Chapter 12 Electrons in Atoms
Dec 13, 2015
Chapter 12
Electrons in Atoms
Greek Idea
Democritus and Leucippus
Matter is made up of indivisible particles
Dalton - one type of atom for each element
Thomson’s ModelDiscovered electrons
Atoms were made of positive stuff
Negative electron floating around
“Plum-Pudding” model
Rutherford’s ModelDiscovered dense positive piece at the center of the atom
Nucleus
Electrons moved around
Mostly empty space
Bohr’s Model
Why don’t the electrons fall into the nucleus?
Move like planets around the sun.
In circular orbits at different levels.
Amounts of energy separate one level from another.
Bohr’s Model
Nucleus
Electron
Orbit
Energy Levels
Bohr’s ModelIn
crea
sing
ene
rgy
Nucleus
First
Second
Third
Fourth
Fifth
} Further away from the nucleus means more energy. There is no “in between” energy Energy Levels
The Quantum Mechanical ModelEnergy 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.
Schrödinger derived an equation that described the energy and position of the electrons in an atom
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.
The Quantum Mechanical Model
Has energy levels for electrons.
Orbits are not circular.
It can only tell us the probability of finding an electron at a certain distance from the nucleus.
The Quantum Mechanical Model
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 %
The Quantum Mechanical Model
Atomic Orbitals
Principal Quantum Number (n) = the energy level of the electron.
Within each energy level the complex math of Schrödinger's equation describes several shapes.
These are called atomic orbitals
Regions where there is a high probability of finding an electron.
There is 1 s orbital for every energy level Spherical shaped
Each s orbital can hold 2 electrons.
Called the 1s, 2s, 3s, etc.. orbitals.
S orbitals
P orbitalsStart 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 electrons
F orbitals Start at the fourth energy level Have seven different shapes 2 electrons per shape
By Energy Level First Energy
Level only s orbital only 2 electrons 1s2
Second Energy Level
s and p orbitals are available
2 in s, 6 in p 2s22p6
8 total electrons
By Energy Level Third energy level s, p, and d
orbitals 2 in s, 6 in p, and
10 in d 3s23p63d10
18 total electrons
Fourth energy level
s,p,d, and f orbitals
2 in s, 6 in p, 10 in d, and 14 in f
4s24p64d104f14
32 total electrons
By Energy Level Any more than
the fourth and not all the orbitals will fill up.
You simply run out of electrons
The orbitals do not fill up in a neat order.
The energy levels overlap
Lowest energy fill first.
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5fIn
crea
sing
Ene
rgy
Electron 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
Electron Configuration Hund’s Rule- When electrons occupy
orbitals of equal energy they don’t pair up until they have to .
Let’s determine the electron configuration for Phosphorus
Need to account for 15 electrons
1s2s 2p3s 3p 3d4s 4p 4d 4f5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f
1s2 2s2 2p6 3s2 3p6 4s23d104p4
What is it?
Selenium
Orbitals 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
Write these electron configurations Titanium - 22 electrons 1s22s22p63s23p64s23d2
•Vanadium - 23 electrons• 1s22s22p63s23p64s23d3
Chromium - 24 electrons 1s22s22p63s23p64s23d4 is expectedBut this is wrong!! Why??
Light 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
All move at 3.00 x 108 m/s = C
Parts of a wave
Wavelength
AmplitudeOrigin
Crest
Trough
Parts of 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 Greek
letter lambda.
Frequency The number of waves that pass a
given point per second. Units are cycles/sec or hertz (hz) Abbreviated the Greek letter nu
c =
Frequency and wavelength 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 The whole range is called a spectrum
Radio waves
Microwaves
Infrared
Ultra-violet
X-Rays
Gamma Rays
Low energy
High energy
Low Frequency
High Frequency
Long Wavelength
Short WavelengthVisible Light
Atomic Spectrum
How color tells us about atoms
Prism White light is
made up of all the colors of the visible spectrum.
Passing it through a prism separates it.
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.
Wave-Particle DualityJJ Thomson won the Nobel prize for describing the electron as a particle.
His son, George Thomson won the Nobel prize for describing the wave-like nature of the electron.
The electron
is a particle!
The electron is an energy
wave!
Confused??? You’ve Got Company!
“No familiar conceptions can be woven around the
electron; something unknown is doing we
don’t know what.”
Physicist Sir Arthur Eddington
The Nature of the Physical World
1934
The Wave-like Electron
Louis deBroglie
The electron propagates through space as an energy
wave. To understand the atom, one must
understand the behavior of
electromagnetic waves.
c = C = speed of light, a constant (3.00 x 108 m/s) = frequency, in units of hertz (hz, sec-1)
= wavelength, in meters
Electromagnetic radiation propagates through space as a wave moving at the speed of light.
Types of electromagnetic radiation:
E = h
EE = Energy, in units of Joules = Energy, in units of Joules (kg·m(kg·m22/s/s22))hh = Planck’s constant (6.626 x 10 = Planck’s constant (6.626 x 10-34-34 J·s)J·s)
= frequency, in units of hertz (hz, sec= frequency, in units of hertz (hz, sec-1-1))
The energy (E ) of electromagnetic radiation is directly proportional to the frequency () of the radiation.
Long Wavelength
=Low Frequency
=Low ENERGY
Short Wavelength
=High Frequency
=High ENERGY
Wavelength Table
…produces all of the colors in a continuous spectrum
Spectroscopic analysis of the visible spectrum…
…produces a “bright line” spectrum
Spectroscopic analysis of the hydrogen spectrum…
This produces bandsof light with definitewavelengths.
Electron transitionsinvolve jumps of
definite amounts ofenergy.