1、普通化学（全英文）教学大纲 Syllabus for General Chemistry (I) Class Webpage: http:/ 1 Introduction (Chapter 1 3) 1.1 Elements, compounds, and mixtures 1.2 Atoms, molecules and ions 1.3 SI units and converting units 1.4 Extensive and intensive properties 1.5 Amount of substance 1.6 Law of conservation of matter 1
2、.7 Mass relationships in chemical reactions 1.8 Balancing the chemical reaction 1.9 Limiting reactants 1.10 Empirical formula 2 Electronic Structures and the Periodic Table (Chapter 7 8) 2.1 Atomic structure, Isotope 2.2 The periodic table (a) Main groups, metals and non-metals (b) s-block, p-block,
3、 d-block elements 2.3 Electron cloud and Atomic orbitals 2.4 Electron configurations (a) 4 quantum numbers: n, l, m, ms (b) Pictures of the orbitals: s, px, py, pz (c) How to fill electrons in different orbitals (d) Be able to draw the electron configuration of the first 20 elements (e) Unpaired ele
4、ctrons, Spin, Paramagnetic and Diamagnetism 2.5 Comparing the atomic orbitals with Bohrs model of atoms (Classical but incorrect) 2.6 Electrons states (a) Each state is related to specific energy (b) e- will transit between states when absorbing/giving off energy, in a form of light (c) Ground state
5、, Excited states, the first excited state (d) Energy of light: relationship among wavelength, frequency, and photons energy (e) Absorption spectrum and Emission spectrum (f) Red shift and blue shift (g) Materials color 2.7 Atomic and ionic radii 2.8 Ionization energy: the first, second, ionization e
6、nergy 2.9 Electron affinity (Note here we use a different definition of the electron affinity from the textbook.) 2.10 Exothermic and Endothermic reactions 3 Chemical Bonds, Valence Bond, and Molecular Shape (Chapter 9) 3.1 Valence electrons 3.2 Ionic bond 3.3 Covalent bond 3.4 Lewis structures (a)
7、Octet rule (b) Single, double, and triple bonds (c) Bond order, bond length, and bond energy: Be able to calculate the heat of a reaction from the given bond energies (d) Lone pair electrons, Coordinate bond (e) Be able to write Lewis structures of common molecules (f) Exception of the octet rule: m
8、ore than 8 e-; fewer than 8 e-; odd number of e- (g) Resonance structures (h) Formal charges a way to select the most effective Lewis structure(s) (i) Failure of Lewis structure: octet rule exception; resonance structure; paramagnetism 3.5 Electronegativity 3.6 Nonpolar and polar covalent bonds 3.7
9、Valence Bond (VB) method (a) Definition of the valence bond method (b) bond and bond How bond and bond are made from s, px, py, pz orbitals? How bond and bond relate to single/double/triple bonds? (c) Hybrid orbitals (only applied to central atoms): sp, sp2, sp3, sp3d (d) Bond angle and the shape of
10、 molecules (e) Cis- and Trans- isomerism (sp2 hybrid) (f) VSEPR method: predict the electron pair shapes and molecular shapes How to calculate the hybrid type of the central atom(s) How to arrange lone pair electrons and bonding electrons (g) Advantages and disadvantages of the VB method 3.8 Molecul
11、es polarity: sum of all bonds dipoles 3.9 Brief introduction to organic molecules 4 Molecular Orbital (MO) Method (Chapter 10) 4.1 Introduction to the MO method (a) Bonding, Antibonding, and Nonbonding MOs (b) bond and bond Compare bond and bond in VB method and in MO method 4.2 Bond orders, unpaire
12、d electrons (a) Compare Bond order definitions in Lewis structure, VB, and MO method (b) Compare Lone pair electrons in Lewis structure, VB, and MO method (c) Predict the stability of molecules/ions (d) Predict paramagnetism/diamagnetism of a molecule 4.3 Homonuclear diatomic molecules (1st and 2nd
