Download Electron configuration PPT - River Dell Regional School District

Chapter 5
Arrangement of Electrons
I.
II.
III.
IV.
V.
VI.
Electromagnetic Waves
Dual Nature of Light
Bohr Model of the Atom
Quantum Model
Quantum Numbers
Determining Number of Orbital
Types and Electrons
VII. Electron Configurations
Review: History of the Atomic
Theory
1803
1897
1909
1913
1935
Today
solid
particle
electron
proton
e- orbit
nucleus
neutron
Quantum
Atom
theory
Dalton
Thomson
Rutherford
Bohr
Chadwick
Schrodinge
r and others
I. Electromagnetic Waves
A. Definition of a Wave
1. method by which energy is transferred
from one point to another
B. Definition of Electromagnetic Wave
1. a form of energy that exhibits wavelike behavior as it travels through
space
Origin - the base line of the energy
Crest - high point on a wave
Trough - low point on a wave
Amplitude - distance from origin to crest
Wavelength - distance from crest to crest
abbreviated l Greek letter lambda.
C. Properties of Electromagnetic
Waves
1. Travels at 3 x 1010 cm / second
(or 3.00 x 108m/s) in a vacuum
Known as the “Speed of Light”
2. Vary in wavelength and frequency
a. wavelength – distance between
corresponding points on waves
b. frequency – the number of waves
that pass a point in a given amount
of time (usually one second)
• The number of waves that pass a given
C. Frequency
point per second
is frequency
• SI units are hertz (hz) or cycles/sec
• Abbreviated n the Greek letter nu
• Relationship between wavelength and
frequency is expressed by c
= ln
Electromagnetic Wave
Disturbance in a magnetic field is perpendicular to
a disturbance in an electric field
D. Examples of Electromagnetic Waves
1. radio waves
2. microwaves
3. infrared
4. white light (visible spectrum)
5. ultraviolet light
6. X-rays
7. gamma radiation
EMS
High
Low
energy
energy
Radio Micro Infrared
Ultra- XGamma
waves waves .
violet Rays Rays
Low
High
Frequency
Frequency
Long
Short
Wavelength
Wavelength
Visible Light
Continuous Electromagnetic Spectrum
---------> increasing wavelength ----------->
E. Wavelength, Frequency and Energy
1. wavelength and frequency
a. the frequency increases as the
wavelength decreases
b. inverse relationship
c = ln
2. frequency and energy
a. as the frequency increases the
energy increases
b. direct relationship
E = hn
---------------- > decreasing energy ---------------------
----------------> decreasing frequency ---------------->
---------------> increasing wavelength ---------------->
Diagram Showing Wavelength and Frequency
F. Types of Spectra
1. Continuous – all wavelengths within a
given range are included
2. Electromagnetic – all electromagnetic
radiation arranged according to
increasing or decreasing wavelength
a. unit for wavelength ranges from
meters to nanometers
b. unit for frequency is hertz (Hz)
(# waves per second)
3. Visible spectrum - light you can see
(ROY-G-BIV)
a. red has the longest wavelength and
the smallest frequency
b. violet has the shortest wavelength
and the greatest frequency
4. Bright Line spectrum (emission spectrum)
a. bands of colored light emitted by
excited electrons when they return to
the ground state
Passing Light Through a Prism
• White light is made
up of all the colors
of the visible
spectrum.
• Passing it through a
prism separates the
colors in white light
If the light is not white
• By heating a gas
with electricity we
can get it to give off
colors.
• Passing this light
through a prism
does something
different.
Producing an Emission Spectrum
• Each element gives
off its own
characteristic colors.
• Can be used to
identify the atom.
• This is how we know
what stars are made
of.
G. Spectroscopy
1. emission spectra of a substance is
studied to determine its identity
2. spectroscope – instrument that
separates light into a spectrum
3. spectral lines – represent wavelength
of light emitted when excited electrons
fall back to the ground state
How Does a Spectroscope Work?
Emission Spectrum (Line Spectrum)
Emission Spectrum
Max Planck (1858-1947)
