WHAT TO STUDY FOR THE FINAL TEST OF THE FIRST SEMESTER.

Below is a listing of concepts that you need to be familiar with for the final test. The test will consist of mostly multiple choice questions and some short answer questions. Though the book is not used much, this covers most of chapter 10(sections 10.5 through 10.11) and a lot of chapter 11 (sections 11.1, 11.2, 11.4 through11.7) in the book.

• Know the general concept of quantum mechanics. Orbitals are areas of highest probability (90%) of the location of an electron.



• Know that all matter moves in a wave pattern. This is especially true for electrons traveling in an orbital. The direction of the electron traveling is not known, but the path length is. The path length around the nucleus must be a multiple of an integer of the wavelength. An example of this is as follows: If an electron has a wavelength of 20 nanometers (nm), two possible path lengths could be 100 nanometers (nm) and 120 nm. Both path lengths are integers of the path length: (100 nm / 20 nm) = 5 (120 nm / 20 nm) = 6. Since the electron travels as a wave, the path lengths given would cause a crest to meet a crest and a trough to meet a trough. If the path length was 110 nm, the electron could not exist at this path length. The 110 nm path length in not an integer of wavelength ((110 nm / 20 nm) = 5.5). This would cause a crest to meet a trough and destructive interference would occur. The knowledge that the electrons travel as a wave gives credibility to the idea that electrons must exist in certain locations. This is supported by the emission spectrums of elements giving off very specific wavelengths of light when excited.
  


• Know the order of filling orbitals by using a periodic table. You will be supplied with a periodic table for the test.



• Be able to properly write electron configurations and/or orbital diagrams for elements in the ground (lowest energy) state. Also, be able to recognize if an electron is excited from an electron configuration or orbital diagram. This was covered in the electron configuration quiz given in class and can also be found in the notes package. An example of this: electron configuration of sodium (11 electrons) in the ground state is 1s^2 2s^2 2p^6 3s^1 - sodium in the excited state could be 1s^2 2s^2 2p^6 3s^0 4s^1



• Know the concept of penetration by an electron to lower principle energy levels and how this concept dictates the order on how the sub-levels and orbitals of different energy levels are filled. This concept is why the 4s sub-level is filled before the 3d sub-level and the 5s sub-level is filled before the 4d sub-level.



• Know the shape of an "s" orbital (sphere) and a "p" orbital (dumb bell)



• Know the sub-levels that are in each principle energy level. 1st - s only; 2nd - s and p; 3rd - s, p, and d; 4th (and on) s, p, d, and f.



• Know how many orbitals are in each sub-level. You should be able to calculate how many electrons could be held in a particular principle energy. For example: How many electrons could be in the 4th principle energy level? The 4th principle energy level is the first time all of the sub-levels are present. Therefore; 1 orbital for the s, 3 orbitals for the p, 5 orbitals for the d, and 7 orbitals for the f will be a total of 16 orbitals. Each orbital can hold 2 electrons for a total of 32 electrons in the 4th principle energy level.



• Know how to write orbital diagrams (show orbitals with arrows representing the electrons), complete electron configurations, and noble gas electron configurations.



• Know the abnormalities for the electron configurations of chromium and copper.



• Know para-magnetism and di-magnetism based upon unpaired and paired electrons in an orbital. An orbital diagram with many unpaired electrons in orbitals would exhibit magnetic character while an orbital diagram with all paired electrons in orbitals would show little to no magnetic character.



• Know that valence electrons are the electrons in the outer-most s and p sub-levels. Any electron that is not a valence electron in an atom is known as a core electron.



• Know the concept of shielding by the core electrons to minimize the effect of the protons in the nucleus on the valence electrons.



• Know all of the general periodic trends of atomic radius (size), 1st ionization energy, electron affinity, and electronegativity. Be able to define what all periodic trends are. Be able to explain in detail the factors that attribute to a periodic trend. Example #1 - Size of atoms across a period (horizontal row on the periodic table) decrease from left to right. The reason is the number of protons in the nucleus increase as you move to the right across the periodic table, but the number of core electrons remains the same. The increased positive charge of the protons and the shielding staying the same allows the protons to attract the valence electrons in closer to the nucleus. Example #2 - 1st ionization energy of atoms decreases going down a group (vertical columns on the periodic table). The reason is the atoms get larger as successive energy levels are added. The greater distance between the nucleus and valence electrons and increase in the amount of shielding due to more core electrons causes the protons in the nucleus to have less effect on the valence electrons. Because of the diminished effect, less energy is needed to remove an electron from the atom.
  


• Know why successive ionization energies get larger and larger. Know that a substantial increase in an ionization energy value from one electron to the next would be an indication of the removal of a core electron.




• Know that chemical bonds are made to lower the potential energy of the atoms involved in the bond. The lowering of the potential energy of the system makes the atoms in the molecule more stable. An example of this was done numerous times in class looking at the system of the energy input (ionization energy) to take the one valence electron away from a sodium atom compared to the energy release (electron affinity) of a fluorine atom gaining an electron to complete its outer valence level. The total process would be exothermic because the energy release from fluorine would be greater in magnitude than the energy consumption of sodium. Be able to apply this concept to other scenario in which the relative values of ionization energy and electron affinity are known.






• Know why the mole concept is important for comparison of ionization energy values and electron affinity values. Both values are reported in kilojoules per mole. The mole concept allows for equal comparisons because the number of atoms and electrons involved with the either process have been counted using the mole concept.






• Know that valence electrons are the electrons involved in bonding.






• Know how to define the two types of chemical bonds: Covalent bonds - sharing of electrons between atoms to achieve (in most cases) an octet (8 valence electrons) or a duet (2 valence electrons - hydrogen most of the time). Occurs between non-metal and non-metal elements (most of the time). A covalent bond is technically defined as an electronegativity difference between the two elements in the bond of less than or equal to 1.7. Ionic bonds - transfer of electrons between atoms to achieve an octet (or duet) for both atoms. Occurs between metal and non-metal elements (in most cases). An ionic bond is technically defined as an electronegativity difference between the two elements in the bond of greater than 1.7.






• An electronegativity table will be supplied to you for bond type determination.






• An electronegativity difference of less-than or equal to 1.7 does not mean electrons are always shared between the elements. It means the electrons are shared a majority of the time and some transferring of electrons does occur. An electronegativity difference of greater than 1.7 does not mean a complete transfer of electrons from one atom to another. It means electrons are transferred from the less electronegative atom to the more electronegative atom a majority of the time and some sharing of the valence electron(s) does occur.






• Know the term / concept of isoelectronic. This is covered in the second VODCast for the Chemical Bonding / Lewis Dot Diagram worksheet.






• Know how the relative sizes of ions compared to their original atom sizes. Cations (positive ions) will always be smaller than their original atoms and anions (negative ions) will always be larger than their original atoms. The reasoning for the trends is givenin the second VODCast for the Chemical Bonding / Lewis Dot Diagram worksheet.






• Know how to show in a Lewis dot diagram and define: a single covalent bond, a double covalent bond, and a triple covalent bond.






• Know how to draw a Lewis dot diagram and structural diagram of a chemical compound. This is covered in the first and second VODCasts for the Chemical Bonding / Lewis Dot Diagram worksheet.

The test will take roughly 50 minutes of the 90 minute testing period. The point value of the test will be that of some of our larger tests this semester - 50 to 60 points. With the remaining time in the testing period, we will do a lab. Please bring pennies for the lab. You will be converting the copper of the pennies to "gold". (Actually, you will be making brass that looks like gold.)