Transcript Ch.4 Electron Configurations

Electrons
in Atoms
Chapter 5
General Chemistry
Objectives
•
•
Understand that matter has
properties of both particles and
waves.
Describe the electromagnetic
spectrum in terms of wavelength
and energy; identify regions of
the electromagnetic spectrum.
Objectives
• Using Bohr’s model of the atom
interpret changes
(emission/absorption) in electron
energies in the hydrogen atom
corresponding to emission
transitions between quantum
levels.
• Write quantum numbers for
electrons in atoms of elements
• Write the electron configurations
for elements.
Review/Link to Previous
Learning
• Atoms are the smallest part of
an element that contains the
properties of that element
• Atoms consist of protons,
neutrons, and electrons
• Quantum Mechanics is the
currently accepted model of the
atom
Properties of Light
Electromagnetic
Radiation
Electromagnetic
Spectrum
• Electromagnetic (EM) radiation
is a form of energy that exhibits
wavelike behavior as it travels
through space.
Electromagnetic
Spectrum
• EM Spectrum: full range of
frequencies of EM radiation
• Listed in order of increasing
frequencies (decreasing
wavelengths)
EM Spectrum
& frequency
Microwaves
Parts of EM Spectrum
• In order of increasing
frequency:
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Radio waves
Microwaves and Radar
Infrared rays
Visible light
Ultraviolet rays
X rays
Gamma rays
Transverse Waves
• Visible light exhibits properties of
wavelike motion.
• One type of wave is a transverse wave.
• Transverse waves are like taking a
rope and moving it up and down.
•A transverse wave is a wave that
causes the medium to vibrate at right
angles to the direction that the wave
travels.
•The particle in the medium travels
perpendicular to the motion of the
wave.
Trough
Parts of Transverse
Wave
• Crest: highest point of the wave
above the resting position
• Trough: lowest point of the
wave below the resting position
Properties of Transverse
Waves
• common properties of
Transverse Waves:
• Frequency (ν )
• Wavelength (λ)
• Speed (c)
Frequency
• Frequency (ν): the
number of
complete cycles of
a wave passing a
given point in a
certain amount of
time
• Measured in
cycles per second
or Hertz (Hz)
Wavelength
• Wavelength (λ)
is the distance
between two
adjacent crests
of a wave.
• Measured in
meters (m)
Calculate Wave Speed
(Velocity)
• Speed of a any wave can be
calculated by:
v=λ xν
• v = speed (velocity) of wave (m/s)
• λ = wavelength (m)
• ν = frequency (Hz or cycles/s)
Constant Speed of EM
Waves
• All forms of EM radiation travel
at a constant speed in a
vacuum
• c = 3.00 x 10 8 m/sec (speed of
light)
EM Radiation
Speed Equation
c=ν xλ
c = 3.00 x 10
of light)
8
m/sec (speed
ν = frequency (Hz, 1 / sec,
sec-1)
λ = wavelength (meters)
Sample Problem
• Red light has a wavelength of
700. nm. What is the frequency
of red light?
Sample Problem
• Red light has a wavelength of
700. nm. What is the frequency
of red light?
• Answer: 4.29 x 1014 Hz
Photoelectric Effect
Photoelectric Effect
• Solar panel on calculator produced
electrical current when light shines
on it
• Photoelectric effect: emission of
electrons from a metal caused by
light striking the metal
• Experiments show only certain
colors of light (only certain amount
of energy) will allow electrical
current to flow
Photoelectric Effect
Quanta of Energy
• A quantum is the minimum
amount of energy that can be
lost or gained by an atom
• Max Planck proposed
relationship between quantum
of energy and frequency of
radiation (and color of light)
Energy in Quantum of
Energy
E=hxν
• E = Energy (Joules)
• h = Planck’s Const. (6.63 x 10
Joules x sec)
• ν = frequency (cycles/sec or
Hertz or sec -1)
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Sample Problem
• Calculate the energy of a
photon with a frequency of
2.85 x 1012 Hz.
Sample Problem
• Calculate the energy of a
photon with a frequency of
2.85 x 1012 Hz.
