Unit 5
Electrons in Atoms
Chemistry I
Mr. Patel
SWHS
Topic Outline
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Continue Learning Major Ions
Atomic Models (5.1)
Electron Configurations (5.2)
Light and Quantum Mechanics (5.3)
Lewis Dot Structures (7.1)
Atomic Models
• Democritus’s Model
• Thomson’s Plum Pudding Model
• Rutherford’s Model
– Electrons travel in orbit around nucleus
– Could NOT explain chemical properties of
elements
– Need a model for electrons
Bohr Model
• Niels Bohr – electrons is found in a specific
orbit around nucleus
• Each orbit has a specific energy = energy level
• The further away from the nucleus, the higher
the energy
Bohr Model
Bohr Model
• An electron can move between levels
– Can not be between levels
– Think of a ladder
• An electron must gain or lose energy to
change levels
• A quantum of energy – energy to move to
another level
Bohr Model
• More energy between levels
when closer to nucleus
• Less energy between
levels when farther
• Energy levels get closer
together
Bohr Model
• Ground state – lowest energy state for an
electron
• Excited state – any higher energy state
Electron Excitation
http://www.youtube.com/watch?v=4jyfi28i928
&feature=relmfu
Bohr Model
• Each ring on a Bohr Model is labeled as “n”
• n must be a whole number
– n=1, n=2, n=3, etc. (period number)
• Each ring (n) can hold a specific number of
electrons
–
–
–
–
n=1 2 electrons
n=2 8 electrons
n=3 18 electrons
n=4 32 electrons
Drawing Bohr (Rutherford) Diagrams
http://www.youtube.com/watch?v=sKAzHE7A7r
Q&feature=relmfu
Ex: Draw the Bohr Diagram for Hydrogen.
Ex: Draw the Bohr Diagram for Neon.
Ex: Draw the PEL Diagram for Calcium.
Ex: Draw the PEL Diagram for Argon.
Bohr Model
• Correct: Electrons have energy levels and can
move
• Incorrect: Electrons move in orbits
• Matter has a Wave-Particle Duality
Dual Nature of Electrons
Electrons as Particles
Electrons as Waves
Photoelectric
Effect
Young's Double
Slit Experiment
Modern Theory
• Rutherford and Bohr based models on
behavior of large objects
• Small objects behave differently – quantum
mechanics
• Schrödinger Equation solutions  quantum
mechanical model of the atom
Schrödinger Equation
The Cat – A Thought Experiment
• Schrodinger Cat 1
• Schrodinger Cat 2
Quantum Mechanical Model
• Determines the allowed
energies of the electrons
• The probability of where an
electron is – electrons
housed in electron clouds
Atomic Orbitals
• Region in space where there is a high
probability of finding an electron
• Principal quantum number (n) – energy level
– think of the ring labels of the Bohr model
• Each energy level can be made up of sublevels
– orbitals of similar energy but different
shapes
1.
• Shape:
sphere
s orbital
2.
• Shape: Dumbbell
p orbital
3.
• Shape:
clover
(mostly)
d orbital
4.
f orbital
• Shape: multiple clover
Atomic Orbitals
http://www.youtube.com/watch?v=K-jNgq16jEY
Electron Configurations
• Electrons found in orbitals
• Electron configuration – ways in which various
electrons are arranged in orbitals
• 4 orbitals: s
p
d
f
(2 electrons),
(6 electrons)
(10 electrons),
(14 electrons)
Three Rules to find Elec. Config
1. Aufbau Principle
– Electrons occupy orbitals of lower energy first
– For same n, low to high energy: s, p, d, f
Three Rules to find Elec. Config
2. Pauli Exclusion Principle
– Each atomic orbital can have at most 2 electron
– Each electron in an orbital must have opposite
spins
– 2 spins: spin up
or spin down
– How we draw: 1 electron in s orbital: ____
2 electrons in s orbital: ____
– We use arrow with “half head”
Three Rules to find Elec. Config
3. Hund’s Rule
– Electrons occupy orbitals to maximize spin
– For same n, place electrons spin up first then pair
them with spin down
– 1 electron in p orbital ____ ____ ____
– 2 electrons in p orbital ____ ____ ____
– 3 electrons in p orbital ____ ____ ____
– 4 electrons in p orbital ____ ____ ____
– 5 electrons in p orbital ____ ____ ____
– 6 electrons in p orbital ____ ____ ____
Orbital Blocks on PT
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s-block: Groups 1A and 2A (exception: He)
p-block: Groups 3A-8A (exception: He)
d-block: transition metals
f-block: inner transition metals
Remember, the period number is n = principal
energy level
Orbital Blocks on PT
How to write electron configuration
• Ex: What is the electron configuration for O?
