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Quick Review of Skills by Week

Week 1

• Give the seven basic SI Units

• Express a derivative unit such as J in basic units

• Use the one-line method for unit conversions and dimensional analysis

• Describe the constituents of an atom, and the atomic theory

• Describe and give an example for the concepts associated with these terms:
• atom
• electron
• atomic nucleus
• proton and atomic number Z
• neutron and number of neutron in a neucleus, N
• mass number
• isotope and nuclide
• abundance of isotopes
• atomic weight

Week 2

• What are the molecular or formula weight of H2, H2O, H2O2, O3, NaCl, and CO32-?

• What is a mole of substance?
What is the meaning of Avogadro's number?
What are the molar masses of 1.0 mole of 12C, H2, CO32-, and C12H22O11 (sugar)?
How many mole is 1.0 kg of each of the above material?

• What is the mass of a hydrogen atom or any nuclide?
What is the mass of 1.0 atomic mass unit?

• What are the weight percentages and molar percentages of carbon in suggar, (C12H22O11) calcium carbonate, (CaCO3) and urea ((NH2)2CO)?

• A compound containing 92.3 weight percent of carbon and 7.7 weight percent of H. What is the empirical formula?

• In the study of any material, our task is to find out what it is and what its properties are. Regarding a pure substance, the first important question is its elemental composition, or its empirical formula. Therefore, you shall be able to perform elemental analysis:
• Explain the principle of elemental analysis.
• Describe the cumbustion of a sample
• Calculate number of moles and mass of C from mass of CO2
• Calculate number of moles and mass of H from mass of H2O
• Determine the mass percentage of elements in a sample
• Determine the empirical formula

• When the empirical formula is given, you need its molecular weight. Some of these methods will be discussed later when the subjects of solutions and gases are discussed. But given a molecular weight, you should be able to determine the molecular formula.

Based on the properties, a chemist will Propose a structure formula In General Chemistry, some structural formula are given, and you are expected to know them. These formulas provide important information about a substance, and they are vital in the study of materials.

• Regarding stoichiometry of chemical reactions, you are expected to solve problems similar to this example:

Example 1

10 g each of Al and Fe2O3 are used for the thermite reaction, 2 Al + Fe2O3 = 2 Fe + Al2O3. How much Fe should be produced when the limiting reactant is exhausted?

Solution

Molar mass of Fe2O3 = 2 * 55.8 + 3 * 16 = 159.6 g / mol.

Since 10 g Fe2O3 requires 3.4 g Al, we assume Fe2O3 completely reduced.

```            1 mol Fe2O3    2 mol Al  27.0 g Al
10 g Fe2O3 ------------- ----------- --------- = 3.4 g Al
159.6 g Fe2O3 1 mol Fe2O3 1 mol Al

Amount Fe produce is calculated this way:

1 mol Fe2O3   2 mol Fe   55.8 g Fe
10 g Fe2O3 ------------- ----------- --------- = __ g Fe
159.6 g Fe2O3 1 mol Fe2O3 1 mol Fe
```

Example 2

When a 2.00-g mixture of CaCl2 (111) and RbCl (121) was analyzed, the result reports that the mixture contains a total of 0.0241 mol Cl- ions. Find weight percentage of RbCl in the sample.

Hint...

