1
Dual Enrollment
Chemistry 105/106
Welcome to CMC Chemistry II. This is a senior level course designed for students who have completed one full
year of Chemistry Pre-AP or Chemistry 105. This course is designed to expand on knowledge that was attained through
your first chemistry course and to address more advanced topics as included in the Nicholls State Chemistry 105/106
curriculum. Those of you who were dual enrolled last year in Chemistry 105 can now earn three additional credit hours
in Chemistry 106. When coupled with Chemistry Lab (Chemistry 110) you will leave the Career Magnet Center with 8
credit hours of chemistry which completes the Freshman college chemistry requirements at Nicholls State. (Note: those
not enrolled in Chemistry Lab 110 will leave the CMC with 6 credit hours of college chemistry and will likely need to
complete Chemistry 110 or its equivalent at the college level). These credit hours are also transferrable to most
universities. Those students who did not dual enroll in Chemistry 105 last year and has earned the minimum
requirement of 21 in Math and 18 on English on the ACT will have the opportunity of dual enrolling in Chemistry 105 in
August.
The following summer assignments are merely a review of many of the major topics that were discussed in
Chemistry I. There are a total of three assignments. They include a reading passage and corresponding questions to
answer. It is very lengthy, but as you will see, most of it is just reading material. My recommendation is that you print
out the entire packet so that you can highlight important concepts. If printing the entire packet is an issue for you, you
will at least have to print part of the packet that needs to be turned in to me.
If you want to sign up for my remind account, you can text @2chem17-18 to the number 81010 OR… on an
iphone, you can enter rmd.at/2chem17-18 and follow the instructions to sign up. By signing into remind, you will be
able to ask questions over the summer break.
You must also sign in to Moodle. Go to LPSD homepage. Click on More  Students  Moodle  Career
Magnet Center  Chemistry 105/106  enrollment key  7117  enroll me. If you did not dual enroll last year in
Chemistry 105 you must print the application form posted in Moodle and bring it in the first day of school so that I can
turn it in to Nicholls.
If you leave me a message under “post a question or concern”, I’ll give you 2 bonus points on your first test.
look forward to seeing all of you in August. In the meantime, have a great summer. 
Anna Cole
.... Continuous effort- not strength or intelligence- is the key to unlocking our potential….. Liane Cordes
I
2
Assignment 1
Matter and Measurement
I recommend that you print this information as notes to prepare for your first test. Place the packet in a 3 ring binder.
If you choose not to print the entire packet, print only those pages that need to be turned in to me. You will have to
refer to the electronic version of this assignment to study from for test purposes. Read the following review then
answer the questions that follow.
Key Vocabulary
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
Matter is classified as a pure substance or a mixture.
Pure substances are either elements or compounds.
An element is a substance all of whose atoms contain the same number of protons.
A compound is a relatively stable combination of two or more chemically bonded elements in a specific ratio.
Properties of compounds differ from the properties of the elements that make them up.
Mixtures consist of two or more substances that retain their individual identities and properties.
Homogeneous mixtures are uniform throughout. Air, seawater and a nickel coin (a mixture of copper and
nickel metals called an alloy) are examples.
Heterogeneous mixtures vary in texture and appearance throughout the sample. Rocks, wood, polluted air, and
muddy water are examples.
Physical properties can be measured without changing the identity or composition of the substance. Physical
properties include color, density, melting point and hardness.
Chemical properties describe the way a substance changes (reacts) to form other substances. The flammability
of gasoline is a chemical property because the gasoline reacts with oxygen to form carbon dioxide and water.
Intensive properties are physical properties that do not depend on the amount of substance in a sample.
Temperature, density, specific heat capacity and boiling point are intensive properties.
Extensive properties are physical properties that depend on the quantity of the sample. Energy (heat), mass,
and volume are examples of extensive properties.
A chemical change (also known as a chemical reaction) transforms a substance into a different substance or
substances. When the chief component of natural gas (methane) burns in air, the methane and the oxygen
from the air are transformed into carbon dioxide and water. In a chemical change, atoms are never lost as a
result of the reaction. The atoms merely rearrange forming different bonds which result in compounds
exhibiting different properties.
A physical change changes the appearance of a substance but does not change its composition. Changes of
physical state, from solid to liquid or from liquid to gas are examples.
 Differences in physical properties as determined by intermolecular forces are used to separate the
components of mixtures. The following are methods commonly used to separate the components of a
mixture.
a. Filtration separates an insoluble solid from liquid.
b. Distillation separates substances based on their difference in boiling points. This may involve a soluble
(dissolved) solid or two miscible liquids.
c. Chromatography is a technique that separates substances based on their differences in intermolecular
forces (polarity) and their abilities to dissolve in various solvents.
3
Units of Measurement
Chemists often use preferred units called SI units after the French Systeme International d’ Unites. Table 1.1 lists the SI
standards and their symbols.
Table 1.1
Physical quantity
Name of unit Abbreviation
Mass
Kilogram
kg
Length
meter
m
Time
second
s or sec
Temperature
Kelvin
K
Amount of substance mole
mol
Electric current
Ampere
A or amp
Temperature is commonly measured using either Celsius or the Kelvin scale. The formula K = °C + 273.15 allows us to
convert from a Celsius temperature to a Kelvin temperature (or vise versa).
Derived units are units derived from SI base units. Example: Volume
Volume is the space occupied by a substance. This is different from mass. Mass (a base unit) is the amount of matter in
an object. Volume is commonly measured in cubic meters (m3), cubic centimeters (cm3), or cubic decimeters (dm3).
Volume is considered a derived unit because it is a cubic combination of the base unit for length. However in chemistry,
we commonly use the non-SI units of L and mL for convenience. One liter is the volume of a cube measuring exactly 10
cm on a side (dm)
1 L = 1000 mL = 1000 cm3 = 1 dm3
Or
1 cm3 = 1 mL
Density is another example of a derived unit. It is defined as the amount of matter packed into a given space. It is often
measured in g/cm3 for liquids and solids and g/L for gases. It can be calculated using the following formula…
𝑚𝑎𝑠𝑠
Density = 𝑣𝑜𝑙𝑢𝑚𝑒
Uncertainty in Measurement
Exact numbers have an exact value and are usually defined or counted. One liters equals 1000 cm3 describes a defined
number. There are 32 students in this class describes a counted number.
Inexact numbers have some degree of error or uncertainty associated with them. All measured numbers are inexact.
Measured numbers are generally reported in such a way that only the last digit is uncertain. Significant figures are all
digits of a measured number including the uncertain one.
Zeros in measured numbers are either significant, or they are merely there to locate the decimal place. The following
guidelines describe when zeros are significant:
1. Zeros between nonzero digits are always significant. (embedded zeros)
2. Zeros at the beginning of a number are never significant (leading zeros)
3. Zeros at the end of a number are significant only when the number contains a decimal point.(trailing zeros)
4
In calculations involving measured quantities, the least certain measurement limits the certainty of the calculated
quantity and determines the number of significant figures in the final answer. Exact numbers do not limit the certainty.
For multiplication and division, the number of significant figures in the answer is determined by the measurement with
the fewest number of significant figures.
For addition and subtraction, the result has the same number of decimal places as the measurement with the fewest
number of decimal digits. In the case of whole numbers, round the answer to the greatest place value containing the
last significant figure.
For example, 340 + 42 = 382 which would be rounded to 380 (since the last significant digit in 340 is in
the ten’s place which is greater than 42 in which the last significant digit is in the one’s place).
Dimensional analysis
Dimensional analysis is a way of converting one unit of measure into another unit using a conversion factor. A
conversion factor is a fraction in which the numerator equals the denominator.