Answer: 1.89 x 10-21 J
Study Buddy Review
• What is the Photoelectric
Effect?
• How is a energy related to
the color of light?
Energy and EM
Radiation in Atoms
Spectra
• Observations of properties of
light emitted by an atom after it
absorbs extra energy
• Continuous spectrum
• Atomic emission spectrum
Line Spectra
• 1) For white light (sun or incandescent
light bulb), you see continuous
spectrum- rainbow of all colors
• Contains all colors of visible light
• Example: prism separates white light into
rainbow colors.
• 2) Atomic Emission Spectrum (Lineemission spectrum)
• contains only certain colors or wavelengths,
mostly black
• unique for each element
• “fingerprint”
• Example: pink glow from hydrogen gas tube
that can be separated into series of lines
Recall Bohr’s Model of
the Atom
• Thought atom was
mostly empty
space
• Nucleus in center
is dense, positively
charge
• Electrons move in
orbits around the
nucleus
http://images.search.yahoo.com/search/imag
es/
Hydrogen Atom
• Bohr’s model arose from
hydrogen atomic emission
spectra.
• If electrons move in specific
orbits, they have a certain fixed
amount of energy
• Observed as discrete bands of
color
Why Bands of Color?
• Electrons can only absorb a
certain amount of energy to
enter the excited state (next
energy level)
• Like rungs of ladder
• Can only move to specific orbit,
which has a specific energy
Parts of Bohr model
• Ground state: lowest energy
state of an atom
• Excited state: state in which an
atom has a higher potential
energy than it has in its ground
state
How is Radiation
Produced from Atom?
• When an excited electron falls back
from an excited state to its ground
state, it emits a photon of radiation.
(E= hν)
• Since only specific frequencies of
light are emitted from elements, the
energy levels in atom are fixed.
• Electrons exist in only certain
energy states.
Quantum Mechanics
Only for Hydrogen
• Bohr’s model of
atom was
successful only for
elements with one
electron
Heisenberg
• Continued building on Bohr model
• Uncertainty Principle – it is
impossible to know both the
exact position and the velocity of
an electron or other particle
simultaneously
• Act of measuring changes what you
are trying to find
• Thus, no well-defined orbit (like Bohr
had proposed)
• Best we can do is represent
PROBABILITY of finding e- within
a given space
Schrodinger
• Erwin Schrodinger developed a
mathematical equation that
allowed for wavelike behavior of
electrons
• Energy of electrons is quantized
Quantum Mechanical
Model of Atom
• atom is mostly empty space
• Nucleus in center is dense,
positively charge
• Electrons are around the nucleus
• e- do not have a precise orbit
(electron cloud)
• e- moves in wavelike motion with
quantized amounts of energy
Quantum Numbers
• Quantum Numbers: specify the
properties of atomic orbitals
and properties of electrons in
orbitals.
Principle Quantum
Number (n)
• n: Principal energy level or
energy number
• n = 1, 2, 3…(as you get farther
from nucleus)
• maximum # of e- in energy level
= 2n2
Angular Momentum (l)
• l: Shape of orbitals (sublevels)
• named s, p, d, f
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s = spherical
p = dumbbells
d = clover leaf
f = butterflies??
• each sublevel has slightly
different energy
Magnetic (ml)
• ml = Orientation of orbital
around nucleus
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s = 1 orbital
p = 3 orbitals
d = 5 orbitals
f = 7 orbitals
Spin ms
• ms = Magnetic spin
• Only two directions of spin to
create magnetic field
• Like N and S of magnet
Electron Configurations
• Electron configuration: Notation
describing the most stable
arrangement of e- around the
nucleus
• Like writing the address for
electrons around nucleus of
atom
Three Rules for Electron
Configurations
1. The Aufbau Principle e- added one
at a time to the lowest energy levels
2. The Pauli Exclusion Principle
orbital can hold at most 2 emust have opposite spin
3. Hund’s Rule (The “bus seat” rule)
e- like to be unpaired if possible &eenter equal energy orbitals until all
orbitals contain one e- with parallel
spin, then they begin to pair up
Electron Configuration
• Write the electron configuration
for the following:
• Oxygen
• Sodium
• copper