Chapter 12
9.x
Thermodynamics
Energy is
SI unit of energy is the
Endothermic Process:
Exothermic Process:
System and Surroundings
System:
Surroundings:
Universe =
12.1
Electromagnetic Radiation
Wavelength
Chapter 12
Frequency
I.
L
II.
III.
Symbol:
Units:
IV.
Speed of light through a vacuum, c, is
12.2
The Quantum Revolution
Black Body Radiation:
Planck:
Quantum
Energy can only be gained or lost only in
E= nhv
h=
n=
v=
The Photoelectric Effect
Einstein used Planck’s quantum theory to explain the
photoelectric effect.
Each energy packet behaves like a tiny packet of light and is
called a
When
; excess photon energy appears as KE of
electron ejected from the metal:
Chapter 12
Determination of
Planck's Constant
What wavelength of radiation has
photons of energy 8.23x10-19 J?
a) 806 cm
b) 4.67 Mm
c) 1.24 Pm
d) 241 nm
e) I am lost
Duality
Dual nature of light:
Matter?
de Broglie Equation:
Diffraction
12.3
Spectrum
Radiation composed of 1 wavelength =
Radiation can be composed of many wavelengths
When separated a spectrum is produced:
Chapter 12
Bohr Model
“Microscopic solar system”
I.
The lowest energy state,
n=
is called the
II.
When the electron is in a
higher state, n=
the atom is said to be in an
III.
Electrons can “jump” from one allowed energy level by emitting/absorbing a photon of
light.
Hydrogen Emission Spectrum
Chapter 12
12.4
Rydberg Equation
n=
Correctly predicts lines in H-atom emission spectrum
Z=
12.5 and 12.7
Quantum Mechanics
I.
Uncertainty Principle:
II.
Schrodinger’s Equation:
Scrodinger’s Equation
Treat the electron as a standing wave
Differential Equation
Only certain solution satisfy the equation
A specific wave function for a given electron is often called an
Chapter 12
When the amplitude passes zero, we call that a
High probabilities make orbitals
12.8 and 12.9
Quantum Numbers
I.
II.
III.
IV.
V.
VI.
VII.
VIII.
IX.
X.
XI.
Each orbital describes a specific distribution of electron density in space, given by its
probability
The
, n, can
have integer value of:
As n
the electron spends
time
away from the
(has a
energy)
Electron shell
Energy of the electron is only dependant on
The angular momentum (Azimuthal) quantum
number
This describes the
Letters are used instead of numbers:
Subshell
The magnetic quantum number,
including zero
The quantum number describes
S-orbital
P-orbital
, can have integer values between
and
Chapter 12
d-Orbital
Sections 12.10, 12.11, 12.13, 12.14
Energy Diagram For Hydrogen
Spin
Electrons act as magnets: Aligned with or against an external magnetic field.
s=
For an electron s=
= spin magnetic q#
The four q#:
are used to characterize a electron in an atom.
Pauli exclusion principle:
Chapter 12
Recap
Quantum numbers
n=
1=
= orientation, integers values between
= +½ or -½
, including zero
Electron Configuration
The way in which electrons are distributed among the various
orbitals of an atom is called its
Any orbital with the same energy is said to be
The most stable,
electron configuration of an atom is that in
which the electrons are in the
possible energy state
Effective Nuclear Charge
Electron Configuration
Chapter 12
Carbon
Hund’s Rule:
Electron Configuration
Aufbau principle
Pauli Exclusion Principle
Hund’s Rule:
Shorthand
Valence Electrons:
Core Electrons:
Short hand notation:
Chapter 12
D-Block
Periodic Table
Section 12.5
Questions
1. How does the atomic radius change down a group?
2. How does the atomic radius change across a row, left to right?
Chapter 12
Atomic radii
Ionization Energy
Ease in which an electron can be removed
Chapter 12
Summary of Trends
Chapter 12
Review Questions
Section 12.1
1. We will studying electromagnetic radiation (light) and using wave mechanics to
describe properties of it. How is wavelength measured? We categorize the type of light
based on its wavelength. Look up what these categorization are and the wavelength
range.
2. How is frequency and wavelength related to each other? Is the speed of light always
constant?
Section 12.2
We first need to characterize atoms. How do we do this?
The interaction of light and matter helps us understand atomic structure and bonding in
molecules.
For example, Rosalind Franklin, James Watson, Francis Crick, and Maurice Wilkins
used X-rays to discover the structure of DNA.
Nature of Light :
The wave properties of light. See Figure 12.7 shows the diffraction pattern that occurs
when x-rays are scattered from a regular array of objects such as the ions in a NaCl
crystal. The bright and dark spots in the diffraction pattern are due to constructive and
destructive interference of waves.
When we heat a piece of iron, as the temperature increases, we observe the color
changes from red to orange to yellow to white.
The white color indicates it is very hot (white light is also observed when current flows
through a tungsten filament in an incandescent light bulb).
The iron metal atoms oscillate at frequency, v.
The iron radiates energy in "packets" or "quanta" of energy of magnitude hv (integral
multiples of hv). That is, light is dual in nature: wave and particle.
