AP CHEMISTRY – SUMMER ASSIGNMENT

Ms. Glassey eglassey@teaneckschools.org

Goal – To review/expand on material covered in previous chemistry class. To
get the most out of it you should read and study the information provided and work out the problems. You want to feel comfortable with the material and be able to evaluate where you need additional help. First set of problems, have the answers so you will be able to check your work.

Due – First day of class. After two weeks – summer assignment will not be accepted

Assignment:

Section 1 – Important material to memorize and use as reference -also additional links provided to help you review

Section 2 – National Math and Science Initiative Units 1-3

-Read and do the practice problems within. Check your work, answers are provided.

Section 3 – AP Chemistry Workbook Units 1, 2, 7

-Additional text given to help you review and then there is a short Problem Set following each unit.

Section 4 – Naming worksheet/Short Multiple Choice

-Short naming worksheet for additional review along with some practice multiple choice questions

Unit 1 – Introduction: Matter and Measurement

Learning Objectives
1.2 The student is able to select and apply mathematical routines to mass data to identify or infer the composition of pure substances and/or mixtures.
2.10 The student can design and/or interpret the results of a separation experiment (filtration, paper chromatography, column chromatography, or distillation) in terms of the relative strength of interactions among and between the components.
2.25 The student is able to compare the properties of metal alloys with their constituent elements to determine if an alloy has formed, identify the type of alloy formed, and explain the differences in properties using particulate level reasoning.

Chemistry is the study of matter and the changes that matter undergoes. Matter is made up of almost infinitesimally small building blocks called atoms. Atoms can combine together to form molecules. Molecules of a few familiar substances are represented here. In later chapters you will learn more about how atoms combine to form molecules, and how molecules have the shapes and properties that they do.

Essential knowledge: Chemical analysis provides a method for determining the relative number of atoms in a substance, which can be used to identify the substance or determine its purity.
a. Because compounds are composed of atoms with known masses, there is a correspondence between the mass percent of the elements in a compound and the relative number of atoms of each element.
b. An empirical formula is the lowest whole number ratio of atoms in a compound. Two molecules of the same elements with identical mass percent of their constituent atoms will have identical empirical formulas.
c. Because pure compounds have a specific mass percent of each element, experimental measurements of mass percents can be used to verify the purity of compounds.

The scientific community uses SI units for measurement of such properties as mass, length, and temperature. There are seven SI base units from which all other necessary units are derived.

Although the meter is the base SI unit used for length, it may not be convenient to report the length of an extremely small object or an extremely large object in units of meters. Decimal prefixes allow us to choose a unit that is appropriate to the quantity being measured. Thus, a very small object might best be measured in millimeters (1 millimeter = 0.001 meters), while a large distance might best be measured in kilometers (1 kilometer = 1000 meters).

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In order to convey the appropriate uncertainty in a reported number, we must report it to the correct number of significant figures.

Guidelines

1. Nonzero digits are always significant–457 cm (3 significant figures); 2.5 g (2 significant figures).

2. Zeros between nonzero digits are always significant–1005 kg (4 significant figures); 1.03 cm (3 significant figures).

3. Zeros at the beginning of a number are never significant; they merely indicate the position of the decimal point–0.02 g (one significant figure); 0.0026 cm (2 significant figures).

4. Zeros that fall at the end of a number or after the decimal point are always significant–0.0200 g (3 significant figures); 3.0 cm (2 significant figures).

5. When a number ends in zeros but contains no decimal point, the zeros may or may not be significant–130 cm (2 or 3 significant figures); 10,300 g (3, 4, or 5 significant figures).

To avoid ambiguity with regard to the number of significant figures in a number with tailing zeros but no decimal point, such as 700, we use scientific (or exponential) notation to express the number. If we are reporting the number 700 to three significant figures, we can leave it written as it is, or we can express it as 7.00 102. However, if there really should be only two significant figures, we can express this number as 7.0 102. Likewise, if there should be only one significant figure, we can write 7 102.

When measured numbers are used in a calculation, the final answer cannot have any greater certainty than the measured numbers that went into the calculation. In other words, the precision of the result is limited by the precision of the measurements used to obtain that result. In order to report results of calculations so as to imply a realistic degree of uncertainty, we must follow the following rules.

1. When multiplying or dividing measured numbers, the answer must have the same number of significant figures as the measured number with the fewest significant figures.

2. When adding or subtracting, the answer can have only as many places to the right of the decimal point as the measured number with the smallest number of places to the right of the decimal point.

