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WHAT ARE ATOMS? • The smallest units of matter.
• Make up elements
• Contain subatomic particles:
• A nucleus containing protons and
neutrons, and surrounded by electrons
• The number of protons and electrons is
unique for every element.
What’s in an atom?
Subatomic
Particles
Particles Charge Mass (kg) Location in
atom
Proton +1 1.67 x 10 -27 In Nucleus
Neutron 0 1.67 x 10 -27
In Nucleus
Electron -1 9.11 x 10 -31 Moving
around
outside of the
nucleus
What’s in an atom?
THE PERIODIC TABLE
• Horizontal (------) rows are called
periods.
• # of protons and electrons increases
by one as you move from left to right
across a period.
• Vertical (up and down) columns are
called groups or families.
• Atoms in the same group usually have
the same # of valence electrons
THE PERIODIC TABLE
• Atomic # (Z)=the number of protons and
electrons in a neutral atom
• Mass # (A)= the number of protons + the
number of neutrons
• # of Neutrons=A-Z
• # of Electrons= Z
• # of Protons= Z
FORMING IONS
• An ion is formed when an atom gains or
loses one or more electrons
• Cation(+)=When an atom LOSES one
electron, it gets a “+1” charge
• Anion(-)=When an atom GAINS an
electron, it gets a “-1” charge
HISTORICAL
MODELS • 4th century B.C.—Democritus
suggested that atoms existed as indivisible particles.
• 1808—John Dalton proposed the atomic theory
• 1897—Thomson—charge to mass ratio of electrons—plum pudding model
• 1911—Rutherford—nuclear model with electrons outside of the nucleus
DALTON’S THEORY
• 1. All matter is composed of indivisible
particles called atoms.
• 2. All atoms of a given element are
identical. Atoms of different elements have
different properties.
• 3. Chemical reactions involve the
combination of atoms, not the destruction of
atoms
• 4. When atoms react to form compounds,
they react in defined, whole-number ratios
CATHODE RAY
EXPERIMENT
Applying voltage causes a glow
to travel from the negative to
positive end glass tube in which
a partial vacuum exists.
What can we learn from a ray??
CATHODE RAY
OBSERVATIONS AND
CONCLUSIONS
• Ray is deflected by a magnetic
field.
• Has electromagnetic properties
• Ray travels away from negative
and toward positive
• Must be negatively charged
CATHODE RAY
OBSERVATIONS AND
CONCLUSIONS
• Any metal will produce a ray.
• All atoms must contain these particles.
• Atoms are electrically neutral.
• Some positive particle must be present to balance the negative charge.
CHARGE VS. MASS
• Thomson was able to calculate
the charge to mass ratio = 1.7 x
1011
• VERY LARGE charge for a
very small mass
NEW MODEL
• Plum pudding—atoms are spherical masses of positive charge with electrons scattered throughout
• Since electrons have a small mass, the mass must come from something else.
RUTHERFORD • 1911 –”Planetary Model”—Nuclear
atom with a nucleus at the center
• Suggested that electrons orbited the
nucleus like planets orbiting the sun
• Named protons and also suggested
the neutrons existed
MOSELEY
• 1914 – studied x-rays to provide
evidence that the atomic number of
elements is the same as the number
of protons in that element
BOHR
• 1922 –The Bohr Model of the atom
• Electrons only existed in specific energy
states
• They (e-) emit or absorb light at
specific wavelengths
• This accounted for line spectra
NEW MODEL
• Modern Theory—Electrons
orbit the nucleus like waves on
a vibrating string.
• Based on observations of
absorption and emission of
electromagnetic radiation by
atoms
EM RADIATION
• Wave-like properties
• Wavelength—l—the distance between two corresponding points on successive waves
• Velocity—c—speed of light = 3.00 x 108 m/s
EM RADIATION
• Frequency—v—the number
of waves that pass a given
point in one second
• For any wave: c = lv
PROPORTIONALITY
• Since speed is constant (speed of light), frequency and wavelength vary inversely
• Frequency increases; wavelength decreases
• Frequency decreases; wavelength increases
PROPORTIONALITY
• Energy and wavelength vary inversely
• wavelength increases; energy decreases
• Energy and frequency vary directly
• Frequency decreases; Energy decreases
WHICH HAS HIGHER
ENERGY?
