Download Light and Electron Notes

LIGHT and ELECTRONS
Unit 6
Chemistry
Langley
LIGHT and its PROPERTIES
Pre-1900
Issac Newton explained light and its behavior
by assuming light moved in waves
1900 and beyond
Experimental evidence began to convince
scientists that that light consists of particles
(after the 1902 experiment of Max Planck)
1905-Einstein
Dual Wave Particle Theory
LIGHT and its PROPERTIES
 Wavelength: distance between two points on
two adjacent waves, symbol is l (Greek symbol
for lamda)
 Frequency: number of waves that pass a given
point in a given amount of time, symbol is n
(Greek symbol nu). Units for frequency are
cycles per second which SI speaking is a Hertz,
Hz (Hz is also a reciprocal seconds-1).
LIGHT and its PROPERTIES
 The frequency and wavelength of light are inversely
proportional to each other.
As the wavelength of light increases, the frequency
decreases
As the wavelength of light decreases, the frequency
increases
 Amplitude: Wave’s height from zero to crest or wave’s
height from zero to trough (can be positive or negative)
 A complete wave cycle starts at zero goes through its
highest point, back through zero, reaches its lowest
point, and back to zero again.
One wave cycle starts at zero and has one crest
and one trough
LIGHT and its PROPERTIES
According to the Wave Model, light
consists of electromagnetic waves
Electromagnetic radiation: light moving in
waves through space
Radio waves, microwaves, infrared waves, visible
light, ultraviolet waves, X-rays, and gamma
raysElectromagnetic spectrum
Speed of light: depending on the wavlength
and frequency, speed of light changes
C = ln
Speed of light in a vacuum = 3.0 x 108 m/s
SPEED of LIGHT PROBLEMS
EXAMPLE 1:
Determine the speed of light if the wavelength
is 3.5 x 10-9 m/s and the frequency is 3.5 Hz.
SPEED of LIGHT PROBLEMS
EXAMPLE 2:
If light has a speed of 5.6 x 103 m/s and a
frequency of 2.3 Hz, what is the wavelength.
SPEED of LIGHT PROBLEMS
EXAMPLE 3:
What is the wavelength of radiation with a
frequency of 1.5 x 1013 Hz? Does this radiation
have a longer or shorter wavelength than red
light?
SPEED of LIGHT PROBLEMS
EXAMPLE 4:
What frequency is radiation with a wavelength
of 5.00 x 10-8 m? In what region of the
electromagnetic spectrum is this radiation?
PHOTOELECTRIC EFFECT
(supporting work for Atomic Spectra)
 The photoelectric effect is a quantum
electronic phenomenon in which electrons are
emitted from matter after the absorption of
energy from electromagnetic radiation such as
x-rays or visible light. The emitted electrons can
be referred to as photoelectrons in this context.
{Wikipedia.org}
PHOTOELECTRIC EFFECT (supporting
work for Atomic Spectra)
Expected: Since all light is energy moving
in waves, all colors of light should knock
electrons off a metal
Shine different color lights on a metal
Measure the number of electrons knocked off
the metal
Found that no electrons were knocked off when
light was below a certain frequency
MAX PLANCK
(his work used in Atomic Spectra)
 German Physicists, founder of quantum theory
 Studied the way light came off hot objects
(diffusion of hydrogen through heated
platinum)
 Concluded that light comes off in small burst of
particles, NOT WAVES
 Quantum-minimum amount of energy that can
be lost or gained by an atom
 To calculate quantum/energy: E = hn
E = energy of the photon
h = Planck’s constant
n = frequency of the incident photon
ATOMIC SPECTRA
As atoms absorb energy, electrons move
into higher energy levels. When the
atoms release energy (lose the energy),
the electron return to the lower energy
levels.
The frequencies of light emitted by an
element separate to give the atomic
emission spectrum of the element
No two elements have the same emission
spectrum
ATOMIC SPECTRA
Atomic line spectra and its existence was
known before Bohr’s atomic model of
hydrogen was produced. What Bohr did
was explain why hydrogen had the specific
frequencies it had, why it “produced/broke
down” into the colors it did; it predicted the
values that agreed with the experiements.
ATOMIC SPECTRA
Hydrogen Atom Line Emission Spectrum
EXPECTED:
Continuous spectrum of
light to be given off. (Since
e- are moving around
nucleus randomly and using
different levels of energy.)
ACTUAL:
Current passed through tube with
Hydrogen gas.
Pink light is given off.
Light passed through spectrum.
Found only specific frequencies
of light given off.
ATOMIC SPECTRA
Lowest possible energy of the electron is
referred to as its ground state
Normal location of an electron
Electrons circle the nucleus in specific
orbits
If an electron absorbs energy, moves up
an energy level (absorption)
If an electron gives off energy, moves
down an energy level (emission)
QUANTUM MECHANICS
EINSTEIN, AGAIN!!!!!!!!!!!!!!!!
Debate between whether light is waves or
particles
Einstein creates dual waves particle theory
(1905)
Light is small particles (photons) that move in
wave shapes
Thought electrons moved around the nucleus
in wave shapes (since electrson are small
particles like photons)
QUANTUM MECHANICS
 Louis de Broglie: Given that light behaves as waves and
particles, can particles of matter behave as waves?
