Atomic Structure Part II Electrons in Atoms · Atomic Structure Part II Electrons in Atoms Radiant...
Transcript of Atomic Structure Part II Electrons in Atoms · Atomic Structure Part II Electrons in Atoms Radiant...
Atomic Structure Part II
Electrons in Atoms
Radiant energy travels in
the form of waves that
have both electrical and
magnetic properties.
These electromagnetic
waves can travel through
empty space, as you know
from the fact that radiant
energy from the sun
travels to Earth every day.
Electrons in Atoms
I. Properties of Waves
1. Definition:
Energythat exhibits wave-like(or
oscillating) behavior as it travels
through space
Electrons in Atoms
2. Wavelength (λ λλλ) distance from peak
to peak, leng
th of
one com
plete
wave
3. Frequency (ν ννν)
a. number of peaks that pass at a given
point each sec
b. can be called cycles per
second (peak/sec)
c. cps now called 1 Hertz (Hz)
Electrons in Atoms-Cont.
4. Velocity (C = speed of light)
a. distance a given peak moves in
a unit of time
b. velocity (m/s) = frequency x
wavelength
c = ν ννν
x λ λλλ
II. Behavior of Light
A.Newton (1600) thought light consisted of particles
(beam of light is a stream of particles)
B. Maxwell (1864) thought light was a wave
phenomenon.
•Calcu
late
d the velocity of th
epr
opag
ation
of anelectro
magn
etic wave
and fou
nd it
was the sam
e for
light
II. Behavior of Light
1. some say light is like waves, some say
its like particles
2. modern theory says that it behaves as
both "wave/particle duality"
II. Behavior of Light
3. Max Planck (early 1900's) said:
a. light is made up of bundles of energy called
photons (or quanta)
b. the energy of each photon is proportional to
the frequency of the light
(Quantum Theory)
Quantum is the minimum amount of
energy that can be gained or lost by
an atom.
•example: CONTINUOUS SPECTRUM
*** when white light is passed through
a prism, it is separated into a band of
colors from red �violet. It's called a
continuous spectrum
c. the work of Planck & Einstein led to
E=energy,
ν= frequency,
h=planks constants (6.6262x10-34J /sec) J is
the symbol for joule the SI unit for energy
E = h x ν ννν
Energy of a quantum is related to the
frequency of the emitted radiation by
this equation
c. the work of Planck & Einstein led to
According to Planck’s theory, for a given
frequency matter can emit or absorb energy
only in whole-number mulitplesof hv, that is
1hv, 2hv, 3hv, and so on.
E = h x ν ννν
c. The photoelectric effect
In the photoelectric effect, electrons, called
photoelectrons, are emitted from a metal’s
surface when light of a certain frequency shines
on the surface. (exam
ple: solar calculator)
c. The photoelectric effect
Einstein said light can both wavelike and
particle like natures. That is, while a beam of
light has many wavelike characteristics, it
also can be thought of as a stream of tiny
particles, or bundles or energy, called
photons.
APhotonis a particle of electromagnetic
radiation with no mass that carries a
quantum of energy.
III. Bright line spectrum
•A. a spectrum that shows separate bright
lines, each with a specific wavelength
•B. bright-line spectra occur when an
element is heated and the colored light
given off is viewed through a
spectroscope. Each element has a unique
set of lines, characteristic of that element
(like a fingerprint)
Line-Emission Spectrum
ground state
excited state
ENERGY IN
PHOTON OUT
Fireworks? Hmmm…
IV. Electromagnetic Spectrum
•A. visible light (like the continuous spectrum)
is only one type of radiation. All other types
are not visible to the human eye.
H I G H E N E R G Y
L O W E N E R G Y
Electromagnetic Spectrum
L O W E N E R G Y
H I G H E N E R G Y
RO
YG.
BI
V
red
orange
yellow
green
blue
indigo
violet
Electromagnetic Spectrum
B. all forms of electromagnetic radiation
travels at the speed of light.
1. speed of light = 3.00 x 108meters/sec
2. use form
ula:
c = ν νννx λ λλλ
3. each line spectrum has a particular
frequency (ν ). If know wavelength (λ), we
can find ν
using c
as a constant.
C. The energy in a photon of light is directly
proportional to the frequency of the light.
•1.
frequency,
energy
•2. can find the energy of a single
quantum (photon) of radiation at
any given frequency.
C. The energy in a photon of light is directly
proportional to the frequency of the light.
•3. proportionality constant that
relates the two is called Planck's
constant (h).
•4. form
ula:
E = hx f
example: a spectral line has frequency of
3.5x1012hertz. What is the energyof a
photon of radiation of this frequency?
E = hx f
(h=6.6262x10-34J/sec)
E = (3.5x1012Hz) (6.6262x10-34J •sec)
E = (2.3x10-21J)
V. Electron energy levels in
Bohr's Model
A. There are certain different
orbits in which an electron can
travel around a nucleus.
1. each circular orbit (or shell) is at a
fixed distance from the nucleus
V. Electron energy levels in
Bohr's Model
2. the greater the radius of that shell,
the greater the energy of the
electron in that shell.
3. these electron orbits are known as
energy levels
B. When electrons absorb energy firm an
outside source, they jump from lower to
higher energy levels.
�when they fall back to their original levels ,
energy is
emitted (light); the same amount as was
absorbed.