1. Wavelength of 400 nm or 700 nm?
2. Frequency of 108 Hz or 1018 Hz?
3. Wavelength of 1 m or 1 cm?
4. Frequency of 106 Hz or 1012 Hz?
CALCULATIONS
• What is the wavelength of EM
radiation with a frequency of
1016 Hz? (Hz= 1/s)
• What type of radiation is it?
(See electromagnetic
spectrum.)
PRACTICE
QUESTIONS
5. What wavelength of
electromagnetic radiation has
a frequency of 3.27 x 104 Hz?
6. What type of radiation is
this? (See electromagnetic
spectrum.)
CALCULATIONS
• What is the frequency of light with
a wavelength of 400 nm?
(400 nm = 4 x 10-7 m)
• What color is this light?
PARTICLE BEHAVIOR OF
LIGHT • Max Planck
• Objects do not emit light at continuous frequencies; rather the energy comes out in specific amounts or quanta
• Quantum—the minimum amount of energy an atom can gain or lose
ENERGY OF A
QUANTUM
• E = hv • E = energy (J)
• h = Planck’s Constant = 6.626 x 10-34 J s
• v = frequency (1/s)
CALCULATIONS
• What is the energy of a photon
of electromagnetic radiation with
a frequency of 1.56 x 1012 Hz?
CALCULATIONS
• What is the energy of a photon
of electromagnetic radiation with
a wavelength of 4.15 x 10-7 m?
PHOTOELECTRIC
EFFECT
• Emission of electrons from a metal when light shines on the metal
• If light is a wave, it should have enough energy to remove electrons at any frequency, but only certain frequencies of light will remove electrons.
EINSTEIN’S
CONTRIBUTION
• Light behaves as a particle and as a wave—wave particle duality
• Photon—a particle of light that carries a quantum of energy
• Energy of photon=energy of quantum—use same equation
EXCITED
ELECTRONS
• Electrons that have absorbed energy become excited and move to higher energy states in an atom
• When they return to lower energy states, they emit the energy
ATOMIC EMISSION
SPECTRA
• H always and only emits
radiation at specific
wavelengths
• Other elements behave the
same way, but patterns are
more complex due to larger # of
electrons
LINE EMISSION
SPECTRA
• Very specific—Act like
fingerprints by which an
element can be identified
***Emission Tubes/
Spectrometers***
DO THE MATH
• A mathematical relationship
exists between the
wavelengths and the emission
spectra of hydrogen.
• The spectral emission of
elements is always the same.
BOHR MODEL • Energy of
emission equals the difference in energy levels
• Added orbits for electrons to the existing atomic model
http://www.robotplatform
.com/knowledge/Atomic
%20Theory/atom-
bohr_en.jpg
ENERGY LEVELS
• Electrons may only go to certain
places (energy levels) within the
atom.
• Evidence: wavelengths in
emission spectra of atoms
LIMITATIONS OF BOHR’S
MODEL
• Since Hydrogen has only
one electron, Bohr’s model
worked well for it.
• The model did not work as
well when applied to atoms
with more electrons.
DRAWING BOHR
MODELS • Only certain numbers of
electrons are allowed in each level
• Level 1: 2
• Level 2: 8
• Level 3: 18
• Level 4: 32
S orbitals of a
hydrogen
atom—
electrons can
transition
between levels,
but cannot stay
in nodes http://www.avon-chemistry.com/electron_lecture.html
DON’T WRITE THIS!! It requires the same amount of
energy every time you
• walk from your seat to the café.
• walk from your seat to your locker.
• walk from your seat to the door.
If these were the only three places
you could go, you would only use
energy in those amounts.
ASSIGNMENT
• Draw Bohr Models for the 18
elements on page 13 of your
packet and answer the two
following questions.
THE QUANTUM
MODEL
• De Broglie: electrons have properties of both waves & particles—interference & diffraction.
