Transcript Unit 4: Electrons in Atoms

Electrons in Atoms
By: Ms. Buroker
Okay …
We now know that an element’s identity lies
in its number of protons … but there is
another particle which is very important as
well …
The electron
Electrons control behavior and reactivity in
substances.
The Electromagnetic Spectrum
Visible light is a kind of
electromagnetic
radiation … a
form of energy
that exhibits
wavelike
behavior as it
travels through
space.
Electromagnetic Radiation
We measure electromagnetic radiation in
two ways …
1.) Wavelength
2.) Frequency
When multiplied together … these two thing
ALWAYS equal the speed of light!!!
C= 3.00 x 108m/s
Wavelength (l)
Wavelength is the distance from one wave crest or wave
trough to the next
Electromagnetic Spectrum
Wavelength is measured in meters … but more
commonly, nanometers.
Frequency (u)
Frequency is defined as the number of waves that pass
a given point in a specific time … usually, a second.
Typically expressed as a
Hertz (Hz) … one wave
per second
Remember What We Said …
All Wavelengths and their frequency’s
when multiplied are equal to the speed of
light; so …
C = lu
Let’s Re-Visit This …
Violet … shortest wavelength
Red … longest wavelength
and highest frequency
and lowest frequency
The Particle Nature of Light
Matter can loose energy only in small,
specific amounts called quanta … so a
quantum is the minimum amount of
energy that can be gained or lost by an
atom.
Equantum = hn
Energy
Frequency
Plank’s Constant
6.626 X 10-34 J.s
The Photoelectric Effect
Einstein suggested that electromagnetic
radiation can be viewed as a stream of
particles called photons.
Ephoton = hn = hc
l
Photoelectric Effect: refers to
the phenomenon in which
electrons are emitted from the
surface a metal when light
strikes it.
Photoelectric Effect Continued
1.) Each metal has a threshold frequency.
2.) For light with frequency lower than the
threshold, no e- are emitted.
3.) For light with frequency greater than threshold,
the number of e- emitted increases with the
intensity of the light.
4.) For light with frequency greater than the
threshold frequency, the kinetic energy, of the
emitted e- increases linearly with the frequency
of light.
Let’s Practice!!
Tiny water drops in the air disperse the white
light of the sun into a rainbow. What is the
energy of a photon from the violet portion of
the rainbow if it has a frequency of 7.23 x 1014
Hz?
Ephoton = hn
Ephoton = (6.626 x 10-34 J.s)(7.23 x 1014/s)
Ephoton = 4.79 x 10-19J
Atomic Emission Spectra
The set of frequencies of the
electromagnetic waves emitted by atoms
of an element.
* They are unique to the
element!
Quantum Theory and the Atom
Bohr Model of the Atom
* Working with the hydrogen atom, he
proposed that the hydrogen atom has only
certain allowable energy states.
Lowest allowed energy state = ground state
Atoms gain energy = excited state
Bohr Model Continued …
Electrons moving around the
nucleus in only certain allowed
circular orbits
Bohr Model Continued …
Assigned each orbital a quantum
number, n
Smaller the orbital = lower energy
Higher orbitals = higher energy
de Broglie Equation
de Broglie predicted that all moving
particles have wave characteristics.
l= h
mv
Mass
Velocity
Heisenberg Uncertainty Principle
States that it is
impossible to know
the precisely both
the velocity and
position of a particle
at the same time.
Hydrogen’s Atomic Orbitals
n = Principle Quantum Number
Energy sub levels = s, p, d, f
Shape =
s: spherical
p: dumbbell shaped
d,f: don’t all have the
same shape
d orbitals
10 eS
orbitals
2 ep orbitals
6 e-
f orbitals
14 e-
In Summary
Principle
Quantum
Number (n)
Sublevels
(types of orbitals)
Present
Number of
Orbitals related
To sublevel
1
s
1
2
s
p
1
3
3
s
p
d
1
3
5
4
s
p
d
f
1
3
5
7
Electron Configuration
The arrangement of electrons in an atom.
There are three main
rules that govern how
electrons can be
arranged in an atom.
The aufbau principle
e- occupy the lowest
energy level FIRST!
The Pauli exclusion principle
A maximum of two e- may occupy a single
atomic orbital, but only if the electrons
have opposite spins!
orbital
Total # of e-
s
2
p
6
d
10
f
14
Hund’s rule
Single e- with the same spin must occupy
each equal energy orbital before additional
e- with opposite spin can occupy the same
orbital.
Atomic Orbitals and Quantum Numbers
n= The principle quantum # (1, 2, 3, ….)
* relates to the size and
energy of the orbital
l = The angular momentum quantum number
(0 to n-1)
l = 0  s orbital
l = 1  p orbital
l = 2  d orbital
l = 3  f orbital
Atomic Orbitals and Quantum Numbers
Ml = magnetic quantum number (l to –l)
* related to the orientation of the orbital in
space relative to the other orbitals in the
atom.
Ms = Electron spin quantum number
(+1/2 or -1/2)
Valance Electrons
Valence electrons are those electrons that
are in the outer most energy level of an
atom … it is these electrons that are
responsible for the reactivity of an
element.
Example:
Write the electron configuration for Carbon.
1s2 2s2 2p2 … the outer most energy level for Carbon is
n=2, right? So to find out how many valance electrons
carbon has, you simply count how many electrons are
in level 2.
Carbon has 4 valance electrons!
Believe it or not … you can look at the periodic table
and determine how many valence electrons an atom
has and how that affects the properties of that element.
All elements with the exception of the transition metals are
called the representative elements.
The representative elements follow the following rules …
The group number tells you
how many valence electrons
there are!!
The period number tells you
the energy level the valence
electrons are in!!