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Essential knowledge: Solutions are homogenous mixtures in which the physical properties are dependent on the concentration of the solute and the strengths of all interactions among the particles of the solutes and solvent.
a. In a solution (homogeneous mixture), the macroscopic properties do not vary throughout the sample. This is in contrast to a heterogeneous mixture in which the macroscopic properties depend upon the location in the mixture. The distinction between heterogeneous and homogeneous depends on the length scale of interest. As an example, colloids may be heterogeneous on the scale of micrometers, but homogeneous on the scale of centimeters.
b. Solutions come in the form of solids, liquids, and gases.
c. For liquid solutions, the solute may be a gas, a liquid, or a solid.
d. Based on the reflections of their structure on the microscopic scale, liquid solutions exhibit several general properties:
1. The components cannot be separated by using filter paper.
2. There are no components large enough to scatter visible light.
3. The components can be separated using processes that are a result of the intermolecular interactions between and among the components.
e. Chromatography (paper and column) separates chemical species by taking advantage of the differential strength of intermolecular interactions between and among the components.
f. Distillation is used to separate chemical species by taking advantage of the differential strength of intermolecular interactions between and among the components and the effects these interactions have on the vapor pressures of the components in the mixture.

Alloys

Metallic solids are often pure substances, but may also be mixtures called alloys.
1. Some properties of alloys can be understood in terms of the size of the component atoms: — Interstitial alloys form between atoms of different radius, where the smaller atoms fill the
interstitial spaces between the larger atoms. (Steel is an example in which carbon occupies the interstices in iron.) The interstitial atoms make the lattice more rigid, decreasing malleability and ductility.
— Substitutional alloys form between atoms of comparable radius, where one atom substitutes for the other in the lattice. (Brass is an example in which some copper atoms are substituted with a different element, usually zinc.) The density typically lies between those of the component metals, and the alloy remains malleable and ductile.
2. Alloys typically retain a sea of mobile electrons and so remain conducting.
3. In some cases, alloy formation alters the chemistry of the surface. An example is formation of a chemically inert oxide layer in stainless steel.

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Introduction: Matter and Measurement- Problem Set

1. A student carries out a density determination lab to find the density of a small spherical metal pellet. He fills a 100-mL graduated cylinder to 50.0 mL and then drops the pellet in. The water level rises to 51.3 mL. If the pellet had a mass or 8.74 grams, what was the pellet’s density to the correct number of significant figures?

2. An aspirin tablet that has a mass of 3.00 grams only contains 325 mg of acetylsalicylic acid. What is the percent by mass of the acetylsalicylic acid in the tablet to the correct number of significant figures?

3. A necklace has a mass of 9.85 grams and a volume of 0.675 cm3. The necklace contains both gold and silver. Gold has a density of 19.3 g/cm3 and silver has a density of
10.5 g/cm3. Calculate the percent by mass of the gold in the necklace assuming that the total volume of the necklace is the sum of the volumes of its gold and the silver components.

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Unit 2 – Atoms, Molecules and Ions

Learning Objectives
1.14 The student is able to use data from mass spectrometry to identify the elements and the masses of individual atoms of a specific element.
2.19 The student can create visual representations of ionic substances that connect the microscopic structure to macroscopic properties, and/or use representations to connect the microscopic structure to macroscopic properties (e.g., boiling point, solubility, hardness, brittleness, low volatility, lack of malleability, ductility, or conductivity).
2.23 The student can create a representation of an ionic solid that shows essential characteristics of the structure and interactions present in the substance.
2.24 The student is able to explain a representation that connects properties of an ionic solid to its structural attributes and to the interactions present at the atomic level.
2.31 The student can create a representation of a molecular solid that shows essential characteristics of the structure and interactions present in the substance.
2.32 The student is able to explain a representation that connects properties of a molecular solid to its structural attributes and to the interactions present at the atomic level.

Essential knowledge: An early model of the atom stated that all atoms of an element are identical. Mass spectrometry data demonstrate evidence that contradicts this early model.
a. Data from mass spectrometry demonstrate evidence that an early model of the atom (Dalton’s model) is incorrect; these data then require a modification of that model.
b. Data from mass spectrometry also demonstrate direct evidence of different isotopes from the same element.
c. The average atomic mass can be estimated from mass spectra.

Many compounds consist of ions rather than molecules. Such compounds are said to be ionic. An ion is an electrically charged "package" consisting of one (monatomic ion) or more (polyatomic ion) atoms. An ion with a positive charge is called a cation (CAT-ion), while an ion with a negative charge is called an anion (AN-ion).

Ionic compounds do not exist as molecules and so do not have molecular formulas. Rather, ionic substances such as sodium chloride and magnesium chloride have only empirical formulas—NaCl and MgCl2, respectively.

The charges on many atomic ions can be predicted using the periodic table. In general, for a nonmetal to form an ion, it will gain as many electrons as it needs in order to have the same number of electrons as a noble gas. Metals will lose electrons to become cations, while nonmetals will gain electrons to become anions.

