Download Specific Objectives:

CHE 100/102
Specific Objectives:
The student should be able to:
Chapter 1 - What is Chemistry?
1. Define and use the terms in bold print.
2. Explain the difference between science and technology; the difference between science and
chemistry.
3. Compare the relationships of matter as pure substances and mixtures. Use the illustrations on page 7
as a starting point.
4. Know and use the atomic symbols for common elements and to understand the significance of
chemical formulas.
5. Distinguish between chemical and physical properties and changes.
6. Describe and use the scientific method and its components.
Chapter 2 - The Numerical Side of Chemistry.
1. Distinguish between precision and accuracy.
2. Define and use significant figures appropriately to show uncertainty in measurements and
mathematical results.
3. Convert numbers between normal and scientific (exponential) notation.
4. Distinguish between mass and weight.
5. Know and use the common SI Units in Chemistry ( page 43) and the derived SI units (page 44)
6. Know and use the Greek prefixes on page 44.
7. Use dimensional analysis to make conversions between metric units and between metric and English
units.
8. Convert between Kelvin, Celsius (Centigrade), and Fahrenheit temperatures.
9. Define and make calculations involving density (specific gravity). Distinguish between intensive and
extensive properties.
10. Know and use the energy units, understand specific heat and be able to calculate energy changes.
Chapter 3 - The Evolution of Atomic Theory
1. State and apply the Law of Conservation of Matter.
2. State and apply the Law of Constant Composition using percent composition.
3. State and apply the Law of Multiple Proportions, using the ratios of small whole numbers.
4. Outline and discuss Dalton's postulates in his Atomic Theory.
5. Name and discuss the attributes of the major subatomic particles (protons, neutrons, and electrons) of
the atom.
6. Discuss the set-up for and the major conclusions from Ernest Rutherford's alpha particle experiment.
7. Define isotope and discuss the make-up of atoms in terms of atomic number, atomic mass numbers,
and atomic weights. Calculate the number of protons, electrons, and neutrons in atoms or ions and
the weighted average atomic weight of an element from the atomic masses and abundances of
isotopes of an element.
8. Discuss the importance of Mendeleev ordering the elements into the Periodic Table.
9. Define and discuss the important areas of the modern periodic table with respect to the properties of
the elements. [Groups, periods, representative elements, transition metals, metals, nonmetals,
metalloids, lanthanides, actinides, noble or rare gases, alkali metals, alkaline earth metals, halogens]
10. Define and discuss the ordering of the elements according to atomic size, ionization energy, and
electron affinity across a period and down a group. Define cation and anion.
11. Discuss electromagnetic radiation with respect to wavelength, frequency, and energy. Calculate
wavelength or frequency of electromagnetic radiation given the other. Calculate the energy of
electromagnetic radiation given its wavelength or frequency.
12. Explain the difference between a continuous and line spectrum and the difference between and
absorption and emission spectrum.
SPECIFIC OBJECTIVES continued.
The student should be able to:
CHAPTER 4 - The Modern Model of the Atom.
1. Explain the concept of quantized energy and its relationship to Classical and Quantum physics.
2. Explain and use Bohr's ideas of quantized energy levels (orbits) for electrons in atoms and relate this
idea to the principal quantum number - n (shells) for an electron in an atom.
3. Discuss the properties of Bohr's orbits including number of electrons and energy of electrons. Use
these ideas to write "gross" (number of electrons in shells) electron configurations.
4. Determine the number of valence electrons in atoms of the representative elements and use the idea to
explain the periodicity of the elements.
5. Distinguish between ground state and excited states for electrons in an atom and calculate energy
differences between energy levels for a hydrogen atom. Use this information to explain the
absorption and emission line spectra for hydrogen (and other elements)
6. Discuss the existence of subshells for electrons in atoms and generally explain the energy overlap of
subshells for different shell levels.
7. Write complete electron configurations and rare gas abbreviation electron configurations for atoms
and ions, that do not show anomalous behavior. Identify valence electrons.
8. Identify elements from the electron configurations of atoms or ions.
9. State the octet rule and use it to predict the chemical formulas of simple ionic compounds.
10. Redefine metals and nonmetals on the basis of valence electron behavior.
11. Discuss atomic size using the ideas developed in this chapter.
12. Explain the basic idea of orbitals, uncertainty, and probability that distinguish the quantum
mechanical model of the atom from the Bohr model.
Additional Material and Objectives for Chapter 4 - The Modern Model of the Atom.
quantum mechanical picture of the atom emphasizes the wave properties of the very small particle, the
electron. λ = h / mv, where λ is the wavelength exhibited by the particle, h is Planck's constant, m is the
mass of the particle, and v is the velocity of the particle. For large particles, such a baseball, the
wavelength is so small that it is completely out of the range of possible observation. For an electron, the
wavelength is approximately of nm length (similar to the wavelength of x-rays).
