CHM1045 Review Topics Test 3
Chapter 11: 11-4 to 11-6 only
1. Solution stoichiometry using redox. Same as other solution stoichiometry. The reaction must be balanced according to the redox balancing method.
2. Be able to identify which substance is being oxidized, reduced. Also the oxidizing agent is the substance being reduced and the reducing agent is the substance being oxidized. You can identify them either by balancing the half reactions or by looking at the change in oxidation numbers of the elements involved in each half reaction
Chapter 12: (All except for enrichment on p. 431, and 12-15 only what we covered in class)
1. Be familiar with the relationships of P, V, T and n from:
a. Boyle's law, b. Charles' Law, c. Avogadro's law, d. Gay-Lussac’s Law
Be able to do calculations when any of the variables are changing and some of the variables are remaining the same (Be able to determine what happens to volume if pressure increases and the temperature and moles of gas remain constant, etc.). In other words, know whether the variables are directly or inversely related.
2. Combined Gas Law: Three variables are changing, only n remains constant.
3. Ideal Gas Equation including calculation of molar mass.
4. Relationship of density of gases and molar mass at STP (22.4 L/ mol ) and also not at STP.
5. Know what STP is.
6. Dalton's Law of Partial Pressure for when you have two different gases together the total pressure is the sum of the individual partial pressures of the two gases. This relates to the sum of the moles of the two gases in the ideal gas law if you want to calculate total pressure. Also this applies to when you have water vapor combined with a gas. (You have to subtract the partial pressure of the water vapor from the total pressure (usually barometric or atmospheric pressure) to get the partial pressure of the gas).
7. Graham's Law of effusion
8. Stoichiometry of gases:
a. Know that the coefficients in a balanced equation can represent gas volumes as long as the pressure and temperature remain constant.
b. Be able to use the ideal gas law to convert from one of the gas variables such as liters to another one such as moles when given temperature and pressure. If at STP you can convert between liters and moles of a gas (22.4L=1mol).
c. Be able to calculate molarity in stoichiometry problems that also involve gases.
d. Be able to determine %purity of a sample which decomposes to produce a gas or to use % purity for other calculations involving gases.
e. Be able to determine molecular formula given enough information to determine empirical formula and also enough information to determine molar mass using the ideal gas law.f. Be familiar with the kinetic molecular theory of gases and what is an ideal gas: points in space, negligible volumes, move at fast speeds in straight lines, undergo elastic collisions, very little attractive forces between particles, low pressures and high temperatures usually required to behave like ideal gases, if temperature increases the speed and kinetic energy increases.
g. Know that gases behave as ideal gases as long as the temperatures are high and the pressures are low. If there are a lot of attractive forces due to polar molecules or if the volumes occupied by the particles are relatively large (mostly due to large molar masses but there could be other reasons dealing with the geometry of the molecule), then Van der Waals equation applies where a coefficient of a represents the attractive forces and affects the P term and a coefficient b represents the molecular volume and affects the V term. If a and b are 0 or very small then the result is the normal ideal gas equation..
Chapter 4 (only 4-11, 4-13, 4-15 to 4-20):
1. Formula of frequency x wavelength = c
3. Formula of E= h x frequency. h will be given. It follows that E = hc/l
4. Rydberg's equation R will be given (Rydberg's constant, not the ideal gas constant)
5. The maximum number of electrons in an energy level: 2n2 (two times n squared)
6. The four quantum numbers and how to determine quantum numbers from an electron configuration and vice versa. The first quantum number, n, is the coefficient of the configuration, corresponds to the period and is the principal energy lever. The second one, l, is the sublevel (l=0 for s, l=1 for p, l=2 for d and l=3 for f). Remember that to determine the ml (third quantum number) and the ms (fourth quantum number) you have to draw the orbitals and place the arrows first. The ml starts at –l and goes through 0 to +l. The ms is +1/2 for arrow pointing up (first one placed) and -1/2 for arrow pointing down (second one placed in an orbital). Remember that an orbital can only hold 2 electrons maximum. The ml corresponds to the individual orbitals. Also remember that when placing the electrons in the orbitals you spread them out before you pair them up.
7. Be able to determine if a set of quantum numbers is not allowed.
8. Shapes of the orbitals and the corresponding name and quantum numbers (the l quantum number) associated with each shape.
9. Electronic configurations for all elements and ions including the exceptions and the configurations for the ions. Remember that Cr, Mo, Cu, Ag and Au are exceptions which end in s1d5 or s1d10. For representative metal ions (Sn, Pb, Bi, Ga, In, Tl) other than IA, IIA or Al, you remove from p first then s. For transition metal ions, including the exceptions, you remove from the s first then d.
10. Noble gas abbreviations.
11. Paramagnetic and diamagnetic and number of unpaired electrons.
12. Isoelectronic species.
13. Remember that when an electron absorbs energy it can go to a higher energy level and release it when it goes to a lower energy level. This energy corresponds to a quantum of energy (a chunk of energy) and is a photon of energy.
14. Remember that energy levels, sublevels and orbitals are just regions in space with 90% probability of finding an electron of that energy content there.
15. Heisenberg Uncertainty Principle indicates it is impossible to know the energy and the momentum of an electron at the same time.