3. Light is dual in nature: particle and wave.
a) What are the particles called?
b) What is the energy of these particles?
c) Which form of electromagnetic radiation has the longest wavelength?
4. The photoelectric effect (Einstein) is discussed. After you read this section in your
book, answer the following questions:
a) Is the following statement is true or false: An electron can be removed from a piece
of metal if the light source is sufficiently intense.
Chapter 12
b) Define all terms in the following equation: KE(electron) = 1/2 mv2 = hv - hvo
hv = ?
hvo = ?
c) Using this equation, how was Planck's constant, h, determined?
5. The energy required to remove one mole of electrons from the surface of rubidium
metal, Ru (s), is 208.4 kJ/mol. If rubidium metal is irradiated with 254-nm light, what is
the maximum kinetic energy the released electrons can have? 1nm = 1 x 10–9 m?
6. What wavelength of radiation has photons of energy 8.23x10-19 J?
Section 12.3
7. Consider the following electronic transitions in the hydrogen atom:
n = 2 -----> n = 3
n = 2 -----> n = 4
a) Is light emitted or absorbed in these transitions?
b) For these transitions is ΔE > 0 or is ΔE < 0?
c) For which transition is the wavelength of light longer?
8. Read Ch. 12 section 3. Figure 12.8 shows that when white light (sunlight) passes
through a slit and a prism, we observe a continuous spectrum of all the colors in white
light. Figure 12.8 also shows the hydrogen line spectrum observed when the following
experiment is done. A high voltage is applied to H2 gas, this forms H-atoms with their
electron in a high energy state. When the electron goes back to a lower energy state, it
emits light. When this light passes through a slit and a prism, a line spectrum is
observed.
a) What do the lines in the H-atom spectrum correspond to?
b) Why don't we see a continuous spectrum?
c) What energy states are allowed for the electron in a hydrogen atom? Are all energy
states allowed?
9. Describe the Bohr Model of the atom.
Chapter 12
10. Consider the following electronic transitions in the H-atom. For which transition is
the wavelength of light emitted shorter?a) n = 3 → n = 2; b) n = 4 → n = 2; c)
n = 2 → n = 3; d) n = 2 → n = 4
Section 12.4
11. What does the Rydberg equation model?
12. Consider the hydrogen atom with the electron in n = 6 state going down to n = 2
state. Calculate the change in energy.
Section 12.5 and 12.7
13. The Bohr model correctly predicts the lines in the hydrogen atom emission
spectrum. Why is the Bohr model fundamentally incorrect for the hydrogen atom? See
the last paragraph of section 12.4 .
14. What does the Heisenberg Uncertainty Principle state? How does Schrödinger’s
equation treat an electron differently from the previous interpretations?
16. Read section 12.5, do Example 12.5.
Sections 12.8 and 12.9
If you haven't already done so, read Section 7 and 8 in Chapter 12. When we solve the
Schrodinger wave equation for the electron in a hydrogen atom we obtain
wavefunctions which correspond to orbitals (1s, 2s, 2p, 3s, 3p, 3d, etc.). The square of
the wavefunction gives the probability distribution for the electron (it tells us where the
electron is likely to be found). By solving the Schrodinger wave equation, we also
obtain the energies of the orbitals. The wavefunctions, Ψ, along with the three quantum
numbers n, l, and ml, are shown in table 12.1.
The quantum numbers are the solutions to Schrodinger's equation.
17. What does each of the quantum numbers n, l, and ml tell you about the electron?
18. What is the total number of orbitals in a shell where n = x?
19. How many ml values would an f orbital have?
Sections 12.10, 12.11, 12.13 and 12.14
Chapter 12
For atoms with many electrons we can not solve the Schrodinger equation exactly
because of numerous electron-electron interactions that must be taken into account.
However, we can use the hydrogen atom orbitals as a templet for atoms with more than
one electron. We do observe a change in the energies of the orbitals. As we fill orbitals
with electrons, the energies of the orbitals shift. For atoms with many electrons, orbitals
with the same principle quantum number "n" do NOT have the same energy (due to
electron-electron repulsions). The general order of the orbitals from low to high energy
for multi-electron atoms are as follows:
High energy
3d __ __ __ __ __
4s __
3p __ __ __
3s__
2p __ __ __
2s__
1s__
Low energy
Rules for filling orbitals:
Aufbau Principal: Electrons are added to the lowest energy orbital first and then fill up
higher energy orbitals in order.
Hund’s Rule: Maximize number of unpaired electrons in degenerate orbitals (orbitals
that have the same energy). For example, there are three degenerate 2p orbitals. If we
have three electrons we place one electron in each orbital, each with the same spin. If
we have more than three electrons, we add an electron to each orbital, with the
opposite spin.
Pauli Exclusion Principal: No two electrons in an atom can have the same four
quantum numbers; an orbital can hold two electrons that have opposite spins.
Chapter 12
Electronic Configurations:
20. Write the electronic configuration (and the short-hand notation) for sulfur (S),
cadmium (Cd) and hafnium (Hf). Identify the core electrons and the valence electrons.
Help: See Example 12.8 text
21. Write the electronic configuration for sulfur ion, S2- and the sodium ion, Na+