For example: How many microseconds are there in one year? Try this problem. The answer is provided. Refer to
your notes if necessary. Show all work. Answer: 3.15× 1013μs
5
Assignment 1 Review Questions
PRINT THESE QUESTIONS TO BE TURNED IN…. Refer to the information provided in Assignment 1 to answer the
following questions. Feel free to use supplementary material as well.
1. What is an alloy?
2. Distinguish between intensive and extensive properties and give examples of each.
3. Distinguish between a mixture and a compound.
4. Distinguish between a chemical and physical change.
5. Determine whether the following properties are intensive or extensive
A. ________________state of matter D. ________________density
B. ________________melting point
E. ________________solubility
C. ________________ mass
6. Determine a separation technique to separate the following
A. _____________________salt water
B. _____________________acetone and water (miscible)
C. _____________________sand in water
D. _____________________ pigments in a mixture of food dye
E. _____________________rubbing alcohol
F. ________________heat
G._________________ temperature
6
Using your notes or memory… Fill in the following table with metric prefixes and numerical values. The first one has
been done for you.
Prefix
Abbreviation
Meaning
Example
Tera
T
1012
1 Terameter (Tm) =
1× 1012 m
Main/base unit
Femto * (look this
one up)
7. ____________________What Greek letter does μ represent? (look this up)
8. _____________________What is the unit of density when describing solids and liquids?
9. _____________________What is the unit of density when describing gases?
10. Distinguish between an exact number and an inexact number.
11. What is a significant figure?
12. Determine the number of significant digits in the following:
A. ______0.003 03
B. ______100.0
C. ______100
D. ______10.040
7
13. Calculate the following and round to the correct number of significant digits.
A. _________________________ (1.40 ×10-3)(5.440× 10-4) / 1.0 ×103
B. _________________________343.34 + 23.56
C. _________________________1560 + 232
D. ______________________Convert 34.5 °C to K
E. ______________________Convert 299.4 K to Celsius
F. ___________________The density of a solution was experimentally determined. The mass of an empty
graduated cylinder has a mass of 89.40 g. After the solution is placed in the cylinder the mass of the cylinder
and the solution measured 120.40 g. The volume of the solution in the cylinder measures 29.9 mL.
Calculate the density to the correct number of significant digits.
G. _________________Convert 34.53 kg/L to mg/mL. Show work.
H. _________________Convert 3.00× 108 m/sec to mi/hr . Show work.
Selected Answers:
12a. 3
b. 4
c. 1
d. 5
13a. 7.6 ×10-10 b. 366.90 c. 1790 d. 307.6 K e. 26.2 °C f. 1.04 g/mL g. 34530 mg/mL h. 6.71×108 mi/hr
(Make sure your answers have the same number of significant figures as those provided here)
8
Assignment 2: Atoms, Molecules, and Ions
Read the following passage, highlighting important ideas. Then answer the questions that follow:
Theories and Laws
Scientific knowledge is empirical- it is based on observation and experiment. Scientists observe and
perform experiments on the physical world to learn about it. These observations and experiments over the
course of time have led us to the development of scientific laws and scientific theories. It is through the
formation of theories and laws that all scientific knowledge has been acquired. A scientific law originates with
a series of similar and consistent observations. It is a brief statement that summarizes past observations and
predicts future ones. Scientific laws describe how nature behaves. We have seen examples of scientific laws
in our study of the gas Laws (Charles’s Law, Boyle’s Law, the Ideal gas law), the Law of Conservation of Matter,
the Law of Definite Proportion, etc. These laws are intended to merely state facts. A scientific theory on the
other hand tries to explain why nature behaves the way it does. Theories are broad explanations of scientific
concepts often consisting of several “postulates” or “assumptions” which may include scientific laws which are
embedded into these postulates. Well established theories are the pinnacle of scientific knowledge, often
predicting behavior far beyond the observations from which they were developed. However they cannot ever
be conclusively proven because some new observation or experiment always has the potential to reveal a
flaw. If a theory is proven wrong by an experiment, it must be revised or tested with new experimentation.
Over time, poor theories and laws are eliminated or corrected, and good theories and laws – those consistent
with experimental results- remain. In fact the more a theory is refined, the closer to truth we get. Established
theories with strong experimental support are the most powerful pieces of scientific knowledge and are the
foundation from which future scientific ideas are built. The development and refinement of the atomic theory
of matter proposed by the English Chemist John Dalton over the course of time is a perfect example of how
“old” ideas are refined to incorporate new observations and data.
Dalton’s atomic theory
All matter is composed of tiny indivisible particles called atoms.
Dalton’s postulates
1. Each element is composed of tiny, indestructible particles called atoms.
2. All atoms of the same element have the same mass and other properties that distinguish them from the atoms
of other elements.
3. Atoms combine in simple, whole number ratios to form compounds.
4. Atoms of one element cannot change into atoms of another element. In a chemical reaction, atoms are neither
created nor
destroyed; they simply change the way they are
bonded together.
9
Dalton’s theory includes two fundamental laws:
The law of definite composition (Law of definite proportion) states that the relative numbers and kinds of atoms in a
given compound are constant. Since every type of atom has its own specific mass and these atoms bond with other
atoms in specific ratios (Postulate 3), then we can also say that on a macroscopic level a compound will have a definite
composition. For example 25 grams of sodium chloride would have the same percentage of sodium and chlorine as a
125 g sample.
The law of conservation of matter states that the mass of reactants will always equal the mass of products in a reaction.
Dalton explains these observations in Postulate 4. All matter is composed of small, indestructible particles. Atoms
simply rearrange during a chemical reaction to form a different compound. Since the particles are only rearranged,
matter is conserved.
Today we accept three of Dalton’s four ideas. Only the second idea is incorrect. We now know that atoms of a given
element are not identical. Evidence from mass spectra (to be discussed later) clearly demonstrates that atoms of the
same element can be composed of different isotopes, each having different masses due to different numbers of
neutrons.
The Evolution of the Atomic model
The Dalton model
As a result of Dalton’s atomic theory, scientists were able to envision what the atom would perhaps look like.
This “picture” is referred to as an atomic model. Being able to picture an atom helps us to comprehend many
concepts that we believe to exist at the atomic level. Dalton’s model depicted the atom as a solid, indivisible
sphere. However, Dalton’s model of the atom had to be refined as new evidence was discovered. The
following is a brief synopsis of how our concept of the atom has changed over the course of time.
The Thomson Model (Plum Pudding Model)
J.J. Thomson is credited with discovering the electron. With this
discovery, Dalton’s atomic theory had to be modified to incorporate the
existence of electrons. Thomson proposed that an atom was a solid
sphere as Dalton suggested, but the electrons were embedded in the
sphere like plums in a pudding. (Figure 2.2)
Figure 2.2 Plum Pudding Model
10
The Rutherford Model (Nuclear Model)
Ernest Rutherford’s gold foil experiment, illustrated in Figure 2.3, provided experimental evidence that contradicted
Thomson’s assumption that atoms are solid particles. In this experiment positively charged alpha particles were fired
at a thin sheet of gold foil. Most of the particles went straight through the foil, but some were severely deflected.
Rutherford concluded from these observations that the atom is mostly empty space but contains a tiny very dense core
of positive charge which he referred to as the nucleus. Consequently, Rutherford’s model is referred to as the nuclear
model. He could not explain however why the negatively charged electrons did not collapse into the positively charged
nucleus.
Figure 2.3
The Bohr Model (Planetary model)
Niels Bohr provided an explanation of Rutherford’s dilemma. He postulated that electrons did not collapse into the
nucleus because they existed in “energy states”. His model described the atom with its electrons orbiting the nucleus
much like the planets orbit the sun. Bohr used evidence from the emission spectrum of elemental hydrogen to support
his theory.
11
Probing into the structure of the atom:
Atomic Spectroscopy and the Bohr Model
An atomic emission spectrum (or line spectrum) is a pattern of discrete lines of different wavelengths when the
light energy emitted from energized atoms is passed through a prism or diffraction grating. Figure 6.3 shows the
atomic emission spectrum of hydrogen. This pattern is unique for hydrogen.