13、period elements) 4.4 Heteronuclear diatomic molecules 4.5 Conjugated bond and electron delocalization (a) Be able to judge whether a molecule has a conjugated bond (b) Draw the molecular shape, and judge the p-orbitals and p-electrons (c) Calculate the nm 4.6 HOMO and LUMO 4.7 Band theory of bonding
14、 in solids: (a) Valence band, conduction band, band gap (b) Metal, semiconductor and insulator 4.8 Inter-molecular interactions (a) van der Waals interaction: polar vs. nonpolar molecules (b) Hydrogen bond (c) Bond energy: chemical bond hydrogen bond van der Waals interaction (d) Relationship of Int
15、er-molecular interactions with Melting, boiling point, hardness (e) Hydrophobic and hydrophilic interaction 4.9 Introduction to proteins and nucleic acids (a) Amino acid Peptide Protein: which interactions are involved in each step? (b) Oligonucleotide Single strand DNA Double strand DNA: which inte
16、ractions? 5 Acids and Bases (Chapter 15, 16.1 16.4) 5.1 The Bronsted-Lowry definitions ( H+, OH- transfer) (a) Strong acid (base) and Weak acid (base) (b) Conjugated acid-base pair Note in each acid-base reactions, there are two conjugated acid-base pairs (c) Poly-protic acid (base), and amphoteric
17、substance 5.2 Lewis acid-base definition ( transfer of lone pair electrons to an empty orbital) 5.3 Ion product of water, Kw = H3O+ OH- = 10-14 ( 25 oC) (a) Kw is a constant in all aqueous solution (under a constant temperature) (b) H3O+ 10-7 M acidic; H3O+ = 10-7 M neutral; H3O+ 10-5 M? (b) weak ac
18、id only Consider the Ka1 only (for poly-protic acid) If Cacid/Ka 100 H3O+ = (Cacid Ka) If Cacid/Ka 100 need to solve the quadratic equation (c) one strong acid + one weak acid Let the strong acid dissociate completely first, and then setup the ICE equation of the weak acids dissociation. (d) one wea
19、k acid + its conjugated base (it is a buffer!) In a buffer: H3O+ = Ka (Cacid / Cbase) (e) one acid + one base Let the acid-base reaction go into completion first. (Need to find out the limiting reactants!) Then the remaining species will belong to one of the previous categories. 5.5 Calculation of O
20、H- in base solution(s) 5.6 Hydrolysis of salt (a) How to quickly judge whether a salt is acidic, neutral or basic (b) Calculate the pH value of a salt solution 5.7 Buffer solutions Note for a good buffer, its acid/base ratio should be 0.110. 5.8 Titration (a) Titration curve: use a strong base to ti
21、trate a strong acid (b) Titration curve: use a strong base to titrate a weak acid Outline for General Chemistry (II) 6A Solution (Chapter 4) 1 class 6.1 Solute, solvent, and solution 6.2 Solubility (S) and molarity (C) (a) Solubility: # grams of solute / 100 g solvent; Molarity: # moles of solute /
22、1 L solution (b) Be able to convert between solubility and molarity of a given solution 6.3 Electrolyte (a) Strong electrolyte: fully dissociate into cations and anions in solution or melt (b) Weak electrolyte: partially dissociate (with a dissociation factor between 01), and majority remains as mol
23、ecules 6.4 Acid, base, and salt Note that insoluble salt is still a strong electrolyte 6.5 Reaction/calculation between ions in solution (a) Water solubility of different acids, bases and salts (b) Find out the limiting reactants for calculation 6.6 Ionic equation 6B Precipitate (Chapter 16.5 16.8)
24、1 class 6.7 Solubility product: (a) Ksp = cationm anionn, (where m, n are coefficients of ions in the precipitates dissociation reaction) (b) Convert between solubility product (Ksp) and solubility (S) 6.8 Predict the solubility formation Compare the magnitude of ion product, cationm anionn, with th