A standing wave is a stationary wave, as on the string of a musical instrument. In a simple picture of the
quantum mechanical model, the electron in the atom can "occupy" energy levels that correspond to an
integral number of wavelengths of the electron but not levels that correspond to fractional numbers of
wavelengths (because of destructive interference).
Erwin Schroedinger solved a wave equation that incorporated both the wavelike and particlelike behavior
of the electron. In solving the equation, he obtained three quantum numbers that described the radial
and angular momentum properties of acceptable energy levels for the electron in an atom. These
quantum numbers are
n = principal quantum number. As the Bohr quantum number it roughly describes the distance of the
electron from the nucleus of an atom. It may have the integral values of 1, 2, 3, 4, … infinity.
l = azimuthal or angular momentum quantum number. This quantum number describes the shape of the
orbital. It may have the integral values of 0, 1, 2, 3, …. n-1. Thus l = 0 corresponds to s orbitals
(spherical), l = 1, to p orbitals ("dumb bells" = 2 lobes), l = 2 to d orbital (mostly 4 equivalent lobes), l = 3
to f orbitals (generally 8 lobes), etc.
ml = magnetic quantum number. This quantum number describes the orientation of the orbitals in space
and in doing so, gives the number of each type of orbital. It may have integral values from -l ..0…+l.
Thus for l = 0, ml must be 0 and there is only one s orbital for a given energy shell. For ml = 1, the
possible values are -1, 0, +1; thus there are three p orbitals. For ml = 2, the possible values are -2, -1, 0,
+1, +2 or 5 d orbitals. For ml = 3, the possible values are -3, -2, -1, 0, +1, +2, +3 or 7 f orbitals.
A fourth quantum number needed to describe the electrons in the atom is
ms = spin quantum number. This quantum number describes whether the spin angular momentum of an
electron is aligned with or against the orbital angular momentum of the electron. It is often viewed as a
clockwise or counterclockwise spin motion. The possible values for this quantum number are +/- 1/2.
Three rules or principles that govern the electrons in an atom are
1. The Pauli Exclusion Principle, which says that no two electrons in an atom may have the same four
quantum numbers. This means that the maximum number of electrons in any single orbital is 2.
2. The Aufbau Principle, which says that in the ground state, electrons occupy the lowest energy level
available. This corresponds to the Bohr model as well.
3. Hund's Rule, which states that for degenerate (same energy) orbitals, electrons will fill first with the
same spin and will "pair" only when all of the degenerate orbitals already have one electron. ("Bunk
beds" and detailed orbital diagrams of the atoms.)
Additional Specific Objectives:
1. Use quantum numbers to determine the number of each type of orbital that are found in an atom and
the number of electrons that can occupy a set of orbitals.
2. Explain the kind of and number of each type of orbital at the various principal energy levels for an
electron in an atom.
3. Draw "orbital diagrams" for detailed electron configurations of the elements.
4. Draw the general shapes of the various kinds of orbitals in an atom.
Chapter 4 - One last additional objective.
Discuss why ionization energies decrease down a group and generally increase across a period. This
includes explaining why Group IIA is generally higher than Group IIIA and why Group VA is generally
higher than Group VIA. Also to discuss how starting to fill a new, higher energy subshell or to pair
electrons in a subshell affects the electron affinity of atoms.
Chapter 5 - Bonding Between Atoms
1. Discuss the nature and properties of molecules and the covalent bonds that hold them together.
2. Use valence electrons and the octet (and duet) rules to draw electron dot structures for simple
molecules.
3. Define and recognize lone pairs of electrons.
4. Define and use multiple bonds to make acceptable electron dot structures.
5. Define and use the concept of resonance when appropriate for electron dot structures.
6. Distinguish ionic and covalent bonds on the basis of whether electrons are transferred or shared.
7. Define the concept of electronegativity and use it to characterize molecules as having ionic or
covalent bonding and for covalent bonds, to define them as polar or non-polar.
8. Distinguish "simple" compounds as binary or ternary, ionic or covalent and name them.
Chapter 6 - The Structure of Molecules
1. Describe the basis for and use the VSEPR (Valence Shell Electron Pair Repulsion Model) with Lewis
dot structures to determine the geometry of the electron pairs and the molecules (ions) for up to and
including 4 electron pairs.
2. Describe the bond angles in the structures detailed in objective number 1 using the idea that lone pair
electrons normally require more "room" than bonded pair electrons.
3. Describe the dipole moment of molecules (ions) using the polarity of the bonds and the geometry of
the molecule (ion) - for which the dipole vectors may cancel. Categorize molecules/ions as polar or
nonpolar.
4. Describe how the polarity of molecules affect their intermolecular attractions for one another.
5. Describe the relative strength of covalent bonds, dipole-dipole attractions, and nonpolar-nonpolar
attractions.