 Referred to the wavelike behavior of particles as matter
waves
 Came up with an equation that predicts all moving objects
have wavelike behavior:
 mv/l = h
 Thanks to experiments conducted by 2 scientists, his theory
was proven correctNobel Prize
 Waves Waves have specific frequencies and electrons have
specific orbits/energy levels
 Waves and electrons can both be bent (diffraction)
 Waves and electrons can both overlap and interfere with
each other (interference)
 Creator of Wave Mechanics
QUANTUM MECHANICS
 DeBroglie’s equation combines Einstein
and Planck’s equations
 mv/l = h
 (Anything with mass and velocity has a
wavelength, so electrons have wavelengths)
 DeBroglie Problems:
 What is the wavelength of an electron that has a
mass of 1.5 X 10-30 kg and a velocity of 2.5 X 104
m/s?
QUANTUM MECHANICS
 DeBroglie Problems:
 What is the velocity of an electron with a mass of
8.3 X 10-29 kg and a wavelength of 400 nm? (Hint:
convert nm to m)
 What is the mass of an electron with a velocity of
4.6 X 103 m/s and a wavelength of 5.6 X 10-2
meters?
 What is the wavelength of an electron that has a
mass of 2.8 X 10-31 kg and a velocity of 3.0 X 108
m/s?
QUANTUM MECHANICS
 Heisenberg
2 Goals in Life:
 find the location of an electron
 find the velocity of an electron
Problem: Electrons cannot be seen under a microscope
Only way to find an electron is to shoot a photon
(particle of light) at the electron
Problem: when the photon hits the electron, it knocks
the electron off course
So with this photon method, you can only know the
position of an electron for a split second, but you still
don’t know the velocity
QUANTUM MECHANICS
Heisenberg
DeBroglie: Tries to help Heisenberg and offers
his equation l = (mv)/h
If you know mass and wavelength of an
electron, equation could help you find velocity
Problem: Equation does not show location!
Equation method will only tell you velocity NOT
location
QUANTUM MECHANICS
Heisenberg
Heisenberg Uncertainty Principle: It is
impossible to know both the position and
velocity of an electron at the same time.
QUANTUM MECHANICS
 Schrodinger
Working with Hydrogen atom that only has 1 electron
Wants to find general location/area of the one electron in
Hydrogen
Creates quantum theory
Quantum theory – uses math to describe the wave
properties of an electron (frequency, wavelength, etc)
Once he plugged his data into the quantum theory, he
found that electrons do not travel in nice, neat orbits
(Bohr model)
Instead, found that electrons travel in 3D regions around
the nucleus
QUANTUM MECHANICS
Schrodinger
Schrodinger’s equation is used to find the
greatest probable location/area of the Hydrogen
atom electron (in the ground state)
QUANTUM MECHANICS
 Quantum Theory
Ground State-normal location of an electron
Excited State-one ring up from the normal location
When excited electron falls back to the ground state, a
photon is given off
Energy of the photn is equal to the difference in energy
between the excited state and ground state
Hydrogen gives off specific colors because its electrons
move from ring 2 to ring 1; Neon gives off a different
color because its electrons move from ring 3 to ring 2
LIGHT AND ELECTRONS REVIEW
 Light was first thought to be wavelike
 Equation for the speed of light is c = ln
 Photoelectric effect challenges this because only certain
frequencies of light could knock off electrons
 Max Planck’s experiment proved that light could be a
particle
 Einstein’s dual wave particle theory says that light is
ACTUALLY small particles (photons) that move in wave
like patterns
 Equation for energy of a photon is E = hn
 Bohr found that electrons orbit the nucleus in specific
orbitals/energy levels
LIGHT AND ELECTRONS REVIEW
 Electrons as Waves:
1924 – Louis de Broglie asked “Could electrons have a
dual wave particle nature like light?”
