Below is a graph of concentration vs time for reaction:
molarity
time
a. Does the graph show the change in reactants or products? How do you know?
b. The slope of the graph at any time gives the instantaneous rate. Is the rate (slope) constant?
c. When is the magnitude of the slope the greatest?
d. What is the sign of the slope?
4. Given the following reaction: A + 2B → C
The rate of reaction in terms of A is -0.5 mol dm-3 sec-1.
a. What is the rate in terms of B?
b. What is the rate in terms of C?
5.
Molarity
1.0
Time
Shown is the graph for the reaction A + 2B → C. The line on the curve represents the concentration of A
over time. Draw lines representing the concentrations of B and C over time and label each on the graph.
6. We do not have a device that directly measures molarity. Name 4 things (that are directly related to
concentration) you can measure in the lab that would allow you to determine a rate?
7. What methods could you use to measure the rate of reaction for the following:
a. Calcium carbonate slowly dissolves in acid to make carbon dioxide gas and a dissolved salt.
b. A sample of mineral ore slowly dissolves in water to make aqueous ions.
c. A piece of zinc is placed in a solution of copper (II) sulfate. (You did this in your types of reactions
lab.)
149
Kinetics Worksheet
1. Kinetics is the study of the ___________ of a reaction.
2. In order for particles to react, they must ___________.
3. In order for collisions to be successful and produce products, the particles must possess what two
characteristics?
a.
b.
4. Name four ways to speed up a reaction.
a.
b.
c.
d.
5. Name two reasons increasing the temperature speeds up a reaction. Put a star next to the primary reason.
6. Draw a potential energy diagram for an exothermic reaction. Label the reactants, products and activated
complex.
PE
Reaction coordinate
7. On the diagram above, draw a dotted line representing the effect of a catalyst.
8. On the diagram above, draw a double line representing the effect of an inhibitor.
9. Explain why the addition of a catalyst increases the rate of reaction.
10. Draw a curve of kinetic energy vs. number of particles for a substance at a given temperature (MaxwellBoltzmann curve). Draw a line to represent the minimum threshold energy for a reaction to occur. Draw a
dotted line that shows how the distribution would change if the temperature were raised. Draw a double
line to show the effect of a catalyst on the distribution.
Number of
Particles
kinetic energy
150
Discovering Equilibrium Activity
To learn a little bit about equilibrium observe the demonstration done by your teacher which involves
two 1-L beakers. One is completely filled with water and represents reactants (A). The second
beaker represents products (B). We will represent the reaction with the following equation. Notice
the double headed arrow. This means the reaction is “reversible”. That means not only can the
reactants make products, but the products can turn around and make reactants.
A ↔ B
Colorless
red
Now let’s start the reaction. (Remember both the forward and reverse reaction occur at the same
time!) To simulate the reaction we will use two 50 mL beakers. Dip each little beaker into our
products and reactants. You should be able to fill the little beaker in the reactants (water) and only
partially fill the beaker in the products (red water). Now start the reaction—turn the reactants into
products by pouring the contents of the little beaker into the products. At the same time pour the
products into the reactants.
1 L beaker
50 mL
beaker
A
↔
B
What volume of A did you put in the little beaker? ______ mL
What volume of B did you put in the other little beaker? _____ mL
Is this system at equilibrium? How do you know?
If we continue dipping the little beakers simultaneously into the big beakers and emptying them, what
will eventually happen to the water levels in the big beakers? Write your hypothesis below:
151
Let’s test our hypothesis.
Continue the reaction by dipping the little beakers into the big beakers and empty their contents into
the opposite side. Repeat this for 20 times. Stop the reaction and look again.
What volume of A did you put in the last little beaker? ______ mL
What volume of B did you put in the other little beaker? ______ mL
Is the system at equilibrium? How do you know?
Equilibrium exists when the rate of the forward reaction equals the rate of the reverse reaction. This
means you are making as many reactants as you are using up. The result is that the amounts of
products and reactants do not change. It does NOT mean that you have equal amounts of products
and reactants.
Summarize your understanding of equilibrium by answering the following true or false questions.
_____ 1. At equilibrium the reaction stops.
_____ 2. To be an equilibrium system, the reaction must be reversible.
_____ 3. At equilibrium the forward and reverse reactions occur at the same rate.
_____ 4. At equilibrium the concentrations of products and reactants become equal.
_____ 5. At equilibrium the concentrations of products and reactants become constant.
Which of the following systems are examples of equilibrium systems where both forward and reverse
reactions occur at the same rate? Place a checkmark in front of any equilibrium system.
_____ 6. Grass growing in a field.
_____ 7. Water evaporating and condensing in a sealed jar.
_____ 8. Wood burning in a fire.
_____ 9. A saturated solution in contact with undissolved solute.
152
Equilibrium Constants
The law of mass action represents a mathematical relationship in equilibrium systems. It is expressed
by the ratio of product concentrations divided by reactant concentrations. Note that only aqueous and
gaseous substances are included in the expression. To write the equilibrium constant expression,
know as K or Keq, look at the following example:
2A(g) + 3B(s) ↔ C(g) + 3D(g)
[C] · [D]3
K=
[A]2
Notice that products are in the numerator and reactants are in the denominator. The brackets
represent molarities. Coefficients become exponents. The reactant, B, was not included because it is
a solid and its molarity is a constant.

1.
Write the equilibrium constant expressions for the following equations.
N2O4(g) ↔ 2NO2(g)
3.
N2(g) + 3H2(g) ↔ 2NH3(g)
2.
PbCl2(s) ↔ Pb2+(aq) + 2Cl-(aq)

4.
3A(g) + B(g) ↔ 2C(s) + D(g)
What does the value of K represent?
5. If the value of K > 1, there is a greater concentration of (products/reactants) at equilibrium.
6. If the value of K< 1, there is a greater concentration of (products/reactants) at equilibrium.

Solve the following problems using this equilibrium system:
N2(g) + 3H2(g) ↔ 2NH3(g)
7. If the concentration of N2 is 0.50 M, the concentration of H2 is 0.40 M and the concentration of
NH3 is 1.0 M, what is the value of K? (You must use the equilibrium constant expression from
problem number 3 above.)
8. At 500 K, the equilibrium constant, K, equals 1.0 × 10-3 M-2. If [H2] = 0.10 M and [NH3] = 0.10
M, what is the concentration of N2 at this equilibrium position?
153
Le Chatelier's Principle
1.