* Wavelengths
are measured in
nanometers
The Bohr model of the atom (developed by
the Danish physicist Niels Bohr) explains
the origin of the lines of the atomic
emission spectrum of hydrogen in Figure
6.3. Bohr proposed that electrons move in
a circular, fixed energy orbits around the
nucleus. He postulated that each circular
orbit corresponds to an “allowed” stable
energy state. The ground state is the
lowest energy state. Excited states are
states of higher energy than the ground
state. When an external energy source is
applied to an element, the electrons in
these atoms jump to a higher energy level
and become “excited”. However these
electrons quickly drop back down to their
ground state and release this energy which
nucleus
is in the form of a photon of light. (See
Figure 6.4) Each photon has a specific wavelength that is dependent on the energy differences between each
energy state. This light energy is what is seen on an emission spectrum of an element. Larger energy differences in
the atom will result in wavelengths of light with higher energy on the emission spectrum. The spectrum shown in
Figure 6.3 shows only the visible portion of hydrogen’s emission spectrum. These colors represent electron
transitions that all originate from energy level energy 2. (See Figure 6.4) Electron transitions from other energy
levels will result in higher energy waves or lower energy waves in the range of ultraviolet and infrared waves.


Match each color in Figure 6.3 with the appropriate electron transitions in illustrated in Figure 6.4. (Hint: how
does wavelength relate to energy? All transitions originate from energy level 2.)
Transitions from which energy level would likely result in ultraviolet radiation?
12
The Modern View of Atomic Structure
Atoms consist of a tiny, dense, positively charged nucleus surrounded by a cloud of negative electrons.
The nucleus contains most of the mass of the atom and consists of positively charged protons and neutral neutrons.
Atoms are electrically neutral because each atom contains the same number of protons as electrons.
Atoms gain electrons to form negatively charged ions called anions or they can lose electrons to become positively
charged ions called cations.
The atomic number is the number of protons in the nucleus.
An element is a substance all of whose atoms contain the same number of protons. Each element is defined by its
atomic number.
The mass number is the number of protons and neutrons in the nucleus of an atom.
Isotopes are atoms containing the same number of protons but different mass numbers. Since the protons in an atom
will never change, any difference in mass is due to differences in the neutron number.
Symbols are often used to distinguish one isotope from another. For example, the isotope of carbon containing six
protons and six neutrons is designated as:
12
6𝐶
where 12 is the mass number and 6 is the atomic number. Because carbon is the only element with 6 protons, often the
atomic number designation is omitted. This isotope of carbon can be represented by any of the following:
12
6𝐶
12
𝐶
=
=
carbon – 12 =
C – 12
Carbon has several isotopes, and each is distinguished from one another by its mass number. Their respective number
of neutrons is calculated by subtracting the atomic number from the mass number. Here are the symbols for various
isotopes of carbon with the number of neutrons in the isotope.
11
6𝐶
5 neutrons
12
6𝐶
6 neutrons
13
6𝑐
7 neutrons
14
6𝑐
8 neutrons
The atomic mass unit, amu, is a convenient way to express the relative masses of tiny atoms. One amu is defined as
one-twelfth of the carbon-12 isotope. Protons and neutrons have a mass of approximately 1 amu each. To put this into
perspective, one amu also equals 1.66054 ×10-24 g which is an extremely small mass.
The atomic mass of an element is provided on any periodic table along with the atomic number of each element.
Elements as they are found in nature typically exist as a mixture of several different isotopes each having the same
number of protons but different number of neutrons. (Recall: the proton number is constant for all atoms of the same
element). Therefore, some atoms in this element are slightly heavier than other atoms. As a result, these atoms have
13
different mass numbers. The atomic mass that is provided on the periodic table is actually a weighted average of all the
isotopes of an element based on the abundance of each isotope found on earth. Because this is a weighted average,
isotopes that are more abundant in nature will contribute more to the average mass that is reported on the periodic
table. It is not a coincidence that the most common isotope of carbon is carbon-12 and the mass reported on the
periodic table for carbon is 12.011 amu. The atomic mass of magnesium is 24.3040 amu. This means that the most
common isotope of magnesium would have a mass number of 24. Other isotopes of each element must be determined
using a variety of other sources such as the internet. It is important to note that NO individual atom of magnesium has a
mass of 24.3040 amu. This is merely an average mass of all of the combined isotopes of magnesium. We must also
remember that the mass number must be a whole number (no decimal digits) since it is the combined number of
protons and neutrons. In our magnesium example above, magnesium’s atomic mass of 24.3040 amu would correspond
to a mass number of 24 amu for magnesium’s most common isotope.
The Periodic Table
The periodic table is an arrangement of elements in order of increasing atomic number with elements having similar
properties placed in vertical columns. The vertical columns are called groups or families and the horizontal rows are
called periods. Figure 2.15 shows the periodic table with the symbol, atomic number, and atomic mass of each element.
It also shows two commonly used numbering systems for the groups. The “A” group elements are also commonly
designated with roman numerals. Elements are also classified as metals, nonmetals, and metalloids based on their
position on the periodic table. See Figure 2.15.
What are the metalloids listed in Figure 2.15?
14
Molecules and Molecular compounds
Although the atom is the smallest representative particle of an element, most matter is composed of molecules or ions,
which are combinations of atoms.
A molecule is an assembly of two or more atoms tightly bonded together. For example, a molecule that is made up to
two atoms is called a diatomic molecule. If the two atoms are identical, we have a diatomic element. Figure 2.18
shows the names, formulas, and pictorial representations of some simple molecules. Notice that both H2 and CO are
diatomic. However, H2 is a diatomic element because both atoms are identical. CO is a diatomic molecule. CO has
different properties than both the carbon and oxygen that make up the molecule.
15
Compounds are substances consisting of two or more different elements. Generally there are two types of compounds:
molecular compounds and ionic compounds.
Molecular compounds are composed of molecules and usually contain only nonmetals. A molecular formula indicates
the actual number and type of atoms in the molecule and is the most often used formula for molecular compounds.
Figure 2.19 shows various ways chemists represent molecular compounds.
Figure 2.19
16
Ions and Ionic compounds
Ions are charged particles composed of single atoms (called monatomic ions) or aggregates of atoms (called polyatomic
ions). A cation is a positive ion, and an anion is a negative ion.
Ionic compounds are composed of ions and usually contain both metals and nonmetals.
An empirical formula gives only the relative number of atoms (lowest ratio of atoms) of each type in the compound.
Empirical formulas are usually used for ionic compounds.
Unlike molecular compounds, ionic compounds do not consist of discrete molecules. Ionic compounds are a collection
of many ions arranged in a regular pattern known as a crystal lattice as shown in Figure 2.21. Rather than draw the ionic
arrangement of sodium chloride, chemists use the much simpler formula, NaCl to represent the compound.
Names of monatomic cations
Monatomic cations (single atom ions with a positive charge) have the same name as the metal. If the metal forms
cations of different charges, a Roman numeral in the name indicates the charge.
K+ = potassium ion
Ca2+ = calcium ion
Co2+ = cobalt (II) ion
Co3+ = cobalt (III) ion
Most transition metals form cations with more than one charge. However there are four common exceptions. These
are the four transition elements that have only a single charge. You must memorize these charges because there will
not be a roman numeral present to indicate their charge.
Sc3+ = scandium ion Ag+ = silver ion Zn2+ = zinc ion
Cd2+ = cadmium ion
17
Names of monatomic anions
When naming a monatomic anion, change the ending to “ide”.