Essential knowledge: Ionic bonding results from the net attraction between oppositely charged ions, closely packed together in a crystal lattice.
a. The cations and anions in an ionic crystal are arranged in a systematic, periodic 3-D array that maximizes the attractive forces among cations and anions while minimizing the repulsive forces.
b. Coulomb’s law describes the force of attraction between the cations and anions in an ionic crystal.
1. Because the force is proportional to the charge on each ion, larger charges lead to stronger interactions. 2. Because the force is inversely proportional to the square of the distance between the centers of the ions
(nuclei), smaller ions lead to stronger interactions.

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Some transition metals can form more than one ion. Iron, for example, forms both Fe2+ and Fe3+. To name such an ion unambiguously, we use the name of the element, a Roman numeral in parentheses to denote the charge, and the word "ion." Fe2+ and Fe3+ would be iron(II) ion and iron(III) ion, respectively. An older method, though still widely used, is to apply the endings –ous for the smaller ionic charges and –ic for the larger ionic charges. For example:

Fe2+ is a ferrous ion

Cu+ is a cuprous ion

Fe3+ is a ferric ion

Cu2+ is a cupric ion

Essential knowledge: Ionic solids have high melting points, are brittle, and conduct electricity only when molten or in solution.
a. Many properties of ionic solids are related to their structure.
1. Ionic solids generally have low vapor pressure due to the strong Coulombic interactions of positive and negative ions arranged in a regular three- dimensional array.
2. Ionic solids tend to be brittle due to the repulsion of like charges caused when one layer slides across another layer.
3. Ionic solids do not conduct electricity. However, when ionic solids are melted, they do conduct electricity because the ions are free to move.
4. When ionic solids are dissolved in water, the separated ions are free to move; therefore, these solutions will conduct electricity. Dissolving a nonconducting solid in water, and observing the solution’s ability to conduct electricity, is one way to identify an ionic solid.
5. Ionic compounds tend not to dissolve in nonpolar solvents because the attractions among the ions are much stronger than the attractions among the separated ions and the nonpolar solvent molecules.
b. The attractive force between any two ions is governed by Coulomb’s law: The force is directly proportional to the charge of each ion and inversely proportional to the square of the distance between the centers of the ions. 1. For ions of a given charge, the smaller the ions, and thus the smaller the distance between ion centers, the
stronger the Coulombic force of attraction, and the higher the melting point.

2. Ions with higher charges lead to higher Coulombic forces, and therefore higher melting points.

Common Cations

Charge Formula	Name
	Charge Formula	Name
	
+1	 H+	Hydrogen ion Li+                      Lithium ion Na+                    Sodium ion K+                      Potassium ion
	+1	 Cs+	Cesium ion Ag+                   Silver ion
Cu+	Copper (I) or Cuprous ion NH +              Ammonium ion
H3O+	Hydronium ion
	
+2	 Mg2+	Magnesium ion Ca2+                   Calcium ion Sr2+                    Strontium ion Ba2+                   Barium ion Zn2+                   Zinc ion
Cd2+	Cadmium ion
Pb2+	Lead or Plumbous ion Sn2+	Tin(II) or Stannous ion
	+2	 Co2+	Cobalt or Cobaltous ion Cu2+                  Copper (II) or Cupric ion Fe2+	Iron (II) or Ferrous ion Hg 2+                Mercury (I) ion
Hg2+	Mercury (II) ion Ni2+                   Nickel (II) ion Mn2+                 Manganese (II)

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2
3
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-
2
-
-
-
-
2
2     3
3
4
2     7
3
4
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Common Anions

Charge Formula	Name
	Charge Formula	Name
	
1-	H-	Hydride ion F-	Fluoride ion Cl-                    Chloride ion Br-                    Bromide ion I-                         Iodide ion NO -               Nitrite ion NO -               Nitrate ion
MnO4-	Permanganate ion HCO3	Hydrogen carbonate ion
(or bicarbonate ion)
	1-	CN-	Cyanide ion OH-                  Hydroxide ion C2H3O2-        Acetate ion
ClO	Hypochlorite ion ClO -               Chlorite ion ClO3-               Chlorate ion ClO4                 Perchlorate ion
H2PO4	Dihydrogen phosphate ion SCN           Thiocyanate ion
HSO4-	Hydrogen sulfate ion N3                  Azide ion
	
2-	O2-	Oxide ion
O 2-	Peroxide ion S2-	Sulfide ion
S O 2-	Thiosulfate ion
HPO42-	Hydrogen phosphate ion
	2-	CO 2-	Carbonate ion CrO 2-	Chromate ion Cr O 2-            Dichromate ion SO 2-                 Sulfite ion
SO 2-                 Sulfate ion C2O42-              Oxalate ion
	
