Transcript Syllabus Chemistry 101 Fall 2010 Professor: Dr. Earle G. Stone

Syllabus Chemistry 101 Fall 2010
Sec. 501 (MWF 8:00-8:50) Sec 505 (MWF 12:40-1:30) RM 100 HELD
Professor: Dr. Earle G. Stone
Office: Room 123E Heldenfels (HELD)
Telephone: 845-3010 (no voice mail) or leave a message at 845-2356
email: [email protected]
(put CHEM 101-Sec. # + subject in subject line of your email)
Office Hours: HELD 123E: Tue. And Thur. 8:00-10:50 AM
I.A. TBA
S.I. Leader: TBA
CHEM 101 and 102 are the first-year chemistry sequence in the core
curriculum. The lecture component of Chemistry 101 covers stoichiometry,
atomic and molecular structure, chemical bonding, basic acid/base chemistry,
solution chemistry, and the gas laws.
In addition this class will cover foundation work in preparation for upper
division course work in inorganic and organic nomenclature and structure,
including discussion of major biological polymers; carbohydrates, sugars, fats,
oils, and lipids, as well as peptides and proteins.
More importantly as most in this class are pre-something, it is the goal of
my lecture section to help you develop the skill set to not only successfully
complete your undergraduate degree but to also prepare for professional school
entrance exams and enable you to succeed in your choice of professional school.
All
College
BIMS
Science
GEST and other
Agriculture other
Engineering
Ag BICH, NUSC, GENE
Education
Liberal Arts
Geosciences
Architecture
Business and other
501
College
Science
GEST and other
BIMS
Engineering
Agriculture other
Ag BICH, NUSC, GENE
Liberal Arts
Education
Geosciences
Architecture
Business and other
The instruction was applicable to my degree
478
128
120
75
42
28
22
18
18
11
9
7
27%
25%
16%
9%
6%
5%
4%
4%
2%
2%
1%
228
75
39
25
20
19
16
10
9
7
6
2
33%
17%
11%
9%
8%
7%
4%
4%
3%
3%
1%
2014
2013
2012
2011
graduate
379
73
21
5
0
478
80% Pre-med, Pre-vet, Pre-Pharm, Pre-dent, Nursing, Physician
Asst. Physical Therapist, Scientist, Eng.
20% degree requirement and thought this would be a fun class
505
College
250
BIMS
103 41%
Science
45
18%
GEST and other
36
14%
Agriculture other
23
9%
Education
9
4%
Engineering
8
3%
Liberal Arts
8
3%
Ag BICH, NUSC, GENE
6
2%
Business and other
5
2%
Geosciences
4
2%
Architecture
3
1%
Kotz and Treichel 6th or 7th ed.
TEXTBOOKS
Chang
New Hardbound ~$200
paperback ~ $150
Solution Manual ~ $60
Online Tutor ~ $45
Total ~$250 - 305
Ebook $45 per semester
Includes
Text
Solution manual
Online tutorial
Optional Used Texts
$5 - $50 + shipping
no sell back
Helpful
Online Dictionary of Chemistry
Useful
As A Second Language
General Chemistry I and
Organic Chemistry I
(There are O-chem II and Physics
books in this series if you find these
useful and will have to take those classes.
Chang’s Essentials
http://slc.tamu.edu/
Tutoring
Supplemental Instruction
Courses
Texas Success Initiative
About Us
Contact Us
118 Hotard North of Sbisa, between
Neeley Hall and the Northside Post Office
(979) 845-2724
The Student Learning Center has won the 2008 National College Learning Center Association
Frank L. Christ Outstanding Learning Center Award!
The award recognizes the center's commitment to supporting and strengthening the
Academic experience of students at Texas A&M University by providing a variety of programs and
services that promote retention and success. Read more...
The Student Learning Center provides Supplemental Instruction and tutoring free of charge to all
Texas A&M University students. The SLC oversees the STLC courses (formerly CAEN), which teach
students how to improve their study skills and prepare for the job market. The SLC manages
Developmental programs for students who have not yet passed the assessment tests required by the state.
Study Tips
•General
•Time Management
•Reading Textbooks
•Setting Goals
•Preparing for Exams
•Success Tips from Fellow Aggies
Tutoring
During the Fall 2008 semester, drop-in tutoring will be offered Sunday nights 5-8pm and Monday through Thursday
nights from 5-10pm. Tutoring will begin on Monday, September 1st. Tutor Zones are currently planned for Studio 12 of The
Commons. Look for our table and tutors and just ask for help! See the schedule...
Drop-in tutoring is available for most lower level math and science courses on the first floor of Hotard Hall. Tutors
are also available to help out with many other courses. If you need help in a particular course and would like to check to see if
a tutor is available for that course, you can contact our tutor coordinator, Linda Callen, at 845-2724.
Grading:
Your grade will be based on
•Four one-hour examinations (each worth 200 points)
•A final examination (400 points)
There are no bonuses or soft points
There will be no extra credit
Major Examination Schedule Fall 2010:
Fri. Sept. 17 Major Exam No.1
Wed. Oct. 6 Major Exam No.2
Mon. Nov. 1 Major Exam No.3
Fri. Dec. 3 Major Exam No. 4
Fri. Dec. 10 Section 501 Final Exam 10:00 to 12:00
Mon. Dec. 13 Section 505 Final Exam 10:30 to 12:30
How grades are determined
The way the real world works
Individual Mastery compared
to a large population
What you are used to and I will
report
1) Raw scores are determined.
2) Small dataset is bootstrapped to
500,000+ exams
3) Individual scores are normalized.
