Download Atomic Structure Electrons in Atoms

Survey
yes no Was this document useful for you?
   Thank you for your participation!

* Your assessment is very important for improving the work of artificial intelligence, which forms the content of this project

Document related concepts

X-ray fluorescence wikipedia, lookup

Astronomical spectroscopy wikipedia, lookup

Reflection high-energy electron diffraction wikipedia, lookup

Ultrafast laser spectroscopy wikipedia, lookup

Rutherford backscattering spectrometry wikipedia, lookup

Mössbauer spectroscopy wikipedia, lookup

Electron scattering wikipedia, lookup

Atomic orbital wikipedia, lookup

Electron configuration wikipedia, lookup

Photoelectric effect wikipedia, lookup

Ion wikipedia, lookup

Marcus theory wikipedia, lookup

Ionization wikipedia, lookup

X-ray photoelectron spectroscopy wikipedia, lookup

Chemical bond wikipedia, lookup

Metastable inner-shell molecular state wikipedia, lookup

Molecular orbital wikipedia, lookup

Heat transfer physics wikipedia, lookup

Bremsstrahlung wikipedia, lookup

Auger electron spectroscopy wikipedia, lookup

Degenerate matter wikipedia, lookup

Transcript
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
QuickTime™ and a
decompressor
are needed to see this picture.
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!!!
QuickTime™ and a
TIFF (Uncompressed) decompressor
are needed to see this picture.
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