1. Chapter 2 and Chapter 4 Review
CHM1111 Section 04
Instructor: Dr. Jules Carlson
Class Time: M/W/F 1:30-2:20
Wednesday, October 5th

2. For the Mid-Term Wednesday, October 12th at 1:30 PM
There will be 10 multiple choice and 2 medium length answer questions
Half of the test is on Chapter 2 ( ), half of the test is on Chapter 4 ( )
You will be given a periodic table and constants but no equations
You are allowed to use most calculators (just no high end programmables TI 83+)

3. Test-taking strategies
Always read the test from beginning to end before starting to answer questions.
If you do not know how to answer a question, skip it and come back to it at the end.
If you finish the test early, look over your answers before submitting your test.
Above all: DON’T PANIC!!!

4. Math and Science Tutor Centre Review
Introduction to the Chemical Properties of Matter
Test 1 Review
Friday, Oct. 7, 2011
2:30 - 4:30
in Room 4M31 (Theatre A)
Andrew Bendor-Samuel Director, Math & Science Tutoring Centre The University of Winnipeg

5. Key Concepts – Chapter 2 How is pressure measured? Unit conversions
Gas Laws
Ideal Gas Law
Concept of Mole Fractions (also ppm, ppb)
Dalton’s Law of Partial Pressures
Gas Stoichiometry Problems
Gas Density
Kinetic Molecular Theory
Molecular speeds, energies
Graham’s Law

6. Mercury Barometer
Originally done with water, fill a tube with water, place it inverted in water basin
Water would drop to 10.3 m above the height in the basin
Later performed with mercury (760 mm at sea level)
Pressure at the top of the column and atmospheric pressure are equal
1 mm of Hg was named 1 torr of pressure

7. Unit Conversions Pascal: 1 Pa = 1 N/m2 = 1 kg/ms2
Units show pressure dependent on mass, position, and time
Bar: 1 bar = 100 kPa = 750 torr
Atmosphere: 1 atm = bar = Pa = 760 torr
Torr: 1 torr = 1 mm Hg = Pa
Standard Temperature and Pressure (STP)
Defined as Pressure = 1 bar, Temperature = K
Molar volume of a gas at this pressure is L
Note: to convert from Celcius to Kelvin
⁰K = ⁰C

8. Relationships Between Gas Properties

9. Ideal Gas Law Ideal Gas Law Constant: R = 8.314 L kPa mol-1 K-1
Ideal Gas Law Constant:
R = L kPa mol-1 K-1
R = L bar mol-1 K-1

10. Relationships we can use

11. Gas concentrations
Gas concentrations can be described in the following ways:
Moles per unit volume (typically mol m-3)
Partial pressures (torr, Pa, kPa, Bar)
Mole fractions
Fraction of 1
Percent, PerMille (Per Thousand)
Parts per million (ppm) – 1 molecule in 106 molecules
Parts per billion (ppb) – 1 molecule in 109 molecules

12. Gas Stoichiometry Problems
To solve stoichiometry problems, follow these steps.
Understand the problem and what physical process is happening, if it helps draw a picture.
Write down all pertinent information. Also, look for assumptions that will simplify a question, or make information not pertinent.
Think about what you need to solve for, what information you have, and what equations are available.
Balance the relevant equation(s).

13. Gas Stoichiometry Problems
Determine the limiting reagent.
Build an ICE (textbook uses ICF) table.
Determine the final number of moles or concentrations of all compounds
Solve for the measurement needed (check for correct units).
Look at your answer, is it reasonable?

14. Gas Density

15. Speed and Energy
Kinetic energy distribution increases with increasing temperature
Velocity decreases with increasing molar mass.

16. Most probable Kinetic Energy and velocity

17. Average Kinetic Energy and Velocity
Called root-mean-square speed

18. Graham’s Law

19. Concepts for Chapter 4 Characteristics of Light
Electromagnetic Spectrum
Photoelectric Effect
Energy-Wavelength-Frequency Relationships
DeBroglie Wavelengths
Momentum (mass x velocity)
Heisenberg’s Uncertainty Principle
Describing Electron Orbitals
4 quantum numbers
Ways of depicting orbitals
Orbital Shapes

20. Characteristics of Light
Light behaves both as a wave and a particle
Light is composed of photons, which have particle properties
Light as a wave can be described by:
Frequency (ν, Greek letter nu): number of cycles the wave undergoes in units of s-1.
Wavelength (λ, Greek letter lambda): the distance between any point on the wave and the corresponding point on the next wave, measured in nm (10-9 m).
Amplitude (A): height of the wave from the axis of propagation, a measure of intensity, measured in m.