1
23
456
�Energy of photon
depends on the
difference in energy
levels
�Bohr’s calculated
energies matched
the IR, visible, and
UV lines for the H
atom
C. In energy atom in its norm
al
state, all electrons are in the
lowest energy levels available
(energetically stable)
VI. Atoms and
Radiation
•A. When all of
the lowest
energy levels
are occupied,
the atom is in
the ground
state
(unexcited).
VI. Atoms and
Radiation
•B. When
electron moves
to higher
energy level,
atom is in the
excited state,
and is
energetically
unstable.
VI. Atoms and
Radiation
C. Bright line
spectrum of an
element
represents the
energy levels in
its atoms.
�problems with Bohr's Model:
� ���only explained some of the lines in the
bright line spectrum
� ���really only worked for hydrogen
� ���need sublevels and electron cloud
model to account for all of the lines.
VII. The Modern Model of the Atom
A. Mechanics
1. Classical Mechanics -Newton's Laws of
Motion (Newtonian Mechanics)
Describes the behavior of visible objects
traveling at ordinary velocities.Bohr’s basis
for his model, but couldn’t explain why
electrons would stay at on energy level or
another. When looking at H-spectral lines,
noticed more one (several closely spaced).
VII. The Modern Model of the Atom
2. Quantum Mechanics –(wave
mechanics)
Describes the behavior of extrem
ely
small particles traveling at velocities at
or near the speed of light
a. Louis de Broglie -particles could
have properties of waves
Planks quanta gave wave properties,
deBrogliesaid electron streams are like
waves of light and have properties of
both particles and waves (matter behaves
as waves)
•b. Schrodinger -described the
behavior of electrons in terms of
quantized energy changes
"quantum mechanics"
Describe a wave equation used to
determine the probability of finding an
electron in any give place or orbital
Schrodinger’sCat
Radial Distribution Curve
Orbital
c. Heisenberg -uncertainty principle
�Region of space where there is a
probability of finding an electron is
called an orbital
"The more
precisely
the POSITIO
Nis
determined,
the less precisely
the
MOMENTUM is
known"
B. Principal Energy Levels
1. Energy Levels
•Bohr -
High Energy
(outer level)
Low Energy
1 2 3 4
Principal Quantum Numbers (N)Number of electrons
2
8 18
32
Corresponds to
energy level
2. Sublevels
Principal Quantum Numbers (N) Sublevel
Present
11s
22s2p
33s3p3d
44s4p4d4f
•Orbital -
–Region of space where an
electron is probably found
•Electron spin
–An orbital can hold 2 electrons that spin in
opposite directions.
Electrons are
represented by arrows
Rules: 1. PauliExclusion Principle
–Each orbital can hold TWO
electrons with opposite
spins.
–No two electrons in an atom
can have the same 4
quantum numbers.
–Each e
-has a unique
“address”:
2. AufbauPrinciple
Electrons fill the
lowest energy
orbitalsfirst.
Electrons to be
added must be
placed in unfilled
orbitalsof lowest
energy for stable
configuration.
3. Hund’sRule
–Within a sublevel, place one e-
per orbital before pairing them.
–“Empty Bus Seat Rule” RIGHT
WRONG
Energy Level Diagram
orbital
Orbital -Place where electrons are probablyfound
Electrons have “up”and
“down”spin
c
Shapes of electron orbitals
The s orbital
The p orbitals
py
px
The p orbitals
pz
The d
The d orbitals
orbitals
The d
The d orbitals
orbitals
The f
The f orbitals
orbitals
f
Click here for orbital viewer
View the grand table
orbital views
C. Electron Configurations
1s2= Helium
1s22s1= Lithium
1s22s22p63s23p6
4s23d104p6
= Krypton
Energy Level
Sub Level # of Electrons
O 8e-
•Orbital Diagram
•Electron Configuration
1s2 2s22p4
Electron Configuration Notation
1s
2s
2p
•Shorthand Configuration
S 16e-
Valenc
e E
lectro
nsCor
e E
lectro
ns
S16e-[Ne]3s2 3p4
1s22s22p63s23p4
Notation
•Longhand Configuration
©1998 by Harcourt Brace & Company
sp
d (n-1)
f (n-2)
1 2 3 4 5 6 7
6 7
Periodic Patterns
Periodic Patterns
•Period #
–energy level (subtract for d & f)
•A/B Group #
–total # of valence e
-
•Column within sublevel block
–# of e-in sublevel
s-block
1st Period
1s1
1st colum
n of
s-block
1
2
3
4
5
6
7
Periodic Patterns
•Example -Hydrogen
1
2
3
4
5
6
7
Periodic Patterns
•Shorthand Configuration
–Core e- :Go up one row and over to the
Noble Gas.
–Valence e- :On the next row, fill in the #
of e-in each sublevel.
[Ar]1
2
3
4
5
6
7
4s23d104p2
Periodic Patterns
•Example -Germ
anium
•Full energy level
1
2
3
4
5
6
7
•Full sublevel (s, p, d, f)
•Half-full sublevel
Stability
1
2
3
4
5
6
7
Stability
•Ion Formation
–Atoms gain or lose electrons to become more
stable.
–Isoelectronicwith the Noble Gases.
1
2
3
4
5
6
7
Stability
•Ion Formation
–Atoms gain or lose electrons to become more
stable.
–Isoelectronicwith the Noble Gases.
Feeling overwhelmed?
Try a few!
Mg =
Fe =
Ru=
Ir=
Ca+2 =
Cl-1 =