Assume x g RbCl are present in the mixture, you have
```   x     2 (2.00-x)
----- + -----------  =  0.0241
121      111```

Solving for x, x = 1.23 g (Practice your solution)

Percentage = 1.23 / 2.00 = 61.4 % RbCl.

Discussion

• A thorough understanding of chemistry is required to make an assumption such that the problem solvable.

• Please generalize and make up a similar problem and work out its solution.

• Instruction and more examples appear in topics under the general heading of Stoichiometry in the Menu. You may also review the skills using a text book on General Chemistry under this heading.

Week 3: Solutions and Reactions

Solutions

• Explain these terms and concepts associated with solutions
• solution, solute, solvent
• solubility, unsaturated solution, saturated solution, supersaturated solution
• polar substance, non-polar substance
• gaseous solution, liquid solution, solid solution

• Distinguish solutions from mixtures and colloids.

• Describe various types of solutions.

• Distinguish unsaturated, saturated, and supersaturated solutions.

Electrolytes

• Explain these terms and concepts associated with electrolytes
• ionic solid, cation, anion
• electrolyte, strong electrolyte, weak electrolyte
• acid, base, salt
• equilibrium, equilibrium constant, ion product
• ionization, autoionization

• Identify what electrolytes are.

• Distinguish between strong and weak electrolytes.

• Explain what happens when electrolytes dissolve in water.

• Give the equilibrium constant expression for ionizaton.

• Explain ion product of water, autoionization of water, and pH.

• Calculate ionization percentage of weak electrolytes

Metathesis Reactions

• Explain metathesis reactions.
• precipitation,
• neutralization,
• gas formation

• Apply metathesis reactions in gravimetric analysis and volumetric analysis

Example 1:

To determine the percentage of MgSO4 in epsom salts, you treat it with BaCl2, because of the following reaction: MgSO4 + BaCl2 ® MgCl2 + BaSO4(s) You can dry the solid substance BaSO4 formed and weigh the resulting solid to determine the quantity of MgSO4 (mol. wt. 120.37) formed. Suppose you started with 1.0000 g of epsom salts, and got 0.5000 g of BaSO4 (mol. wt. 233.39). Calculate the percentage of MgSO4 in the sample.

Hint -
see Metathesis Reactions for its solution.

Example 2:

A sample weighing 3.77 g containing CaCl2 and AlCl3 dissolved in water was treated with AgNO3, and the dry AgCl collected weighs 13.07 g. Calculate the weight and mole percentages of CaCl2 in the sample.

Formula wt: CaCl2, 111.1; AlCl3, 133.5; AgCl, 163.4.

Hint -
see Metathesis Reactions for its solution.

• Apply metathesis reactions for making chemicals

Making volatile HCl:

• H2SO4 + 2 NaCl ® 2 HCl (g) + Na2SO4

Making volatile HNO3:

• H2SO4 + 2 NaNO3 ® 2 HNO3 (g) + Na2SO4

Making volatile and poisoneous HCN gas in the gas chamber.

• H2SO4 + 2 NaCN ® 2 HCN (g) + Na2SO4

Solution Stoichiometry

• Expressing concentrations
• Converting concentrations between various units: g/L, percent, weight percentate, mole percent, mol/L (M), ppm, etc.
• Calculating amounts of solute in a given volume of solution Amount = Concentration * Volume Applying this formula to solve titration problems
• Preparing a solution of prescribed concentration
• Solving any problem involving solution stoichiometry

Balancing Redox Reaction Equations

The skill to balance redox reaction equation is rather complicated, but it can be broken down into some steps.

The first step is the skill to identify Oxidation states of any element in a chemical formula. This link helps you to acquire that skill, if you have not already done so.

The second step is to Balance Half Reaction Equations. a skill to acquire from this link.

As a review, here are the guidelines for balance reaction equations:

1. Identify the elements that are oxidized and reduced by examining their Oxidation states
2. Write the oxidation and reduction half-reactions and Balance Half Reaction Equations. In an acid solution, use H+ and H2O to balance the charges and other atoms. In a basic solution, use OH- and H2O to balance the charges and other atoms.
3. Add the two half-reactions algebraically such that the electrons in the two half-reaction equations cancel completely. Cancel other species such as H+, OH-, and H2O common to the two sides, if necessary.
4. Check your equation and make certain that numbers of atoms and charge are equal on both sides.

Week 4: The Gaseous State

• Explain Avogadro's law, Boyles law, Charles law, Daltons law of partial pressure (including mass and mole fractions) and ideal gas law.

• Apply these laws to solve problems.

• Correlate Ideal gas law with other laws.

• Describe the postulates of the kinetic theory of gas. Show how gas laws are related to these postulates.

• Compare molecular speeds according to kinetic theory of gas, and explain the diffusion and effusion phenomena.

• Discuss the difference between real gases and ideal gas.

Week 5: The Gaseous State, continue

Ideal gas law and Dalton's law of partial pressures

From the Ideal Gas law to Dalton's law for a system containing gases 1, 2, 3, ... P V = n R T
= n
Total R T
= (n1 + n2 + n3 + ...) R T
= (P1 + P2 + P3 + ...)

Mole fraction of gas 1, X1

n1
X1 = -----
nTotal
Example: A 1.0-L cylinder contains 0.5 g each of N2, O2 and CO2 at 273 K. Calculate the total pressure, partial pressure, and mole fractions.
Hint

See the table of results

Gas Amount /mol P = n R T /atm Mole fraction
N2 0.5/28 = 0.179 P = 0.40 0.40
O2 0.5/32 = 0.156 P = 0.35 0.35
CO2 0.5/44 = 0.114 P = 0.2540.25
Total 0.449 1.004 1.00

Collecting Gases over Water

The vapour pressure of water is a function of temperature, and a plot is shown.
```|Vapor Pressure
|
|     Illustrated during the lecture
|
|___________Temperature

```
The vapor-pressure vs. temperature curve is part of the Phase Diagram of water, and the vapor pressure is related to humidity and dew points. Example: On July 1, the atmosphere pressure was 745 mmHg kPa at 30 C and the air is saturated with water vapor, 31.8 mmHg. Estimate the mole fraction of water vapour.
Hint

X = 31.8 / 745 = 0.043 (4.3 %)

The partial pressure of dry air would have been 745 - 31.8 = 713 mmHg.

Kinetic Theory of Gases

Postulates:
• Molecular size, attraction, & repulsion negligible
• Average kinetic energies of gases are the same at the same T
• Pressure due to kinetic energy
For 1.0 mole of ideal gas, N = Avogadro's No., V = molar volume.

P V = N m u2/2
= (3/2) R T

u = (3 R T / N M)1/2
= (3 k T / M)1/2

Comparing 2 gases