21. Characteristics of Light
The height of a wave is the amplitude
Brighter light has a wave of larger amplitude.
Light of higher frequency has a shorter wavelength.

22. Electromagnetic Spectrum
Lower Energy
Higher Energy
Visible light is a very small portion of the electromagnetic spectrum
Wavelength is inversely proportional to frequency

23. The Photoelectric Effect
Experiments performed by Phillipp Lenard, and Albert Einstein
Shine light onto a metal surface, and electrons would be ejected
Electrons hit a detector, can measure a current
Electrons would only be ejected if the light had a certain minimum frequency

24. Photoelectric Effect 2
Observations of ejected electrons were as follows:
Below the threshold frequency (ν0), no electrons are observed, regardless of light intensity
Above ν0, the maximum kinetic energy increases linearly with light intensity.
Above ν0, the number of emitted electrons increases with light intensity, but the energy of each electron is independent of light intensity.
All metals show the same pattern, but each metal has a different ν0.

25. Energy-Wavelength-Frequency Relationships

26. Quantization of Energy
n = 5
n = 4
n = 3
n = 2
n = 1
Ground state

27. Atomic Spectra
Atoms absorb specific and characteristic frequencies/wavelengths of light.
Depends upon the energy differences between ground and excited states for the atoms electrons.
The pattern of absorbed photons is an absorption spectrum.
The pattern of emitted photons is an emission spectrum.

28. Atomic Spectra
When a photon is absorbed, it has to have a wavelength or frequency which matches the difference between the energy levels.
Sometimes, the electron can drop from a higher energy state to a lower energy state producing a photon - the photon produced also has a wavelength or frequency that matches the difference between the energy levels.

29. Momentum of Particles

30. Summary of Particle and Wave Equations
Table 4-1 p. 215

31. Heisenberg’s Uncertainty Principle
Heisenberg’s uncertainty principle states the following:
The position and momentum of an electron are linked. The better we know the position, the less certain we are of its momentum and vice-versa

32. Describing Electron Orbitals
The spatial distribution of an electron around a nucleus is described by a 3-dimensional wave.
These 3-dimensional waves are called orbitals.
The quantized properties of electron orbitals can be identified using quantum numbers.
There are 4 different quantum numbers that describe electron orbitals, comes from solution to the Schrödinger equation.
These quantum numbers and their properties will be useful to explain bonding and magnetic properties.

33. What do electron orbitals look like?
Electrons can be found in particular regions where electrons may be found described by their wavefunction solution to the Schrödinger equation - describes an electron orbital.
These orbitals can have different shapes and sizes.

34. Principle Quantum Number
The Principal Quantum Number (n) indexes energy.
No simple equation in multi-electron systems for energy of levels like there is in Hydrogen atoms.
Values are positive integers.
Tells you something about the size of an orbital (more energy an electron has, the more it will move)
As n increases, energy increases, orbital size increases.

35. Azimuthal Quantum Number
Value of l 1 2 3
Orbital Letter s p d f

36. Magnetic Quantum Number

37. Magnetic Quantum Number

38. Spin Quantum Number
+1/2
-1/2

39. Ways of depicting orbital shapes
Electron density plot: shows electron density with varying distance from the nucleus
Orbital density picture: Two-dimensional picture showing a cross-section of orbital
Electron contour drawing: Shows a contour surface that encloses almost all (often 90%) of the electron density

40. Orbital Size As n gets larger, orbital size increases
All orbitals with the same n are similar in size
However, orbitals decrease in size with increasing nuclear charge

41. Orbital Shapes – s orbitals
s orbitals are spherical, and have one node fewer than n.
Nodes occur where the wave has a minimal amplitude (see circles).
Waves can also have different phases (see with p orbitals)

42. Orbital shapes – p orbitals
p orbitals are lobed-shaped
The three p orbitals are oriented across the three axes
There is a node which separates each p-orbital into 2 phases, the electron occupies both phases
Oriented along z-axis
Oriented along y-axis
Oriented along x-axis

43. Orbital shapes – d orbitals

44. Energies of orbitals
For Hydrogen (and other atoms): When hydrogen absorbs a photon, the electron goes from the ground state to an excited state. One excited state may be a 2p orbital. If the electron is given sufficient energy, the electron can escape all bound states and the atom ionizes. This energy is called the ionization energy (IE).