4) Normalized scores are transformed.
5) Letter grades are assigned
Sum of points assigned to correct responses
A context-free evaluation of relative performance
An absolute score is assigned to a defined scale
>89.501 A
>59.501 D
>79.501 B
<59.501 F
>69.501 C
Raw % Score Distribution
20
Results of Bootstrap
15
5000Repetitions
549Secs
Sheet1!$B$112
Sheet1!$B$113
Average
117.136Average
35.319
SD
3.3410SD
1.8824
Max
131.739Max
41.707
Min
104.090Min
29.000
Simulation
Stats
10
5
0
Probability Density
1.E-02
8.E-03
6.E-03
4.E-03
2.E-03
0.E+00
30
50
100
150
Sheet1!$B$112
Sheet1!$B$113
Cumulative Probability
1
0.9
0.8
0.7
0.6
0.5
0.4
0.3
0.2
0.1
0
1.E-02
0
Histogram of Sheet1!$B$112
And Sheet1!$B$113
100%
90%
80%
70%
60%
50%
40%
30%
20%
10%
0%
501
200
T-Score Distribution - Course
25
20
15
10
5
0
60%
70%
80%
90%
100%
29
Student X
Student Y
Student X
Exam 2
Student Y
Student X
Exam 3
Student Y
Student X
Exam 4
Student Y
Final Student X
Exam Student Y
Avg Exam Student X
Diff
Student Y
Student X
Avg Raw
Student Y
Avg T Student X
Score Student Y
Course Student X
Grade Student Y
Exam 1
79
129
Form Diff Raw
T
Grade
A
1 146.00 89.19
B
C
4 138.00 85.08
B
B
2 130.00 84.50
B
C
4 131.00 80.41
B
A
2 75.00 77.61
C
C
4 87.00 79.74
B
B
1 25.00 57.83
F
D
4 108.00 79.57
B
B
1 234.00 84.54
B
C
3 230.00 76.62
C
1.4
3.8
50.8%
57.8%
79.7%
79.7%
B
B
Overall
T-score
A
B
C
501 15.5% 37.0% 38.8%
503 13.3% 37.6% 37.3%
All
14.3% 37.3% 38.0%
16.0% 35.0% 35.0%
Exam 1
T-score
A
B
C
501 18.5% 33.8% 30.2%
503 17.6% 38.3% 30.1%
All
18.0% 36.2% 30.1%
16.0% 35.0% 35.0%
Exam 2
T-score
A
B
C
501 20.1% 29.7% 34.7%
503 16.5% 39.2% 27.5%
All
18.1% 34.8% 30.8%
16.0% 35.0% 35.0%
Exam 3
T-score
A
B
C
501 17.4% 33.5% 36.2%
503 17.0% 31.6% 36.4%
All
17.2% 32.5% 36.3%
16.0% 35.0% 35.0%
Exam 4
T-score
A
B
C
501 19.5% 34.9% 32.6%
503 19.3% 37.3% 26.5%
All
19.4% 36.2% 29.3%
16.0% 35.0% 35.0%
Final
T-score
A
B
C
501 19.8% 29.5% 35.9%
503 18.7% 31.3% 35.7%
All
19.2% 30.5% 35.8%
16.0% 35.0% 35.0%
D
6.8%
10.6%
8.9%
14.0%
F
1.8%
1.2%
1.5%
2.0%
D
15.8%
10.2%
12.8%
14.0%
F
1.8%
3.9%
2.9%
2.0%
D
12.8%
13.3%
13.1%
14.0%
F
2.7%
3.5%
3.2%
2.0%
D
12.8%
14.6%
13.8%
14.0%
F
0.0%
0.4%
0.2%
2.0%
D
9.8%
13.7%
11.9%
14.0%
F
3.3%
3.2%
3.2%
2.0%
D
14.3%
13.1%
13.6%
14.0%
F
0.5%
1.2%
0.9%
2.0%
Maximum
avg raw score
Cumulative
extra raw needed on every
Points after current avg desired avg score points remaining exam
grade
grade
Last Exam
needed for to earn desired
Taken
50.000%
79.501% desired grade
grade
200
100
400
200
318
118
170.80
600
300
477
177
188.50
800
400
636
236
218.00
1200
900
954
54
277.01
85.401%
94.252%
109.002%
138.503%
Maximum
avg raw score
Cumulative
extra raw needed on every
Points after current avg desired avg score points remaining exam
grade
grade
Last Exam
needed for to earn desired
Taken
90.000%
79.501% desired grade
grade
200
180
400
360
318
-42
154.80
600
540
477
-63
148.50
800
720
636
-84
138.00
1200
1080
954
-126
117.01
77.401%
74.252%
69.002%
58.503%
Day 1
And we begin:
The mere formulation of a problem is far more often essential than
its solution, which may be merely a matter of mathematical or
experimental skill. To raise new questions, new possibilities, to
regard old problems from a new angle requires creative imagination
and marks real advances in science.
~Albert Einstein
Problem - A situation that presents
difficulty, uncertainty, or perplexity:
Question - A request for data:
inquiry, interrogation, query.
Answer - A spoken or written reply,
as to a question.
Solution - Something worked out to
explain, resolve, or provide a method for dealing
with and settling a problem.
1. Numbers – Significant Figures, Rounding Rules, Accuracy, Precision,
Statistical Treatment of the Data
2. Units – 5 of the 7
1. Time – seconds
Conversion factors – realtionship
2. Length – Meters
between two of the basic units
3. Mass – grams
Molecular Weight (Mass)
4. Amount – Moles
Molar Mass, Mole Ratio, Molarity,
5. Temperature – Kelvins
molality, Density?