```        M1 u12  =  M2 u22.
Thus,

M1     u2
-- = (---)2
M2     u1```

The Graham's law of effusion is

k             k '
rate of effusion = ------ = ------
d1/2       M1/2

Applications:

• Ratio of effusion rates

• Balloon business

• Separation of isotopes

Real Gases

van Der Waals Equation (P + (n/V)2 a) (V - nb) = n R T

nb - Volume of n moles molecules.

(n/V)2 a - Pressure correction term.

Week 6: Thermochemistry

Explain what energy is and expand the discussion on the following:

• work = F . s;   1 N m = 1 J

• potential energy = m g h

• kinetic energy = ½ m v2

• temperature scale, kinetic energy of molecules, temperature measurements

• heat, 1 cal = 4.182 J
Measurement of heat = C DT
= s m Dt

• electric energy E = q V (charge q times voltage V)

• electromagnetic radiation E = h v

• matter and energy E = m c2

• chemical energy - energy in substances by virtue of state, bonding, and composition

Thermochemistry (Chemical energy)

• Define a system and surroundings for the following discussion.

• What is zeroth law of thermodynamics?

• What is the first law of thermodynamics?

• Explain first law as the conservation of energy by defining internal energy E (Some textbook uses U) as the increase in internal energy of a system being the amount of heat q plus the amount of work w. DE = q + w
= Efinal - Einitial

• Define what is enthalpy = DH, a special term required for reaction taking place at constant pressure. DH = DE + DP V
= Hfinal - Hinitial

• Define and evaluate enthalpy of reaction, DH for a given amount of substances (mole, gram, ton)

• Define and evaluate standard enthalpy of reaction, DHo for a given amount of substances (mole, gram, ton)

• Define and evaluate standard enthalpy of formation, DHof for a given amount of substances (mole, gram, ton)

• Use energy level to express chemical energy.

• Define Hesses' law and apply it to calculate energy of processes.
Work on some examples and problems.

Week 7: Hess's Law and Introduction to Quantum Theory

Hess's law

Hess's law states that energy changes are state functions. C + 1/2 O2 -> CO,         dH° = -110 kJ/mol
CO + 1/2 O2 -> CO2,     dH° = -283 kJ/mol.
+ ___________
C + O2 -> CO2,         dH° = (-110)+(-283) = -393 kJ/mol.
From which, we write
```
0 kJ ------------ C(graphite) + O2
|       |
-110 kJ |       |
V       |
CO + /2 O2 -----     |
|       | -393 kJ
|       |
-283 kJ |       |
|       |
V       V
------------ CO2

```

Example 1

The enthalpy of combustion for H2, C(graphite) and CH4 are -285.8, -393.5, and -890.4 kJ/mol respectively. Calculate the standard enthalpy of formation dHf for CH4.