3-	 N3-	Nitride ion P3-                  Phosphide ion
	3-	PO 3-	Phosphate ion PO33-                 Phosphite ion

Note that polyatomic ions that contain oxygen have names that end in either –ite for the smaller number of oxygens or –ate for the larger number of oxygens. For Example:

NO2-	Nitrite ion

NO3-	Nitrate ion

SO32-	Sulfite ion

SO42-	Sulfate ion

Prefixes are used when the series of oxyanions of an element extends to four members. The prefix per- indicates one more oxygen, while the prefix hypo- indicates one less oxygen. For example:
ClO-	Hypochlorite ion	BrO-	Hypobromite ion
2	2
3	3
ClO -               Chlorite ion                  BrO -               Bromite ion ClO -               Chlorate ion                 BrO -               Bromate ion
ClO4-	Perchlorate ion	BrO4-	Perbromate ion

Names and formulas of acids follow naturally from the naming of ionic compounds. (Even though acids are not ionic!) A working definition of acid for this exercise will be: an anion with enough hydrogen ions attached to make it neutral. Thus, Cl– requires one hydrogen ion to become HCl, a familiar acid. SO42– requires two hydrogen ions to become H2SO4, another familiar acid.

To name an acid derived from an atomic anion, remove the ide ending from the anion, replace it with ic, and surround the new name with the prefix hydro and the word acid. Example: The anion in HCl is

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the chloride ion. Remove the ide ending, and replace it with ic. Surround the new word with hydro and acid, and you have hydrochloric acid.

For acids derived from polyatomic anions, simply replace the suffix of the anion name and add the word acid. The ending ate gets replaced with ic; the ending ite gets replaced with ous. The acid derived from nitrate ion becomes nitric acid. That derived from nitrite ion becomes nitrous acid. (Some of the anion roots change slightly for acid names. Example: The acids derived from sulfate and sulfite ions are sulfuric and sulfurous acids, respectively.)

Essential knowledge: Molecular solids with low molecular weight usually have low melting points and are not expected to conduct electricity as solids, in solution, or when molten.
a. Molecular solids consist of nonmetals, diatomic elements, or compounds formed from two or more nonmetals.
b. Molecular solids are composed of distinct, individual units of covalently bonded molecules attracted to each other through relatively weak intermolecular forces.
1. Molecular solids are not expected to conduct electricity because their electrons are tightly held within the covalent bonds of each constituent molecule.
2. Molecular solids generally have a low melting point because of the relatively weak intermolecular forces present between the molecules.
3. Molecular solids are sometimes composed of very large molecules, or polymers, with important commercial and biological applications.

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Atoms, Molecules and Ions- Problem Set

1. Write the chemical formulas for the following compounds:
a) copper (I) oxide	b) ferrous sulfide	c) sodium peroxide
d) ferric carbonate	e) iron (III) sulfate	f) lithium hypobromite g) calcium hydride	h) magnesium nitride i) sodium perchlorate

2. Name the following compounds:
a) Fe(OH)2	b) Ca(C2H3O2)2 d) LiClO4	e) NaOCl
g) ZnHPO4	h) Mg3N2

c) Mg(ClO2)2 f) K2Cr2O7
i) BaSO3

3. Write the chemical formulas for the following acids:
a) hypochlorous acid b) iodic acid	c) sulfurous acid d) hydrobromic acid e) perchloric acid	f) sulfuric acid
g) carbonic acid	h) phosphoric acid	i) phosphorous acid

4. Name the following acids:
a) HClO3	b) HOBr d) HC2H3O2	e) HNO2 g) HI	h) HBrO4

c) H2SO3 f) HNO3 i) HCl

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Unit 7 – Electronic Structure

Learning Objectives
1.5 The student is able to explain the distribution of electrons in an atom or ion based upon data. 1.6 The student is able to analyze data relating to electron energies for patterns and relationships.
1.7 The student is able to describe the electronic structure of the atom, using PES data, ionization energy data, and/or Coulomb’s Law to explain of how the energies of electrons within shells in atoms vary.
1.8 The student is able to explain the distribution of electrons with Coulomb’s Law to analyze measured energy. 1.12 The student is able to explain why a given set of data suggests, or does not suggest, the need to refine the
atomic model from a classical shell model with the quantum mechanical model.
1.13 Given information about a particular model of the atom, the student is able to determine if the model is consistent with specified evidence.
1.15 The student can justify the selection of a particular type of spectroscopy to measure properties associated with vibrational or electronic motions of molecules.

