Download Atomic Structure Electrons in Atoms

ATOMIC STRUCTURE
Electrons in Atoms
• In the early 1900’s scientists studying the behavior of
atoms observed that certain elements emitted visible
light when heated by a flame or if subjected to a high
voltage spark
– Much of our understanding of how electrons behave
in atoms comes from the study of light
– In order to understand this, you must know about the
nature of light…
LIGHT
• Visible light (the light we see with our
eyes) is a type of electromagnetic
radiation
• All other electromagnetic radiation is
invisible
ELECTROMAGNETIC RADIATION
• Electromagnetic (EM) radiation is the
transmission and emission of energy in the form of
electromagnetic waves
– Visible light, X-rays, microwaves , infrared waves (IR),
ultraviolet waves (UV), radio waves, etc.
ELECTROMAGNETIC SPECTRUM
• The electromagnetic spectrum
encompasses all forms of electromagnetic
radiations
CHARACTERISTICS OF WAVES
• A wave can be thought of as a vibrating
disturbance by which energy is transmitted
• Waves can be characterized by their:
– Wavelength  (lambda): distance between
identical points on a successive wave
• nanometers (nm)
– Frequency  (nu): the number of waves that
pass through a particular point in 1 second
• hertz (Hz); 1 Hz = 1 cycle/s
– Amplitude: the vertical distance from the
midline to the peak or trough
High Frequency
Low Frequency
SPEED OF LIGHT
• Another important property of waves is their
speed
– dependent on type of wave and medium
it’s traveling through
• The speed of a wave is a product of its
wavelength and frequency:
c = 
• Notice  and  are inversely
proportional
• In a vacuum, All electromagnetic waves
travel at the speed of light (3.00 x 108
m/s)
PLANCK’S QUANTUM THEORY
• According to the theory atoms and molecules can
emit or absorb energy only in discreet quantities,
which Planck called “quantum”
• quantum (meaning “fixed amount”) is the smallest
quantity of energy that can be gained or lost by an
atom
• The energy E of a single quantum of energy is given
by
E =h
where h is called Planck’s constant (6.63 x 10-34
J•s) and  is the frequency of radiation
• According to Planck’s theory energy is always
released or absorbed by matter in whole-numbers of
THE PHOTOELECTRIC EFFECT
• Albert Einstein used
Planck’s quantum theory to
explain the photoelectric
effect
• When light shines on a
clean metal surface the
surface emits electrons (can
be converted to electrical
energy)
• For each metal there is a
minimum frequency of light
needed to emit electrons
– red light in incapable of
releasing electrons from
sodium metal even if intense
– faint violet light releases
electrons easily
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DUAL NATURE OF LIGHT
• In explaining the photoelectric effect Einstein proposed
that light (electromagnetic radiation) has dual nature:
– Can behave like a wave
– Can behave like a stream of particles (photons)
• PHOTON - a particle of electromagnetic radiation
– A photon has no mass
– A photon carries a quantum of energy
• Einstein calculated that a photon’s energy
depends on its frequency
Ephoton = h
ATOMIC EMISSION SPECTRA
• Light of a neon sign is produced by
passing electricity through a tube of
neon gas
• The atoms in the tube absorb energy
and become excited and unstable
• They become stable by releasing the
energy as light
ATOMIC EMISSION SPECTRA
• If the light emitted from neon is passed
through a prism neon’s atomic
emission spectrum is produced
ATOMIC EMISSION SPECTRA
• The atomic emission spectrum of an element
is its “fingerprint”
• It’s the set of frequencies EM waves the
element’s atoms emit (give off)
• Each element’s atomic emission spectrum is
unique and can be used to identify the element
– Hydrogen’s emission spectrum:
violet
410 blue-violet
nm
434 nm
blue-green
486.1 nm
red
656.2 nm
Light Bulb
(white light)
Hydrogen Bulb
THE BOHR MODEL OF THE ATOM
• Bohr proposed:
– Electrons move around the nucleus in circular
orbits (“rings”) with distinct energy levels
• smaller orbits have lower energy, larger orbits
higher energy
– In other words, electrons found closer to the
nucleus has less energy than electrons found
at greater distances from the nucleus
• Bohr assigned a quantum number (n) to each
level
BOHR’S MODEL OF THE ATOM
Energy
n=7
+
n=6
-
n=5
n=4
n=3
n=2
n=1
This model is often called the planetary model
BOHR’S ATOM CONTINUED
• The lowest energy state of an atom is its
ground state
• When an atom gains energy (through heating
for example) it is in an excited state
– in an excited state the electron absorbs the
energy & jumps to higher energy level when it
jumps back down to its ground state it releases
excess energy in the form of light
• Even though hydrogen contains only one
electron, it can have many excited states
BOHR MODEL CONTINUED
• Because electrons jump between orbitals that have specific
energy levels only certain frequencies of electromagnetic
radiation can be given off (only certain colors can be
emitted):
• If an excited electron drops from n=3 to n=2 red light is
emitted
• If an excited electron drops from n=4 to n=2 blue-green
light is emitted
• 5-2: blue light
• 6-2: violet light
• This is how Bohr explained hydrogen’s emission spectrum
E = h
E = h
Wait!
• Bohr’s model explained the emission spectrum of
Hydrogen, but it did not explain the emissions of any
other element!
• It was eventually found that Bohr was incorrect:
– Electrons do not travel in circular orbits around the
nucleus
• In 1924, a young French scientist Louis de Broglie
(1892-1987) proposed an idea that eventually
accounted for the fixed energy levels of Bohr’s model
and better explained the behavior of electrons
THE QUANTUM MECHANICAL MODEL
OF THE ATOM
• If waves can act like particles, particles can act as
waves!
• Electrons behave like waves
• Can’t know electrons position/path around the nucleus:
– Electrons move about in a cloud around the nucleus in what
appears to be a random pattern
• The Quantum Model only predicts where an electron is
likely to be found
HEISENBERG UNCERTAINTY
PRINCIPLE
• According to this principle it is fundamentally
impossible to know the exact position and velocity
of a particle at the same time
• locating an electron produces uncertainty in the
position and motion of the electron
• It’s like trying to locate a helium filled balloon in a
completely darkened room:
– If you locate it by touch, you change it’s
position-Once you find it, it’s already
somewhere else!
SO WHERE CAN ELECTRONS BE FOUND?
• In the quantum model, the nucleus is not
surrounded by orbits, but by atomic orbitals
• Atomic Orbital: a three-dimensional pocket
of space around the nucleus that the electron
is most likely to be found
– An electron has a 90% chance of being
found in the atomic orbital
– That is the best we can do!
Electron probability
density for hydrogen
Where 90% of the
e- density is found
for the 1s orbital
e- density (1s orbital) falls off rapidly
as distance from nucleus increases
ATOMIC ORBITAL ORGANIZATION
1. Principal energy level (n: 1-7)
• (n) indicates relative size and energy of orbital. As n
increases so do energy and size
2. Energy sublevels (s, p, d, f)
– sublevels labeled according to shape
– s: spherical; p: dumbbell; d/f: varied
3. Orbitals: Each sublevel has a specific number of
orbitals:
–
–
–
–
s: 1orbital
p: 3 orbitals
d: 5 orbitals
f: 7 orbitals Each orbital can hold two electrons!!!
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increasing
energy
Each orbital can hold two electrons
4d
5s
4p
Energy
3d
4s
3p
3s
2p
2s
1s