Solution:
Lets interpret the information about enthalpy of formation by writing out the equations as shown here.
dH°f
/(kJ/mol)
(1) H2(g) + 0.5 O2(g) -> H2O(l) -285.8
(2) C(graphite) + O2(g) -> CO2(g) -293.5
(3) CH4(g) + 2O2(g) -> CO2(g) + 2H2O(l) -890.4
From the above equations, derive
C + 2H2 -> CH4
Answer: C + 2H2 -> CH4 -74.7
Hint: 2*(1) + (2) - (3), Thus,
dHf = 2 * (-285.8) + (-393.5) - (-890.4) = ?

Discussionn:

```
===C(graphite) + 2 H2(g) + 2 O2(g)===
- 74.7 kJ |              |
== CH4 (g) + 2 O2(g)== |
|              |
|              |
|              |
|              |-965.1 kJ
-890.4 kJ |              | [(-2*285.8-393.5) kJ]
|              |
|              |
|              |
|              |
V              V
==========CO2(g) + 2 H2O(l)==========
```

Example 3

The standard enthalpies of formation of SO2 and SO3 are -297 and -396 kJ/mol respectively. Calculate the standard enthalpy of reaction for the reaction: SO2 + 1/2 O2 -> SO3.

Solution:
In order to show how the chemical reactions take place, and for a better appreciation of the technique of problem solving, we write the equations according to the data given:

SO2(g) -> S(s) + O2(g)       dH = 297 kJ
S(s) + 3/2 O2 -> SO3           dH = -396 kJ
Add the two equations to give SO2(g) + 1/2 O2 -> SO3       dH = -99 kJ

Your turn to work:
Sketch an energy level diagram for the combinations of substances.

Calorimeter Measurements

Calorimetry the science of heat measurements, already mentioned. What is a Calorimeter?
Do you know how to apply the following: q = C dT = s m dT
C = q / dT

dH = dE + d(PV) = dE + P dV = dE + dn R T

Example 5 (see Calorimetry for other examples)

A table of thermodynamic data gives dHf = -285.8 kJ/mol for water. A bomb calorimeter measurement gives the heat of combustion for H2 as -282.0 kJ/mol. Estimate the error of the enthalpy measurement.

Solution
Reinterpret the problem, we have

H2(g) + 0.5 O2(g) = H2O(l),   dE = 282.0 kJ/mol.
Furthermore,
dn = -1.5
dH = dE + dn R T,
= - 282.0 + (-1.5 mol * 8.314 J/(mol K) * 298 K)
= (-282.0 - 3.72) kJ
= -285.7 kJ
The error is (285.8-285.7)/285.8 = 0.03%

Discussion
More heat is giving of if the reaction is carried out at constant pressure, since the P-V work (1.5 R T) due to the compression of 1.5 moles of gases in the reactants would contribute to dH.

If 1.0 mole water is decomposed by electrolysis at constant pressure, we must supply an amount of energy equivalent to enthalpy change, dH, a little more than internal energy, dE. More energy must be supplied to perform the P-V work to be done by the products (H2 and O2).

Introduction to Quantum Theory

Stories told:
• Discovery of electrons by J.J. Thomson, who determined the charge to mass ratio of electrons.
• Electromagnetic radiation formulas, c = l n
l
= c / n
c: velocity of light (3e8 m s-1 l: wavelength
n: frequency
• Discovery of X-rays by W. C. Roentgen while he performed experiments on cathode rays.
• Discovery of radioactivity by H. Becquerel from urnaium salt while trying to see if X-rays will be emitted by other material exposed to Sun rays.
• Rutherford's alpha scattering experiments revealed the size of the nucleus and the structure of the atoms. Among his conclusion is that the radius of the nucleuss is (1/100,000)th of that of the atom. Since the radius of the atom is 1e-10 m, the radius of the nucleus is 1e-15 m.