• Wave-particle duality
• Waves could only exist with wavelengths that corresponded to Bohr’s orbits
HEISENBERG
UNCERTAINTY
PRINCIPLE
• If electrons are like both waves and particles, where are they in an atom?
• Electrons are located by their interaction with photons, but that interaction knocks them off course—as soon as you locate it, it’s somewhere else!
HEISENBERG
UNCERTAINTY
PRINCIPLE
It is impossible to know both the
position and velocity of an
electron or any other particle at
the same time.
SCHRODINGER’S WAVE
EQUATION
• Only certain frequencies of waves
solve the equation
• These frequencies correspond to
quanta of energy & Bohr’s orbitals
The Shrodinger equation is:
SCHRODINGER’S
WAVE EQUATION
• Schrodinger’s wave function describes a probable space in which electrons can be found
• Orbital—a 3-dimensional region around the nucleus that indicates the probable location of an electron
• An electron configuration is a way of showing
the arrangement of electrons around the
nucleus.
• The first # indicates the energy level
• The letter indicates the type of orbital: s, p,
d, f
• A superscript indicates the # of electrons in
the orbital.
• Ex: 1s2 = 1st energy level, ‘s’ orbital, 2
electrons
• Electrons occupy the lowest energy orbital first
• An atom will gain or lose electrons to get 8
electrons in the outer shell (the “Octet” rule)
• ‘s’ and ‘p’ orbitals are the outer shell
LOCATION ON THE PERIODIC
TABLE ALSO TELLS ELECTRON
CONFIGURATION
http://chemed.chem.wisc.edu/chempaths/GenChem-
Textbook/Electron-Configurations-and-the-Periodic-Table-
564.html
DIAGONAL RULE This diagram
shows the order
in which
electrons fill
orbitals. Begin
with 1s and
follow the
diagonal lines. http://www.wyzant.com/Help/Sc
ience/Chemistry/Electron_Confi
guration/
SPECIAL
DIRECTIONS
• “D” orbital: The energy level for
the d-orbital is always one less
than the row #
• Ex: Zinc ( )
• “F” orbital: The energy level for
the f-orbital is always two less
than the row #
• Ex: Cerium ( )
EXCEPTIONS
Cr
we would predict:
1s2 2s2 2p6 3s2 3p6
4s2 3d4
But it is actually:
1s2 2s2 2p6 3s2 3p6
4s13d5
Cu
we would predict:
1s2 2s2 2p6 3s2 3p6 4s2
3d9
But it is actually:
1s2 2s2 2p6 3s2 3p6 4s1
3d10
NOBLE-GAS
NOTATION • AKA—abbreviated ground state
• Noble gases=elements in group 18
• Electron Configurations of Noble gases
which have eight electrons in their
outermost orbitals
• Uses brackets: [Ne] or [He]
• [Ne]=1s22s22p6
• Can be used to write shorthand versions of
other electron configurations:
• Ex: S=
ORBITAL
SHAPES
http://chemwiki.ucdavis.edu/Physical_Chemistry/Quantum
_Mechanics/Atomic_Theory/Electrons_in_Atoms/Electroni
c_Orbitals
ORBITALS • Each orbital holds 2 electrons
• Different numbers of orbitals for each sublevel
• 1 s orbital—holds 2 __
• 3 p orbitals—hold 6 __ __ __
• 5 d orbitals—hold 10 __ __ __ __ __
• 7 f orbitals—hold 14 __ __ __ __ __ __ __
THE AUFBAU
PRINCIPLE
• Each electron occupies the
lowest energy orbital
available
•Ex: An electron will be in
2s before 2p and 4s
before 6s
HUND’S RULE • Single electrons will occupy all
orbitals of the same sublevel
before pairing.
• All single electrons must have
the same spin.
PAULI EXCLUSION
PRINCIPLE
No two electrons in a single
atom can have the same set
of four quantum numbers.
Two electrons in a single
orbital must have opposite
spins
ORBITAL NOTATION
(ORBITAL DIAGRAM)
• Shows electrons in orbitals
• __ __ __ = unoccupied orbitals
• __ __ __ = orbitals occupied with
electrons of the same spin
• __ __ __ = orbitals occupied with
electrons with different spins