A few common diatomic anions also end in ide: OH1- hydroxide CN1- cyanide O22- peroxide
Names of Polyatomic cations
Polyatomic cations (positive ion containing more than one atom): ammonium NH4+
Hg22+ mercury (I)
Polyatomic oxyanions (anions containing oxygen)
Some polyatomic anions end in ate or ite. These are called oxyanions because they contain oxygen. Because there is no
logical way to predict their formulas or charges, they must be memorized. Some polyatomic oxyanions have hydrogen
attached. The word hydrogen or dihydrogen is used to indicate anions derived by adding H+ to an oxyanion. An added
hydrogen changes the charge by +1. See Table 2.3.
The ending –ate denotes the most common oxyanion of an element. The ending –ite refers to an oxyanion having the
same charge but one fewer oxygen. For example: SO42- is sulfate, but SO32- is sulfite. Other patterns are seen by
examining the chlorate family.
ClO- = hypochlorite; hypo denotes one fewer oxygen than chlorite
ClO2- chlorite
ClO3- chlorate
ClO4- perchlorate; Per- denotes one more oxygen than chlorate
*the other halogen oxyanions follow the same pattern.
18
Naming ionic compounds
Review the following summary or refer to your Chemistry I Notes
 Metals in Groups IA, IIA, and IIIA will form ions that have charges of 1+, 2+, and 3+ respectively. These positive
ions have only one possible charge and are named by using the name of the element.
 Metals (or nonmetals) that have several positive oxidation states must indicate the charge of the ion by writing
a Roman numeral in parentheses following the name of the element. Common transition metals with only one
oxidation state are: Zinc (+2), Cadmium (+2) and silver (+1) . Most others will have a roman number.
 Negative, monatomic (one atom) ions are named by using the root word of the nonmetal and adding the suffix
-ide. Nonmetals in Groups VIIA, VIA, VA and IVA will form ions that have charges of 1-, 2-, 3-, and 4- respectively
in binary (two element) compounds.

The charges of the –ate and –ite ions are usually the same as that of the –ide ions. We can determine the “ide”
ions from their positions on the periodic table. For example Sulfide has a -2 charge and both sulfate and sulfite
have a -2 charge. Other examples: phosphide = -3 and phosphate and phosphite = -3, chloride ( or any
halogen) = -1 and chlorate, chlorite, hypochlorite and perchlorate are all -1. Exceptions: Nitrate, carbonate.
There is no correlation to the periodic table for these polyatomic ions. These must be memorized.

*Note: you must memorize the common –ate polyatomic ions: sulfate, carbonate, nitrate, phosphate,
chlorate, bromate, and iodate. You can then use the pattern described above to determine other
polyatomic ions. Other ions that need to be memorized are hydroxide, acetate, cyanide, and
ammonium. Make flashcards.
Names of ionic compounds consist of the cation followed by the anion name.
MgBr2 = magnesium bromide
Fe2O3 = iron (III) oxide
Ca3(PO4)2 = calcium phosphate
FeO = iron (II) oxide
*Remember that the Roman numbers III and II refer to the +3 and +2 charges of the iron cations, respectively. Notice
the charge on the iron atom in each case is inferred by the ratio in which iron combines with the -2 charged oxide ion.
Hydrogen will take on a -1 charge when bonded to groups 1 or 2. These compounds are known as hydrides.
Names of binary molecular compounds
A binary molecular compound contains two nonmetals. The rules for naming binary molecular compounds are the
following:
1. Name the first element
2. Name the 2nd element
3. Use a prefix to denote how many of each element is present in the formula.
Ex. SO3 = sulfur trioxide
N2O5 = dinitrogen pentoxide
P4O10 = tetraphosphorus decoxide
19
Names and formulas of Acids
The names and formulas of some common acids, an
important class of hydrogen containing compounds
are listed in Table 2.4. Acids are molecular
compounds (even though they may contain a
polyatomic ion). In all cases, the formulas of acids
are composed of one or more hydrogens added to a
common monatomic anion or oxyanion. Acids of
monatomic ions are called binary acids and acids of
oxyanions are called oxyacids.
To name binary acids (Hydrogen + a nonmetal),
replace the –ide ending of the anion with –ic acid
and add the prefix hydroEx. Bromide, Br- becomes hydrobromic acid, HBr.
To name oxyacids (Hydrogen + oxyanion), replace the –ate ending of the oxyanion with –ic acid or the –ite ending of the
oxyanion with –ous acid
Ex. Nitrate, NO3- becomes nitric acid, HNO3
Hypochlorite, ClO- becomes hypochlorous acid, HClO
Note: Phosphate, PO43- becomes phosphoric acid, H3PO4
Sulfate, SO42- becomes sulfuric acid, H2SO4
You must be fluent in writing the formulas of various acids.
Naming hydrates??? Know your prefixes….
Name the compound as is. Designate the number of water molecules with a prefix, then add the word
“hydrate”.
Example: CuSO4 • 5 H2O is called copper (II) sulfate pentahydrate
20
Organic compounds
Organic chemistry is the study of carbon compounds.
Hydrocarbons are molecular compounds containing only carbon and hydrogen
All carbon atoms must have a total of 8 electrons or 4 total bonds. (octet)
Organic compounds that have all single bonds are called saturated hydrocarbons; those with multiple bonds are
unsaturated.
Saturated hydrocarbons may consist of alkanes or cycloalkanes. Both alkanes and cycloalkanes have all singe bonds.
Unsaturated hydrocarbons consist of alkenes, cycloalkenes, or alkynes. Alkenes and cycloalkenes contain a double
bond. Alkynes contain a triple bond.
Alcohols are formed when a hydrogen on an alkane is replaced with a hydroxyl group, OH. (…note this is not the
hydroxide ion as is used in inorganic chemistry). Many organic compounds have other elements such as oxygen,
nitrogen, phosphorus, or halogens that are referred to as “functional groups”. Functional groups give organic
compounds a unique set of properties. The presence of the hydroxyl group on an organic compound tells us that the
compound is an alcohol.
21
Assignment 2 Review Questions
You MUST print this assignment to turn in to me.
1. Refer to your notes to fill in the following table with the appropriate family names
Group number
Name of group/family
Family I (1A)
Family II (2A)
Family VI (16 or 6A)
Family VII (17 or 7A)
Family VIII (18 or 8A)
2.
What are the seven diatomic elements? Provide both name and appropriate formula.
3. Consider the four systems in the figure:
A
B
C
D
a. Which of the above system(s) represent a pure substance?______________________________
b. Classify the pure substances that you labeled in “a” as an element or a compound.________________
c. Which system is a mixture?_____________________
4. Explain why Dalton’s 2nd postulate is incorrect.
5. What experimental evidence disproved Thomson’s model that the atom was solid?
6. _______________________What is Thomson’s model called?
22
7. _______________________What is Rutherford’s model called?
8. Describe the atom according to Rutherford’s findings.
9. How did Niels Bohr explain why electrons did not fall/collapse into the nucleus? What experimental evidence
verified his claim?
10. What is atomic mass? How is it different from mass number?
11. How does a molecular formula differ from an empirical formula?
12. Tell which chemical form of chlorine is implied in these two sentences: a) chlorine is a toxic gas b) common
table salt contains the element chlorine.
13. What is an oxyanion?
14. What would be the name for the following oxyanions:
A. __________________ BrO- b. _____________________IO4- c. _____________________ IO215. List the prefixes up to ten.
16. ____________________________________________What is the name of MgSO4 • 7 H2O?
17. ______________________Look up the formula for ammonia.
18. Name the following molecular compounds:,
a. _________________________________ SCl2
b. _________________________________ CF4
c.
______________________________________________________BrI3
d. _________________________________ SF6
19. Name the following acids (practice without polyatomic ion list then check yourself)
a. _________________________HClO4
e. _________________________H3PO3
b.
_________________________HCl
f.
_________________________CH3COOH
c.
_________________________HNO3
g.