Here is an overallview of quantum theory:

• Electromagnetic Radiation
Transmission of energy through space via no medium is electromagnetic radiation. The visible region of the electromagnetic radiation is light, but that is a very small region. There is much more than light to meet the eye.

key equation: c = l n
or velocity of light = wavelength * frequency

• Spectra
Diagrams showing the distribution of intensity versus wavelengths are called spectra. Their study reveals the fundamentals of electromagnetic radiation as well as leading to useful applications. For example, the Hydrogen Spectra study led Bohr and others to develop the quantum theory to describe the atomic structures. For some cool spectrum demonstration, check out the IR Tutor created by Charles Abrams. Here is one of his animated pictures.

Week 8

Review and examples. I hand-wrote the review lecture notes, which cannot be post here. Sorry!

• Quantum numbers
The states of electrons are represented by wavefunctions. Each wavefunction has a set of numbers, called quantum numbers. We often use quantum numbers to describe properties of electrons. This page gives a simplistic but important view of quantum numbers.

• Atomic orbitals
Electronic states, represented by wavefunctions, in an atom are called atomic orbitals. Since we use quantum numbers to describe them, atomic orbitals are labelled by quantum numbers, such as 2s orbital, s represent quantum number l = 0, that implies m = 0. Each atomic orbital accomodates two electrons due to electron spin.

• Periodic table
The beauty of quantum theory is its mathematical results not only explain the arrangement of the elements in the Periodic Table of chemical elements, but they seem rationalize the existence of the elements. Its rationalization lies in the Electronic configuration of atoms with more than two electrons. Based on the electronic configurations, Atomic properties are nicely explained.

• Zests of elements - A delightful look at the chemical elements. Elements review - A review and quiz.

Week 9

(Chapter 8) Electronic Configurations and Chemistry of the Periodic Table,
Midterm test on Friday for chapters 6, 7, and 8.

Review: Explain and apply the following terms

• atomic orbital
rules for quantum numbers
representations K, L, M, etc.
s, p, d, f, ... coressponding to l = 0, 1, 2, 3, ..., and subshells
electron distribution and shape of orbitals
• Applcation of Pauli exclusion principle
application of Hunds rule (for subshells).
• electron configuration the build-up (aufbau principle)
Apply the general rule to give the electronic configuration for any atom (not exceptions).
Figure out the electronic configuration from location of the element in the periodic table.
• magnetic properties of material
• diamagnetism - apply to all substance
• paramagnetism - substance having unpaired electrons
• ferromagnetism - has lined-up domains
• antiferromagnetism - with antiparallel domains
• Explain and point out the trends in
• atomic radii
• 1st IP
• Electron affinity (EA)

Week 10. Ionic and Covalent bonds (Chemical bonding)

Ionic Compounds
• Acids H+(aq) and anions
• Bases Cations and OH-(aq)
• Salts - by products of acid-base reactions or by reaction HCl + NaOH = H2O + NaCl In reality, HCl gas dissolves in water forming H+(aq) and Cl-(aq)
NaOH solid dissoves in water forming Na+(aq) and OH-(aq)
Neutralization reaction, H+(aq) + OH-(aq) = H2O
Spectator ions Na+(aq) + Cl-(aq)
By direct reaction Na ([Ne]3s1) + Cl ([Ne]3s23p5)
=> Na ([Ne]) + Cl ([Ne]3s23p6 same as __)
Using Lewis dot structure: Na. + :Cl.
Write out the Lewis Dot Symbols for the following

H

Li     Be     B     C     N     O     F     Ne

Na     Mg     Al     Si     P     S     Cl     Ar

Electronic configurations of ions

Main groupelements, ionization potentials (IP) of some elements

Element1st IP2nd IP3rd IP4th IP
Na 496 4,562 6,912 9,543
Mg 738 1,451 7,733 10,540
Al 578 1,817 2,754 11,577

Group IA, IIA, and IIIA - ions have pseudo-noble-gas electronic configurations

Groups IIIA to VA heavy metals: charge = (group number - 2)

Tl1+, Sn2+, Pb2+ and Bi3+

Anions of Groups VA to VIIA (pseudo-noble-gas electronic configurations)

N3-, O2-, Br1-

Transition-Metal Ions

• Transition-metals lose n s electrons first, and then they may lose one or more (n-1) d electreons to form ions.
• Transition metal ions are usually colored, due to the d-d transitions. The d-sub shell is split for ions in solutions.
• Electronic configurations for transition metal ions.