3. Vocabulary – Approximately 300 new terms or words and applying new or
more rigid definitions to words you may already own.
4. Principles (Theories and Laws) – Stoichiometry, Quantum Theory,
Bonding, Chemical Periodicity, Solutions, Thermodynamics,
Intermolecular Forces, Gas Laws, Colligative Properties, Kinetics, Equilibrium,
Electrochemistry
cp = q/mDT
DG = DH – TDS
PV = nRT
DT = Kmi
rate = k[A]m[B]n
∆E = q + w
Eocell = Ecathode = Eanode
[C]c[D]d
%yield = actual/theoretical * 100% K = a b
[A] [D]
c (ms-1)
E = n =
l (m)
Dr. Stone’s patent pending chemistry problem solver
1. Write down everything you are given
•Vocabulary
•Numbers
•Units
2. Write down what you want to know
•Vocabulary
•Numbers
•Units
3. Write down mathematical equation(s)
that include(s) these values and units
•Principles
4. Write a balanced stoichiometric equation
•Mole concept
5. Convert everything to moles
•Dimensional analysis
6. Convert everything to the unknown’s
units
•Rounding, significant figures, accuracy and
precision
Observations
•
Can be qualitative or quantitative
• Qualitative observations describe properties or occurrences in ways that do not
rely on numbers.
• Quantitative observations are measurements that consist of a number a unit and
a label
Hypotheses
•
A tentative explanation for the observations that may not be correct, but puts the
scientist’s understanding of the system being studied into a form that can be tested
Experiments
•
•
•
Tests the validity of the hypothesis
Are systematic observations or measurements made under controlled conditions, in
which the variable of interest is clearly distinguished from any others
If experimental results are reproducible, they are summarized in a law.
Law
•
A verbal or mathematical description of a phenomenon that allows for general
predictions that describes what happens and not why and is unlikely to change
greatly over time unless a major experimental error is discovered.
Theory
•
Attempts to explain why nature behaves as it does which is incomplete and
imperfect and evolves with time to explain new facts as they are discovered
Scientific Method - A procedure that searches for
answers to questions and solutions to problems, which
consists of:
http://museum.nist.gov/exhibits/adx2/index.htm
http://www.batteryequaliser.com/behome.html
http://museum.nist.gov/exhibits/adx2/index.htm
http://www.patentstorm.us/patents/5945236-description.html
Natural Laws
• Law of Conservation of Mass – The notion that mass, or
matter, can be neither created nor destroyed.
• Law of Conservation of Energy – A law that states that in
any system not involving nuclear reactions or velocities approaching the velocity of light,
energy cannot be created or destroyed. The First Law of Thermodynamics.
• Law of Conservation of Mass – Energy –
Einstein’s General Theory of Relativity - E=mc2 – with work becomes the special theory of
relativity, which has been verified by experiment, has shown that the mass of a body
changes as the energy possessed by the body changes. Such changes in mass are too small
to be detected except in subatomic phenomena. Matter may be created by the
materialization of a photon into an electron-positron pair; or it may be destroyed, by the
annihilation of this pair of elementary particles to produce a pair of photons.
Natural Laws
Law of Definite Proportions – When two or more elements
combine to form a compound, their masses in that compound are in a fixed and definite ratio.
This data helps justify an atomic view of matter.
Law of Multiple Proportions – When two elements
combine to form more than one compound, the mass of element A which combines in the first
compound with a given amount of element B has a simple whole number ratio with the mass of
element A which combines in the second compound with the same given mass of element B.
Use of Numbers
• Exact numbers
– 1 dozen = 12 things for example
• Accuracy
– how closely measured values agree with the
correct value
• Precision
– how closely individual measurements agree
with each other
• Significant Figures – start at the left and proceed
to the right
– If the number does not have a decimal point
count until there are no more non zero
numbers
– If the number has a decimal point start
counting at the first non-zero number and
continue counting until you run out of
decimal places
• Scientific notation – use it.
•
•
Use of Numbers
Multiplication & Division rule
Easier of the two rules
Product has the smallest number of significant figures of multipliers
4.242
x 1.23
2.7832
x 1.4
5.21766
round off to 5.22
3.89648
round off to 3.9
Addition & Subtraction rule
More subtle than the multiplication rule
Answer contains smallest decimal place of the addends.
3.6923
 2.02
8.7937
 2.123
6.9463
round off to 6.95
6.6707
round off to 6.671
 1.234
When a 5 appears. Is there anything to the right of the 5 greater than zero?
Is the number to the left of the 5 odd?
Is the number to the left of the 5 even? (Treat 0 as even.)
Use of Numbers
How many sig figs?
How many sig figs in the answer?
a)
b)
c)
d)
e)
f)
g)
h)
i)
a)
b)
c)
d)
e)
f)
g)
h)
i)
j)
k)
l)
m)
n)
0.0713200
7843000
1.4800
100
100.0
894.003
89400
0.03000
74.000
Round off to two sig figs
a)
b)
c)
d)
e)
f)
g)
h)
34.78
17.51
48.50
45.50001
24.33
17.50
20.5
45.5000
472x101
4600x0.005
36.0x4752
45.08/36.2
1.003/8500
0.003/472x12
3.003/475.0x0.30/524
0.3005x4.1
23.56+24.983
4.78-2.892
46.83-0.03
34.892+5.0
134.033-0.02
48.2-46
[
]
23.56 – 2.3
=
(1.68)
3
1.248 x 10
Day 2
Vocabulary
• Chemistry - Science that describes matter – its properties-composition-structure, the
changes it undergoes, and the energy changes that accompany those processes
• Matter - Anything that has mass and occupies space.