State Le Chatelier's Principle
2.
Define equilibrium
3.
State three different types of stress you can apply to an equilibrium system.
4.
An increase in temperature always favors the endothermic/exothermic reaction.
5.
An increase in concentration of a substance in an equilibrium system will shift the reaction
towards/away from that substance.
6.
An increase in pressure will shift an equilibrium system towards the side with the fewest/most
moles of gaseous substances.
7.
Given the following hypothetical chemical reaction:
2A(g) + B(g)  2C(g) + heat
a.
b.
c.
d.
e.
f.
8.
Is this reaction endothermic or exothermic?
If we increase the concentration of A, which direction will the reaction shift?
(towards the reactants or towards the products)
If we increase the concentration of C, which direction will the reaction shift?
If we increase the pressure of the system, which direction will it shift?
If we increase the temperature of the system, which direction will it shift?
If we increase the temperature of the system, what will happen to the value of K?
Given the following hypothetical chemical reaction:
heat + A(g) + 3B(g)  5C(g)
a. Is this reaction endothermic or exothermic?
b. If we decrease the concentration of A, which direction will the reaction shift?
(towards the reactants or towards the products)
c. If we decrease the concentration of C, which direction will the reaction shift?
d. If we decrease the pressure of the system, which direction will it shift?
e. If we decrease the temperature of the system, which direction will it shift?
f. If we decrease the temperature of the system, what will happen to the value of K?
154
Lab: Equilibrium and LeChatelier’s Principle
LeChatelier’s principle states that if an equilibrium system is subjected to stress, the system will react
to relieve the stress. To relieve a stress, the system can do one of two things: form more products
using up reactants, or reverse the reaction and form more reactants using up products. In this
experiment you will observe several equilibrium systems. Then by putting different stresses on the
systems, you will observe how equilibrium systems react to a stress.
Pre-Lab Questions
1. Define saturated, unsaturated and supersaturated. Which of these represents an equilibrium
system?
2. What types of stress can be placed on an equilibrium system?
3. Write dissociation reactions for the following:
a. AgNO3
b. NaOH
c. HCl
d. KSCN (SCN- is a polyatomic ion known as thiocyanate)
e. Na3PO4
Materials
Equilibrium solutions
(1) Saturated NaCl solution (colorless)
(2) Acid/Base Indicator solution (green)
(3) FeSCN2+ solution (reddish-orange)
(4) Cobalt complex solution (purple)
**HCl, 12M (extremely corrosive solution… use only in fume hood)
HCl, 0.1 M
AgNO3, 0.1 M
NaOH, 0.1 M
Test tubes
KSCN solid
Beaker
Na3PO4 solid
Hot plate
155
Procedure
1. Equilibrium in a saturated solution.
NaCl(s) ↔ Na+(aq) + Cl-(aq)
Obtain a small quantity of saturated NaCl solution in a test tube. Note its appearance. In the
fume hood, add a few drops of 12 M HCl. (HCl is a source of both H+ ions and Cl- ions but only
the Cl- are part of the equilibrium system.) Note any changes in the test tube. Empty the contents
in the waste container in the fume hood and wash the test tube.
2. Acid/Base indicator equilibrium.
HX(aq) ↔ H+(aq) + X-(aq)
Yellow
blue
An indicator is a weak acid, HX, where X is typically a complicated organic ion. Its formula is
not important for this concept. What is important is that HX is one color and X- is another color.
This allows us to detect shifts in the equilibrium.
Place a small quantity of acid/base indicator solution in a test tube. Note its appearance. Add 5
drops of 0.1 M HCl solution and stir. Note any color change. To the same test tube add 0.1 M
NaOH until no further color change occurs. Again, note the color. (Adding OH- ions causes the
H+ concentration to decrease as the ions combine to form water molecules.) See if you can add
the right amount of acid and/or base to this test tube to cause the solution to be green in color
after it is stirred. Rinse and clean the test tube. You may dispose of the solution by washing it
down the drain with water.
3. Complex ion equilibrium.
Fe3+(aq) + SCN-(aq) ↔ FeSCN2+(aq)
colorless
colorless
red-brown
Obtain a small quantity of FeSCN2+ solution and divide it into three small test tubes. The first
test tube will be your reference. Note its color. To the second test tube add 2-3 crystals of KSCN.
To the third test tube add a few crystals of Na3PO4. (PO43- ions have the ability to form complex
ions with Fe3+, which has the same effect as removing Fe3+ ions from solution.) Note the change.
Rinse and clean the test tubes. You may dispose of the solutions by washing them down the drain
with water.
4. Cobalt complex equilibrium.
Co(H2O)62+ + 4 Cl- → CoCl42- + 6 H2O
ΔH = + 50 kJ/mol
Pink
blue
Obtain a small quantity of Cobalt complex equilibrium solution and divide it into four small test
tubes. The first test tube will be your reference. Note the color. To the second test tube add a
few drops of H2O. (Water is not the solvent in this solution it is a product!) Note the color. To
the third test tube add a few drops of 12 M HCl (do this in the fume hood). Note the color. To
the fourth test tube add a few drops of 0.1 M AgNO3 solution. Note the color.
Obtain a sealed pipet containing the cobalt complex equilibrium solution. Note its color. Place
the pipet into an ice water bath for a few minutes. Shake the pipet to stir the contents. Note the
color. Then place the pipet into a hot water bath (~60ºC) while gently shaking the pipet. Note
the color.
156
Data and Analysis
1. Saturated Solution
NaCl(s) ↔ Na+(aq) + Cl-(aq)
Initial observation of system:
Change observed when HCl added:
Stress: increasing the concentration of ClDirection of Shift to relieve the stress: (right or left)
2. Acid/Base indicator
HX(aq) ↔ H+(aq) + X-(aq)
Yellow
blue
Initial observation of system:
Change observed when HCl added:
Stress :
Direction of shift:
Change observed when NaOH added:
Stress:
Direction of shift:
3.
Complex ion equilibrium.
