Chapter 4 reg Electron Configurations TUESDAY 11315 Learning Target Explain the electromagnetic spectrum Learning Outcome Be able to d escribe a wave in terms of frequency wavelength speed and ID: 730795
Download Presentation The PPT/PDF document "Arrangement of the Electrons" is the property of its rightful owner. Permission is granted to download and print the materials on this web site for personal, non-commercial use only, and to display it on your personal computer provided you do not modify the materials and that you retain all copyright notices contained in the materials. By downloading content from our website, you accept the terms of this agreement.
Slide1
Arrangement of the Electrons Chapter 4 (reg.)
(Electron Configurations)Slide2
TUESDAY 11/3/15
Learning Target:
Explain the electromagnetic spectrum.
Learning Outcome:
Be able to
d
escribe
a wave in terms of frequency, wavelength, speed, and
amplitude.
Slide3
What holds the atom together?
What
is energy
?
Is
energy matter
?
Do you think it’s possible for light energy to cause changes in atoms?Slide4
Spectrum of Light!
Electromagnetic Radiation
-form of
energy
that exhibits wave-like behavior as it travels through space.
Electromagnetic Spectrum
-ordered arrangement by wavelength or frequency for all forms of electromagnetic radiation.Slide5
Parts of the wave
Wavelength
-lambda (
λ
)
The distance between corresponding points on adjacent waves.
Units
: m, nm, cm, or Å
Frequency
-nu (
ν) The number of waves passing a given point in a definite amount of time. Units: hertz (Hz) or cycles/s = 1/s = s-1Slide6
Relationship between λ
and
ν
c =
λ∙ν
λ
= wavelength (m)
ν
= frequency (Hz)
c = speed of light=
3.0 x 108 m/sec (constant)λ and ν
are _______________ related.Slide7
B. EM Spectrum
EX
: Find the frequency of a photon with a wavelength of 434 n
m.Slide8
Practice Problem
Truck-mounted helium-neon laser produces
red
light
whose frequency (
v
) is 4.7 x 10
14 Hz. Determine the wavelength
.
*
Remember that c=3.0x108m/s. *Use the formula c= λ . v
Slide9
c= λ .
v
c =
3.0x10
8
m/s
c=
λ
.
v v=c / λ
λ = 633nm
=
6.33x10
-7
m
v =
3.0x10
8
m/s
= 0.47x 10
15
s
-1
= 4.7x10
14
s
-1
6.33x10
-7
m
Frequency = 4.7x10
14
Hz (cycles
per second)Slide10
WEDNESDAY 11/4/15
Learning Target:
Know how to calculate energy of light using
the formula E =
hv
.
Learning Outcome: I will complete the Energy Calculation worksheet.Slide11
Calculate the frequency for the yellow-orange light of sodium.
Calculate the frequency for violet light.Slide12
Relationship between
Energy and
ν
E = h∙
ν
E = energy (joule)
h = Planck’s constant
=
6.63 x 10
-34
j∙secν = frequency (Hz)E and ν
are ______________ related.
Calculate the energy for the yellow-orange light for sodium.Slide13
Calculate the energy for the violet light.Slide14
When an electric field changes, so does the magnetic field. The changing magnetic field causes the electric field to change. When one field vibrates—so does the other.
RESULT-An
electromagnetic wave.Slide15
Waves or Particles
Electromagnetic radiation has properties of waves but also can be thought of as a
stream of particles.
Example: Light
Light as a wave: Light behaves as a transverse wave which we can filter using polarized lenses.
Light as particles (photons)
When directed at a substance light can knock electrons off of a substance (Photoelectric effect)Slide16
Warm up
1. What
is the frequency of a photon of light whose wavelength is 344 nm?
2. A
red light has a wavelength of 676 nm. What is the energy of the light?Slide17
Thursday 11/5/15
Learning Target:
Know the roles Max Planck and Einstein played in the development of the photoelectric effect.
Learning Outcome:
Be able to
describe the photoelectric effect.
Slide18
Light as waves and particles
(the Particle Theory of light)
2 problems that could not be explained if light only acted as a wave.
1.)
Emission of Light by Hot bodies
:
Characteristic color given off as bodies are heated:
red
yellow
white
If light were a wave, energy would be given off continually in the infrared (IR) region of the spectrum.Slide19
The second problem………
2.)
Absorption of Light by Matter =
Photoelectric Effect
Light can only cause electrons to be ejected from a metallic surface if that light is at least a minimum
threshold frequency
. The
intensity
is
not important. If light were only a wave intensity would be the determining factor, not the frequency!Slide20
Max Planck (1900’s)
Particle Theory of Light
When an object loses energy, it doesn’t happen continuously but in small packages called “
quanta
”
.
“
Quantum
”
-a definite amount of energy either lost or gained by an atom.
“
Photon
”
-a
quantum of light
or a particle of radiation.Slide21
Monday 11/9/15
Learning Target:
Know the Bohr Model of the hydrogen atom.