Essential knowledge: The atom is composed of negatively charged electrons, which can leave the atom, and a positively charged nucleus that is made of protons and neutrons. The attraction of the electrons to the nucleus is the basis of the structure of the atom. Coulomb’s law is qualitatively useful for understanding the structure of the atom.
a. Based on Coulomb’s law, the force between two charged particles is proportional to the magnitude of each of the two charges (q1 and q2), and inversely proportional to the square of the distance, r, between them. (Potential energy is proportional to q1q2 /r.) If the two charges are of opposite sign, the force between them is attractive; if they are of the same sign, the force is repulsive.
b. The first ionization energy is the minimum energy needed to remove the least tightly held electron from an atom or ion. In general, the ionization energy of any electron in an atom or ion is the minimum energy needed to remove that electron from the atom or ion.
c. The relative magnitude of the ionization energy can be estimated through qualitative application of Coulomb’s law. The farther an electron is from the nucleus, the lower its ionization energy. When comparing two species with the same arrangement of electrons, the higher the nuclear charge, the higher the ionization energy of an electron in a given subshell.
d. Photoelectron spectroscopy (PES) provides a useful means to engage students in the use of quantum mechanics to interpret spectroscopic data and extract information on atomic structure from such data. In particular, low-resolution PES of atoms provides direct evidence for the shell model. Light consists of photons, each of which has energy E = hν, where h is Planck’s constant and ν is the frequency of the light. In the photoelectric effect, incident light ejects electrons from a material. This requires the photon to have sufficient energy to eject the electron. Photoelectron spectroscopy determines the energy needed to eject electrons from the material. Measurement of these energies provides a method to deduce the shell structure of an atom. The intensity of the photoelectron signal at a given energy is a measure of the number of electrons in that energy level.
e. The electronic structure of atoms with multiple electrons can be inferred from evidence provided by PES. For instance, both electrons in He are identical, and they are both roughly the same distance from the nucleus as in H, while there are two shells of electrons in Li, and the outermost electron is further from the nucleus than in H.

When we say "light," we generally are referring to visible light—a type of electromagnetic radiation. Visible light constitutes a very small segment of the electromagnetic spectrum, which is composed of various types of electromagnetic radiation in order of increasing wavelength. Other types of electromagnetic radiation in this spectrum include microwave, infrared, and ultraviolet radiation.

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All types of electromagnetic radiation share certain characteristics. All have wavelike characteristics—similar to those of waves moving through water—including frequency and wavelength. All move through a vacuum at the "speed of light," which is 3.00 108 m/s. For any electromagnetic radiation, the wavelength, frequency, and speed of light are related by the equation

Where v (nu) is the frequency of the radiation in reciprocal seconds (s–1),	(lambda) is the wavelength in meters, and c is the speed of light in meters per second. Based on the above equation, we can see that as
wavelength increases, frequency decreases. The visible spectrum of light (ROYGBIV) ranges from red at about 700 nm to violet at about 400 nm. (Note a nanometer (nm) is 1 x10-9 meters).

Although matter was believed to consist of tiny particles called atoms prior to the start of the twentieth century, energy was considered to be "continuous." In 1900 Max Planck proposed that, like matter, energy is quantized; in other words, it consists of tiny, discrete amounts. He termed the smallest "particle" of energy that can be emitted or absorbed as electromagnetic radiation the quantum. Planck's theory was revolutionary, but it enabled scientists to reconcile some observations that could not be explained by classical physics. According to Planck's theory, the energy of radiation is directly proportional to its frequency.

E is the energy of a single quantum, and h is Planck's constant (6.63 10 34J s). Albert Einstein coined the term photon to refer specifically to a quantum of radiant energy (electromagnetic radiation carries energy through space, and is therefore also known as radiant energy).

For many years scientists have used flame tests to verify the presence of certain elements in compounds. Sodium, for example, gives off a characteristic yellow color when burned. Potassium burns with a violet flame.

The colored light given off in these tests can be separated with a prism into its component wavelengths, just as sunlight can be separated into a rainbow or visible spectrum. An important difference is that the rainbow is a continuous spectrum—a spectrum containing light of all wavelengths—while the colored light of a flame separates into a line spectrum, which consists of only a few specific wavelengths.

Essential knowledge: The interaction of electromagnetic waves or light with matter is a powerful means to probe the structure of atoms and molecules, and to measure their concentration.
a. The energy of a photon is related to the frequency of the electromagnetic wave through Planck’s equation (E = hν). When a photon is absorbed (or emitted) by a molecule, the energy of the molecule is increased (or decreased) by an amount equal to the energy of the photon.
b. Different types of molecular motion lead to absorption or emission of photons in different spectral regions. Infrared radiation is associated with transitions in molecular vibrations and so can be used to detect the presence of different types of bonds. Ultraviolet/visible radiation is associated with transitions in electronic energy levels and so can be used to probe electronic structure.

The line spectrum of hydrogen contains four visible lines.

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Although the existence of these lines had been known for many years, they were not explained until early in the twentieth century. Building on the work of Planck and Einstein, Niels Bohr applied quantum theory to explain the line spectrum of hydrogen in terms of the behavior of the electron in a hydrogen atom.