Fe3+(aq) + SCN-(aq) ↔ FeSCN2+(aq)
colorless
colorless
red-brown
Initial observation of system:
Change observed when KSCN added:
Stress:
Direction of shift:
Change observed when Na3PO4 added:
Stress:
Direction of shift:
157
4. Cobalt complex equilibrium.
Co(H2O)62+ + 4 Cl- → CoCl42- + 6 H2O
Pink
blue
ΔH = + 50 kJ/mol
Initial observation of system:
Change observed when H2O added:
Stress:
Direction of shift:
Change observed when HCl added:
Stress:
Direction of shift:
Change observed when AgNO3 added:
Stress:
Direction of shift:
Change observed when placed in ice water:
Stress:
Direction of shift:
Change observed when placed in hot water:
Stress:
Direction of shift:
Questions
1. When the concentration of a reactant was added to an equilibrium system, in what direction did
the system shift? How was the effect different if a product was added?
2. In the cobalt complex equilibrium, neither Ag+ nor NO3- ions were part of the equilibrium
equation yet its addition changed the system. Explain what happened when the silver nitrate
solution was added and why it affected the equilibrium system.
3. Look at the cobalt complex equilibrium equation. Is this reaction endothermic or exothermic?
Which direction did the equilibrium shift when the temperature was increased? Explain why
increasing the temperature of the system shifted the system in the direction it did.
158
Study Guide: Topics 6 & 7 Kinetics and Equilibrium
1.
Which quantities in the enthalpy level diagram are altered by the use of a catalyst?
Enthalpy
I
II
III
Time
A.
B.
C.
D.
I and II only
I and III only
II and III only
I, II and III
(Total 1 mark)
2.
Which statement is correct for a collision between reactant particles leading to a reaction?
A.
Colliding particles must have different energy.
B.
All reactant particles must have the same energy.
C.
Colliding particles must have a kinetic energy higher than the activation energy.
D.
Colliding particles must have the same velocity.
(Total 1 mark)
3.
Which changes increase the rate of a chemical reaction?
I.
Increase in the concentration of an aqueous solution
II.
Increase in particle size of the same mass of a solid reactant
III. Increase in the temperature of the reaction mixture
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
(Total 1 mark)
159
4.
The sequence of diagrams represents the system as time passes for a gas phase reaction in which
reactant X is converted to product Y.
Diagram 1
t = 7 seconds
Diagram 2
t = 5 minutes
Diagram
t = 10 minutes
Diagram 4
t = 5 days
Time, t
X=
Y=
Which statement is correct?
A.
At t = 5 days the rate of the forward reaction is greater than the rate of the backward
reaction.
B.
At t = 7 seconds the reaction has reached completion.
C.
At t = 10 minutes the system has reached a state of equilibrium.
D.
At t = 5 days the rate of the forward reaction is less than the rate of the backward reaction.
(Total 1 mark)
5.
Which statement concerning a chemical reaction at equilibrium is not correct?
A.
The concentrations of reactants and products remain constant.
B.
Equilibrium can be approached from both directions.
C.
The rate of the forward reaction equals the rate of the reverse reaction.
D.
All reaction stops.
(Total 1 mark)
6.
What is the equilibrium constant expression, Kc, for the reaction below?
N2(g) + 2O2(g)
2NO2(g)

NO2 
A.
Kc =
N 2 O 2 
2NO 2 
B.
Kc =
3N 2 O 2 
C.
Kc =
D.
Kc =
NO2 2
N 2 O 2 2
NO2 2
N 2   O 2 2
(Total 1 mark)
160
7.
Sulfur dioxide and oxygen react to form sulfur trioxide according to the equilibrium.
2SO2(g) + O2(g)
2SO3(g)
How is the amount of SO2 and the value of the equilibrium constant for the reaction affected by
an increase in pressure?
A.
The amount of SO3 and the value of the equilibrium constant both increase.
B.
The amount of SO3 and the value of the equilibrium constant both decrease.
C.
The amount of SO3 increases but the value of the equilibrium constant decreases.
D.
The amount of SO3 increases but the value of the equilibrium constant does not change.
(Total 1 mark)
8.
What will happen to the position of equilibrium and the value of the equilibrium constant when
the temperature is increased in the following reaction?
Br2(g) + Cl2(g)
2BrCl(g)
∆H = +14 kJ
A.
B.
C.
D.
Position of equilibrium
Shifts towards the reactants
Shifts towards the reactants
Shifts towards the products
Shifts towards the products
Value of equilibrium constant
Decreases
Increases
Decreases
Increases
(Total 1 mark)
9.
In the reaction below
N2(g) + 3H2(g)
2NH3(g)
∆H = –92 kJ
which of the following changes will increase the amount of ammonia at equilibrium?
I.
Increasing the pressure
II.
Increasing the temperature
III. Adding a catalyst
A.
B.
C.
D.
I only
II only
I and II only
II and III only
(Total 1 mark)
10.
The manufacture of sulfur trioxide can be represented by the equation below.
2SO2(g) + O2(g)
2SO3(g)
∆Hο = –197 kJ mol–1
What happens when a catalyst is added to an equilibrium mixture from this reaction?
A.
The rate of the forward reaction increases and that of the reverse reaction decreases.
B.
The rates of both forward and reverse reactions increase.
C.
The value of ∆Hο increases.
D.
The yield of sulfur trioxide increases.
(Total 1 mark)
161
11.
The reaction between two substances A and B
A+BC+D
has the following rate expression:
rate = k [A]
Draw the graphical representation of:
[A] against time
[B] against time
[B]
[A]
time
time
rate against [A]
rate against [B]
rate
rate
[A]
[B]
(Total 3 marks)
12.
When excess lumps of magnesium carbonate are added to dilute hydrochloric acid the following
reaction takes place.
MgCO3(s) + 2HCl(aq) → MgCl2(aq) + CO2(g) + H2O(l)
(a)
Outline two ways in which the rate of this reaction could be studied. In each case sketch a
graph to show how the value of the chosen variable would change with time.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
162
(4)
(b)
State and explain three ways in which the rate of this reaction could be increased.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(6)
(c)
State and explain whether the total volume of carbon dioxide gas produced would increase,
decrease or stay the same if
(i)
more lumps of magnesium carbonate were used.
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
(2)
(ii)
the experiments were carried out at a higher temperature.
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
(2)
(Total 14 marks)
163
13.
(i)
Draw a graph that shows the distribution of molecular energies in a sample of a gas at two
different temperatures, T1 and T2, such that T2 is greater than T1.
(2)
(ii)
Define the term activation energy.
(1)
(iii)
State and explain the effect of a catalyst on the rate of an endothermic reaction.
(2)
(Total 5 marks)
14.
The table below gives information about the percentage yield of ammonia obtained in the Haber
process under different conditions.
