Download Chapter 4 Modern Atomic Theory

Chapter 4
Modern Atomic Theory
Lesson 4.1: Emission Spectra
Learning Target: I will understand why the Hydrogen
atom emission line spectrum led to the development of
the Bohr model.
Success Criteria: I can describe the hydrogen atom
emission line spectrum and how it led to the concept
of energy levels in atoms.
Electromagnetic Spectrum
• Electromagnetic radiation
• Form of energy that exhibits wavelike behavior as it travels
through space
• All travel at a constant speed of 3.00 x 108 m/s through a
vacuum (speed of light), slightly slower through matter
Ground vs. Excited States
• Ground State: Lowest energy state of an atom is the
ground state
• Excited state: is state where atom has a higher
potential energy than ground state
• Lots of excited states but only one ground state
• When excited atom returns to its ground state, it gives
off the energy it gained
What does light have to do with it?
• When energy is added to an atom (Heat or
Electricity) electrons move to higher energy levels
• Electrons move farther away from the nucleus
• When they move back to the lower energy level, light
is emitted
• Different colors for different elements
• Different electron configurations
result in different colors
Hydrogen-Atom Emission Line
Spectrum
• Electric current passed
through vacuum tube of
H2 gas at low pressure,
emission of pinkish glow
• Shone pink light through
a prism and it separated
into 4 specific colors of
the visible spectrum
• Called emission line
spectra
Oh MY Gosh!!
• Classical theories predicted that
hydrogen atoms would be
excited by all energies of light
and form a continuous
spectrum
• Should see a continuous range of
frequencies of electromagnetic
radiation, no lines of color
Excited Hydrogen Atom
• Excited hydrogen atom falls to ground
state or lower energy state, it emits a
photon of radiation equal to difference in
energy between atom’s initial state and its
final state
• Emits specific frequencies (energies) of light
• i.e. hydrogen atom exists only in very specific
energy states
Lesson 4.2: Bohr Model
Learning Target: I will understand the Bohr model and
its development.
Success Criteria: I can describe why the Bohr model
was created. I can draw a Bohr model for the first 20
elements.
Enter Niels Bohr…
• Bohr Model of the Atom
• Explained hydrogen emission spectra and only certain
frequencies (energy) allowed
• Electron can circle the nucleus only in allowed orbits
• When in an allowed orbit, electron has a definite fixed
energy
• Electron (and therefore atom) is in lowest energy state when
it is in orbit closest to the nucleus
• Orbit is separated from nucleus by large empty space where
electron
Think
No steps
Between rungs
Drawing Bohr Models:
1. Find the element on the periodic table
2. Determine the number of electrons – it is the same as
the atomic number for a neutral atom. This is how
many electrons you will draw
3. Find out which period (row) your element is in. This
will be the number of energy levels (orbits) you will
draw.
4. Draw a nucleus in the center, and then draw the number
of orbits (shells) that corresponds to the energy levels
Drawing Bohr Models:
1. Add the electrons from the inside out, according to
the chart below:
Electron Shell
Maximum Number of Electrons
1
2
2
8
3
8
4
18
5
18
6
32
7
32
Valence Electrons
• Valence Electrons: the electrons in the outermost
shell (farthest from the nucleus)
• The shell containing electrons that are furthest from
the nucleus is called the valence shell
• Noble gases have a full outer shell, all others have
partially filled outer shells
• Atoms will try to gain or lose electrons to have a full
valence shell
Bohr’s Model explained Line spectra
• Electron can gain energy equal to difference in
energy between orbits and move to higher energy
orbit
• Called absorption
• When electron falls back to lower energy level, a
photon is emitted equal to difference between orbits
• Called emission
Bohr’s Model Continued
• Bohr calculated allowed energy level differences and
matched them to the experimentally observed values
for the line spectra in each series
• Unfortunately, it only worked for hydrogen
• Couldn’t explain spectra for atoms with > 1 electron
Lesson 4.3: Quantum Theory
Learning Target: I will understand the development of
the quantum model of the atom.
Success Criteria: I can name the 4 quantum numbers,
and discuss how they relate to the “address” of an
electron.
Could electrons Behave Like Waves?
• 1924 Louis DeBroglie
• If light can behave like a particle, can electrons behave like waves?