• Energy - The capacity to do work or transfer heat.
• Chemical Properties - chemical changes - describes the characteristic ability of a
•
•
•
substance to react to form new substances (flammability and corrosion).
– rusting or oxidation, chemical reactions
Physical Properties - physical changes - Characteristics that scientists can measure
without changing the composition of the sample under study (mass, color, volume, amount
of space occupied by the sample).
– changes of state, density, color, solubility
Extensive Properties - depend on quantity
a. Vary with the amount of the substance
b. Include mass, weight, and volume.
Intensive Properties - do not depend on quantity
a. Include color, melting and boiling point, electrical conductivity, and physical state at a
given temperature
b. Determine a substance’s identity,
density = mass
d = m
c. Have an important intensive property called
volume
V
density (d), a ratio of two extensive properties,
mass and volume
Vocabulary
• Three distinct states of matter:
1. Solids — relatively rigid and have fixed shapes and volumes
— volumes of solids independent of temperature and pressure
2. Liquids — have fixed volumes but flow to assume the shape of their containers
— Volumes of liquids independent of temperature and pressure
3. Gases — have neither fixed shapes nor fixed volumes and expand to fill their
containers completely
— Depends strongly on temperature and pressure
Vocabulary – Isotopes and Atomic Masses
 Atoms of different elements exhibit
different chemical behavior.
 Identity of an element is defined by
its atomic number.
 (Z) is the number of protons in the
nucleus of an atom of the element.
 The atomic number is therefore
different for each element.
 Known elements are arranged in
order of increasing Z in the
periodic table.
Vocabulary – Isotopes and Atomic Masses
 The chemistry of each element is determined by its number of protons and
electrons.
 In a neutral atom, the number of electrons equals the number of protons.
 Symbols for elements are derived directly from the element’s name.
 Nuclei of atoms contain neutrons as well as protons.
 The number of neutrons is not fixed for most elements, unlike protons.
 Atoms that have the same number of protons, and hence the same atomic
number, but different numbers of neutrons are called isotopes.
 Isotopes - All isotopes of an element have the same number of protons and
electrons, which means they exhibit the same chemistry. Isotopes of an element
differ only in their atomic mass.
Vocabulary – Isotopes and Atomic Masses
• Atomic mass
1. The mass of any given atom is not simply the sum of the masses of its electrons,
protons, and neutrons.
2. Most elements exist as mixtures of several stable isotopes. The weighted average is of
the masses of the isotopes is called the atomic mass. Atoms are too small to measure
individually and do not have a charge.
3. The arbitrary standard that has been established for describing atomic mass is the
atomic mass unit (amu), defined as one-twelfth of the mass of one atom of 12C.
4. Electrons added or removed from an atom produce a charged particle called an ion,
whose charge is indicated by a superscript after the symbol for the element.
Vocabulary – Essential Elements
• Elements that are absolutely required in the diets of humans are called
essential elements (highlighted in purple).
• Essential elements are restricted to the first four rows of the periodic table
with only two exceptions (Mo and ).
• An essential element is one that is required for life and whose absence
results in death.
• An element is considered to be essential if a deficiency consistently causes
abnormal development or functioning and if dietary supplementation of
that element and only that element prevents this adverse effect.
Classification of the Essential Elements
• Most living matter consists primarily of bulk elements—
oxygen, carbon, hydrogen, nitrogen, and sulfur. They are
the building blocks of the compounds that make up our
organs and muscles; they also constitute the bulk of our
diet.
• Six elements—sodium, magnesium, potassium, calcium,
chlorine, and phosphorus—are called macrominerals and
provide essential ions in body fluids and form the major
structural components of the body.
• Remaining essential elements called trace elements and are
present in small amounts (V, Cr, Ni, F, Sn, Si, As).
The Trace Elements
• It is difficult to detect low levels of some of the essential
elements, so the trace elements were relatively slow to be
recognized.
• Many compounds of trace elements are toxic.
• Dietary intakes of elements range from deficient to optimum
to toxic with increasing quantities; the optimum levels differ
greatly for the essential elements.
Chemical Compounds
Atoms in all substances that contain more than one atom are held together by
electrostatic interactions—interactions between electrically charged particles such as
protons and electrons.
Compounds & Molecules
• COMPOUNDS are a combination of 2 or more elements in
definite ratios by mass. (Law of Definite Proportions)
• The character of each element is lost when forming a
compound.
• MOLECULES are the smallest unit of a compound that retains
the characteristics of the compound.
The composition of molecular compounds is given by a
• CHEMICAL FORMULA
WRITING FORMULAS
Chemical Formula – chemical symbols and number of each representing composition
Empirical Formula – simplest ratio of elements that does not represent the actual
number and is non-positional
Molecular Formula – chemical symbols and number of each representing
composition representing actual number but not position
Structural Formula – chemical symbols and number of each representing
composition representing actual number and position
The structural chemical formula for glycine can be
written as
H2NCH2COOH
to show atom ordering
or in the bond-line structural formula
H H O
H N C C O H
H
Enables chemists to create a three-dimensional
model that provides information about the physical
and chemical properties of the compound
WRITING FORMULAS
Empirical and molecular formulas are precise and informative
but have disadvantages:
– inconvenient for routine verbal communications
– many compounds have the same empirical and molecular formulas
but different arrangements of atoms, which results in different
chemical and physical properties
Compounds & Molecules
STRUCTURAL
FORMULA
BOND-LINE
FORMULA
Day 3
MOLECULAR FORMULAS
• In one molecule of glycine there are
– 2 C atoms
– 5 H atoms
– 1 N atom
– 2 O atoms
Empirical Formula _______________
Molecular Formula _______________
Structural Formula _______________
Representations of Molecular Structures
 Different ways of representing the structure of a
molecule
1. Molecular formula gives only the number of each kind of atom
present.