Bohr proposed that the electron in a hydrogen atom could circle the nucleus only in specific orbits designated by a quantum number n. The quantum number can have integer values, with n = 1 corresponding to the orbit closest to the nucleus. He showed the relationship between the value of n and the energy of an electron is

RH is the Rydberg constant (2.18 10–18J). The energy of an electron is, by convention, a negative number. When an electron resides in the orbit designated by n = 1, it is said to be in the ground state. This is the lowest possible energy level in which hydrogen's electron can exist. If hydrogen's electron is in a higher energy (less negative) orbit, with n greater than 1, the atom is said to be in an excited state. Bohr assumed that the electron could "jump" from one allowed energy state to another by absorbing or emitting photons of radiant energy of certain frequencies. He described the lines in the hydrogen spectrum as the energy given off when an electron in an excited state returns to the ground state. A flame or the application of high voltage imparts energy to the electron in a hydrogen atom and promotes it to an orbit of higher n value. When the excited state electron returns to the ground state, it releases the excess energy in the form of visible light. The frequency of this emitted radiant energy corresponds exactly to the energy difference between the two states.

Using his equation for the energy of an electron, Bohr calculated the energy change and the frequency associated with changing values of the quantum number n.

Using this relationship, Bohr was able to show that the visible line spectrum of hydrogen was due to the transitions of electrons in hydrogen atoms from n = 6 to n = 2, n = 5 to n = 2, n = 4 to n = 2, and n = 3 to n = 2.

For the transition from n = 5 to n = 2

The negative sign indicates that this frequency light is emitted when an electron goes from n = 5 to n = 2. (If an electron were to go from n = 2 to n = 5, light of the same frequency would be absorbed.) This corresponds to the 434.8-nm line in the hydrogen line spectrum. Thus, energy must be absorbed for an electron to move to a higher energy state, and radiant energy is emitted when an electron drops to a lower energy state.

The visible lines in the hydrogen line spectrum are known as the Balmer series, in honor of Johann Balmer who first developed an equation by which their frequencies could be calculated. Electron transitions ending in n = 1 and n = 3 are called the Lyman and the Paschen series, respectively.

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Essential knowledge: The electronic structure of the atom can be described using an electron configuration that reflects the concept of electrons in quantized energy levels or shells; the energetics of the electrons in the atom can be understood by consideration of Coulomb’s law.
a. Electron configurations provide a method for describing the distribution of electrons in an atom or ion. b. Each electron in an atom has a different ionization energy, which can be qualitatively explained through
Coulomb’s law.
c. In multielectron atoms and ions, the electrons can be thought of as being in ―shells‖ and ―subshells,‖ as indicated by the relatively close ionization energies associated with some groups of electrons. Inner electrons are called core electrons, and outer electrons are called valence electrons.

d. Core electrons are generally closer to the nucleus than valence electrons, and they are considered to ―shield‖ the valence electrons from the full electrostatic attraction of the nucleus. This phenomenon can be used in conjunction with Coulomb’s law to explain/rationalize/predict relative ionization energies. Differences in electron-electron repulsion are responsible for the differences in energy between electrons in different orbitals in the same shell.

Heisenberg stated in his ―Uncertainty Principle‖ that it is inherently impossible to know simultaneously both the exact momentum of the electron and its exact location in space. In Bohr's model of the atom a single quantum number described an orbit in which an electron could exist. In the quantum mechanical model of the atom, four quantum numbers are required to describe what is now called an orbital.

1. The principal quantum number, n, can have positive integer values. (1, 2, 3, . . . ). The principal quantum number determines the size of the orbital or the energy level.

2. The second or azimuthal quantum number, l, can have integer values from 0 to n–1. The value of l determines the shape of the orbital. Each value of l has a letter associated with it to designate orbital shape.

Value of l                0      1     2     3 Letter used              s      p     d     f

3. The magnetic quantum number, ml, can have integer values from –l through +l. The magnetic quantum number determines the orbital's orientation in space.

Each value of n defines an electron shell. Within a shell, each value of l defines a subshell. Within a subshell, each value of ml defines an individual orbital.

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1. The shell with principal quantum number n will consist of exactly n subshells. Each subshell corresponds to a different allowed value of l from 0 to n–1. Thus, the first shell (n = 1) consists of only one subshell, the 1s (l = 0); the second shell (n = 2) consist of two subshells, the 2s (l = 0) and 2p (l = 1); the third shell consists of three subshells, 3s, 3p, 3d, and so forth.

2. Each subshell consists of a specific number of orbitals. Each orbital corresponds to a different allowed value of ml. For a given value of l, there are 2l + 1 allowed values of ml, ranging from –l to l. Thus, each s (l = 0) subshell consists of one orbital; each p (l = 1) subshell consists of three orbitals; each d (l = 2) subshell consists of five orbitals, and so forth.