Sc [Ar]4s2 3d1
Sc2+ [Ar] 3d1
Sc3+ [Ar] (colorless)

Ti [Ar]4s2 3d2
Ti4+ [Ar] (colorless)

V [Ar]4s2 3d3
V2+ [Ar] 3d3

Cr [Ar]4s1 3d5
Cr3+ [Ar] 3d3

Fe [Ar]4s2 3d6
Fe2+ [Ar] 3d6
Fe3+ [Ar] 3d5

Co [Ar]4s2 3d7
Co2+ [Ar] 3d7

Ni [Ar]4s2 3d8
Ni2+ [Ar] 3d8

Cu [Ar]4s1 3d10
Cu1+ [Ar] 3d10 (colorless)
Cu2+ [Ar] 3d9

Zn [Ar]4s2 3d10
Zn2+ [Ar] 3d10 (colorless)

Ionic Radii
• Ionic radii of the same charge increase for a group as the atomic numbers increase.
• Isoelectronic ions are ions with the same electronic configuration.
• Across a period, the cations decrease in radius, even if they have the same charge. The higher the charge, the smaller the ion.
• Sizes of anions are much larger than their isoelectronic cations.

Born-Haber Cycle for the Calculation of Lattice Energy

Since the concepts of lattice energy and energy of crystallization are sound, and lattice energy is a useful quantity to judge the chemical properties of the salt, we can and usually calculate it by applying the law of conservation of energy, in a manner similar to the Hess's law. For the calculation of lattice energy, this process is called the Born-Haber cycle.

In 1919, M. Born and F. Haber separately devised a method to calculate the lattice energy from known thermodynamic data such as,

 Hsub Enthalpy of sublimation, Hmelt Enthalpy of melting, Hvap Enthalpy of vaporization, Hf Enthalpy of formation, U Lattice energy, Ecryst Energy of crystallization, H Enthalpy of reaction, D Bond (dissociation) energy IP Ionization potential, or ionization energy, IE EA Electron affinity,
For example, to calculate the lattice energy of NaBr, we start with Na metal and Br2 liquid. To make NaBr, each of the reactants need to go through a series of processes in order to calculate the lattice energy:

2 Na (s) --2 Hsub ® 2 Na (g) --2 IP ® 2 Na+(g)

Br2(l) --Hvap ® Br2 (g) --D ® 2 Br(g) --2EA ® 2 Br-(g)

One other reaction is required

2 Na + Br2 ® 2NaBr, 2Hf

The lattice energy corresponds to the reaction

2 NaBr(s) ® 2Na+(g) + 2 Br-(g)

To put all the information together, we may now draw the cycle:

```          THE BORN-HABER CYCLE
-----------2Na+ + 2 Br(g)--------
­                       |
|2IP + D            |2 EA
|                       ¯
2Na(g) + Br2(g)     2 Na+(g) + 2Br-(g)
­                       |
|2Hsub+Hvap    |
|                       |
2Na(s) + Br2(l)        | 2Ecryst
|                       |
|2Hf                  |
¯                      ¯
----------2 NaBr(s)---------------
```
Thus,
2Hf - (2Hsub + Hvap + 2IP + D) = 2EA + 2Ecryst

Note that we have used two moles of NaBr in the above diagram.

This scheme shows that we can calculate the lattice energy of NaBr from some known thermodynamic data. The same can be calculated from reaction equations and their associated energies. This is illustrated below

 2 Na(s) + Br2(l) ® 2 NaBr(s) 2 Hf 2 Na(g) ® 2 Na(s) or 2 Na(s) ¬ 2 Na(g) - 2 Hsub 2 Na+(g) + 2e ® 2 Na(g) or 2 Na(g) ¬ 2 Na+ + 2e - 2 IP Br2(g)®Br2(l) or Br2(l) ¬ Br2(g) - Hvap 2 Br(g) ® Br2(g) or Br2(g) ¬ 2 Br(g) - D 2 Br-(g) ® 2 Br(g) + 2 e or 2 Br(g) + 2 e ¬ 2 Br-(g) - 2 EA Add all the above equations leading to 2 Na+(g) + 2 Br-(g) ® 2NaBr(s) 2 Ecryst
Thus, 2 Hf - 2Hsub - 2 IP - Hvap - D - 2 EA = 2 Ecryst
Ecryst = Hf - Hsub - IP - (Hvap + D) / 2 - EA
The same result as shown in the diagram