2. Structural formula shows which atoms are present and how they
are connected.
3. Ball and stick model shows the atoms as spheres and the bonds
as sticks.
4. A perspective drawing, called a wedge-and-dash representation,
attempts to show the three-dimensional structure of the
molecule.
5. The space-filling model shows the atoms in the molecule but not
the bonds.
6. The condensed structural formula is the easiest and most
common way to represent a molecule—it omits the lines
representing bonds between atoms and simply lists the atoms
bonded to a given atom next to it. Multiple groups attached to
the same atom are shown in parentheses, followed by a subscript
that indicates the number of such groups.
1.
2.
3.
4.
5.
6.
Chemical Compounds
• Chemical bonds – two different kinds
1. Ionic — ionic compounds consist of positively and negatively charged
ions held together by strong electrostatic forces. (Formula Mass)
2. Covalent — covalent compounds consist of molecules, which are
groups of atoms in which one or more pairs of electrons are
shared between bonded atoms. Atoms are held together by
the electrostatic attraction between the positively charged
nuclei of the bonded atoms and the negatively charged
electrons they share. (Molecular Mass)
Ionic Chemical Compounds
• Electrostatic attraction between oppositely charged particle species
(positive and negative) results in a force that causes them to move toward
each other.
• Electrostatic repulsion between two species that have the same charge
(either both positive or both negative) results in a force that causes them to
repel each other
• When the attractive electrostatic interactions between atoms are stronger
than the repulsive interactions, atoms form chemical compounds and the
attractive interactions between atoms are called chemical bonds.
Ionic Chemical Compounds
• IONS are atoms or groups of atoms with a net positive or negative charge.
• Taking away an electron from an atom gives a CATION with a positive charge
• Adding an electron to an atom gives an ANION with a negative charge.
• Ionic compounds contain both cations and anions in a ratio that results in no net
electrical charge.
CATION + ANION → COMPOUND
• A neutral compound requires equal number of + and - charges.
In general
metals (Mg) lose electrons to become
cations
nonmetals (F) gain electrons to
become anions
Ionic compounds are held together by the
attractive electrostatic interactions between
cations and anions.
Cations and anions are arranged in space to
form an extended three-dimensional array
that maximizes the number of attractive
electrostatic interactions and minimizes the number of repulsive electrostatic interactions.
Physical Properties of Ionic Compounds
 Ionic compounds
– Usually form hard crystalline solids that melt at high
temperatures and are very resistant to evaporation
– Properties stem from the characteristic internal structure of an
ionic solid, which is a three-dimensional array of alternating
positive and negative ions held together by strong electrostatic
attractions
PbS
FeS
SiO
Binary Ionic Compounds
 An ionic compound that contains only two elements, one
present as a cation and one as an anion, is called a binary
ionic compound.
 For such compounds, the subscripts in the empirical
formula can also be obtained using the absolute value of
the charge on one ion as the subscript for the other ion and
then reduce the subscripts to their simplest ratio to write
the empirical formula.
Polyatomic Ions
 Groups of atoms that bear a net electrical charge
 Atoms that make up a polyatomic atom are held together
by the same covalent bonds that hold atoms together in
molecules
 Many more kinds of polyatomic ions than monatomic ions
and polyatomic anions are more numerous than
polyatomic cations
 Method used to predict empirical formula for ionic
compounds that contain monatomic ions can be used for
compounds containing polyatomic ions. Overall charge on
the cations must balance the overall charge on the anions
in the formula unit.
Table of Common Ions
Common Positive Ions (Cations)
Monovalent
Hydronium
(or hydrogen)
Lithium
Sodium
Potassium
Rubidium
Cesium
Francium
Silver
Ammonium
Thalium
Copper I
H3O+
H+
Li+
Na+
K+
Rb+
Cs+
Fr+
Ag+
NH4+
Tl+
Cu+
Divalent
Magnesium
Calcium
Strontium
Beryllium
Manganese II
Barium
Zinc
Cadmium
Nickel II
Palladium II
Platinum II
Copper II
Mercury II
Mercury I
Iron II
Cobalt II
Chromium II
Lead II
Tin II
Mg2+
Ca2+
Sr2+
Be2+
Mn2+
Ba2+
Zn2+
Cd2+
Ni2+
Pd2+
Pt2+
Cu2+
Hg2+
Hg22+
Fe2+
Co2+
Cr2+
Pb2+
Sn2+
Trivalent
Aluminium
Antimony III
Bismuth III
Al3+
Sb3+
Bi3+
Iron III
Cobalt III
Chromium III
Fe3+
Co3+
Cr3+
Table of Common Ions
Common Negative Ions (Anions)
Monovalent
Hydride
Fluoride
Chloride
Bromide
Iodide
Hydroxide
Permangante
Cyanide
Thiocynate
Acetate
Nitrate
Bisulfite
Bisulfate
Bicarbonate
Dihydrogen phosphate
Nitrite
Amide
Hypochlorite
Chlorite
Chlorate
Perchlorate
HFlClBrIOHMnO4CNSCNC2H3O2NO3HSO3HSO4HCO3H2PO4NO2NH2ClOClO2ClO3ClO4-
Divalent
Oxide
Peroxide
Sulfide
Selenide
Oxalate
Chromate
Dichromate
Tungstate
Molybdate
tetrathionate
Thiosulfate
Sulfite
Sulfate
Carbonate
Hydrogen phosphate
phosphite
O2-
Trivalent
Nitride
N3-
O22S2Se2C2O42CrO42Cr2O72WO42MoO42S4O62S2O32SO32SO42CO32HPO42-
Phosphate
PO43-
Hydrates
 Ionic compounds that contain specific
ratios of loosely bound water
molecules, called waters of hydration.