 Similarities between waves and electrons
 Waves have specific frequencies and electrons have
specific orbits/energy levels
 Waves and electrons can both be bent (diffraction)
 Waves and electrons can both overlap and interfere with
each other (interference)
 DeBroglie’s equation combines Einstein and Planck’s
equations
 mv/l = h
 (Anything with mass and velocity has a wavelength, so
electrons have wavelengths)
QUANTUM NUMBERS and ATOMIC ORBITALS
 REVIEW
Energy levelsSpecific energies electrons can have
Quantum of energyamount of energy required to
move an electron from one energy level to another
energy level
The amount of energy an electron gains or loses in an
atom is not always the same
Energy levels in an atom are not equally spaced
Higher energy levels are closer together
Modern description of the electrons in atoms, quantum
mechanical model, comes from the mathematical
solutions to the Schrodinger equation
The quantum mechanical model determines the allowed
energies an electron can have and how likely it is to find
the electron in various locations
QUANTUM NUMBERS and ATOMIC ORBITALS
QUANTUM NUMBERS
Quantum numbers are used to describe the
location and behavior of an electron (zip code
for electrons)
First quantum number = Principal = n
Second quantum number = Angular Momentum
Third quantum number = Magnetic Quantum
Fourth quantum number = Spin Quantum
QUANTUM NUMBERS and ATOMIC ORBITALS
 Principal (first) quantum number = n
 Main quantum number
 Describes the size of the electron cloud (the smaller the number, the
smaller the cloud)
 ALSO, shows the distance from the nucleus, the smaller the
number, the closer the cloud is to the nucleus
 Called energy levels or shells
 Positive integers (1,2,3,4,…)
 Symbol is n
 Each energy level has a maximum number of electrons it can hold
n
1
2
3
4
# Electrons
2
8
18
32
 Example: Energy level 1
2 electrons
close to the nucleus
small electron cloud
QUANTUM NUMBERS and ATOMIC ORBITALS
Second Quantum Number:
Each energy level has sublevels
The number of sublevels is equal to n
Example: Energy level 1 has 1 sublevel
Sublevels are called: s,p,d,f
QUANTUM NUMBERS and ATOMIC ORBITALS
 Third Quantum Number
 Divides sublevels into orbitals
 Also tells the shape the electron is moving in
 The number of orbitals for each level is:
 S has 1
 P has 3
 D has 5
 F has 7
 The number of orbitals for an energy level is equal to n2
 Example: 2nd Energy level
 n2 = 4
 1s, 3p
 Each orbital can only hold a maximum of 2 electrons
 Shapes of orbitals:
 S is spherical
 P is peanut shaped
 D is daisy shaped
 F is unknown
QUANTUM NUMBERS and ATOMIC ORBITALS
Fourth Quantum Number:
Describes the electron spin
Both electrons in an orbital are negative, so
they repel each other and spin in opposite
directions
Use arrows to represent electrons
QUANTUM NUMBERS and ATOMIC ORBITALS
Pauli Exclusion Principle:
No two electrons in an atom can have the
same set of 4 quantum numbers because
electrons repel each other
Example: 2 electrons may both be:
in the first energy level (same first number)
sitting in an s sublevel (same second #
moving in a sphere shape (same third #)
BUT one electron spins clockwise and one
spins counter clockwise ( which means they
have different fourth #s)
ELECTRON CONFIGURATIONS
Example 1: Map out the quantum
numbers for all the electrons in Hydrogen
Find the # of electrons in hydrogen (atomic
number will give you this number)
What order do you fill in s, p, d, f in the rings?
ELECTRON CONFIGURATIONS
 Diagonal RulePattern that shows the order the
electrons fill in the orbitals: Some People Do Forget
1s
2s
3s
4s
5s
6s
7s
2p
3p
4p
5p
6p
7p
3d
4d
5d
6d
Notice that the electrons do
not fill in all of the level 3
first (3s, 3p, 3d) and then
move to level 4
4f
5f
Instead, electrons fill in the
orbitals in the order that is
easiest to them (easier for
an electron to fill in a 4s
before it fills in a 3d)
Aufbau Principle: Electrons have to fill in the lowest
(easiest) energy level or orbital first
ELECTRON CONFIGURATIONS
 Hund’s Rule:
Every orbital must get
one electron first,
before you double up.
“Cookie Rule”
 Example 2: Helium
ELECTRON CONFIGURATIONS
 Example 3: Lithium
 Example 4: Fluroine
ELECTRON CONFIGURATIONS
 Orbital Notation
drawing out
configurations with
arrows
 Electron
Configuration
Notation/Superscript
Notation:
writing configurations
with superscripts to
represent electrons
ELECTRON CONFIGURATIONS
 Do Orbital Notation and Electron
Configuration for the following:





Zn
I
Cl
Mg
As
NOBLE GAS CONFIGURATIONS
 Noble Gas Configurations:
 Write out the superscript notations for:
Neon:
Sulfur:
Sulfur has the same configuration as Neon plus a 3s23p4
So, you could use the noble gas as a shortcut and write
Sulfur’s configuration as
[Ne] 3s23p4 OR
[Ne]
 Noble gas configuration: write the noble gas (group 18) that comes
directly before the element in question and then add the rest of the
configuration
 Practice:
 Write the noble gas superscript notation for the following elements.
C
W
Np
Sn
DOT DIAGRAMS
Lewis Dot Diagrams:
Way to show the number and position of
electrons on the outermost energy level
Since the energy levels all overlap and cover
one another, only the outermost energy level is
able to bond with other elements
The electrons involved in bonding are called the
valence electrons (to get these electrons look at the
column number)
DOT DIAGRAMS
Lewis Dot Diagrams:
Chemical symbol + Number of valence
electrons
The rules for orbitals still apply, so no side can
have more than two dots, and each “p” orbital
side gets one dot, before you double up
p1
s orbital
p orbitals
p2
X
p3
s
DOT DIAGRAMS
Noble gases have a full valence
There are no empty spaces so the
element does not need any more electrons
Stable octet – 8 electrons in the valence
so the element does not want to bond (this
means it is stable)
Only the noble gases have a stable octet
DOT DIAGRAMS
Practice: Write the noble gas superscript
notation and then draw the dot diagram for
the following:
V
Br
Al
K