Quick Review of Skills by Week
- 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
- 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
- What are the molecular or formula weight of H2,
H2O, H2O2, O3, NaCl, and
- 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
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
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:
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?
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
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.
Assume x g RbCl are present in the mixture, you have
x 2 (2.00-x)
----- + ----------- = 0.0241
Solving for x, x = 1.23 g (Practice your solution)
Percentage = 1.23 / 2.00 = 61.4 % RbCl.
- 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
- 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.
- 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
- Explain metathesis reactions.
- gas formation
- Apply metathesis reactions in gravimetric analysis and volumetric analysis
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.
Metathesis Reactions for its solution.
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,
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
The skill to balance redox reaction equation is rather complicated, but it
can be broken down into some steps.
- 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
The first step is the skill to identify
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:
- Identify the elements that are oxidized and reduced by examining their
- 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.
- 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.
- 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
= nTotal R T
= (n1 + n2 +
n3 + ...) R T
= (P1 + P2 +
P3 + ...)
Mole fraction of gas 1, X1
X1 = -----
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
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 = 0.254||0.25
|Total ||0.449 ||1.004 || 1.00
Collecting Gases over Water
The vapour pressure of water is a function of temperature, and a plot
| Illustrated during the lecture
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.
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.
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
For 1.0 mole of ideal gas, N = Avogadro's No.,
V = molar volume.
- Molecular size, attraction, & repulsion negligible
- Average kinetic energies of gases are the same at the same T
- Pressure due to kinetic energy
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.
-- = (---)2
The Graham's law of effusion is
k k '
rate of effusion = ------ = ------
- Ratio of effusion rates
- Balloon business
- Separation of isotopes
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
Thermochemistry (Chemical energy)
Work on some examples and problems.
- Define a system and surroundings for the following
- 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,
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
Week 7: Hess's Law and Introduction to Quantum Theory
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 | |
CO + /2 O2 ----- |
| | -393 kJ
-283 kJ | |
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.
Lets interpret the information about enthalpy of formation
by writing out the equations as shown here.
|(1) H2(g) + 0.5 O2(g) -> H2O(l)
|(2) C(graphite) + O2(g) -> CO2(g)
|(3) CH4(g) + 2O2(g) -> CO2(g) + 2H2O(l)
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) = ?
===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]
==========CO2(g) + 2 H2O(l)==========
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.
SO2(g) -> S(s) + O2(g)
dH = 297 kJ
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:
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:
Calorimetry the science of heat measurements, already mentioned.
What is a Calorimeter?
Sketch an energy level diagram for the combinations of substances.
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
H2(g) + 0.5 O2(g) = H2O(l),
dE = 282.0 kJ/mol.
Reinterpret the problem, we have
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%
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
- 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
- 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:
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
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
Review and examples. I hand-wrote the review lecture notes, which cannot be
post here. Sorry!
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.
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.
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
of atoms with more than two electrons. Based on the electronic configurations,
Atomic properties are
Zests of elements
- A delightful look at the chemical elements.
- A review and quiz.
(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
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)
Write out the Lewis Dot Symbols for the following
- Acids H+(aq) and anions
- Bases Cations and OH-(aq)
- Salts - by products of acid-base reactions or by reaction
HCl + NaOH = H2O + NaCl
HCl gas dissolves in water forming H+(aq) and Cl-(aq)
NaOH solid dissoves in water forming Na+(aq) and OH-(aq)
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.
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
|Element||1st IP||2nd IP||3rd IP||4th 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-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 of the same charge increase for a group as the atomic numbers
- 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,
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:
|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,
2 Na (s) --2 Hsub
® 2 Na (g) --2 IP
® 2 Na+(g)
® 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
2Hf - (2Hsub +
Hvap + 2IP + D)
= 2EA + 2Ecryst
-----------2Na+ + 2 Br(g)--------
|2IP + D |2 EA
2Na(g) + Br2(g) 2 Na+(g) + 2Br-(g)
2Na(s) + Br2(l) | 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 Hf - 2Hsub - 2 IP - Hvap
- D - 2 EA = 2 Ecryst
|2 Na(s) + Br2(l) ® 2 NaBr(s)
|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
or Br2(l) ¬ Br2(g)
|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)
Ecryst = Hf - Hsub
- IP - (Hvap + D) / 2 - EA
The same result as shown in the diagram
Describe the hydrogen molecule in light of the following:
- Valence bond theory of H2
- Molecular orbital theory of H2
- Electron configuration of molecules
- 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
- 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.
Identify the central atom in a molecule containing more than two
atoms as a start.
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.
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.
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
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.
Predict the values of bond angles and describe the
used by the central atoms in the molecules or ions.
Dipole moment and molecular geometry
Mostly hand-written notes, not available here.
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
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.
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
|CO ||CN- ||NO ||SO2
Week 13. (Monday only)
Review: stoichiometric calculations, chemical energy, Hess's law,
Born-Habor cycle, etc.
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.