 Waters of hydration can be removed
by heating.
 Compounds that differ only in the
numbers of waters of hydration can
have very different properties.
CuSO4.5H2O
Prefixes for indicating numbers of
species in chemical names
Prefix
Number
Prefix
Number
mono -
1
hepta-
7
di-
2
octa-
8
tri-
3
nona-
9
tetra-
4
deca-
10
penta-
5
undeca-
11
hexa-
6
dodeca-
12
Day 4
Naming Polyatomic Ionic Compounds
Many compounds have more than one name:
1. Common name — have historical origins
2. Systematic name — write structure of the compound from its name and vice versa
Procedure for naming binary ionic compounds, which contain only two elements, uses the
following steps:
1) Place the ions in their proper order: cation and then anion
2) Name the cation
1) Metals that form only one kind of positive ion. These metals are in Groups 1–3, 12, and 13.
The name of the cation of a metal that forms only one kind of positive ion is the same
as the name of the metal
2) Metals that form more than one cation. These metals are transition metals, actinides, and the
heaviest elements of Groups 13–15. Positive charge on the metal is indicated by a
Roman numeral in parentheses following the name of the metal.
3) Name the anion
1) Monatomic anions — named by adding the suffix
–ide to the root of the name
of the parent element
2) Polyatomic anions
1) Have common names that must be learned
2) Polyatomic anions that contain a single metal
or nonmetal atom plus one or more oxygen
atoms are called oxoanions. Relationship
between names of oxoanions and number
of oxygen atoms present is:
Naming Polyatomic Ionic Compounds
Write the names of the compound as the name of the cation followed by the
name of the anion.
– It is not necessary to indicate the number of cations or anions present per
formula unit in the name of an ionic compound because information is implied by
the charges on the ions
– When writing the formula for an ionic compound from its name the charge of the
ions must considered.
Covalent Compounds
 Physical Properties of Covalent compounds
– Can be gases, liquids, or solids at room temperature
and pressure, depending on the strength of the
intermolecular interactions
– Covalent molecular solids tend to form soft crystals
that melt at low temperatures and evaporate easily
– Consist of discrete molecules held together by
comparatively weak intermolecular forces (the forces
between molecules) even though the atoms within each
molecule are held together by strong intramolecular
covalent bonds (the forces within the molecule)
Covalent Compounds
• Covalent compound are represented by a molecular formula, which gives the
atomic symbol for each component element, in a prescribed order,
accompanied by a subscript indicating the number of atoms of that element
in the molecule
C8H10N4O2
C4H10OBCl3
C3H6N6O6
Covalent Compounds
• Inorganic compounds
– Compounds that consist primarily of elements other than
carbon and hydrogen
– Include both covalent and ionic compounds
– Formulas are written when the component elements are listed
beginning with the one farthest to the left in the periodic table
with those in the same group listed alphabetically
• Organic compounds
– Covalent compounds that contain predominantly carbon and
hydrogen
– Formulas of organic compounds written with carbon first,
followed by hydrogen and then by other elements in
alphabetical order
Inorganic Covalent Compounds
• Some pure elements exist as covalent molecules
• Hydrogen, nitrogen, oxygen, and the halogens
occur as diatomic molecules and contain two atoms
• A few pure elements, such as elemental phosphorus
and sulfur, are polyatomic molecules and contain
more than two atoms
Binary Inorganic Compounds
 Binary covalent compounds — covalent compounds that contain
only two elements
 The procedure for naming them uses the following steps:
1. Place the elements in their proper order.
a. Element farthest to the left in the periodic table is named first. If both elements
are in the same group, the element closer to the bottom of the column is named
first.
b. Second element is named as if it were a monatomic anion in an ionic compound
with the suffix –ide attached to the root of the element name
2. Identify the number of each type of atom present.
a. Prefixes derived from Greek stems are used to indicate the number of each type
of atom in the formula unit.
b. If the molecule contains more than one atom of both elements, then prefixes are
used for both.
c. With some names, the final a or o of the prefix is dropped to avoid awkward
pronunciation.
3. Write the name of the compound.
a. Binary compounds of the elements with oxygen are named as “element oxide”
with prefixes that indicate the number of atoms of each element per formula unit.
b. Certain compounds are always called by their common names assigned long ago
when names rather than formulas were used.
Binary Inorganic Compounds
Organic Covalent compounds
Hydrocarbons and Biologically Important Compounds
 Hydrocarbons
 Consist entirely of carbon and hydrogen
 Four major classes of hydrocarbons:
1. Alkanes — contain only carbon-hydrogen and carboncarbon single bonds
2. Alkenes — contain at least one carbon-carbon double bond
3. Alkynes — contain a least one carbon-carbon triple bond
4. Aromatics — contain rings of six carbon atoms that can be
drawn with alternating single and double bonds
To Be Covered in Weeks 6 and 7
Biologically Important Compounds
1. Carbohydrates –Sugars, starches
2. Lipids, Fats, Oils
3. Biopolymers
1. Proteins
2. Nucleic Acids
To Be Covered in Week 8
Coordinate Inorganic Compounds
To Be Covered
in Week 9

Acids and Bases
Acids
– A substance with at least one hydrogen atom that can dissociate
to form an anion and an H+ ion (a proton) in aqueous solution,
thereby forming an acidic solution

Bases
– Compounds that produce hydroxide ions (OH–) and a cation
when dissolved in water, thus forming a basic solution

Neutral
– Solutions that are neither basic nor acidic
To Be Covered
in Weeks 10 and 11
Counting Atoms
Chemistry is a quantitative science—we need a
“counting unit.”