_________________________H2S
d.
_________________________H2SO4
23
20. Write the formulas for the following:
h.
_______________Nitrous acid
l.
_______________Hydrobromic acid
i.
_______________Phosphoric acid
m. _______________Perbromic acid
j.
_______________Hydroiodic acid
n.
k.
_______________Sulfurous acid
_______________Bromous acid
21. Name the compounds in problems a-g using a simple periodic table and no polyatomic ion list. You
can check your work afterwards with your polyatomic ion sheet. For examples h – n write the
correct formulas.
a. _____________________________Na3PO4
h. _______________Mercury I chloride
b. _____________________________FeSO4
i.
_______________Sodium sulfate
c. _____________________________Al(ClO4)3
j.
_______________Copper II nitrate
d. _____________________________MgCO3
k. _______________Iron II hydroxide
e. _____________________________Zn(OH)2
l.
f.
m. _______________Sodium nitrate
_____________________________CuC2H3O2
g. _____________________________Sr(OH)2
_______________Ammonium acetate
n. _______________Aluminum sulfite
24
Refer to your chemistry I notes or research the answers to questions 21-26
22. What are the 10 simplest Alkanes? Write their names, molecular formulas, condensed structures, and line
structures on a separate sheet of paper. (Refer to your notes for questions 26 -30)
23. What are the 6 simplest alkenes? (Names, molecular formulas, condensed structures and line structures.)
24. What are the 6 simplest alkynes? (Names, molecular formulas, and condensed structures.)
25. What are the 6 simplest cycloalkanes?
26. What are the three simplest alcohols?
27. The following symbols represent atoms of carbon, hydrogen, and oxygen (in no particular order).
a. Using your knowledge of periodic trends, determine which atom is which.
b. Draw a particulate representation (model) of a methane molecule, an oxygen molecule, carbon dioxide
molecule, and a water molecule
c. Draw a particulate representation describing the reaction of the combustion of methane. Be sure to make sure
that all mass is conserved.
25
Assignment 3: Quantum Mechanical Model of the atom
The review questions for this assignment are embedded throughout the reading passage. You will have to
answer these questions on a separate sheet of paper. Be sure to clearly label each assignment.
Many great philosophers since 400 B.C believed that matter was composed of tiny atoms. However this concept
was based solely on intuition rather than experimentation. The evidence to support the atomic theory has only been
available since the early 1900’s through the study of chemical reactions. This early experimental evidence led John
Dalton to put together the first atomic theory. Although Dalton’s theory had to be modified, it has withstood the test
of time extremely well.
Once the scientific world came to believe in the idea of the atom, it was logical to ask: What is the nature of an
atom? Does an atom have parts, and if so, what are they? In this assignment we will explore how the atomic theory has
evolved into its present state. We will see that the modern theory of atomic structure accounts for the periodic trends
(periodicity) that occur because of the electron arrangement in atoms.
However before we examine atomic structure, we must consider the revolution that took place in physics during
the first 30 years of the 1900’s. During this time, experiments were conducted that could not be explained by classical
physics. Isaac Newton and all of this followers thought they had it all figured out. Their theories could explain
phenomenon as diverse as the motion of the planets and the dispersion of visible light through a prism. It was rumored
that students were being discouraged from pursuing physics as a career because it was felt that all the major problems
had been figured out by the current theories of the time. Through various mathematical formulas, physicists were able
to accurately describe the behavior of matter on a macroscopic level. However, when scientists attempted to explain
atomic structure using basic knowledge of physics, the experimental observations did not agree. Consequently, a new
branch of physics was necessary in order to account for the behavior of the atom. This “new physics” became known as
“quantum physics”.
Contradictions with classical physics came through the study of light energy (electromagnetic radiation). An
electromagnetic wave consists of various parts and propagates through space at a speed of 3× 108 m/s. Based on the
properties of a wave, various mathematical formulas were used to convert the frequency of a wave to its wavelength (or
vice versa) as discussed in chemistry I. (recall: c = λν)
Example: Calculate the following for practice: What is the frequency of a light wave having a wavelength of 555 nm?

When calculating frequency from wavelength, be sure that the units for the wavelength match the units for the
speed of light. In this case we must convert nm to meters.
Answer: frequency = 5.41 × 1014 Hz
26
Assignment 3a
Concept review…. Google Khan Academy: atomic structure, introduction to light, electromagnetic waves and the
electromagnetic spectrum to review these basic concepts. Then answer the following questions on a separate sheet of
paper.
1) Why is light energy also referred to as electromagnetic radiation?
2) Describe the properties of a wave in terms of amplitude, wavelength, energy, frequency, and speed. Show an
illustration in your discussion.
3) Describe the major divisions of the electromagnetic spectrum and their relative energy, frequency, and wavelength.
4) The brilliant red colors seen in fireworks are due to the emission of light with wavelengths around 650 nm when
strontium salts such as Sr(NO3)2 and SrCO3 are heated. Calculate the frequency of red light with a wavelength of 650
nm. Answer: 4.61× 1014 Hz. Show all work…
At the end of the 19th century, it was assumed that there was no intermingling of matter and light. Matter was
thought to consist of particles, whereas energy was described as a wave. Particles were described as entities that had
mass and whose position in space could be specified. Waves, on the other hand, were described as mass-less and
delocalized: that is, their position in space could not be specified. Waves also exhibited two properties known as
interference and diffraction.
Interference and Diffraction
Waves interact with each other in a characteristic way called interference: they can either cancel each other out or
build each other up, much like ocean waves. If waves coming from two sources align so that their crests overlap, they
are said to be in phase. Consequently, the amplitude is larger. (See Figure below) This is called constructive
interference. On the other hand, if the waves are completely out of phase (if they align so that the crest from one
source overlaps the trough from the other source) the waves cancel by destructive interference.
Constructive Interference
Destructive Interference
27
In addition, as is the case with ocean waves, when light waves encounter an obstacle or a slit they will bend around it, a
phenomenon called diffraction. Since light exhibits these characteristics, it has long been assumed that light was
thought to be purely a wave phenomenon.
At the beginning of the twentieth century, however, experimental evidence began to reveal that light can also
exhibit particle-like nature. This concept was suggested by a German physicist named Max Planck. Planck’s
experimentation suggested that light energy can come in the form of little packets/bundles of energy rather than one
continuous wave. Planck referred to this bundle of energy as a quantum of energy. Planck’s findings were a surprise.
Now it seemed clear that energy is quantized, that is, it could be described as having a discrete value. Each little
“packet” of energy was a whole number multiple of hν where h is a constant called Planck’s constant which was
experimentally determined to be 6.626× 10-34 J/s and ν is the frequency of the electromagnetic radiation absorbed or
emitted. Through his work, Planck was able to calculate a value of a quantum of energy using the equation: E = hv
Assignment 3b: The blue color in fireworks is often achieved by heating copper (I) chloride (CuCl) to about 1200°C. Then
the compound emits blue light having a wavelength of 450 nm. What is the increment of energy (the quantum of
energy) that is emitted at 4.5× 102 nm by CuCl. Answer: 4.41× 10-19 J. Show all work on a separate sheet of paper.
From the calculation above, we can see that a sample of CuCl emitting light at 450 nm can lose energy only in
increments of 4.41× 10-19 J, which is the size of a quantum in this case (or any whole number multiple of this value).
28
Albert Einstein expanded the idea of light acting as a particle through his analysis of the photoelectric effect.
The photoelectric effect refers to a phenomenon in which electrons are emitted from the surface of a metal when light
strikes it. These electrons are then used to generate an electric current. It was previously thought that light will
dislodge electrons from a metal much like an ocean wave might dislodge a rock from a cliff. With enough energy, the
rock will be broken off. This energy might come in the form of one huge wave, or many successive smaller waves. In
other words, even if the waves were smaller, eventually there would be enough energy to cause the rock to break off.