• O = oxygen, atomic number 8 = 8 electrons
• Draw spaces:
____
____
____ ____ ____
1s
2s
2p
• Fill spaces according to rules:
____
____
____ ____ ____
1s
2s
2p
• Write: 1s22s22p4
How to write electron configuration
• Ex: What is the electron configuration for C?
• C = carbon, atomic number 6 = 6 electrons
• Draw spaces:
____
____
____ ____ ____
1s
2s
2p
• Fill spaces according to rules:
____
____
____ ____ ____
1s
2s
2p
• Write: 1s22s22p2
3 ways to write electron configurations
1. Using boxes and arrows
____
____
____ ____ ____
1s
2s
2p
2. Long EC: Cl: 1s22s22p63s23p5
3. Short EC: Cl: [Ne] 3s23p5
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Put last noble gas in brackets and write electrons from there
Writing EC
• This is much easier than
it looks.
• Simply, start at
hydrogen and walk to
the desired element
counting all the
elements you pass
Ex. Write EC (all three ways) for Boron.
Ex. Write EC (all three ways) for Mg.
Ex. Write EC (all three ways) for V.
Ex. Write EC (long and short) for Fr.
A Look Back…
• So far we have covered (and mastered):
– Evolution of the Atomic Model
• Democritus, Thomson, Rutherford, Bohr, QM
– Bohr Model and Bohr Diagram
– Quantum Mechanical Model and Orbitals
– Rules of Electron Configuration
– Writing Electron Configurations
Electrons
• Chemical reactions are the breaking and
forming of bonds
• There are two types of bonds: covalent and
ionic (and metallic) = next unit
• Bonding involves the movement of electrons
Valence Electrons
• Valence Electron: electrons in the highest
occupied level
• These are the electrons that participate in
bonding!!!
Valence Electrons
• You do not have to draw a Bohr Model every
time you need to determine the VE’s
• The valence electrons (valency) for an atom is
the same as the group number
• Note: In general, transition metals have two
valence electrons.
Determine the valence electrons for:
1.
2.
3.
4.
5.
6.
7.
Ca
Be
O
Si
H
Ne
Ar
1.
2.
3.
4.
5.
6.
7.
2
2
6
4
1
10
10
Lewis Dot Structures
• Show bonding electrons
• These structures show only
valence electrons.
• How to draw:
– Write Symbol for element
– Determine group number
– Place that many (group number)
dots around symbol
N
B
Lewis Dot Structures
1
2
3
4
5
6
7
8
Lewis Structures (Future)
Draw the Lewis Dot Structure for:
Cs
Al
Ge
Br
Practice Time!!!
The Octet Rule
• Remember that Noble Gases were very stable
– They all have 8 valence electrons (2 for He)
– FULL outer shell of electrons
• Every element will try to become like a noble
gas
• The Octet Rule – atoms will try to have a full
outer shell (= 8 electrons) when bonding
Cations
• Metals tend to lose electrons to have a full
outer shell
• Cation – positively charge ion
– Results from metals losing electrons
• Naming: element name + ion
– Ex: Na = Sodium
but Na1+ = sodium ion
Anions
• Nonmetals/Metalloids tend to gain electrons to
have a full outer shell
• Anion– negatively charged ion
– Results from nonmetals gaining electrons
• Naming: element name with –ide ending + ion
– Ex: Br = Bromine
but Br1- = bromide ion
Cations/Anions
• To determine the charge of an element’s ion,
look at the group/column that it is in
– Group 1: 1+
– Group 2: 2+
– Group 3: 3+
– Group 4: 0
– Group 5: 3– Group 6: 2– Group 7: 1-
Lewis Dot Structures for Ions
• Draw the normal Lewis Dot structure for the
neutral element
• Add electrons if gained or remove electrons if
lost
• Place the appropriate charge
Nitrogen:
N
3-
Nitride
ion:
N
1. Draw the Lewis Dot Structure for Phosphorus.
2.
