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Transcript
Arrangement of Electrons in
Atoms
Chapter 4
The New Atomic Model


Investigations  relationship between light
and atom’s electrons
How are electrons arranged? Why don’t
they fall into the nucleus?
Light a wave or particle?

Wave Description:


Electromagnetic Radiation: energy that
acts like a wave in space
All forms create Electromagnetic Spectrum
Electromagnetic Spectrum
Electromagnetic Spectrum


All forms move at speed of light, c,
3.00x108 m/s
Forms identified by:


wavelength, , the distance b/ corresponding
points on adjacent waves. Units: nm, cm, or
m
Frequency, , # of waves that pass a given
point in a specific time, 1 sec. Unit: 1/s =
Hertz, Hz
Wavelength and Frequency
Wavelength and Frequency
c = 
speed of
light, m/s

wavelength,
m
Inverse proportion equation!!
Frequency,
1/s
Calculation


Calculate the wavelength of a radio wave
with a frequency of 102.7 x 106s-1
Determine the frequency of light whose
wavelength is 5.267 nm.
Particle Nature of Light

Photoelectric
Effect:
emission of
electrons from
a metal when
light shines on
the metal
Photoelectric Effect





Light had to be certain frequency to knock
e- loose
Wave theory  any frequency should work
(just might take a while)
Light must also be a particle!
Max Planck(1900) explanation: objects
emit energy in small packets called quanta
Video - 16
Max Planck

Quantum of energy is the smallest amount
of energy that can be lost or gained by an
atom
E = h
Energy of
quantum,
in joules, J
Planck’s
constant,
6.626x10-34 Js
Frequency,
s-1
Energy Calculation

What is the energy of green light, with a
wavelength of 500. nm?
Albert Einstein



Light is both wave and particle!
Particle of light = photon, having zero
mass and a quantum of energy
Photons hit metal and knock e- out, but
photon has to have enough energy
H-atom Emission Spectrum



Pass a current through gas at low pressure
it excites the atoms
Ground state: lowest energy state of an
atom
Excited state: atom has higher potential
energy than it has in ground state
H – Atom Spectrum

When atom jumps from excited state to
ground state it gives off energy  LIGHT!
E2
Ephoton = E2 – E1 = hv
E1
Bohr Model of H-atom
H-atom Line Emission Spectrum
Element Emission Spectras
Helium – 23 lines
Neon – 75 lines
Argon - 159 lines
Xenon – 139 lines
Mercury – 40 lines
H-atom Line Emission Spectrum



More lines in UV (Lyman series) and
IR(Paschen series)
Why did H-atom only emit certain colors
of light?
Explanation led to new atomic theory 
Quantum Theory
Bohr Model of H-atom





1913 – Niels Bohr
e- circles nucleus in certain paths, orbits or
atomic energy levels
e- is higher in energy the farther away from
nucleus
e- cannot be between orbits
Video - 23
Bohr Model of H-atom
Bohr Model of H-atom


From wavelengths of emission spectrum
Bohr calculated energy levels of H-atom
Model worked ONLY for H-atom
Quantum Model of Atom

Can e- behave as a wave?



Yes!
To find e- use a photon, but photon will
knock the e- off course
Heisenberg Uncertainty Principle:
impossible to determine position and
velocity of a particle at the same time.
Schrödinger Wave Equation



1926 – developed equation and only ewaves of certain frequencies were solutions
Quantization of e-  probability of finding
e- in atom
No neat orbits  probability clouds or
orbitals
Electron Configurations
Atomic Orbitals


Def: 3-D region around nucleus that
indicates the probable location of an
electron
Energy levels or shells:


Numbered 1-7
Smaller number = closer to nucleus, lower
energy
Sublevels


Each shell has sublevels
s


p



1 – s orbital
3 – p orbitals
d

5 – d orbitals

7 – f orbitals
f
Shells and Sublevels






Shells and sublevels together:
1s
2s, 2p
3s, 3p, 3d
4s, 4p, 4d, 4f, etc.
s is the lowest energy and f is the highest
Orbitals

Each orbital in a sublevel can
hold a maximum of 2 e



1 – s 2 e- max.
3 – p orbitals 6 e- max.
5 – d orbitals 10 e- max.
7 – f orbitals 14 e- max.
Electron Configurations




Arrangement of e- in atom
Orbital Notation:
H has 1eRules:
1.
Aufbau Principle: electron occupies lowest
energy level that can receive it
Electron Configurations
2. Pauli Exclusion Principle: no two e- in
an sublevel orbital can have the same
spin
3. Hund’s Rule: orbitals of equal energy are
occupied by one e- before pairing up e-. All
single occupied orbitals must have same
spin.

He – 2e-
Energy of sublevels
Electron Configurations




N
S
Ti
I
Electron Configuration Notation



B
Ni
Hg
Noble Gas Notation

Use noble gas from previous row


Al
Pb
Special Cases



d sublevel more stable with half-filled or
completely filled sublevel
Cr
Cu