Transcript Arrangement of Electrons in Atoms

Arrangement of
Electrons in Atoms
4-1 The Development of a New Atomic Model
The Development of a New
Atomic Model
How do electrons surround the nucleus of the atom?
“The emission of light is fundamentally related to the
behavior of electrons.”
Properties of Light
Scientists early in the 1900’s believed that light
behaved like a wave.
Light also behaves like a particle. Confused? We’ll
discuss this later…
Electromagnetic Radiation – a form of energy that
exhibits wavelike behavior as it travels through
space.
Properties of Light
Electromagnetic Spectrum – all the forms of
electromagnetic radiation.
Fig. 4-1 on pg. 92.
* You need to know all these forms of light, including
approximate wavelength and frequency.
Properties of Light
All these forms of light have the same speed.
c = 3 x 108 m/s
Speed of light in a vacuum. Speed of light in other
mediums is more.
Properties of a Wave:
Wavelength () – the distance between
corresponding points on adjacent waves.
Wavelength values can be very small or very large.
Red light = 700 nm
Radio wave = 100 m
1 nm = 1 x 10-9 m
Properties of a Wave:
Frequency () – the number of waves that pass a
given point in a specific time, usually 1 second.
Unit is wave/second.
* 1 wave/second = 1 hertz (Hz)
Fig. 4-2 pg. 92
Properties of a Wave:
Amplitude – is the height of the wave measured from
the origin to its crest.
The brightness or intensity of the light depends
on the amplitude.
The greater the amplitude, the brighter the light.
Properties of a Wave:
Important Equation #1: c = 
Speed of Light = wavelength x frequency
 is inversely proportional to 
As  increases  decreases. As  decreases 
increases.
Photoelectric Effect
Refers to the emission of electrons from a metal
when light shines on the metal.
The photoelectric effect explains phenomenon such
as solar powered calculators, automatic street lights,
automatic doors, …
Photoelectric Effect
For a given metal, no electrons were emitted if the lights
frequency was below a certain minimum.
The brightness of a light won’t necessarily cause
electrons to flow.
Ex. Red light will not cause electrons to flow in a sheet
of sodium metal, no matter how long or bright the
source is. Violet light will cause electrons to flow.
Violet light has a greater frequency, and a greater
amount of energy per photon. Photon?
Different metals require different amounts of energy
depending on how tightly the electrons are bound to
the metal.
Light as Particles
Max Planck said an object emits energy in small
amounts called quanta.
Quantum – the minimum quantity of energy that can
be lost or gained by an atom.
Follow closely now, this is weird.
Planck’s Theory
Max Planck predicted
accurately how the
spectrum of radiation
emitted by an object
changes with
temperature.
Light as Particles
Energies absorbed or emitted by atoms are quantized,
which means that their values are restricted to certain
quantities. Energy is not continuous.
Ramps vs. Stairs = Continuous vs. Quantized
Ex. Imagine a car’s fundamental quantum of energy
corresponds to a speed of 10 km/hr. If the car has 7 quanta
of energy, it will have a speed of 70 km/hr. If the car has 9
quanta of energy, it will have a speed of 90 km/hr. This
shows that a car can only move in multiples of 10 km/hr (in
this case). Speeds such as 88 km/hr and 41 km/hr are
impossible.
Light as Particles
Quantized
Continuous
Light as Particles
Important Equation #2: E = h
Planck’s Constant (h) = 6.6262 x 10-34 Js (Joule x
seconds)
Einstein showed their was a wave-particle duality
(meaning wave and particle)
The Hydrogen-Atom LineEmission Spectrum
Ground State – the lowest energy state of an atom
Excited State – state at which an atom has a higher
potential energy than its ground state.
See fig. 4-8 on pg. 96
The Hydrogen-Atom LineEmission Spectrum
When an excited electron returns to the ground
state, it gives off energy in the form of
electromagnetic radiation.
The energy of a photon is equal to the difference
in energy between the atom’s initial state and its
final state.
See fig. 4-5, 4-6 and 4-7 on pg. 95
The Hydrogen-Atom LineEmission Spectrum
Bohr Model of the Hydrogen
Atom
According to the Bohr’s Model, electrons orbit the
nucleus in allowed paths, called orbits. Similar to
planets orbiting the sun.
According to this model, electrons cannot exist
between energy levels.
Bohr Model of the Hydrogen
Atom
The farther the electron is from the nucleus, the
more energy it possesses.
I’ll explain line-spectrum’s more in class.
Bohr’s Model is OK, but it isn’t complete. It is only
good for explaining the hydrogen atom, not multielectron atoms.
Arrangement of
Electrons in Atoms
4-2 The Quantum Model of the Atom
Electrons as Waves
Last section we learned that light can behave as both
a particle and a wave. What about electrons?
Louis De Broglie stated that electrons could be
considered waves confined to a space around an
atomic nucleus.
Electron waves can exist, but only at specific
frequencies corresponding to specific
frequencies.
Electrons as Waves
Experiments showed that electrons (like light)
could be bent, or diffracted. Also, electron beams
could interfere with each other.
Diffraction – bending of light when passed through a
crystal.
Interference – overlapping of waves, reducing energy in
some areas.
Heisenberg Uncertainty
Principle
The position and momentum of a moving object can
not simultaneously be measured and known exactly.