3.
4.
5.
6.
Will this element for a cation or anion?
What charge will it have?
What will be the name of the ion?
What noble gas is it similar to?
Draw the Lewis Dot Structure for the ion.
1. Draw the Lewis Dot Structure for Barium.
2.
3.
4.
5.
6.
Will this element for a cation or anion?
What charge will it have?
What will be the name of the ion?
What noble gas is it similar to?
Draw the Lewis Dot Structure for the ion.
Light – A Wave
• Newton tried to prove light to be a particle
• However, experimental data showed that light
was actually behaving as a wave
• The study of light led to the quantum
mechanical model of the atom
Properties of Waves
• Wavelength (λ)
– “lambda”
– Distance between
crests
Crest
• Frequency (ν)
– “nu”
– Cycles per second
– Hertz (Hz)
• Amplitude
– Height from zero
to the crest
Trough
Properties of Waves
• The Wave Equation:
c = λ∙ν
– c = speed of light = 3 x 108 m/s
– λ = wavelength – must be in meters
– ν = frequency – in Hertz
• Energy of Light:
E = h∙ν
– E = energy – in Joules
– h = Planck’s constant = 6.626 x 10-34 J∙s
– ν = frequency – in Hertz
Electromagnetic Spectrum
• When light passes through a prism, it is
separated into different frequencies
• You need to know:
– Name and order of each regions
– Order based on wavelength
– Order based on frequency
– Order based on energy
– Details of the Visible Region
Electromagnetic Spectrum
EM Spectrum Song
Long Wavelength
Low frequency
Low Energy
Short Wavelength
High frequency
High Energy
Review: Properties of Waves
• The Wave Equation:
c = λ∙ν
– c = speed of light = 3 x 108 m/s
– λ = wavelength – must be in meters
– ν = frequency – in Hertz
• Energy of Light:
E = h∙ν
– E = energy – in Joules
– h = Planck’s constant = 6.626 x 10-34 J∙s
– ν = frequency – in Hertz
Equations:
c = λ∙ν
E = h∙ν
• c = speed of light = 3 x 108 m/s
• h = Planck’s constant = 6.626 x 10-34 J∙s
Ex: What is the frequency of a wave with a
wavelength of 3.68 x 10-9 m? The energy?
Equations:
c = λ∙ν
E = h∙ν
• c = speed of light = 3 x 108 m/s
• h = Planck’s constant = 6.626 x 10-34 J∙s
Ex: What is the frequency of a wave with a
wavelength of 700 nm? The Energy?
Practice!!!
Atomic Spectra
• When atoms absorb energy, they move into
higher energy levels
• These electrons then return back to a lower
level and release energy as light
• Each atom releases light in a special way
Atomic Spectra
• Atomic Emission
Spectrum – the
frequencies of light
released by an element
split into separate
discrete lines (unlike
light)
Hydrogen Spectrum
• Balmer Series
– Visible Region
• Lyman Series
– UV Region
Light – A Particle
• Light deserves a quantum mechanical
treatment
• Light also behaves as a particle and a wave
(Particle-Wave Duality)
• Light particles called photons – packets or
quanta of light (E = h∙ν)
deBroglie Relation
• Louis deBroglie determined that all matter
that is moving can be considered as waves
• Large object have such a small wavelength
that it can not be observed
• His math showed that as mass decreases, the
wave function becomes more important
(λ=h/mv)
Heisenberg
• Heisenberg Uncertainty Principle – it is
impossible to know the exact velocity and
position for a particle at the same time
Heisenberg Uncertainty principle