Temperature/°C
Pressure/
atmosphere
10
100
200
300
400
600
(a)
200
50.7
81.7
89.1
89.9
94.6
95.4
300
14.7
52.5
66.7
71.1
79.7
84.2
400
3.9
25.2
38.8
47.1
55.4
65.2
500
1.2
10.6
18.3
24.4
31.9
42.3
From the table, identify which combination of temperature and pressure gives the highest
yield of ammonia.
……………………………………………………………………………………….
(1)
(b)
The equation for the main reaction in the Haber process is
N2(g) + 3H2(g)
∆H is negative
2NH3(g)
Use this information to state and explain the effect on the yield of ammonia of increasing
(i)
pressure: …………………………….………………………………………..
……………………………………………………………..………………….
………………………………………………………………………………..
………………………………………………………………………………..
(2)
164
(ii)
temperature: ………………………………………………………………….
…………………………………………………………………………….….
………………………………………………………………………………..
………………………………………………………………………………..
………………………………………………………………………………..
(2)
(c)
In practice, typical conditions used in the Haber process are a temperature of 500 °C and a
pressure of 200 atmospheres. Explain why these conditions are used rather than those that
give the highest yield.
……………………………………………………………………………………….
……………………………………………………………………………………….
……………………………………………………………………………………….
……………………………………………………………………………………….
(2)
(d)
Write the equilibrium constant expression, Kc, for the production of ammonia.
……………………………………………………………………………………….
……………………………………………………………………………………….
(1)
(Total 8 marks)
15.
(a)
An industrial gas mixture is produced by the catalytic reforming of methane using steam.
CH4(g) + H2O(g)
H = +206 kJ
CO(g) + 3H2(g)
By circling the appropriate letter(s) below, identify the change(s) that would shift the
position of equilibrium to the right.
A
increasing the temperature
B
decreasing the temperature
C
increasing the pressure
D
adding a catalyst
E
decreasing the pressure
F
increasing the concentration of H2
(2)
165
(b)
The following graph represents the change of concentration of reactant and product during
a reaction.
0.7
0.6
0.5
Product
0.4
[reactant] or
[product] / 0.3
mol dm –3
Reactant
0.2
0.1
0.0
0
(i)
10
20
30
Time / s
40
50
60
Calculate the average rate of reaction over the first 15 s, stating the units.
............................................................................................................................
............................................................................................................................
............................................................................................................................
............................................................................................................................
............................................................................................................................
............................................................................................................................
(3)
(ii)
After 19 s the concentrations of the reactant and product do not change. State what
this indicates about the reaction.
............................................................................................................................
............................................................................................................................
(1)
(Total 6 marks)
166
Topic 8: Acids and Bases
8.1 Theories of Acids and Bases
8.1.1 Define acids and bases according to the Bronsted-Lowery and Lewis theories.
8.1.2 Deduce whether or not a species could act as a Bronsted-Lowery and/or a Lewis acid or base.
8.1.3 Deduce the formula of the conjugate acid (or base) of any Bronsted-Lowery base (or acid).
8.2 Properties of Acids and Bases
8.2.1 Outline the characteristic properties of acids and bases in aqueous solution.
8.3 Strong and Weak acids and Bases
8.3.1 Distinguish between strong and weak acids and bases in terms of the extent of dissociation,
reaction with water and electrical conductivity.
8.3.2 State whether a given acid or base is strong or weak.
8.3.3 Distinguish between strong and weak acids and bases, and determine the relative strengths of
acids and bases using experimental data.
8.4 The pH Scale
8.4.1 Distinguish between aqueous solutions that are acidic, neutral or alkaline using the pH scale.
8.4.2 Identify which of two or more aqueous solutions is more acidic or alkaline using pH values.
8.4.3 State that each change of one pH unit represents a 10- fold change in the hydrogen ion
concentration [H+(aq)].
8.4.4 Deduce changes in [H+(aq)] when the pH of a solution changes by more than one pH unit.
167
Topic 8 NOTES
168
Acids and Bases Worksheet
1.
Properties of Acids
a.
Acids have a pH value less than ___________.
b.
Acids have a ________ taste.
c.
Acids react with metal to form ____________ gas.
d.
Acids react with carbonate compounds to form ______________ gas.
e.
When a strong acid reacts with a strong base it forms ____________ and
____________.
f.
Litmus turns ___________ in the presence of an acid solution.
2.
Properties of Bases
a.
Bases have a pH value greater than ____________.
b.
Bases have a ____________ taste.
c.
Bases feel ______________ .
d.
Litmus turns __________ in the presence of a basic solution.
e.
Phenolphthalein turns _____________ in the presence of a basic solution.
3.
Definitions
a.
Define an acid and base according to Bronsted-Lowry.
4.
b.
How does a conjugate acid/base pair differ from one another?
c.
Water undergoes "autoionization." What does this mean? Write a chemical equation
showing this process.
d.
What is a hydronium ion? Write its formula.
e.
Define "electrolyte". Give an example of a nonelectrolyte, strong electrolyte and a
weak electrolyte.
f.
Name the three common strong acids.
g.
What is the difference between a strong and weak acid?
The pH scale.
a.
Write the mathematical equation for pH.
b.
What is the pH of a solution with a [H+] concentration of 10-5?
c.
What is the pH of a solution with a [H+] concentration of 10-11?
d.
What is the [H+] of a solution with a pH of 8?
e.
A solution has a pH of 2. Is it basic, acidic, or neutral?
f.
A solution has a pH of 12. Is it basic, acidic, or neutral?
g.
A solution has a pH of 3. A second solution has a pH of 5. Which is more acidic?
h.
In part "g" how many times more acidic (in terms of [H+] is the most acidic solution
compared to the least acidic solution?
169
Brønsted-Lowry Acids and Bases
1. Label the Brønsted-Lowry acids and bases in the following reactions and show the direction of
proton transfer.
H+
H2O + Cl- ↔ OH- + HCl
Acid base
base
acid
a. NH3 + H2O ↔ NH4+ + OHb. H2O + H2O ↔ H3O+ + OHc. HCO3- + HSO4- ↔ H2CO3 + SO42d. H3PO4 + F- ↔ H2PO4- + HF
e. NO2- + H2O ↔ HNO2 + OH-
5. Fill in the table of the following acid base conjugate pairs.
Acid
H2SO4
Base
Acid
HPO42-
NO3-
Base
NH3
SO32-
H2O
HCO3-
H2O
6. Name the following acids.
a. HCl
b. H2SO4
c. HBr
d.
e.
f.