MOLE - 1 mole is the amount of
substance that contains as many
particles (atoms, molecules) as there
are in 12.0 g of 12C.
Avogadro’s Number
6.02214199 x 1023
There is Avogadro’s number of
518 g substance.
of Pb, 2.50 mol
particles in a mole of any
Amedeo Avogadro
1776-1856
Molar Mass
1 mol of 12C
= 12.00 g of C
= 6.022 x 1023 atoms of C
12.00 g of 12C is its MOLAR MASS
Taking into account all of the
isotopes of C, the molar mass of
C is 12.011 g/mol
One-mole
Amounts
Day 5
Molar mass
• Flowchart for converting between mass, number of moles,
and number of atoms, molecules, or formula units
Molar mass
PROBLEM: What amount of Mg is represented by
0.200 g? How many atoms? Mg has a molar mass of 24.3050 g/mol.
PROBLEM: How many hydrogen atoms in 0.036 moles
of propane C3H8? In 2.7 moles?
MOLECULAR MASS & MOLAR MASS
Molecular weight = sum of the atomic weights of all
atoms in the molecule.
Molar mass = molecular weight in grams
Problem: What is the molar mass
of ethanol – C2H6O?
1 mol contains
2 mol C (12.01 g C/1 mol C/1 mol C2H6O) =
6 mol H (1.01 g H/1 mol H/1 mol C2H6O) =
1 mol O (16.00 g O/1 mol O/1 mol C2H6O) =
TOTAL = Molar Mass =
MOLECULAR MASS & MOLAR MASS
• Molecular mass of a substance
– Sum of the average masses of the atoms in one molecule of the substance
– Calculated by summing the atomic masses of the elements in the
substance, each multiplied by its subscript in the molecular formula
– Units of molecular mass are atomic mass units (amu)
• Formula = C8H9NO2
• Molar mass =
Tylenol
Molar Mass
PROBLEM: How many moles of alcohol are there in a
“standard” can of beer if there are 21.3 g of C2H6O?
(a) Molar mass of C2H6O = 46.08 g/mol
(b) Calc. moles of alcohol
PROBLEM: How many molecules of alcohol are there in
a “standard” can of beer if there are 21.3 g of C2H6O?
(a) 21.3 g of C2H6O is 0.462 mol of C2H6O.
(b) Calc. molecules of alcohol
PROBLEM: How many atoms of C from ethanol are
there in a “standard” can of beer if there are 21.3 g of
C2H6O?
(a) 21.3 g of C2H6O is 0.462 mol of C2H6O which is 2.78 E23 .
(b) Calc. Atoms of carbon.
Calculating Mass Percentages
A pure compound always consists of the same elements combined in the same
proportions by weight.
Therefore, we can express molecular composition as PERCENT BY WEIGHT
Ethanol, C2H6O
52.13% C
13.15% H
34.72% O

Law of definite proportions states that a chemical compound always contains

Percent composition — the percentage of each element present in a pure

Calculation of mass percentage
the same proportion of elements by mass
substance—is constant
1. Use atomic masses to calculate the molar mass of the compound
2. Divide the mass of each element by the molar mass of the compound and then
multiply by 100% to obtain percentages
3. To find the mass of an element contained in a given mass of the compound,
multiply the mass of the compound by the mass percentage of that element
expressed as a decimal
Percent Composition
Consider NO2, Molar mass = ?
What is the weight percent of N and of O?
What are the weight percentages of N and O in NO?
Percent Composition
A compound of B and H is 81.10% B. What is its empirical formula?
1.
Calculate the number of moles of each element in 100.0 g of sample.
2.
Take the ratio of moles of B and H. Always divide by the smaller number.
3.
Now, recognize that atoms combine in the ratio of small whole numbers.
4.
Find the ratio of moles of elements in the compound.
We need to do an EXPERIMENT to find the MOLAR MASS. Experiment gives 53.3 g/mol
Compare with the mass of B2H5 = 26.66 g/unit , i.e. Find the ratio of these masses.
Determining the Empirical Formula of
Penicillin
 Can use the empirical formula of a substance to
determine its percent composition
 Can use the percent composition of a sample to
determine its empirical formula, which then can
be used to determine the molecular formula—a
procedure used to determine the empirical and
molecular formulas of penicillin
Determining the Empirical and Molecular
Formulas of Penicillin
The combustion analysis of Penicillin G
C = 53.9%
H = 4.8%
N = 7.9%
S = 6.5%
Na = 6.5%
Total = 82.1%
Day 6
Chemical Equations
• A chemical equation is an expression that gives the
identities and quantities of the substances in a chemical
reaction
• Chemical formulas and other symbols are used to indicate
the starting material(s) or reactant(s), which are written on
the left side of the equation, and the final compound(s) or
product(s), which are written on the right side. An arrow,
read as yields or reacts to form, points from the reactants to
the products.