Likewise, even light of low intensity should be able to knock off electrons if the metal was exposed long enough to the
energy. Einstein however proved this not to be the case. It was found that each metal required a specific minimum
frequency of light, known as the threshold frequency, that allowed the electrons to be ejected from its surface.
Anything above this minimum frequency caused the electrons to be ejected. Lower frequency light would not eject
electrons regardless of how long the metal was exposed to the energy. For example, in viewing Figure 1 , we can see
that a wavelength of light at 700 nm will eject no electrons, but wavelengths of 550 nm and 400 nm (which are higher in
energy) are sufficient enough to eject the electrons. Moreover, as long as the minimum frequency is met, shorter
wavelengths of light will produce more kinetic energy. (i.e.. 400 nm of light produces a greater velocity of kinetic
energy than 550 nm…. Recall the shorter the wavelength the higher the energy).
Figure 1
The findings of Albert Einstein reinforced the findings of Planck in that light energy must come in
packets/bundles of energy rather than in the form of one continuous wave. These little packets can be compared to
little BB’s fired from a BB gun. Each BB has its own specific amount of energy. Likewise each little packet of light energy
must have certain minimum amount of energy needed to knock off electrons. Einstein used Planck’s idea and called
these packets of mass-less particles photons. Einstein suggested that if each photon had sufficient energy to dislodge an
electron, even low intensity light could be used since each photon had the minimum frequency needed for ejection.
From Einstein’s perspective, a beam of light may not always act like a wave propagating through space, but instead a
shower of particles. (See Figure2 ) The energy of each photon was directly related to its frequency. The photoelectric
effect is utilized today in solar panels where radiation from the sun knocks electrons away from metal sheets. Einstein
eventually won the Nobel Prize for this explanation of the photoelectric effect.
Figure 2: Einstein’s view of light: light consisted of distinct “particles” of electromagnetic
radiation. Each little packet of energy was called a photon.
29
According to Einstein, each little packet/bundle of energy had a specific wavelength and consequently a specific
amount of energy. The energy of each photon can be calculated using the expression: E =
𝒉𝒄
𝝀
where E is the energy of
a photon of light, h is Planck’s constant, and λ is the wavelength of the radiation.
Assignment 3c:
Google Khan’s Academy: photoelectric effect. Then answer the following questions. Turn in when complete.
1) Describe the photoelectric effect (in your own words)
2) Explain how the wave theory was not able to explain this phenomenon (in your own words)
3) Describe Einstein’s explanation of this phenomenon and how it implied the particle nature of EM radiation. (in your
own words.
Through Einstein’s observations, the following conclusions were made:
1) Each different type of metal required a specific minimum frequency of light to discharge the electrons
2) For light with frequencies lower than the threshold frequency, no electrons are emitted regardless of the intensity of
light (no matter how long the light was applied)
3) For light with frequency greater than the threshold frequency, the number of electrons emitted increases with the
intensity of light.
4) For light with frequency greater than the threshold frequency, the kinetic energy of the emitted electrons increases
linearly with the frequency of light. (light with higher frequencies could emit more electrons than light with lower
frequencies as long as the threshold frequency was met. i.e. violet light could emit more electrons than red light as
long as both frequencies were above the minimum frequency required for that metal) See Figure 1.
In a related development, Einstein derived his famous equation: E = mc2 as part of the theory of relativity. Although we
won’t go into the dynamics of this theory, the main significance of this equation is that energy has mass. This is more
𝐸
apparent if we rearrange the equation in the form: m = 𝑐 2. From this arrangement, we can see that if we know energy,
mass can be calculated by dividing by the speed of light squared. So in terms of light energy, if we know the wavelength
of light, we can calculate the energy of that photon using E =
ℎ𝑐
𝜆
. The energy can then be used to calculate its mass.
This tells us that a photon of energy acts like a particle because it has mass. However, it is also very clear that the mass
of a photon is VERY, VERY, VERY small in a classical sense.
However despite the work of Planck and Einstein, we cannot completely abandon our view of light as a wave. Many
experiments continue to validate its wavelike properties. One such experiment is the double slit experiment which is
illustrated in Figures 3 and 4.
30
Figure 3 : Light as Particle in the double slit
experiemnt
Figure 4: Light as a wave in the double slit experiment
If light were to exist only as little particles, these particles would pass through a screen with two slits as illustrated in
Figure 3 above. This however is not what happens. The findings that we actually get are illustrated in Figure 4. The
resulting pattern occurs because of the diffraction and interference of light waves as they interact with each other.
Assignment 3d: Google Khan’s Academy: Interference, and double slit experiment. View the videos provided and
answer the questions below. Answer all questions in your own words. Turn in as you go.
1) Explain diffraction and interference in terms of waves.
2) Discuss the differences between constructive and destructive interference patters
3) Explain what happens when light is actually passed through two slits
4) What is meant by the term diffraction patterns?
5) What do the dark areas indicate on a diffraction pattern?
6) What do the bright spots indicate on a diffraction pattern?
7) Explain how a particulate view of EM radiation could not explain the concepts of diffraction and interference.
We can summarize the work of Planck and Einstein as follows:
1) Energy is quantized. It can occur only in discrete units called quanta.
2) Electromagnetic radiation, which was previously thought to exhibit only
wave properties, seems to show certain characteristics of a particle-like
state as well. This phenomenon is referred to as the dual nature of light.
See Figure 5.
Figure 5: dual nature of light
31
Thus light, which previously was thought to be purely wavelike in terms of Classical Physics was now found to
have certain characteristics of a particle as presented by Planck and Einstein. This breakthrough began to revolutionize
the thinking of scientists at the time. Now scientists began to ask the question: Is the opposite also true? Can matter
also exhibit wavelike properties? This question was raised by a young French physicist named Louis de Broglie. De
Broglie found that if we combine the two equations provided by Einstein and Planck’s work we can derive the following
ℎ
equation: m = 𝜆𝜈. Rearranging for wavelength, we would get
ℎ
λ =𝑚𝑣. This equation, known as the de Broglie
equation, allows us to calculate the wavelength of a moving particle with a given mass.
Assignment 3e: Compare the wavelength for an electron (mass = 9.11× 10 -31 kg) traveling at a speed of 1.0 ×107 m/s
with that for a ball (mass = 0.10 kg) traveling at 35 m/s using the de Broglie equation. Show work on a separate sheet.
Wavelength of the electron… show work
Wavelength of the ball… show work
Answer: electron = 7.27× 10-11 m; ball = 1.9 × 10-34 m
Notice that the wavelength of a moving ball is incredibly short…. Impossible to detect. The electron on the other
hand, although still quite small has a wavelength in proportion to its submicroscopic size. We can test de Broglie’s
equation by studying diffraction patterns.
Diffraction patterns can only be explained in terms of waves. Thus, if electrons possess wavelike properties,
they should also possess the ability to be bent in the same way that light waves do. Experimental observations show
that this is exactly the case. Electrons do indeed produce a diffraction pattern similar to other forms of light energy.
This finding concludes the fact that electrons, which are particle in nature, do exhibit wavelike characteristics. Wow…
we have now come full circle. Electromagnetic radiation thought to be clearly a wave, was found to have particulate
properties. And electrons, only thought of as tiny particles, also exhibit wavelike characteristics. However, it still must
be said that larger chunks of matter have such small wavelengths that they are impossible to verify experimentally.
This type of matter (such as our ball) will exhibit predominantly particulate properties since their wavelength’s are
impossible to observe. Photons, which are very small “bits of matter”, while showing some particulate properties,
exhibit predominantly wave properties. But the electrons are just the right size to exhibit both wavelike and particle
like behavior. (Whew…. That was deep…. )
32
Back to the atom…..
Assignment 3f: Use the following reading passage to answer the questions below on a separate sheet of paper. (in
your own words).