HNO3
HNO2
H3PO4
7. Write formulas for the following acids.
a. hydroiodic acid
b. sulfurous acid
c. carbonic acid
d.
e.
f.
acetic acid
hydrosulfuric acid
chloric acid
170
pH and pOH
1.
Define pH.
2.
Define pOH.
3.
pH + pOH = __________
4.
If the pH of a solution is 3, the solution is
5.
If the pOH of a solution is 5, the solution is
6.
Kw is called the "ionization constant" for water. What is the equation for Kw and what is its
numerical value at 25C?
7.
In the symbol [H+], what do the brackets represent?
8.
If the [H+] of a solution is 10-8, what is the [OH-]?
9.
If the pOH is 9, what is the [OH-]?
10.
If the pH of a solution is 3, what is the pOH? the [H+]? the [OH-]?
11.
In a strong acid solution, the [H+] concentration is (greater than, equal to, less than) the
original concentration of the acid.
12.
In a weak acid solution, the [H+] concentration is (greater than, equal to, less than) the
original concentration of the acid.
13.
If 1.0% of a 0.10 M solution of weak acid HF ionizes, what is the concentration of the [H+]?
What is the pH?
14.
5% of a 0.20 M weak base ionizes. What it the [OH-] of the solution? What is the pOH and
the pH?
15.
What is the pH of a 10-2 M solution of HCl? What is the pOH?
(acidic/basic/neutral)__ .
171
(acidic/basic/neutral)__ .
Lewis Acids and Bases
1. Define an acid according to Lewis theory.
2. Define a base according to Lewis theory.
3. Define dative bond.
4. Draw Lewis dot diagrams for NH3 and BH3.
5. Using the diagrams you drew in #4, use an arrow to show the electron pair that will be used to
make the dative bond. (Draw the arrow from the base to the acid to show the electrons
involved.)
6. Identify the Lewis acids and bases in the reactants of the following reactions:
a. Ag+ + 2NH3 → [Ag(NH3)2]+
b. B(OH)3 + H2O → B(OH)4- + H+
c. Fe3+ + 6CN- → [Fe(CN)6]3-
7. Boron trifluoride, BF3, and ammonia, NH3, react to produce a product. A dative bond is formed
between the boron atom on BF3 and the nitrogen atom on NH3. Write the equation for this
reaction, using Lewis electron-dot formulas. Label the Lewis acid and the Lewis base. Determine
how many grams of product are formed when 10.0g of each reactant are placed in a reaction
vessel, assuming that the reaction goes to completion.
172
pH and pOH
pH is an indication of how acidic or basic a solution is. pH = -log[H+]. A logarithm (log) is just the
exponent to which ten is raised to obtain a given value. For example if [H+] = 10-5, then the log of
10-5 = -5 and the pH would be 5. pOH = -log[OH-].
pH + pOH = 14
[H+][OH-] = 10-14
Using the relationships above, fill in the following table.
[H+]
10-4
[OH-]
pH
pOH
8
-4
10
4
-11
10
7
-2
10
1
10-6
Calculate the pH of the following solutions.
1. 0.01 M HCl
2. 0.001 M NaOH
3. 0.10 M HNO3
4. 0.050 M Ca(OH)2
5. 0.01 M HF (assume 1% ionization)
6. 0.0001 M HC2H3O2 (assume 10% ionization)
7. 2.00 M NH4OH (assume 5% ionization)
173
Acidic or basic
EXTEND- Titration to Determine the Molarity of Vinegar
Procedure
Student’s trial run–
Put approximately 10 mL of vinegar (RECORD EXACT AMOUNT) and 15 mL of distilled water into an
Erlenmeyer flask. Add 3 drops of phenolphthalein indicator. Titrate to the color change with the
standardized NaOH solution. Once the solution has changed color, stop the titration. Wait for all of the lab
groups to complete their first run. Compare your solution with the class.
Remember: Do not let NaOH go below the 50mL mark on the buret
Data Collection –
Collect data for three additional trials. Record data and observations in the table below.
Data
REMEMBER UNCERTAINTIES!!
Volume of
Initial volume of
Trial
Vinegar (mL)
0.500 M NaOH*
(mL)
Trial Run
1
2
3
*The NaOH solution is 0.500 ± 0.005 mol dm-3
Qualitative observations made during lab:
174
Final volume of
0.500 M NaOH*
(mL)
Volume of 0.500 M
NaOH* used
(mL)
Use your data from the activity above to answer the following questions.
1. Calculate the number of moles of NaOH added in each trial. Show a sample calculation and record
all results in the table.
2. At the equivalence point in any titration, how do the moles of base compare to the moles of acid?
3. Based on your answer to number 2, record the number of moles of acid present in each trial in the
table below.
4. Calculate the molarity of vinegar for all 3 trials. Show a sample calculation and record all results in
the table.
5. Report the average molarity of vinegar.
Trial
Moles of NaOH
Moles of vinegar
Molarity of vinegar
(M or mol/L or mol dm-3)
1
2
3
Average
6. Calculate the % error for your data. (Obtain theoretical value from teacher)
7. Describe at least two possible sources of error. Suggest one way that you could improve your lab
techniques to lower your percent error the next time you do this lab.
8. One method of determining molarities during titration is to use the equation M1V1 = M2V2, where
M1V1 equals moles of acid and M2V2 equals moles of base. Using your data from Trial 1 solve for the
molarity of the vinegar, M1. Compare this value to your results in the table above.
9. Write a balanced chemical equation for the neutralization of sodium hydroxide and vinegar,
HC2H3O2.
175
EVALUATE
Write balanced equations for the following neutralization reactions:
1. H3PO4 + KOH →
2. HI + Mg(OH)2→
Titrations
1. What is the formula for determining molarity using titration?
2. If we neutralize 15.0 cm3 of 1.00 mol dm-3 HCl with 12.5 cm3 of NaOH, what is the molarity of the
NaOH?
3. What volume of a 0.50 M HCl solution do we need to neutralize 35 mL of 0.80 M NaOH?
4. What is the molarity of KOH if we used 50.0 mL of 0.100 mol dm-3 HNO3 to neutralize 25.0 mL of the
KOH?
5. What is the purpose of an indicator solution in a titration?
176
Electrolytes
Using a conductivity probe or meter, test the conductivity of the following substances in water
solution and categorize into one of the following categories (your teacher may do this for you).
