Title: Chapter 4 Arrangement of Electrons in Atoms
1Chapter 4 Arrangement of Electrons in Atoms
- Section 4.1
- The Development of a New Atomic Model
2Light Is it a Wave?
- The study of light led to the development of the
quantum mechanical model. - Light is a kind of electromagnetic radiation.
- Electromagnetic radiation includes many kinds of
waves (X-rays, UV rays, microwaves, radio waves).
- They all move at the same speed, 3.00 x 108 m/s,
which is called symbol (c)
3Parts of a Wave
Origin
4Parts of a Wave
- 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
- Wavelength - is abbreviated l, Greek letter
lambda.
5Frequency
- The number of waves that pass a given point per
second. - Units are cycles/sec or hertz (hz)
- Abbreviated n, the Greek letter nu
- c ln
6Frequency and Wavelength
- They are inversely related
- As one goes up the other goes down.
- Different frequencies of light is different
colors of light. - There is a wide variety of frequencies, which is
called a spectrum.
7Radio waves
Microwaves
Infrared .
Ultraviolet
X-Rays
Gamma Rays
Long Wavelength
Short Wavelength
Visible Light
8Light Is it a Particle?
- Studied by Max Planck in 1900
- He wondered why a hot object does not emit (give
off) electromagnetic radiation all the time (it
should if light were just a wave) - He proposed that a hot object emits energy in
small packets (chunks) called quanta. - A quantum is the minimum quantity of energy that
can be lost or gained by an atom.
9Light Is it a Particle?
- Planck devised a formula E h?, where v is
frequency and h 6.626 x 10-34 Js (Plancks
constant). - This was expanded on by Einstein in 1905 and
refers to the duality of light (its both a
wave a particle). - Einstein called the particles of light photons.
- A photon has zero mass at rest and its energy
depends on the frequency. - Einstein explained the photoelectric effect
(emission of electrons by a metal when light
strikes it) by referring to these photons.
10An Explanation of Atomic Spectra
11Where the electron starts
- The lowest energy level of an electron is called
its ground state. - When an electron has higher energy than in its
ground state, its called an excited state.
12Changing the Energy
- Lets look at a hydrogen atom
13Changing the Energy
- Heat or electricity or light can move the
electron up energy levels
14Changing the Energy
- As the electron falls back to ground state it
gives the energy back as light
15Changing the Energy
- May fall down in steps
- Each with a different energy
16Visible Ultraviolet Ultraviolet Ultraviolet
Infrared
Visible
Ultraviolet
17Ultraviolet
Visible
Infrared
- Further they fall, more energy, higher frequency.
- This is simplified
- The orbitals also have different energies inside
energy levels - All the electrons can move around.
18Bohrs Model
- Why dont the electrons fall into the nucleus?
- Move like planets around the sun.
- In circular orbits (orbitals) at different
levels. - Amounts of energy separate one level from
another. - A good model, but did not explain everything.
19Bohrs Model
Nucleus
Electron
Orbit
Energy Levels
20Bohrs Model
Fifth
- Further away from the nucleus means more energy.
- There is no in between energy
- Energy Levels
Fourth
Third
Increasing energy
Second
First
Nucleus
21Chapter 4 Arrangement of Electrons in Atoms
- Section 4.2
- The Quantum Model of the Atom
22Electrons as Waves
- If light could behave as a particle and wave,
what about electrons? - Louis de Broglie in 1924 did this and found
electrons could be diffracted (bent) and can
interfere with each other. - This lead to the question, if electrons are both
waves and particles, then where are they in the
atom?
23Heisenberg Uncertainty Principle
- Studied by Werner Heisenburg in 1927
- He said, that it is impossible to know exactly
the speed and velocity of a particle. - The better we know one, the less we know the
other. - The act of measuring changes the properties.
24After
Before
Photon
Photon changes wavelength
Moving electron
Electron changes velocity
25The Quantum Mechanical Model
- Developed by Erwin Schrodinger derived an
equation that described the energy and position
of the electrons in an atom - Energy is quantized. It comes in chunks.
- A quanta is the amount of energy needed to move
from one energy level to another. - Since the energy of an atom is never in between
there must be a quantum leap in energy.
26The Quantum Mechanical Model Cont
- Things that are very small behave differently
from things big enough to see. - The quantum mechanical model is a mathematical
solution - It is not like anything you can see.
27The Quantum Mechanical Model
- Has energy levels for electrons.
- Orbits are not circular.