3. The total number of orbitals in a shell is n2, where n is the principal quantum number of the shell.

4. In the quantum mechanical model of the atom, each orbital can accommodate two electrons. Electrons possess an intrinsic property known as electron spin. In order to distinguish the two electrons in a single orbital, we must use a fourth quantum number, the electron spin quantum number. The electron spin quantum number, ms, can have values of + ½ and – ½.

Essential knowledge: The currently accepted model of the atom is based on the quantum mechanical model. a. Coulomb’s law is the basis for describing the energy of interaction between protons and electrons.
b. Electrons are not considered to follow specific orbits. Chemists refer to the region of space in which an electron is found as an orbital.
c. Electrons in atoms have an intrinsic property known as spin that can result in atoms having a magnetic moment. There can be at most two electrons in any orbital, and these electrons must have opposite spin.
d. The quantum mechanical (QM) model addresses known problems with the classical shell model and is also consistent with atomic electronic structures that correspond with the periodic table.
e. The QM model can be approximately solved using computers and serves as the basis for software that calculates the structure and reactivity of molecules.

In a many-electron atom, repulsions between electrons result in differences in orbital energies within an electron shell. In effect, electrons are shielded from the nucleus by other electrons. The attraction between an electron and the nucleus is diminished by the presence of other electrons. Instead of the magnitude of attraction being determined by the nuclear charge, in a many-electron system the magnitude of attraction is determined by the effective nuclear charge Zeff. This is the diminished nuclear charge that is felt by the electrons in orbitals with principal quantum number 2 and higher.

Z is the nuclear charge (atomic number), and S is the average number of electrons between the nucleus and the electron in question. The result is that orbital energies in many-electron atoms depend not only on the value of n, but also on the value of l. Within a shell, subshells of higher l value have higher energies.

The Pauli exclusion principle states that no two electrons in an atom can have the same set of four quantum numbers: n, l, ml, and ms. Thus, for two electrons to occupy the same orbital, one must have ms = + ½ and the other must have ms = – ½.

The arrangement of electrons within the orbitals of an atom is known as the electron configuration. The most stable arrangement is called the ground-state electron configuration. Hund's rule states that for orbitals with the same energy, the electrons will occupy single orbitals before pairing up.

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Essential knowledge: As is the case with all scientific models, any model of the atom is subject to refinement and change in response to new experimental results. In that sense, an atomic model is not regarded as an exact description of the atom, but rather a theoretical construct that fits a set of experimental data.
a. Scientists use experimental results to test scientific models. When experimental results are not consistent with the predictions of a scientific model, the model must be revised or replaced with a new model that is able to predict/explain the new experimental results. A robust scientific model is one that can be used to explain/ predict numerous results over a wide range of experimental circumstances.
b. The construction of a shell model of the atom through ionization energy information provides an opportunity to show how a model can be refined and changed as additional information is considered.

The Aufbau principle states that electrons must fill up from lowest energy orbitals that can receive them. That is why the 4s orbitals come before the 3d orbitals. 4s orbitals are slightly lower in energy then 3d orbitals, even the though the 4s is considered a higher energy level. There are also some exceptions to the Aufbau principle:

1. Cr, Mo and W, according to Aufbau, would have a d4s2. However, it has been found that ½ filled sublevels are more stable and therefore there electron configurations are d5s1.
2. Cu, Ag, and Au, according to Aufbau, would have a d9s2. However, it has been found that a completely filled d-sublevel is more stable and therefore there electron configurations are d10s1.

Lastly, whenever a atom or ion has at least one unpaired electron it is called paramagnetic. These types of substances are attracted to magnetic fields. If all of an atom or ions electrons are paired, it is called diamagnetic. These types of substances are repelled by magnetic fields.

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Unit 7 – Electronic Structure Problem Set

Given the following list of atomic and ionic species, find the appropriate match for questions 1 – 4. a) Fe2+	b) Cl	c) K+	d) Cs	e) Hg+

1. Has the electron configuration 1s22s22p63s23p63d6. 2. Has a noble gas electron configuration.
3. Has electrons in f orbitals 4. Is isoelectronic with gold

5. Which of the following combinations of particles represents an ion net charge of -1 and has a mass number of 82?
a) 46 neutrons, 35 protons, 36 electrons b) 46 neutrons, 36 protons, 35 electrons c) 46 neutrons, 36 protons, 36 electrons d) 47 neutrons, 35 protons, 35 electrons e) 47 neutrons, 35 protons, 36 electrons

6. For a neutral As atom in the ground state, how many electrons have quantum numbers n = 4, l = 1? a) 2                             b) 3                 c) 4                             d) 5                             e) 6