Due to the duel nature of matter and energy
Only important with small scale objects
The Schrödinger Wave
Equation
Erwin Schrödinger developed an equation, which treated
electrons in atoms as waves.
Solutions to wave equation are known as wave
functions.
Don’t worry about wave functions, we do a little more with it
in AP
Coupled with Heisenberg Uncertainty Theory, lead to
Quantum Theory
Quantum Theory – describes mathematically the wave
properties of electrons and other very small particles.
The Schrödinger Wave
Equation
Most Important Idea: We can only know the
probability of finding an electron, not its exact
location.
Orbital – a 3-dimensional region around the
nucleus that indicates the probable location of an
electron.
Fig 4-11
Review
Energy is quantized ( found in specific amounts)
Electrons have wavelike behavior
Impossible to know electron position and
momentum.
Can predict the probability of electron location
Called the Quantum-mechanical model
Probability and Orbital
The density of an electron cloud is called the electron
density.
Higher density – more likely to find electron
Lower density – less likely to find electron
An orbital is the region where a given electron is
likely found.
There are different types of orbitals….s, p, d, f which
we will talk about more later.
Orbitals and Energy
 To describe orbitals, scientists use quantum
numbers.
Quantum Number – specify the properties of atomic
orbitals and the properties of electrons in orbitals.
Principal Quantum
Number
indicates the main energy level occupied by the electron.
Sometimes considered the shell.
n are positive integers (n = 1, n=2, n=3, …)
As n increases, energy and distance from nucleus increases.
n = 1 is the lowest energy level, closest to the
nucleus.
More than one electron can have the same value of n.
The total number of orbitals that exist in a given shell is equal
to n2.
Angular Momentum
Quantum Number (l)
indicates the shape of an orbital
Also considered the sublevel.
The number of orbital shapes possible is equal to n
l can have values of 0 and all positive integers less than
or equal to n-1
If n = 1, l = 0: (l = n – 1 = 1 –1 = 0)
If n = 2, l = 1 and 0: (l = n – 1 = 2 – 1 = 1)
Each orbital is assigned a letter, which corresponds
to a shape
s orbital – see figure 4-25 pg 144
p orbital- see figure 4-26 in book
d orbital – see figure 4-27 in book
Each atomic orbital is designated by the principal
quantum number followed by the letter of the
sublevel.
Ex. 1s sublevel is the s orbital is in the first main energy
level
Ex. 2p sublevel is the set of p orbitals in the second
energy level
Ex. 3d sublevel is the set of d orbitals in the third
energy level
Magnetic Quantum
Number (ml )
indicates the orientation of an orbital around the nucleus
ml = +/- l and every integer in between
Ex. If n = 1, l = 0, ml = 0
This means there is a single s orbital in the first energy level
If n = 2, l = 1, ml = -1, 0, +1
In the second energy level there are three p orbitals
If n = 4, l = 2, ml = -2, -1, 0, +1, +2
In the fourth energy level there are five d orbitals.
If n = 4, l = 0, ml = 0
In the fourth energy level there is 1 s orbital
Spin Quantum Number
(ms)
indicates the spin states of an electron in an orbital, either +1/2,
or –1/2.
o
Electrons spin on an internal axis either clockwise or
counterclockwise.
A single orbital can hold a maximum of two electrons, which must
have opposite spins.
Summary of Energy
Levels, Sublevels, and
Orbitals
Principal Energy
Level
n=1
Sublevels
Orbitals
1s
1s (one)
n=2
2s, 2p
2s (one) + 2p
(three)
n=3
3s, 3p, 3d
3s(one) +
3p(three)+3d(five)
n=4
4s, 4p, 4d, 4f
4s(one)+4p(three)+4d(five)+
4f(seven)
Max Number of
Electrons in Each
Sublevel
Sublevel
# of Orbitals
s
1
Max # of
Electrons
2
p
3
6
d
5
10
f
7
14
Arrangement of
Electrons in Atoms
4-3 Electron Configurations
Electron Configuration
Describes where the electrons are found and what
energies they possess.
They are determined by distributing the atoms
electrons among levels, sublevels, and orbitals based
on a set of principles.
Determining Electron
Configurations
They fill using the Aufbau Principle
Electrons in atoms want to assume the lowest
possible energy.
Ground-state electron configuration
Order Orbitals are filled
* Note 4sfills before 3d,
this is what is found in
nature. The energy
levels get closer
together farther from
nucleus.
Pauli Exclusion
Pauli Exclusion Principle
An orbital can hold a maximum of 2 electrons.
2 electrons in the same orbital must have opposite
spins.
An electron is "paired" if it is sharing an orbital with
another electron with an opposite spin.
An electron is "unpaired" if it is alone in an orbital
Hunds Rule
Electrons occupy equal-energy orbitals so that a
maximum number of unpaired electrons results.
You must put a single electron in each equal-energy
orbital before you begin to pair.
Short Cut to remembering
order….either one works
Sample 1
Sample 2
You try
Write the configurations for A) magnesium and B)
Nickel. How many unpaired electrons does each
possess?
A) Mg: 1s22s22p63s2 No unpaired
B) Ni: 1s22s22p63s23p64s23d8 Two Unpaired
We will learn this chart
later