- It can only tell us the probability of finding
an electron a certain distance from the nucleus.
28The Quantum Mechanical Model Just watch
- The atom is found inside a blurry electron
cloud - A area where there is a chance of finding an
electron. - Draw a line at 90
29Atomic Orbitals Quantum Numbers
- Within each energy level the complex math of
Schrodingers equation describes several shapes. - These are called atomic orbitals.
- Regions where there is a high probability of
finding an electron. - They specify the properties of the orbitals and
the electrons in each orbital. - It is like the chemical address or seating chart
for an electron.
30Principal Quantum Numbers
- Principal Quantum Number (n) is the energy level
of the electron. - The maximum energy level for an atom is the same
as the period it is on in the Periodic Table. - The total number of electrons per energy level is
found by 2n2.
31Angular Momentum Quantum Number
- Indicates shape of orbital.
- Symbolized by l
- For every n, the number of orbital shapes l is 0
and all positive integers ? n-1
32S orbitals
- 1 s orbital for every energy level
- Spherical shaped
- Each s orbital can hold 2 electrons
- Called the 1s, 2s, 3s, etc.. orbitals.
- Their l 0
33P orbitals
- Start at the second energy level
- 3 different directions
- 3 different shapes
- Holds total of 6 electrons
- Their l 1
34P Orbitals
35D orbitals
- Start at the third energy level
- 5 different shapes
- Holds total of 10 electrons
- Their l 2
36F orbitals
- Start at the fourth energy level
- Have seven different shapes
- Holds a total of 14 electrons
- Their l 3
37F orbitals
38Summary
of shapes
Max of electrons
Starts at energy level
s
1
2
1
p
3
6
2
5
10
3
d
7
14
4
f
39Magnetic Quantum Number
- Indicates orientation of orbital.
- Symbolized by m
- It is found by all integers from l to l
- For example, if l 1, then m -1, 0, and 1
40Spin Quantum Number
- Indicates spin of orbital.
- Since every orbital holds 2 electrons, there are
only 2 values for spin, 1/2 and 1/2
41Chapter 4 Arrangement of Electrons in Atoms
- Section 4.3
- Electron Configurations
42Electron Configurations
- The way electrons are arranged in atoms.
- Aufbau principle- electrons enter the lowest
energy first. - This causes difficulties because of the overlap
of orbitals of different energies. - Pauli Exclusion Principle- at most 2 electrons
per orbital - different spins (one 1/2, other
1/2)
43Electron Configurations
- Hunds Rule- When electrons occupy orbitals of
equal energy they dont pair up until they have
to . - Lets determine the electron configuration for
Phosporus - Need to account for 15 electrons
44- The first two electrons go into the 1s orbital
- Notice the opposite spins
- only 13 more
45- The next electrons go into the 2s orbital
- only 11 more
46- The next electrons go into the 2p orbital
- only 5 more
47- The next electrons go into the 3s orbital
- only 3 more
48- The last three electrons go into the 3p orbitals.
- They each go into separate shapes
- 3 unpaired electrons
- 1s22s22p63s23p3
49The easy way to remember
50Fill from the bottom up following the arrows
51Fill from the bottom up following the arrows
52Fill from the bottom up following the arrows
53Fill from the bottom up following the arrows
- 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2
54Fill from the bottom up following the arrows
- 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2
55Fill from the bottom up following the arrows
- 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2
4f14 5d10 6p6 7s2
56Fill from the bottom up following the arrows
- 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2
4f14 5d10 6p6 7s2 5f14 6d10 7p6
57This is the Diagonal Rule!
58Orbitals fill in order
- Lowest energy to higher energy.
- Adding electrons can change the energy of the
orbital. - Half filled orbitals have a lower energy.
- Makes them more stable.
- Changes the filling order
59Write these electron configurations
- Titanium - 22 electrons
- 1s22s22p63s23p64s23d2
- Vanadium - 23 electrons 1s22s22p63s23p64s23d3
- Chromium - 24 electrons
- 1s22s22p63s23p64s23d4 is expected
- But this is wrong!!
60Chromium is actually
- 1s22s22p63s23p64s13d5
- Why?
- This gives us two half filled orbitals.
- Slightly lower in energy.
- The same principal applies to copper.
61Coppers electron configuration
- Copper has 29 electrons so we expect
- 1s22s22p63s23p64s23d9
- But the actual configuration is
- 1s22s22p63s23p64s13d10
- This gives one filled orbital and one half filled
orbital. - Remember these exceptions