• Diffraction experiments showed that electron beams can
show interference patterns as they pass through a
diffraction grating
Heisenberg and the Dawn of the
Uncertainty Principle
• 1927 Werner Heisenberg
• Electrons are detected by interaction with photons
• Since electrons have same energy as photons, any attempt to detect
them, throws the electron off course
• Always uncertainty in trying to locate an electron
• Heisenberg Uncertainty Principle
• It is impossible to determine simultaneously both the position and velocity
of an electron or any other particle
Schrodinger Wave Equation
• 1926 Erwin Schrodinger
• Developed an equation that treats electrons in atoms as waves
• Only waves of specific energies (and frequencies) provide solutions
to the equation
• Together with the Uncertainty principle, the Schrodinger
equation lays the foundation for the modern quantum
theory of the atom
Quantum Theory
• Describes mathematically the wave properties of electrons
and other very small particles
• Solutions are known as wave functions
• Wave functions give the probability of finding an electron at a given
place around the nucleus
• NO NEAT ORBITS as BOHR suggested
• Electrons exist in regions called orbitals
• Orbitals: 3D regions of space around the nucleus that indicates the
probable location of an electron
Quantum Numbers
• Quantum numbers: specify the properties of atomic orbitals
and the properties of electrons in atomic orbitals
• n=principal quantum number
• Indicates the main energy level (think period number
• Electron’s energy and average distance from the nucleus increase
• l -Angular Momentum Quantum Number
• Indicates the shape of the orbital (s, p, d, f)
Quantum Numbers
• m =Magnetic Quantum Number
• Indicates the orientation of an orbital around the nucleus
• ms = spin quantum number
• Two possible fundamental spin states of an electron in an
orbital
• A single orbital can hold a maximum of 2 electrons but the
two electrons must have opposite spin states (+1/2 and 1/2
Orbital Shapes and Orientation
•
•
•
•
s orbitals have 1 orientation
p orbitals have 3 orientations
d orbitals have 5 orientations
f orbitals have 7 orientations
Each orbital orientation can
hold a pair of electrons.
Lesson 4.4: Electron Configuration
Learning Target: I will understand how to write
electron configurations.
Success Criteria: I can write the electron configurations
for the first 20 elements using the periodic table.
Electron Configurations
• Electron Configuration: shows the
distribution of electrons in an atom. It
shows the number of electrons in each
sublevel of each energy level of the
atom.
1
1s
Number of electrons
Sublevel
Principle energy level
Rules Governing Electron
Configurations
• Aufbau Prinicple: An electron occupies the lowest energy orbital that
can receive it
• i.e. start with energy level 1, s orbital
• Pauli Exclusion Principle: No two electrons in the same atom can
have the same set of four quantum numbers
• i.e. electrons in the same orbital must have different spin
• Hund’s Rule: Orbitals of equal energy are each occupied by one
electron before any orbital is occupied by a second electron and all
electrons in singly occupied orbitals must have the same spin state
• i.e. put one up spin electron in each orbital before adding down spin electrons
Using the Periodic Table and its Blocks to
get Electron Configurations
Exceptions
• Cr electron configuration
• would expect [Ar] 4s2 3d4
• Actually [Ar]4s1 3d5
• Half filled d orbitals are more stable so one of the s electrons
goes to a d orbital
• Cu electron configuration
• Would expect [Ar] 4s2 3d9
• Actually [Ar]4s1 3d10
• Filled d orbitals are more stable so one of the s orbitals goes to
a d orbital
Lesson 4.5: Orbital Notation and
Noble Gas Notation
Learning Target: I will understand how to write an
electron configuration in orbital notation and noble gas
notation
Success Criteria: I can use lines and arrows to visually
represent the electron configuration. I can determine
the proper noble gas to use in noble gas notation.
Electron Configuration Notation
• Orbital Notation: Uses the lines and up and down
arrows to represent electron configuration
• Each line represents an orbital, each arrow represents an
electron
• Energy levels are represented by distance from the bottom
of the paper
• __ orbital notation for He
1s2
Electron Configuration Notation
• Noble Gas Notation: Takes closest noble gas symbol
without going over the atomic number and continues
electron configuration from that point on
• This method allows you to write and abbreviated electron
configuration
• Ie. If you have Sr, use [Kr] 5s2