Learning Outcome:
Be able to
describe the Bohr Model and how it helped describe the behavior of electrons.
Slide22
Line Spectrums
Excited State
: Higher energy state than the atom normally exists in.
Ground State
: Lowest energy state “happy state”
Line Spectrum
: Discrete wavelengths of light emitted.
2 Types
:
1.)
Emission Spectrum: All wavelengths of light emitted by an atom.
2.)
Absorption Spectrum
: All wavelengths of light that are
not
absorbed by an atom.
This is a continuous spectrum with wavelengths removed that are absorbed by the atom. These are shown as
black lines
for absorbed light.
Continuous Spectrum
: All wavelengths of a region of the spectrum are represented (i.e. visible light)Slide23
Hydrogen line Spectrum & niel’s
Bohr
Hydrogen’s spectrum can be explained with the wave-particle theory of light.
Niel’s
Bohr
(1913)
1.) The electron travels in orbits (energy levels) around the nucleus.
2.) The orbits closest to the nucleus are lowest in energy, those further out are higher in energy.
3.) When energy is absorbed by the atom, the electron moves into a higher energy orbit. This energy is released when the electron falls back to a lower energy orbit. A
photon
of light is emitted.Slide24
Hydrogen Spectrum
Lyman Series
-electrons falling to the 1
st
orbit, these are highest energy, _____ region.
Balmer
Series
- electrons falling to the 2
nd
orbit, intermediate energy, _______ region.
Paschen Series-electrons falling to the 3rd orbit, smallest energy, ______ region.Slide25
Bohr’s equation for Hydrogen
E
n
= (-R
H
) 1/n
2
E
n
= energy of an electron in an allowed orbit (n=1, n=2, n=3, etc.)
n = principal quantum number (1-7)RH = Rydberg constant (2.18 x 10-18 J)When an electron jumps between energy levels: ΔE =Ef – EiBy substitution: ΔE = h
ν
= R
H
(1/n
i
2
- 1/n
f
2
)
When n
f
> n
i
then
Δ
E = (+)
When n
f
< n
i
then
Δ
E = (-)Slide26
THURSDAY 11/13/15
Learning Target:
Know what quantum numbers represent in the quantum mechanical model of the atom.
Learning Outcome:
Determine the electron configuration and orbital diagram of atoms.
Slide27
New Theory Needed to explain more complex atoms!
DeBroglie (1924)
-Wave properties of the electron was observed from the diffraction pattern created by a stream of electrons.
Schrodinger (1926)
-Developed an equation that correctly accounts for the wave property of the electron and all spectra of atoms. (very complex)Slide28
Quantum Theory
(current theory of the atom)
Rather than
orbits
we refer to
orbitals
. These are 3-dimensional regions of space where there is a high probability of locating the electron.
Heisenberg Uncertainty Principle
-it is not possible to know the exact
location and momentum (speed) of an electron at the same time.Quantum Numbers-4 numbers that are used to identify the highest probability location for the electron.Slide29
Quantum numbers (reg. chem.)
1.)
Principal Quantum Number
(n)
States the main energy level of the electron and also identifies the number of sublevels that are possible.
n=1, n=2, n=3, etc. to n=7
2.)
Orbital Quantum Number
Identifies the shape of the orbital
s
(2 electrons) sphere 1 orbitalP (6 electrons) dumbbell 3 orbitalsd
(
10 electrons) 4-4 leaf clovers & 1-dumbbell w/doughnut
5 orbitals
f
(14 electrons)
very complex 7 orbitalsSlide30
Quantum numbers (cont.)
3.)
Magnetic Quantum Number
Identifies the
orientation
in space (x, y, z)
s 1 orientation
p 3 orientations
d 5 orientations
f 7 orientations
4.) Spin Quantum Number
States the spin of the electron.
Each orbital can hold at most 2 electrons with opposite spin.Slide31
Quantum numbers
1.)
Principal Quantum Number
(n)
States the main energy level of the electron and also identifies the number of sublevels that are possible.
n=1, n=2, n=3, etc. to n=7
2.)
Azimuthal Quantum Number
(l)
Values from 0 to n-1
Identifies the shape of the orbitall = 0 s sphere 1 orbitall = 1 p dumbbell 3 orbitalsl = 2 d 4-4 leaf clovers & 1-dumbbell w/doughnut5 orbitals
l = 3 f very complex 7 orbitalsSlide32
Quantum numbers (cont.)
3.)
Magnetic Quantum Number
(m
l
)
Values from –l
l
States the
orientation
in space (x, y, z)ml = 0 s only 1 orientationm
l
= -1, 0, +1 p 3 orientations
m
l
= -2,-1,0,+1,+2 d 5 orientations
m
l
= -3,-2,-1,0,+1+2,+3 f 7 orientations
4.)
Spin Quantum Number
(m
s
)
Values of +1/2 to -1/2
States the spin of the electron.
Each orbital can hold at most 2 electrons with opposite spin.