Make sure you test pure water to see what its conductivity is before you test the solutions.
Substance (0.1 M
solution)
Excellent Conductor
Poor Conductor
Nonconductor
NaCl
Sugar (C12H22O11)
HCl
Vinegar (HC2H3O2)
NaOH
NH3
Alcohol (C2H5OH)
Based on the results above, answer the following questions.
Salts, like NaCl, are strong/weak/non electrolytes. (circle one)
HCl is a strong/weak acid and is a strong/weak electrolyte.
Vinegar is a strong/weak acid and is a strong/weak electrolyte.
NaOH is a strong/weak base and is a strong/weak electrolyte.
NH3 is a strong/weak base and is a strong/weak electrolyte.
Sugar and alcohols are strong/weak/non electrolytes.
Sports drinks, like Gatorade, contain electrolytes. What ingredient do these drinks contain that
makes them a good source of electrolytes?
177
178
Study Guide: Topic 8 Acids and Bases
1.
Which is a conjugate acid-base pair in the following reaction?
HNO3 + H2SO4
A.
B.
C.
D.
H2NO3+ + HSO4–
HNO3 and H2SO4
HNO3 and H2NO3+
HNO3 and HSO4–
H2NO3+ and HSO4–
(Total 1 mark)
2.
Which equation represents an acid-base reaction according to the Lewis theory but not the
Brønsted-Lowry theory?
A.
NH3 + HCl
NH4Cl
H3O + OH–
B.
2H2O
C.
NaOH + HCl
D.
CrCl3 + 6NH3
+
NaCl + H2O
[Cr(NH3)6]3+ + 3Cl–
(Total 1 mark)
3.
In which reaction is H2PO4–(aq) acting as a Brønsted-Lowry base?
A.
B.
C.
D.
H2PO4–(aq) + NH3(aq) → HPO42–(aq) + NH4+(aq)
H2PO4–(aq) + OH–(aq) → HPO42–(aq) + H2O(l)
H2PO4–(aq) + C2H5NH2(aq) → HPO42–(aq) + C2H5NH3+(aq)
H2PO4–(aq) + CH3COOH(aq) → H3PO4(aq) + CH3COO–(aq)
(Total 1 mark)
4.
An aqueous solution of which of the following reacts with magnesium metal?
A.
Ammonia
B.
Hydrogen chloride
C.
Potassium hydroxide
D.
Sodium hydrogencarbonate
(Total 1 mark)
5.
Which acids are strong?
I.
HCl(aq)
II.
HNO3(aq)
III. H2SO4(aq)
A.
B.
C.
D.
I and II only
I and III only
II and III only
I, II and III
(Total 1 mark)
179
6.
Solutions of hydrochloric acid (HCl(aq)) and ethanoic acid (CH3COOH(aq)) of the same
concentration reacted completely with 5.0 g of calcium carbonate in separate containers. Which
statement is correct?
A.
B.
C.
D.
CH3COOH(aq) reacted slower because it has a lower pH than HCl(aq).
A smaller volume of CO2(g) was produced with CH3COOH(aq) than with HCl(aq).
A greater volume of CO2(g) was produced with CH3COOH(aq) than with HCl(aq).
The same volume of CO2(g) was produced with both CH3COOH(aq) and HCl(aq).
(Total 1 mark)
7.
When the following 1.0 mol dm–3 solutions are listed in increasing order of pH (lowest first),
what is the correct order?
A.
B.
C.
D.
HNO3  H2 CO3  NH3  Ba(OH)2
NH3  Ba (OH)2  H2 CO3  HNO3
Ba (OH)2  H2 CO3  NH3  HNO3
HNO3  H2 CO3  Ba (OH)2  NH3
(Total 1 mark)
8.
Define the terms Brønsted-Lowry acid and Lewis acid. For each type of acid, identify one
example other than water and write an equation to illustrate the definition.
(Total 5 marks)
9.
Identify one example of a strong acid and one example of a weak acid. Outline three different
methods to distinguish between equimolar solutions of these acids in the laboratory. State how
the results would differ for each acid.
(Total 5 marks)
180
10.
Vinegar has a pH of approximately 3 and some detergents have a pH of approximately 8. State
and explain which of these has the higher concentration of H+ and by what factor.
(Total 1 mark)
11.
The pH values of solutions of three organic acids of the same concentration were measured.
acid X
acid Y
acid Z
(i)
pH = 5
pH = 2
pH = 3
Identify which solution is the least acidic.
(1)
(ii)
Deduce how the [H+] values compare in solutions of acids Y and Z.
(2)
(iii)
Arrange the solutions of the three acids in decreasing order of electrical conductivity,
starting with the greatest conductivity, giving a reason for your choice.
(2)
(Total 5 marks)
181
Topic 9: Oxidation and Reduction
9.1 Introduction to Oxidation and Reduction
9.1.1 Define oxidation and reduction in terms of electron loss and gain.
9.1.2 Deduce the oxidation number of an element in a compound.
9.1.3 State the names of compounds using oxidation numbers.
9.1.4 Deduce whether an element undergoes oxidation or reduction in reactions using oxidation
numbers.
9.2 Redox Equations
9.2.1 Deduce simple oxidation and reduction half-equations given the species involved in a redox
reaction.
9.2.2 Deduce redox equations using half-equations.
9.2.3 Define the terms oxidizing agent and reducing agent.
9.2.4 Identify the oxidizing and reducing agents in redox equations.
9.3 Reactivity
9.3.1 Deduce a reactivity series based on the chemical behavior of a group of oxidizing and reducing
agents.
9.3.2 Deduce the feasibility of a redox reaction from a given reactivity series.
9.4 Voltaic Cells
9.4.1 Explain how a redox reaction is used to produce electricity in a voltaic cell.
9.4.2 State that oxidation occurs at the negative electrode (anode) and reduction occurs at the
positive electrode (cathode).
9.5 Electrolytic Cells
9.5.1 Describe, using a diagram, the essential components of an electrolytic cell.
9.5.2 State the oxidation occurs at the positive electrode (anode) and reduction occurs at the negative
electrode (cathode).