1) What is atomic spectroscopy?
2) What is a bright line emission spectrum?
3) Describe atomic emission spectroscopy and how an atomic spectrum differs from the familiar white light spectrum.
4) Describe how Bohr was able to use the experimental evidence of atomic spectroscopy to construct his model of the
atom (in your own words).
5) Discuss the differences between the Balmer series, the Paschen series, and the Lyman series.
In the early days, it was clearly evident that when white light is passed through a prism, a continuous spectrum
of visible colors are seen (ROY G BIV) where one color blends into another color as the wavelength becomes longer and
longer. The study of light as it is passed through a prism is referred to as atomic spectroscopy. When studying specific
elements, it was found that when a source of energy was applied, the element gave off light energy. However, when
this resulting light was passed through a prism, it did not result in a continuous spectrum as did white light. Instead an
incontinuous spectrum was produced. This spectrum of incontinuous light is referred to as a bright line emission
spectrum (or just “line” spectrum or emission spectrum for short) The earliest breakthrough came in the study of
hydrogen. Passing the light emitted by excited hydrogen atoms resulted in only 4 visible lines each having its own
unique wavelength. See Figure 6.
Figure 6: Emission spectrum of hydrogen
You may be wondering, so what?? What is so significant of the
line spectrum of hydrogen? Well then…. Here it is…..These
incontinuous line spectrums tell us that only certain energies
are allowed for the electrons in a hydrogen atom. In other
words, the energy of the electron in the hydrogen atom is
quantized. This observation ties in perfectly with the postulates
of Max Planck. The lines seen are evidence of electrons
changing energy levels as they become excited and then drop
back down to their ground state. The ground state is the
lowest energy state of an electron. Excited states are states of
higher energy than the ground state. Energy in the form of a
photon is emitted or absorbed by an electron only when it
changes from one allowed energy state to another. The lines of
the atomic emission spectrum of hydrogen result when an electron falls from a higher allowed state to a lower allowed
state (such as n = 3 to n = 2). The closer the energy states are to each other in an atom, the lower the energy of the
photon produced (i.e. the longer the wavelength of the light wave).
Since only certain lines are observed, there are only certain energies possible to make these transitions occur. If
any amount of energy would be allowed, a continuous spectrum would be observed rather than an incontinuous
spectrum. Atomic spectroscopy and the line spectrum of hydrogen led to the development of the Bohr model.
33
The Bohr model
In 1913, a Danish physicist name Neils Bohr used the evidence described above to develop a “quantum” model of the
hydrogen atom. Bohr proposed that the electron in a hydrogen atom moves around the nucleus only in certain allowed
circular orbits. Each orbit had a distinct energy that was represented by the line spectrum of hydrogen. See Figure 7.
Figure 7: The Bohr model
The principal energy levels or energy states that an electron can
exist in are represented as n. As electrons become excited they
can only absorb a specific wavelength of light to transition from
one energy state to another. When these electrons drop to a
lower energy state they give off photons of light with
wavelengths that correspond to the experimental evidence
presented by the emission spectrum of hydrogen.
The four visible lines that we see on hydrogen’s emission
spectrum have just the amount of energy needed for electrons
to transition back down to energy level 2. These wavelengths of
light are referred to as the Balmer series and only show
transitions that are visible to the eye. Going from energy state
3 to energy state 2 would give off the least amount of energy. Since red light has the least amount of energy, it would
have the longest wavelength of light. An electron in energy level 6 transitioning to energy level 2 would give off the
most energy. This associated energy would give off a photon of light in the violet range of visible light (since violet light
has the highest energy of all visible colors). See Figures 8 and 9.
Figure 9: Bright Line Spectrum and Bohr model
Figure 8: The Balmer Series
Examine Figure 8. Notice how close energy levels 4-6 are to each other. Also notice how far energy level 1 is
from energy level 2. Transitions that occur when electrons drop to energy level 1 would give off much more energy.
34
This quantity of energy would have wavelengths shorter than those of visible light resulting in higher energy. The
resulting light released in this case is in the ultraviolet region. (Recall: ultraviolet light has more energy than violet)
Consequently, these transitions were not noticed for a while because UV waves are undetectable by the naked eye.
Transitions that show the ultraviolet transitions are referred to as the Lyman series. The same holds true for transitions
to energy level 3. The energy states are so close to each other giving off less energy; these transitions result in
wavelengths longer than red light and the resulting emissions are in the infrared region. This is referred to as the
Paschen series. (see Figure 10)
Figure 10: Paschen, Balmer, and Lyman Series of a hydrogen atom
At first Bohr’s model was very promising. The energy levels calculated by Bohr closely agree with the values
obtained from hydrogen’s emission spectrum. However, when Bohr’s model was applied to atoms other than hydrogen,
it did not work at all. Some attempts were made to adapt the model by
Figure 11:
constructing elliptical orbits. The result of this work is often used as an
illustration of the current model as seen in Figure 11, but we know that
this model is also fundamentally incorrect. (although it looks pretty
cool…)
Even though the Bohr model is fundamentally incorrect, it was very important historically because it showed that the
observed quantization of energy in atoms paved the way for later theories.
35
As time went on, de Broglie along with two other scientists: Werner Heisenberg and Erwin Schrodinger developed an
approach using the electron’s wavelike properties. This model is known as the wave mechanical model or the quantum
mechanical model. De Broglie originated the idea that the electrons possessed wavelike properties, when it was
previously thought of as a particle. These waves could exhibit constructive and destructive interference patterns as all
waves do. Schrodinger went on to develop a series of equations (much too complicated to go into here), and the
solutions to his equations were referred to as wave functions. Each wave function is characterized by a particular
value of energy. According to the wave mechanical model, the electron’s position was arranged around the nucleus in a
3-D arrangement along the x, y, and z axes of a Cartesian graph. These wave functions are referred to as orbitals. An
orbital is a region in space along the axis of a graph where the probability of finding an electron is the greatest. An
orbital is NOT the same as a Bohr “orbit”. Electrons do not move in a circular orbit around the nucleus. Well then.. how
are they moving you might ask???? Well the answer to this is idk…. We don’t know. When we view matter
macroscopically, we are able to predict their pathways. For example when two billiard balls with a known velocity
collide, we can predict their motions after collision using classical physics. However we cannot predict the pathway of
an electron. The wave function gives no detailed pathway on the movement of the electron. This is quite disturbing
since we want to know the exact location and movement of the electron.
To help us understand the nature of an orbital, we need to consider a principle discovered by Werner
Heisenberg. Heisenberg’s analysis led him to a surprising conclusion: There is a fundamental limitation to just how
precisely we can know both the position and momentum (velocity) of the electron at any given time. We may be able to
establish its position at one point in time, but in order to do this we must apply some sort of disturbance. When this
happens, its momentum is altered. This theory/statement is referred to as the Heisenberg Uncertainty principle. The
more accurately we know a particle’s position, the less accurately we know its momentum. (and vise versa) This
limitation is so small for macroscopic objects such as baseballs that it is unnecessary. However when applied to the
electron, the uncertainty principle states that since it is impossible to know the exact motion of the electron as it moves
around the nucleus, we cannot assume that it is moving in a circular, well defined orbits as implied by the Bohr model.
The Physical meaning of a Wave Function
Given the limitations set forth by Heisenberg, what then is the physical
meaning of a wave function for an electron. That is, what is an atomic orbital?
Without getting into too much detail here, the wave function gives us no visual
meaning, but if we square the wave function, what we can determine is the
probability of finding an electron near a particular point in space. We have no
information of how the electrons move in space between two different
positions, but at least we can determine the probability of finding an electron in
a certain position. The probability of finding hydrogen’s electron is seen in
Figure 12. If we plot the solutions to Schrodinger’s equations on a graph, we see
areas of high density and low density. The areas of high density is where the
probability of finding the electron is the greatest.
