Transcript Chapter 2 and Chapter 4 Review

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
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 (2.1-2.6), half
of the test is on Chapter 4 (4.1-4.6)
• 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+)
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!!!
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
786-9866
[email protected]
Key Concepts – Chapter 2
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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
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
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 = 1.01325 bar = 101 325 Pa = 760
torr
• Torr: 1 torr = 1 mm Hg = 133.3 Pa
• Standard Temperature and Pressure (STP)
– Defined as Pressure = 1 bar, Temperature = 273.15 K
– Molar volume of a gas at this pressure is 22.65 L
– Note: to convert from Celcius to Kelvin
⁰K = ⁰C + 273.15
Relationships Between Gas Properties
Ideal Gas Law
Ideal Gas Law Constant:
R = 8.314 L kPa mol-1 K-1
R = 0.08314 L bar mol-1 K-1
Relationships we can use
Gas concentrations
• Gas concentrations can be described in the
following ways:
1. Moles per unit volume (typically mol m-3)
2. Partial pressures (torr, Pa, kPa, Bar)
3. Mole fractions
1.
2.
3.
4.
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
Gas Stoichiometry Problems
• To solve stoichiometry problems, follow these
steps.
1. Understand the problem and what physical
process is happening, if it helps draw a picture.
2. Write down all pertinent information. Also, look
for assumptions that will simplify a question, or
make information not pertinent.
3. Think about what you need to solve for, what
information you have, and what equations are
available.
4. Balance the relevant equation(s).
Gas Stoichiometry Problems
5. Determine the limiting reagent.
6. Build an ICE (textbook uses ICF) table.
7. Determine the final number of moles or
concentrations of all compounds
8. Solve for the measurement needed (check for
correct units).
9. Look at your answer, is it reasonable?
Gas Density
Speed and Energy
• Kinetic energy distribution
increases with increasing
temperature
• Velocity decreases with
increasing molar mass.
Most probable Kinetic Energy and
velocity
Average Kinetic Energy and Velocity
Called root-meansquare speed
Graham’s Law
Concepts for Chapter 4
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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
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.
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.
Electromagnetic Spectrum
Lower Energy
Higher Energy
• Visible light is a very small portion of the electromagnetic
spectrum
• Wavelength is inversely proportional to frequency
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
Photoelectric Effect 2
• Observations of ejected electrons were as
follows:
1. Below the threshold frequency (ν0), no electrons are
observed, regardless of light intensity
2. Above ν0, the maximum kinetic energy increases
linearly with light intensity.
3. Above ν0, the number of emitted electrons
increases with light intensity, but the energy of each
electron is independent of light intensity.
4. All metals show the same pattern, but each metal
has a different ν0.
Energy-Wavelength-Frequency
Relationships
Quantization of Energy
n=5
n=4
n=3
n=2
n=1
Ground state
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.
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.
Momentum of Particles
Summary of Particle and Wave
Equations
Table 4-1 p. 215
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
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.
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.
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.
Azimuthal Quantum Number
Value of l
0
1
2
3
Orbital Letter s
p
d
f
Magnetic Quantum Number
Magnetic Quantum Number
Spin Quantum Number
+1/2
-1/2
Ways of depicting orbital shapes
• Electron density plot: shows
electron density with varying
distance from the nucleus
• Orbital density picture: Twodimensional picture showing a
cross-section of orbital
• Electron contour drawing:
Shows a contour surface that
encloses almost all (often 90%)
of the electron density
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
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)
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
Orbital shapes – d orbitals
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).