• Abbreviations are added in parentheses as subscripts to
indicate the physical state of each species:—(s) for solid, (l)
for liquid, (g) for gas, and (aq) for an aqueous solution.
• A balanced chemical equation is when both the numbers of
each type of atom and the total charge are the same on both
sides. A chemical reaction represents a change in the
distribution of atoms but not in the number of atoms.
Chemical Equations
• Look at the information an equation provides:
D
Fe 2O 3(s) + 3 CO (g) 
 2 Fe (s) + 3 CO 2 (g)
• Balanced chemical equation
– Provides qualitative information about the identities and
physical states of the reactants and products
– Provides quantitative information because it tells the relative
amounts of reactants and products consumed or
produced in the reaction
– The number of atoms, molecules, or formula units of a reactant
or product in a balanced chemical equation is the
coefficient of that species
– Mole ratio of two substances in a chemical reaction is the ratio
of their coefficients in the balanced chemical equation
Balancing Simple Chemical Equations
 Method for balancing chemical equations
1. Identify the most complex substance.
2. Beginning with that substance, choose an element that
appears in only one reactant and one product. Adjust
the coefficients to obtain the same number of atoms of
this element on both sides.
Optionally - Balance polyatomic ions (if present)
as a unit.
3. Balance the remaining atoms, usually ending with the
least-complex substance and using fractional
coefficients if necessary. If a fractional coefficient is
used, multiply both sides of the equation by the
denominator to obtain whole numbers for the
coefficients.
4. Count the numbers of atoms of each kind on both sides
of the equation to be sure that the chemical equation is
balanced.
Chemical Equations
• Law of Conservation of Matter
– There is no detectable change in quantity of matter in
an ordinary chemical reaction.
– Balanced chemical equations must always include the
same number of each kind of atom on both sides of the
equation.
– This law was determined by Antoine Lavoisier.
• Propane,C3H8, burns in oxygen to give carbon
dioxide and water.
D
C3H 8  5 O 2 
 3 CO 2  4 H 2 O
Law of Conservation of Matter
• NH3 burns in oxygen to form NO & water
• C7H16 burns in oxygen to form carbon dioxide and
water.
KClO3 +
C12H22O11
CaSO4 + CH4 + CO2
KCl + CO2 +
H 2O
CaCO3 + S + H2O
Mass Relationships in Chemical Equations
• A balanced chemical equation gives the identity of the
reactants and products and the accurate number of
molecules or moles of each that are consumed or produced.
• Stoichiometry is a collective term for the quantitative
relationships between the masses, numbers of moles, and
numbers of particles (atoms, molecules, and ions) of the
reactants and products in a balanced reaction.
• A stoichiometric quantity is the amount of product or
reactant specified by the coefficients in a balanced
chemical equation.
Stoichiometry Problems
• Steps in converting between masses of reactants and
product
1.
2.
3.
Convert the mass of one substance (Substance A) to the
corresponding number of moles using its molar mass.
From the balanced chemical equation, obtain the number of
moles of another substance (B) from the number of moles of
substance A using the appropriate mole ratio (the ratio of their
coefficients).
Convert the number of moles of substance B to mass using its
molar mass.
• Converting amounts of substances to moles, and vice versa,
is the key to all stoichiometry problems.
Calculations Based on Chemical Equations
• How many CO molecules are required to react with 25
MOLES of Fe2O3?
• How many iron atoms can be produced by the reaction of
2.50E5 MOLES of iron (III) oxide with excess carbon
monoxide?
• What mass of CO is required to react with 146 g of iron (III)
oxide?
Limiting Reactants
• If one or more of the reactants is not used up completely
but is left over when the reaction is completed, then the
amount of product that can be obtained is limited by the
amount of only one of the reactants
• A limiting reactant is the reactant that restricts the amount
of product obtained. The reactant that remains after a
reaction has gone to completion is present in excess.
Limiting Reactant Concept
• What is the maximum mass of sulfur dioxide that can be
produced by the reaction of 95.6 g of carbon disulfide with
110. g of oxygen?
CS2  3 O2  CO2  2 SO2
• What is the maximum mass of sulfur dioxide that can be
produced by the reaction of 95.6 g of carbon disulfide with
110. g of oxygen?
Percent Yields from Reactions
• The Limiting Reactant determines the maximum amount
of product that can be formed from the reactants when
reactants are not present in stoichiometric quantities.
• The Theoretical Yield is calculated by assuming that the
reaction goes to completion.
– Determined from the limiting reactant calculation.
• Actual Yield is the amount of a specified pure product
made in a given reaction.
– In the laboratory, this is the amount of product that is
formed in your beaker, after it is purified and dried.
• Percent Yield indicates how much of the product is
obtained from a reaction.
actual yield
% yield =
 100%
theoretical yield
Percent Yields from Reactions
• A 10.0 g sample of ethanol, C2H5OH, was boiled with excess
acetic acid, CH3COOH, to produce 14.8 g of ethyl acetate,
CH3COOC2H5. What is the percent yield?
Classifying Chemical Reactions
 Most chemical reactions can be classified into one
or more of only four basic types:
1. Acid-base reactions
acid + base  salt
2. Exchange reactions (Single Displacement, Double Displacement,
Metathesis)
AB + C  AC + B or AB + CD  AD + CB
3. Condensation reactions (and the reverse, cleavage reactions)
(Combination, Decomposition)
Condensation: A + B  AB
Cleavage: AB  A + B
4. Oxidation-reduction reactions
(covered in Chem 102)
oxidant + reductant  reduced oxidant + oxidized reductant
To Be Covered in Weeks 10 and 11
Day 7