Albert Einstein
Light and Electron Arrangement
II. Light Has a Dual Nature!!!!!
A. Light can act like a particle or a wave
1. emission and absorption of light by
matter can not be explained by wave
theory
2. only certain frequencies of light
produce the photoelectric effect
a. emission of electrons by some
metals when they are exposed to
light
3. In 1900 Max Planck observed that a hot
object loses energy in packets called
quanta
a. this energy is directly related to the
wave frequency ( E = hv)
b. in 1905 Einstein said this relationship
held for all electromagnetic radiation
Ephoton = hv
Light Has a Dual Nature (Particle + Wave)
Light Interference Pattern (Wave Nature)
Photoelectric Effect – Particle Nature
Light hits a metal and electrons are released and an
electric current may be produced
Photoelectric Effect – Particle Nature of Light
Only light of a certain frequency or higher
will cause the photoelectric effect
4. Vocabulary
a. quantum – quantity of energy gained
or lost by an atom when electrons are
excited
b. photon – a quantum of light
c. ground state – lowest energy level of
an atom
d. excited state – a heightened state of
energy in an atom
III. The Bohr Model of the Atom
A. Electrons of hydrogen circle the
nucleus in orbits
1. orbits have a fixed amount of energy
in the ground state
2. orbits are a fixed distance from the
nucleus
3. orbits furthest from the nucleus have
the greatest energy
Niels Bohr
(1885 – 1962)
Bohr Model of the Atom
4. Electrons in the ground state can absorb
quanta of energy – become excited- and
move to a higher orbit
5. Electrons emit quanta of energy when
they return to the ground state
6. Model applies only to hydrogen atoms
Bohr’s Model
Nucleus
Electron
Orbit
Energy Levels
Bohr’s Model
Increasing energy
Fifth
Fourth
Third
Second
First
Nucleus
• Further away
from the nucleus
means more
energy.
• There is no “in
between” energy
• Energy Levels
The Electron Becomes Excited
• The energy level and electron starts from
is called its ground state.
• As it absorbs energy it goes up to an
excited state energy level.
• Then what happens?
Ground State of the Electron in the
Hydrogen Atom
The black dot in energy level 1 is the
electron
• Energy in the form of heat, light or
electricity can excite the electron so that it
moves to higher energy levels (it becomes
excited)
• As the electron falls back to ground state it
gives the energy back as light
• Electron may fall down in steps
• Each with a different energy
Lyman, Balmer, Paschen Series for Hydrogen
IV. Quantum Model of the Atom
A. Problem With the Bohr Model – Why
could the electron in hydrogen orbit in
only a small number of allowed paths?
B. Solving the Problem
1. Louis de Broglie – electrons have a
dual nature - they can act like
particles or waves !!!
Diffraction Patterns
x-rays through Al
electrons through Al
2. Schrodinger – developed equations
that treat electrons in atoms like waves
a. describe the shapes of the orbitals
in which electrons have a high
probability of being found
b. quantum theory – mathematical
explanation for the wave properties
of electrons that apply to all atoms
Louis de Broglie
(1892-1987)
Electrons have a dual
nature (particle + wave)
Erwin Schrodinger
(1887-1961)
Schrodinger equation
describes wave
properties of electrons
mathematically
The Quantum Mechanical Model
• The atom is found
inside a blurry
“electron cloud”
• A area where there is
a chance of finding an
electron
C. Principles of the Quantum Model
1. electrons act like waves and particles
2. probability of an electron being found
at various distances from the nucleus
3. orbitals – a 3-D region about the
nucleus where a specific electron may
be found
4. electrons have greater energy as their
distance from the nucleus increases
5. energies of orbitals are quantized within
main energy levels
6. the exact location of electrons can not
be pinpointed – they are found in regions
of high probability called orbitals or
electron clouds
Similarities -Bohr and Quantum Model
Quantum
Atomic
Model
1. The closer the orbital to the nucleus the
lower the energy
2. To move from a lower to a higher level
the energy absorbed must be equal to
the difference between the levels
3. When e- drops from a higher to lower
level electromagnetic radiation is emitted
equal to the difference in energy levels
4. The most probable location of the e- is a
distance equal to the lowest energy level.
S orbitals
• 1 s orbital for
Every energy level
• Spherical
shaped
• Each s orbital can hold 2 electrons
• Called the 1s, 2s, 3s, etc.. orbitals.