Figure 12: 90% probability circle
of a 1s orbital on Cartesian graph
36
Figure 13 Distance of 1s orbital from the nucleus
When we use the term atomic orbital, we are picturing the electron
density- a region where we might find the electron 90% of the time
(which is called a 90% probability circle). Now think of this
probability (density) graph three-dimensionally with the nucleus in
the center. The electrons are likely to be located at some distance
away from the nucleus. Notice in figure 13, that the greatest
probability of finding the electron in a 1s orbital would occur at a
distance of 0.053 nm away from the nucleus. As we move away
from 0.053 nm, the probability of finding the electron diminishes;
but yet it is still quite possible of finding it at other distances… just
not as probable as finding it at a distance of 0.053 nm. In other
words, in an atom, the closer we approach 0.053 nm to more likely
we will find the electron in a 1s orbital. Bohr’s model was very accurate in predicting the energy state of a hydrogen
atom because hydrogen’s electron is found in the 1s orbital which is spherical in shape, just as Bohr predicted.
Quantum numbers
When we solve Schrodinger’s equations, there are many wave functions (orbitals) that satisfy it. Each of these orbitals
are characterized by a series of numbers called quantum numbers, which describe various properties of the orbital. The
probable location of these orbitals are summarized below:
The principal quantum number (n) has integral values: 1,2,3, etc. This quantum number is related to size and energy of
the orbital. As n increases, the orbital becomes larger and the electron spends more time farther from the nucleus. This
means that it has more energy and is less attracted to the nucleus.
l
The angular momentum quantum number ( )…. This quantum number is related to the shape of atomic orbitals.
l are considered to be in the same subshell or sublevel. The value of l for a
particular orbital is commonly assigned a letter: l = 0 is called an s orbital; if l = 1 then we have a p orbital; l = 2 is a d
orbital; and if l = 3 we have an f orbital. All orbitals with the same angular momentum quantum number have the
Electrons having the same value for
same shape as we will see shortly.
The magnetic quantum number (ml) has values between
l and – l.. including 0. The value of ml is related to the
orientation of the orbital in space relative to the other orbitals in the atom.
37
The Quantum numbers for the first four energy levels of orbitals is summarized below:
N
1
2
3
4
l
0
0
1
0
1
2
0
1
2
3
sublevel designation
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
ml (-l - + l )
number of orbitals
0
0
-1, 0, 1
0
-1, 0, 1
-2, -1, 0, 1, 2
0
-1, 0, 1
-2, -1, 0, 1, 2
-3, -2, -1, 0, 1,2,3
1
1
3
1
3
5
1
3
5
7
Orbital shapes
We have seen that the meaning of an orbital is represented most clearly by a probability distribution which
represents a surface that surrounds 90% of the total electron density. In viewing the hydrogen atom, we gain
a visual as seen in Figures 14 and 15. Figure 14 shows the actual probability points. Figure 15 shows more of a
three-dimensional look of the spherical shape of the orbitals.
Figure 14: Probability circles of s orbitals
Figure 15: Shape of s orbital
All of hydrogen’s apparent orbitals are spherical in shape. Remember that although hydrogen’s electron
occupies its ground state 1s orbital, it can move to other orbitals when it is excited. Notice in the diagram
above that there are regions in a hydrogen atom where the probability of finding an electron is zero and there
are regions where we might find the electron 90% of the time. Areas where there is zero probability of finding
the electron are called nodal surfaces (spheres). These are the areas located between the 90% circle for each
energy level. Also notice that the number of nodal surfaces is characterized by n-1. For example, in
describing hydrogen’s 3s orbital we see that there are 2 areas of zero probability. See Figure 16. For our
38
purposes, we will focus in terms of their overall spherical shape which become larger as the value of n
increases.
Figure 16: Nodal surfaces of a s orbitals.
The p orbitals are a bit more complicated. They are not spherical like the s orbitals, but instead have
two lobes separated by a “node” at the nucleus. Figure 17 shows the probability circle of a single p
orbital. There are a total of three p orbitals which lie on different axis of the Cartesian graph.
Figure 17: 90% Probability circle of a “p” orbital
If we draw in the 90% circle we can see the shapes of each p orbital as they lie on each axis. See Figure
18. The p orbitals are labeled according to the axis of the “xyz” coordinate system along which the
lobes lie. For example the 2p orbitals with lobes centered along the x axis is called the 2p x orbital. The
orbitals lying on the y axis are labeled 2py, etc.
39
Although all p orbitals have the same shape, as the value of n
gets larger, the orbitals are larger and further away from the
nucleus. Another point to note: each of the p orbitals, 2p x, 2py
and 2pz in the second energy level are equal in energy. They
only differ on their orientation along the xyz axis. When we
look at the energies of the electrons that occupy these orbitals
they are all equal (meaning that each orbital is equidistant from
the nucleus. When this happens we say that the electrons are
degenerate with each other. Degenerate electrons have equal
energy. So an electron in a 2px orbital would have the same
energy (and is equidistant from the nucleus) as an electron in a
2py orbital. Electrons however in a 2px and a 3px are not
degenerate since they are not in the same energy level.
Likewise electrons in the 2s are not degenerate with electrons
in the 2p orbitals.
Figure 18: Shapes of “p” orbitals
along each axis
40
The d orbitals are even more complex than the p orbitals as seen in Figure 19:
Figure 19: Shapes of “d” orbitals
Four of the 5 orbitals have 4 lobes that lie in a
single plane that is indicated by the orbital label
that differ only in how they are positioned
along the xyz axis. The 5th orbital is unique with
2 lobes along the z axis and a belt centered in
the xy plane.
The f orbitals are even more complex than the
d orbitals. Since they aren’t involved in
bonding, we will not include them here.
41
Electron Spin and the 4th quantum number.
So far we have discussed the first 3 quantum numbers. The concept of the 4th quantum number was
introduced many years after Schrodinger’s work. In studying spectral data it was found that the electron has a
magnetic moment with two possible orientations when the atom was placed in an external magnetic field.
Since it was known from classical physics that a spinning charge produces a magnetic moment, it seemed
reasonable to assume that the electron could have two spin states, producing two oppositely charged
magnetic moments. The new quantum number that was adopted to describe this phenomenon is called the
electron spin quantum number (ms) and has values of + ½ and – ½ . This tells us that the electron can spin
in one of two opposite directions. For our purposes, the main significance of the electron spin is connected to
the postulates of an Austrian physicist named Wolfgang Pauli. Pauli’s contribution was known as Pauli’s
Exclusion Principle which states that in a given atom, no two electrons can have the same set of four quantum
numbers n, l, ml, and ms. Since electrons in the same orbital have the same value of n, l, and ml, then they
must have a different value of ms. And since there are only two values for ms, no two orbitals can hold more
than two electrons. If any orbital does hold two electrons, they must have opposite spins. More review on an
atoms 4 quantum numbers will be included with Ch. 8.
Assignment 3g: Answer the following on a separate sheet of paper.
What are the two properties of waves? Distinguish between each.
2. What is a photon?
1.
3. What does it mean for energy to be quantized?
4. According to Bohr, what is happening in the atom to cause the lines on an emission spectrum?
5. Distinguish between the excited state and ground state of an atom
6. What is a line spectrum? How is it used to identify an element?
7. Why does electron transitions from energy level one result in ultraviolet radiation?
8. What are areas in an atom where there is zero probability of finding an electron and where are they
located?
9. For a 4s electron in a hydrogen atom, how many nodal surfaces would there be?
10. How many p orbitals are there? How can we distinguish one p orbital from another.
11. What is a degenerate orbital? Give an example.
12. Define the Heisenberg’s Uncertainty principle; Pauli’s Exclusion Principle, Wave-particle duality theory
The end…. 