7. Which electron configuration is impossible? a) 1s22s22p6
b) 1s22s22p3
c) 1s22s22p63s1 d) 1s22s22p63s2 e) 1s22s22p62d2

8. Which set of quantum numbers (n,l,ml,ms) represents the outermost electron in an aluminum atom? a) 2, 1, 0, +½	b) 2, 1, -1, +½	c) 3, 0, 0, +½	d) 3, 1, -1, +½

9. Which species is paramagnetic in the gaseous state?
a) Cr+3	b) Zn+2	c) Sn+2	d) Al+3

10. How many unpaired electrons are found in the ground state of a sulfur atom?
a) 0	b) 2	c) 4	d) 6

11. The most stable electron configuration for Mn+2 is:
a) [Ar] 4s23d3	b) [Ar] 3d5	c) [Ar] 4s13d5	d) [Ar] 4s13d4

12. What is the wavelength of radiation emitted from a mercury arc sunlamp if the frequency of the radiation is about 1.2 x 1015 sec-1? (c = 3.0 x 1010 cm/sec)
a) 2.0 x 10-5 cm	b) 4.0 x 10-4 cm	c) 2.5 x 10-6 cm	d) 2.5 x 10-5 cm

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13. The 1st line in the hydrogen spectrum is dark purple with a wavelength of 410.18 nm. Calculate:

a) the frequency of this light

b) the change in energy that produces this color (h = 6.626 x 10-34 J.s/ particle) (c = 3.0 a 108 m/sec)

14. An electron in a hydrogen atom gets excited to the 3rd energy level and then falls to back down to the1st energy level.

a) How much energy was released from this transition?

b) What is the frequency of this transmission?

c) What is the wavelength of this transmission in nanometers?

d) Would you be able to see the color of this transmission?

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Unit 8 – Periodic Table and Properties

Learning Objectives
1.9 The student is able to predict and/or justify trends in atomic properties based on location on the periodic table and/or the shell model.
1.10 Students can justify with evidence the arrangement of the periodic table and can apply periodic properties to chemical reactivity.
1.11 The student can analyze data, based on periodicity and the properties of binary compounds, to identify patterns and generate hypotheses related to the molecular design of compounds for which no data is supplied.

Elements display periodicity in their properties when the elements are organized according to increasing atomic number. This periodicity can be explained by the regular variations that occur in the electronic structures of atoms. Periodicity is a useful principle for understanding properties and predicting trends in properties. Its modern-day uses range from examining the composition of materials to designing new materials. Although a simple shell model is not the currently accepted best model of atomic structure, it is an extremely useful model that can be used qualitatively to explain and/ or predict many atomic properties and trends in atomic properties. In particular, the arrangement of electrons into shells and subshells is reflected in the structure of the periodic table and in the periodicity of many atomic properties. Many of these trends in atomic properties are important for understanding the properties of molecules, and in being able to explain how the structure of the constituent molecules or atoms relates to the macroscopic properties of materials. Students should be aware that the shells reflect the quantization inherent in quantum mechanics and that the labels given to the atomic orbitals are examples of the quantum numbers used to label the resulting quantized states. Being aware of the quantum mechanical model as the currently accepted best model for the atom is important for scientific literacy.

Essential knowledge: Many properties of atoms exhibit periodic trends that are reflective of the periodicity of electronic structure.
a. The structure of the periodic table is a consequence of the pattern of electron configurations and the presence of shells (and subshells) of electrons in atoms.
b. Ignoring the few exceptions, the electron configuration for an atom can be deduced from the element’s position in the periodic table.
c. For many atomic properties, trends within the periodic table (and relative values for different atoms and ions) can be qualitatively understood and explained using Coulomb’s law, the shell model, and the concept of shielding/effective nuclear charge. These properties include:
1. First ionization energy	2. Atomic and ionic radii 3. Electronegativity	4. Typical ionic charges
d. Periodicity is a useful tool when designing new molecules or materials, since replacing an element of one group with another of the same group may lead to a new substance with similar properties. For instance, since SiO2 can be a ceramic, SnO2 may be as well.

The size of an atom is defined by its atomic radius. As a particular group of elements is examined, it is found that the atomic radius of each successive element in that group increases. The reason for this is due to the fact that as you go down a group, the number of energy levels increases, thus the size of the atom increases. However, as a particular period of elements is examined, it is found that the atomic radius of each successive element in that period generally decreases (though there are exceptions due to electron configuration irregularities). The reason for this is due to the fact that as you go across a period, the level of the atom’s electrons remains the same, but the number of protons in the nucleus of the atom increases; thereby increasing the ―attractive force‖ on the electrons. The end result is that the electrons get pulled in closer and the size of the atom gets smaller.

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