25、e Ksp (a) cationm anionn Ksp Over saturated solution, Precipitate is formed 6.9 The common ion effect 6.10 Calculation of ion concentrations in solution When more than one precipitate can be formed, which one will precipitate first? When the 2nd precipitate starts to form, what is the remaining ion
26、concentration for the 1st ion? 7 Oxidation-Reduction Reaction (Chapter 11) 2 class 7.1 Oxidation numbers (Oxidation state) 7.2 Rules for assigning the oxidation number Apply the 5 steps in order 7.3 Oxidation numbers and the periodic table Highest possible oxidation number = # of its valence e-s; Lo
27、west possible oxidation number = # of e-s needed to gain for an octet 7.4 Oxidation-reduction reaction (a) Any oxidation number change Red-ox reaction (b) Oxidation half-reaction (for Reducing agent): Oxidation # , lose e-s, being oxidized (c) Reduction half-reaction (for Oxidizing agent): Oxidation
28、 # , gain e-s, being reduced 7.5 Balance the Red-ox reactions (Half-reaction method) Find out species that are oxidized or reduced (based on their oxidation #s); Balance each half reaction; If in a basic solution, add OH- to each side to neutralize H+; Double check the conservation of mass and charg
29、es 8 Changes in States (Chapter 5, 12, 13.5 13.7) 3 classes 8.1 State of Matters Gas, liquid, and solid 8.2 Gas and Ideal Gas 2 assumptions for the ideal gas molecules: (1) no molecule size; (2) no intermolecular interaction 8.3 Ideal gas law: PV = nRT Select the correct value for the gas constant R
30、, based on the units. 8.4 Avogadros law (a) same T (b) Standard temperature and pressure, 1 mole of any ideal gas takes 22.4 liter. 8.5 Mole fraction (Xi): Xi = ni / ntotal = Pi / Ptotal = Vi / Vtotal Note that all the ns here only include gas species 8.6 Partial pressure (Pi): pressure of a gas com
31、ponent if it solely occupies the whole volume Ptotal = PA + PB + + P i Pi = Ptotal Xi 8.7 Partial volume (Vi): Volume of a gas component under T and Ptotal Vtotal = VA + VB + + V i Pi Vtotal = Ptotal Vi 8.8 Phase, Phase change, and Phase equilibrium (a) Liquid evaporation and the vapor pressure Esti
32、mate the relative magnitude of different molecules vapor pressure, based on their molecular structures (b) Change of the vapor pressure with T: ln (P2/P1) = Hvap / R (1/T2 1/T1) (c) Boiling temperature (Tb): Tb, vapor pressure = 1 atm (d) Melting P = P0, solvent Xsolute For two volatile components:
33、Ptotal = P1,0 X1 + P2,0 X2 Calculate the ratios of two volatile components in both liquid and gas phases the one with higher vapor pressure will be enriched in the gas phase Distillation (b) Boiling point elevation: Increase of Tb: Tb = Tb Tb = Kb m; (Kb is a constant for the solvent, m is the molal
34、ity of solute) Using the boiling point elevation to calculate the solutes Molar mass. (c) Freezing point depresstion Decrease of Tf: Tf = Tf Tf = Kf m; (Kf is a constant for the solvent, m is the molality of solute) (d) Osmotic pressure = Csolute * R * T Using the osmotic pressure to calculate the s
35、olutes Molar mass. Summary and Exercise Class #3 1 class Midterm Exam 1 class 9 Chemical Thermodynamics and Thermochemistry (Chapter 6) 4 classes 9.1 System, surroundings, and universe System + Surroundings = Universe Open system, Closed system, Isolated system 9.2 State and State function (a) State