  
Hydrogen
Beryllium
Nitrogen
Oxygen
Fluorine
Helium
Lithium
Boron
Carbon
Neon
2
1
3
4
5
6
7
8
10
9
He
Ne
Be
Li
O
H
C
N
B
F
4.003
1.008
6.941
9.012
10.81
12.01
14.01
16.00
20.18
19.00
ORDER OF ORBITALS (FILLING) IN MULTIELECTRON ATOM
ELECTRON CONFIGURATION
• An atoms electron configuration is the way an
atom’s electrons are distributed among the
orbitals of an atom
• The most state stable electron configuration
is an atom’s ground state
– Ground state: all electrons are in the lowest
possible energy state
• Electron configuration represented by writing
symbol for the orbital and a superscript to
indicate the number of electrons in the orbital
Li: 1s2 2s1
The Pauli Exclusion Principle
• The two electrons in an orbital must
spin in opposite directions 


  

  
1s
2s
2p
3s
3p
4s
paramagnetic
diamagnetic
unpaired electrons paired electrons
3d
HUND’S RULE
• Negatively charged electrons repel each other, so:
– Electrons won’t pair up unless they have to
– Once there is one electron in every orbital…the
pairing will begin!
1.
  
1s 2s
Add an electron:
2.
2p
   
1s 2s
2p
Add an electron:
3.
    
1s 2s
Add an electron:
4.
2p
    
1s 2s
2p
ELECTRON CONFIGURATION
• The periodic table can be divided into four distinct
blocks based on valence electron configuration
• electron configuration explain the recurrence of
physical and chemical properties
SHORTHAND (NOBLE GAS)
NOTATION
• Shows electron filling starting from
previous noble gas:
– Na: 1s22s22p63s1
– Noble gas configuration: [Ne]3s1
Electron Configuration for
Cation and Anions
• What is the electron configuration for a
sodium atom?
– Na: 1s22s22p63s1
• What is the electron configuration for a
sodium ion (Na+)?
– Na+: 1s22s22p6 or [Ne]
TRANSITION CATIONS
• When transition metals form cations
electrons are removed from the s-orbital
before the d-orbital
– Mn: [Ar]4s2 3d3
– Mn2+: [Ar]3d3
• This happens because d-orbitals are
more stable than s-orbitals in transition
metals
• The s-orbitals are always in a higher
energy level
ISOELECTRONIC
• What do you notice about the following
electron configurations:
– F- : 1s22s22p6 or [Ne]
– O2- : 1s22s22p6 or [Ne]
– N3- : 1s22s22p6 or [Ne]
• They have the same number of
electrons and therefore identical
ground-state electron configurations
– These three ions are isoelectronic
VALENCE ELECTRONS
• Valence electrons are the electrons
found in the outermost orbitals of the
atom (highest energy level)
• They are the electrons involved in
bonding and are responsible for the
chemical properties of an element
• Carbon has 4 valence electrons: 1s2 2s2
2p2
• How many does magnesium have?
VALENCE ELECTRONS & GROUP
NUMBER
• One of the most important relationships
in chemistry:
– Atoms in the same group have similar
chemical properties because they have the
same number of valence electrons!
• For the representative elements (group
A elements)
– group # = number of valence electrons
– Exceptions: He is in group 8 but only has 2
valence electrons
The octet/duet rule
• Atoms will gain, lose or share electrons
to achieve noble gas configuration,
meaning all atoms want a full outer
orbital:
– 2 valence electrons for He
– 8 valence electrons for all other noble
gases
ELECTRON-DOT STRUCTURE
• Chemists often represent valence
electrons in electron-dot structures
• Electron-Dot Structure consists of the
element’s symbol surrounded by dots
representing the atom’s valence electrons
– valence electrons are placed one to each
of the four sides first,
– when each side has one dot, you may
begin doubling up
S