9.5.3 Describe how current is conducted in an electrolytic cell.
9.5.4 Deduce the products of the electrolysis of a molten salt.
182
Topic 9 NOTES
183
Oxidation - Reduction
Rules for determining oxidation numbers:
1. Pure elements have an oxidation number = 0
2. The oxidation number of a monatomic ion = charge of the ion
3. Alkali metals in compounds have an oxidation number = +1
4. Alkaline earth metals in compounds have an oxidation number = +2
5. Fluorine in compounds has an oxidation number = -1
6. Oxygen in compounds has an oxidation number = -2 except in peroxides (-1)
7. Hydrogen has an oxidation number = +1 except in metal hydrides (-1)
8. The sum of all oxidation numbers in a compound = 0
9. The sum of all oxidation numbers in a polyatomic ion = charge of the ion
Assign oxidation numbers to each element in the following.
1. K
2. RbCl
3. Na2O
4. NH3
5. CO2
6. N2O5
7. N2O
8. N2O3
9. CaCO3
10. KNO3
11. Na2Cr2O7
12. KMnO4
13. H3PO4
14. H2O2
15. FeCl3
16. Li2SO4
Define each of the following terms:
17. Oxidation
18. Reduction
19. Oxidizing Agent
20. Reducing Agent
In the following equations, identify the substance oxidized, the substance reduced, the oxidizing
agent, and the reducing agent.
Ox.
21. CH4 + O2 → CO2 + H2O
22. Zn + AgNO3 → Zn(NO3)2 + Ag
23. Al + Cl2 → AlCl3
24. Na + H2O → H2 + NaOH
184
Red.
OA
RA
Writing and Balancing Half Reactions
For each of the reactions below, split the equation into half reactions and balance using the method
demonstrated by your teacher.
1.
Zn(s) + Al3+(aq) → Al(s) + Zn2+(aq)
2.
MnO4-(aq) + Fe(s) → Fe2+(aq) + Mn2+(aq)
(occurs in acidic solution)
3.
Cr(s) + NO3-(aq) → Cr3+(aq) + NO(g)
(occurs in acidic solution)
4.
NO2-(aq) + Al(s) → NH3(g) + AlO2-(aq)
(occurs in basic solution)
5.
CN-(aq) + MnO4-(aq) → CNO-(aq) + MnO2(s)
(occurs in basic solution)
185
Activity Series of Metals and Nonmetals
Metals have different reactivities. For example, the alkali metals all react with water, most metals
react with acid, while a small number react with neither. The metals can be placed in a reactivity
series that helps identify whether one metal is capable of replacing another in a single replacement
reaction.
Reactivity Series of metals:
Li, K, Ca, Na, Mg, Al, Zn, Fe, Ni, Sn, Pb, Cu, Hg, Ag, Au
Most reactive
least reactive
Using the series it is possible to determine if aluminum is capable of replacing copper in the
compound CuCl2. (Al + CuCl2 → ? ) Since Al is higher in the series it will be capable of
replacing the Cu ions in CuCl2. Since Al is higher than Cu we say that it is more easily oxidized and
is a better reducing agent than Cu.
Predict if the following reactions will occur. For those that do, complete and balance the equation,
and label the oxidizing agent. For those that do not, write “no reaction.”
1.
Zn + FeCl3 →
2.
Ag + Cu(NO3)2 →
3.
Na + KCl →
4.
Ca + Mg(NO3)2 →
5.
Al + MgF2 →
6.
Zn + Pb(NO3)2 →
The reactivity of the halogens is a matter of periodic trends, fluorine being the most reactive and
iodine being the least reactive as you travel down the family. Predict whether or not the following
reactions will occur. Complete and balance those that do.
7.
Cl2 + NaI →
8.
Br2 + CaF2 →
9.
F2 + AlCl3 →
10.
I2 + KBr →
186
Electrochemical Cells
An electrochemical, galvanic, or voltaic cell converts chemical potential energy into
electricity. It relies on the difference in activity between the two metals that make up the cathode
and anode or the two metal ions in the solutions.
The electrons flow as a result of a spontaneous reaction. The spontaneity can be somewhat
predicted based on the activity series. Species higher up on the chart will act as the anode while
species lower on the chart will act as cathodes. A more detailed prediction of the output voltage can
be calculated using Standard Reduction Potentials (which we are not going to cover in this class).
Factors that affect the voltage output of a cell are the activity difference in the two metals
(emf), concentration, and temperature. A quantitative description is given by the Nernst equation
(also something we’re going to skip).
At the anode, oxidation takes place, the neutral metal atoms lose electrons which flow
through the wire to the light bulb causing it to light up. The positive ions that form as a result of
oxidation, dissolve in solution.
At the cathode, reduction takes place, the positive ions in solution gain electrons that have
flowed from the anode. The ions become neutral atoms and precipitate out of solution onto the
cathode.
Sometimes an inert but conducting material is used as an electrode when one or both of the
species is non-conducting. (For example, if the following is one of the half-reactions in the cell:
MnO4- + 5e- + 8H+  Mn2+ + 4H2O
Neither MnO4- nor Mn2+ can serve as electrodes since they are not solid so a Platinum electrode
could be used)
Salt bridge to allow
flow of ions so there is
no charge build up in
either solution. Charge
build-up would shift the
equilibrium and the
reaction would cease
e-
Light or
voltmeter
ee-
e-
Anode:
Oxidation,
Negative
electrode
Cathode:
Reduction,
Positive
electrode
Electrolyte: allows for charges to
move from one electrode to the
other.
187
Electrochemical Cells
V
_
+
Given the following redox reaction that takes place in the electrochemical cell above:
3 Cu2+ (aq) + 2 Al (s) → 3 Cu (s) + 2 Al3+ (aq)
1.
2.
3.
4.
Which species is being oxidized?
Write the equation for the oxidation half reaction.
Oxidation takes place at which electrode?
What is the charge of this electrode?
5.
6.
7.
8.
Which species is being reduced?
Write the equation for the reduction half reaction.
Reduction takes place at which electrode?
What is the charge of this electrode?
9. Electrons in the external circuit flow from which electrode to which electrode?
10. Towards which electrode do positive ions in the salt bridge flow?
11. Electrochemical cells convert _____________ energy into _____________ energy.
11. On the diagram above, label the following:
salt bridge, anode, cathode, Cu, Al, Cu2+ solution, Al3+ solution, direction of electron flow in the
external circuit, direction of ion flow in the salt bridge
188
Electrolytic Cells
Electrolytic cells are the opposite of galvanic cells:
They convert electrical energy into chemical potential energy.
Electricity must be supplied in order to force a non-spontaneous reaction to occur. The spontaneity
can be somewhat predicted based on the activity series.
Current flows from the species being oxidized to the species being reduced.