36、 functions: changes that only depends on the initial state and the final state, not on the path of changes (b) Internal energy and Internal energy change (U, U) (c) Heat (Q) and Work (W) they are NOT state functions! Q = n C (Tfinal Tinitial), where C (Cp or Cv) is the molar heat capacity 9.3 Thermo
37、dynamics First Law: U = Q W (a) Signs of U, Q, and W (b) Calculation of the volume work (i.e. expansion, compression) Under constant external pressure (Pex): W = Pex V Under vacuum (Pex = 0): W = 0 V = 0 9.4 Enthalpy and Enthalpy change (H, H) Relationship between H and U: H = U + ngasRT 9.5 Thermal
38、 Chemical Equation H and Q (heat) are extensive properties, and relate to moles of a reaction. Exothermic (q 0): H 0 9.6 Hess Law When a reaction can be written as combination of a few other reactions, those thermodynamic state functions, (e.g. U, H, S, G, K) can be expressed as corresponding combin
39、ation. How to find out the combination coefficients of different sub-reactions? 9.7 Standard States (XoT): For a pure solid or liquid they are standard states, concentration = 1 For a gas Standard states: Ppartial = 1 atm For an ion in solution Standard states: ion = 1 M If non-standard states, we n
40、eed to convert, e.g. GT =GoT + RTln Q 9.8 Calculation of H (a) Using the enthalpy of formation (Hfo): Ho = (Hfo products) (Hfo reactants) (b) Using the enthalpy of combustion (Hco): Ho = (Hco reactants) (Hco products) (c) Using the bond energy (B.E.): Ho = (B.E. reactants) (B.E. products) 9.9 Entrop
41、y and Entropy Change (S, S) (a) Factors that affect entropy: State of matter, Temperature, Pressure, Molecular weight, molecular symmetry (b) Qualitatively, if gas molecules # S (c) So = (So products) (So reactants) 9.10 Thermodynamics Second Law: Sisolated = Ssystem + Ssurroundings 0 Sisolated 0 sp
42、ontaneous; Sisolated = 0 equilibrium; Sisolated 0 spontaneous; G = 0 equilibrium; G 0 spontaneous all T H 0, S 0, S 0 spontaneous high T (T H/S) (e) Calculation of G298Ko: G298Ko = (Gf, 298Ko products) (Gf, 298Ko reactants) (f) Calculation of GTo: GoT = Ho298K T So298K , (assume H, S are not tempera
43、ture dependent) (g) Calculation of G from Go: GT =GoT + RTln Q , (Q = products / reactants) 9.12 Thermodynamics Third Law: for all materials, S0K = 0 10 Chemical Equilibrium (Chapter 14, 17) 3 classes 10.1 Introduction to chemical equilibrium 10.2 Equilibrium constants (K) K only depends on temperat
44、ure! 10.3 Calculations involving equilibrium Set up the chemical equation with ICE (initial, change, and equilibrium) concentrations, pressures, or moles, then solve the equation of K. 10.4 Relationship between K and GoT: GoT = RTln K 10.5 Predict the reaction direction by comparing the magnitude of
45、 Q and K: (Q K) 10.6 Hess law for G and K Multiple ionic equilibria in solution (using Hess Law to find out combination coefficients) Pay attention to K: e.g. if (reaction 1) = 2 * (reaction 2) 3 * (reaction 3) G1 = 2 *G2 3 *G3, K1 = K22 / K33 10.7 Shift of equilibrium (a) Qualitatively, use Le Chte
46、liers principle to predict the shift direction (b) Quantitatively, set up ICE equation to solve the equation of K. (c) Change of K with T: ln (K2/K1) = Ho / R (1/T2 1/T1) 10.8 Learn how to plot Concentration Time graphs 11 Chemical Kinetics (Chapter 18) 3 classes 11.1 Rates of Reactions, Rate and co
47、ncentration v = c/t = dC/dt for aA + bB dD: v = dA/adt = dB/bdt = dD/ddt 11.2 Rate laws How to find rate laws based on reactants concentrations and reaction rates Differential rate laws: v = k Am Bn Integrated rate laws: C t 11.3 Zero-order, First-order, Second-order reaction Be able to write rate laws for Ze