The H-H molecule

Describe the hydrogen molecule in light of the following:
• H-H
• H:H
• Valence bond theory of H2
• Molecular orbital theory of H2
• Electron configuration of molecules

Bondlength and bond energies

• Define bondlength and bond energy and note relationship between the two
• Define bond order explain its relationship to bondlength or bond energy (sigma and pi bond)
• Evaluate enthalpies of reactions using bond energies
• Recognize covalent substances and characterize ionic character as difference in electonegativity
• Describe trends in bondlengths of a series of related compounds

Week 11. Ionic and Covalent bonds (continue) and Molecular Geometry

Lewis Dot Structure

• Draw the Lewis dot structure of a given molecule or ion.
• Draw several possible structures to consider resonance structures.
• Assign formal charge to an atom in a dot structure.
• Assess the stability of a structure by considering formal charges of atoms.
• Give examples for molecules and ions that do not follow the octet rule.

Valence-shell electron-pair repulsion models

1. Identify the central atom in a molecule containing more than two atoms as a start.
2. Identify the number of valence electrons of any element. This concept is important, because you need to know the number of valence electrons in order to write a Lewis dot structure for the molecule in question.
3. Count the number of VSEPR pairs or steric number (SN) for the central atom in a molecule. You need this number in order to describe or predict the shape of the molecule in question.
4. Determine the number of lone electron pairs that are not shared with other atoms. Often, a Lewis dot structure is useful to help you count this number.
5. Predict the shape of molecules or inos as the key concept of VSEPR theory. From the shape and by applying the idea that lone electron pairs takes up more space, you can predict the bond angles withing 5% of the observed values.
6. Predict the values of bond angles and describe the hybrid orbitals used by the central atoms in the molecules or ions.

Dipole moment and molecular geometry

Mostly hand-written notes, not available here.

Hybrid atomic orbitals valence bond theory and VSEPR model.

• Describe the valence bond (VB) approach to chemical bonding
• Demonstrate hybridization of atomic orbitals for VB
• Correlate the molecular shape to the hybrid atomic orbitals of some central atoms.
• Combine the concepts of hybrid orbitals, valence bond theory, VSEPR, resonance structures, and octet rule to describe the shapes and structures of some common molecules.

Week 12. Molecular Geometry

Valence bond theory and hybrid orbitals

Mostly contained in the website above.

Hybrid orbitals in chemical bonds of organic compounds

Describe the hybrid orbitals used in various organic compounds. The hybrid orbitals are described using hand-drawn diagrams and models. If you miss the lectures, it will take you a lot of time to catch up. Your ability to describe the bonding and shape of these molecules is very important. By applying the basic ideas provided here, you will be able to understand the complicated organic molecules in the future.
HCCH H2CCH2 H3CCH3 H3CNH2
H3COH C3CHO H3CCOOH H3CCCCH3
CO CN- NO SO2
C6H6 C6H12 NCS- CO2
H3C-CC-C-NH-CH(OH)-CH=C=O

Molecular orbital (MO) theory. MO of Li2, Be2, B2, C2, N2, O2 and F2

Much more extensive than the coverage of the text book. Because hand drawing and explanation are required, no information is provided on the website. Final Examination Questions will require all concepts introduced during the lectures.

Week 13. (Monday only)

Review: stoichiometric calculations, chemical energy, Hess's law, Born-Habor cycle, etc.

Final Examination

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The final examination is divided into two parts.

Part A (30%) are questions that require you to provide written answers. In providing answers to them, you have to apply several concepts. Most of the exercises and CAcT questions are base on one concept. We have not finalized the exam paper yet, but we are currently considering letting you choose 3 out of 4 questions in Part A. Chemical bonding (Chapters 9 and 10) concepts are important for this part.

Part B (70%) consists of 35 multiple choice questions. You already know the style of these questions.

The average for the final examination is usually lower than those of term tests. The average for the written part is usually lower than the multiple choice part.

E-mail cchieh@uwaterloo.ca