P orbitals
•
•
•
•
Start at the second energy level
3 different directions
3 different shapes
Each can hold 2 electrons
P Orbitals
D orbitals
• Start at the third energy level
• 5 different shapes
• Each can hold 2
electronshttp://www.falstad.com/qmatom/
F orbitals
Orbitals (s, p, d, f)
Orbitals (s, p, d types)
s orbitals
(one type)
p orbitals
(3 types)
d orbitals
( 5 types)
Orbitals in Sodium (Na)
V. Quantum Numbers
A. Principal Quantum Number
1. main energy level
B. Orbital Quantum Number
1. shape of orbital (s,p,d,f)
C. Magnetic Quantum Number
1. orientation of orbital about the nucleus
D. Spin Quantum Number
1.indicates clockwise or counterclockwise spin of the electron (+ or – ½)
VI. Determining Number of Orbital
Types and Electrons
A. If n = the number of the principal energy
level or shell ( 1-7) and there is a maximum
of 2 electrons per orbital then:
1. n = the possible number of orbital types
for that shell
2. n2 = total number of orbitals possible
3. 2n2 = total number of electrons possible
Orbital shapes and maximum number of electrons
in the in the first four energy levels
# of
Max
shapes electrons
Starts at
energy level
s
1
2
1
p
3
6
2
d
5
10
3
f
7
14
4
B. Examples
If n = 3 then in energy level 3:
• 3 orbital types possible (s,p,d)
• 9 orbitals are possible
• 18 electrons are possible
(n)
(n2)
(2n2)
If n = 4 then in energy level 4:
• 4 orbital types possible (s,p,d,f) (n)
• 16 orbitals are possible
(n2)
• 32 electrons are possible
(2n2)
VII. Electron Configuration
A. Rules and Principles
1. Aufbau Principle – an electron
occupies the lowest energy orbital that
can receive it
2. Hund’s Rule – orbitals of equal energy are
each occupied by one electron before
any orbital is occupied by a second
electron
3. Pauli Exclusion Principle – no two
electrons in the same atom can have the
same set of four quantum numbers
4. Heisenberg Uncertainty
Principle
Both the velocity and
position of a particle
(electron) can not be
measured at the
same time
Werner Heisenberg
(1901-1976)
Heisenberg Uncertainty Principle
• It is impossible to know exactly the
position and velocity of a particle at the
same time.
• To measure where a electron is, we use
light.
• But the light moves the electron
• And hitting the electron changes the
frequency of the light.
Before
Photon
Moving
Electron
After
Photon
changes
wavelength
Electron Changes
Velocity
B. Types of Electron Configurations
1. Electron –configuration notation
a. indicates number of the principal
energy level, the orbitals, and
the number of electrons possible
2. Orbital Notation – arrows indicate location
and spin of electrons
3. Electron-dot structure – indicates valence
shell electrons
ASSIGNING CONFIGURATIONS
RULES TO REMEMBER
1) Electron occupies the lowest energy
orbital that can receive it
2) An orbital can hold a maximum of 2 e3) Orbitals of equal energy are each
occupied by one electron before any
orbital is occupied by a second electron
4) No two electrons in the same atom can
have the same set of quantum numbers
5) Always start with 1s
Electron Configuration Notation
Writing Electron Configuration Notation
Examples
•
•
•
•
Hydrogen
•
Helium
• 1s2
Lithium
• 1s22s1
Berylium
• 1s22s2
1s1
Writing Electron Configuration Notation
More Examples
• Boron
• 1s2 2s2 2p1
• Carbon
• 1s2 2s2 2p2
• Nitrogen
• 1s2 2s2 2p3
• Oxygen
• 1s2 2s2 2p4
How to Remember Order to Assign
Electrons (The easy Way)
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
•
2
1s
• 2 electrons
Fill from the bottom up
following the arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
•
2
2
1s 2s
• 4 electrons
Fill from the bottom up
following the arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
•
2
2
6
2
1s 2s 2p 3s
• 12 electrons
Fill from the bottom up
following the arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
•
2
2
6
2
1s 2s 2p 3s
6
2
3p 4s
• 20 electrons
Fill from the bottom up
following the arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
•
2
2
6
2
1s 2s 2p 3s
6
2
10
6
3p 4s 3d 4p
5s2
• 38 electrons
Fill from the bottom up
following the arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
•
2
2
6
2
1s 2s 2p 3s
6
2
10
6
3p 4s 3d 4p
5s2 4d10 5p6 6s2
• 56 electrons
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
3d
3p • The first two electrons
go into the 1s orbital
2p
• Notice the opposite
spins
• only 13 more
5f
4f
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
5f
4f
3d
3p • The next electrons go
into the 2s orbital
2p
• only 11 more
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
3d
3p • The next electrons go
into the 2p orbital
2p
• only 5 more
5f
4f
Increasing energy
7s
6s
5s
4s
7p
6p
6d
5d
5p
4d
4p
3p •
3s
2s
1s
2p •
•
•
5f
4f
3d
The last three electrons
go into the 3p orbitals.
They each go into
separate shapes
3 unpaired electrons
1s22s22p63s23p3
Writing Orbital Notation
Writing Orbital Notation
Writing Electron-dot Notation
(show only valence electrons- outer shell)
Writing Shorthand Configurations
• Chlorine (Cl)
1s22s22p63s23p5
• Calcium (Ca)
1s22s22p63s23p64s2
• Iron (Fe)
1s22s22p63s23p63d64s2
Shorthand Configuration
[Ne] 3s23p5
[Ar] 4s2
[Ar] 3d64s2
Practice Writing Shorthand Configuration
1. Magnesium (z = 12)
[Ne] 3s2
2. Iodine (z = 53)
[Kr] 4d105s25p7
3. Lead (z = 82)
[Xe] 4f14 5d10 6s2 6p2