Major Differences:
 Direction of electron flow changes
 Oxidation still occurs at the anode and reduction still occurs at the cathode. Notice that
cathode is now the negative electrode, it attracts cations
 Notice that the anode is now the positive electrode, it attracts anions
 Notice that there is a power supply (battery) instead of a light bulb or voltmeter
e-
Power
Source
ee-
Salt bridge to allow
flow of ions so there is
no charge build up in
either solution.
e-
Cathode:
Reduction
occurs here
Negative
Electrode
Cations
attracted to
anode
Anode:
Oxidation
occurs here
Positive
Electrode
Anions are
attracted to
cathode
Electrolyte: allows for charges to
move from one electrode to the
other.
Examples of electrolysis in action are:
 making hydrogen and oxygen gas from water
o 2H2O  2H2 + O2
o 2H+ + 2e-  H2 (Reduction at negative electrode/cathode)
o O2-  O2 + 4 e- (Oxidation at positive electrode/anode)


Making sodium and chlorine from molten salt
o NaCl(l)  Na + Cl2
o Cathode: Na+ + 1 e-  Na
o Anode: 2Cl- + 2 e-  Cl2
Electroplating precious metals onto less expensive metals
189
Electrolytic Cells
+
-
1. Molten potassium iodide is placed into the reaction vessel in the above electrolytic cell. What
are the two reactants in this reaction?
2. Write the half reaction for the oxidation that will take place.
3. At which electrode will oxidation take place?
4. What is the charge of this electrode?
5. Write the half reaction for the reduction that will take place.
6. At which electrode will reduction take place?
7. What is the charge of this electrode?
8. Label the anode and cathode on the diagram above showing what is produced at each electrode.
9. Write the overall balanced equation for the reaction.
10. Electrolytic cells convert _____________ energy into ______________ energy.
11. State two ways that electrochemical and electrolytic cells are similar.
12. State two ways that electrochemical and electrolytic cells are different.
190
Topic 9 NOTES
191
Study Guide: Topic 9 Oxidation and Reduction
1.
Which are examples of reduction?
A.
B.
C.
D.
I.
Fe3+ becomes Fe2+
II.
Cl– becomes Cl2
III.
CrO3 becomes Cr3+
I and II only
I and III only
II and III only
I, II and III
(Total 1 mark)
2.
In which change does nitrogen undergo oxidation?
A.
NO2  N2O4
B.
NO3–  NO2
C.
N2O5 NO3–
D.
NH3  N2
(Total 1 mark)
3.
What are the oxidation numbers of the elements in sulfuric acid, H2SO4?
Hydrogen
Sulfur
Oxygen
A.
+1
+6
–2
B.
+1
+4
–2
C.
+2
+1
+4
D.
+2
+6
–8
(Total 1 mark)
4.
Consider the following reaction:
H2SO3(aq) + Sn4+(aq) + H2O(l) → Sn2+(aq) + HSO4–(aq) + 3H+(aq)
Which statement is correct?
A.
H2SO3 is the reducing agent because it undergoes reduction.
B.
H2SO3 is the reducing agent because it undergoes oxidation.
C.
Sn4+ is the oxidizing agent because it undergoes oxidation.
D.
Sn4+ is the reducing agent because it undergoes oxidation.
(Total 1 mark)
192
5.
Which statement is correct about an oxidizing agent in a chemical reaction?
A.
It reacts with oxygen.
B.
It reacts with H+ ions.
C.
It loses electrons.
D.
It undergoes reduction.
(Total 1 mark)
6.
A voltaic cell is made from magnesium and iron half-cells. Magnesium is a more reactive metal
than iron. Which statement is correct when the cell produces electricity?
A.
Electrons are lost from magnesium atoms.
B.
The concentration of Fe2+ ions increases.
C.
Electrons flow from the iron half-cell to the magnesium half-cell.
D.
Negative ions flow through the salt bridge from the magnesium half-cell to the iron halfcell.
(Total 1 mark)
7.
What process occurs at the cathode in a voltaic cell and at the anode in an electrolytic cell?
A.
Cathode of
voltaic cell
Oxidation
Anode of
Electrolytic cell
Reduction
B.
Oxidation
Oxidation
C.
Reduction
Oxidation
D.
Reduction
Reduction
(Total 1 mark)
8.
Which statement is correct for the electrolysis of a molten salt?
A.
Positive ions move toward the positive electrode.
B.
A gas is produced at the negative electrode.
C.
Only electrons move in the electrolyte.
D.
Both positive and negative ions move toward electrodes.
(Total 1 mark)
9.
Iron in food, in the form of Fe3+, reacts with ascorbic acid (vitamin C), C6H8O6, to form
dehydroascorbic acid, C6H6O6.
(i)
Write an ionic half-equation to show the conversion of ascorbic acid to
dehydroascorbic acid in aqueous solution.
.........................................................................................................................
.........................................................................................................................
(1)
(ii)
In the other ionic half-equation Fe3+ is converted to Fe2+. Deduce the overall
equation for the reaction between C6H8O6 and Fe3+.
.........................................................................................................................
.........................................................................................................................
(1)
(Total 2 marks)
193
10.
Deduce the change in oxidation number of chromium in the below reaction. State with a reason
whether the chromium has been oxidized or reduced.
Cr2O72 + 14H+ + 6Fe2+  2Cr3+ + 6Fe3+ + 7H2O
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
(Total 2 marks)
11.
A part of the reactivity series of metals, in order of decreasing reactivity, is shown below.
magnesium
zinc
iron
lead
copper
silver
If a piece of copper metal were placed in separate solutions of silver nitrate and zinc nitrate
(i)
determine which solution would undergo reaction.
……………………………………………………………………………………………
(1)
(ii)
identify the type of chemical change taking place in the copper and write the
half-equation for this change.
……………………………………………………………………………………………
……………………………………………………………………………………………
(2)
(iii)
state, giving a reason, what visible change would take place in the solutions.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(2)
(Total 5 marks)
194
12.
A current is passed through molten sodium chloride. Identify the substance formed at each
electrode and write an equation to represent the formation of each substance. Determine the mole
ratio in which the substances are formed.
……………………………………………………………………………………………………
……………………………………………………………………………………………………
……………………………………………………………………………………………………
……………………………………………………………………………………………………
……………………………………………………………………………………………………
……………………………………………………………………………………………………
(Total 5 marks)
195
196
197
198
199