Download Atomic structure BV

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Dalton’s Atomic Theory (1808)
• All matter is made of atoms. Atoms are indivisible and
indestructible*.
Lecture 1
The structure of atom
• All atoms of a given element are identical in mass* and
properties*
• Compounds are formed by a combination of two or more
different kinds of atoms.
• A chemical reaction is a rearrangement of atoms.
*not exactly! But pretty good for 1808
Atom
Electron
•  nucleus very small and dense, contains most of the mass, electrons circle nucleus, they
have very little mass but take up most of the space, electrons 1/100,000 the diameter of
the nucleus. The equivalent of a taw (2 cm ) on in the middle of the stadium compared to
the whole of Giant Stadium at the Meadowlands…
•  negatively charged subatomic particle located
outside the nucleus.
Proton
•  positively charged subatomic particle located in
the nucleus.
Neutron
•  neutral subatomic particle located in the
nucleus.
•  Um… by the way, how do we know this?
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(1908 Nobel Prize in Chemistry)
α  particle velocity ~ 1.4 x 107 m/s
(~5% speed of light)
1.  atoms positive charge is concentrated in the nucleus
2.  proton (p) has opposite (+) charge of electron (-)
3.  mass of p is 1840 x mass of e- (1.67 x 10-24 g)
2.2
2.2
Subatomic Particles (Table 2.1)
Rutherford’s Model of the Atom
atomic radius ~ 100 pm = 1 to 2 x 10-10 m
nuclear radius ~ 1 to 20 x 10-15 m
mass p = mass n = 1840 x mass e2.2
Atomic number (Z) = number of protons in nucleus
Mass number (A) = number of protons + number of neutrons
= atomic number (Z) + number of neutrons
Isotopes are atoms of the same element (X) with different numbers of
neutrons in their nuclei
Mass Number
A
ZX
Atomic Number
1
1H
235
92
2
1H
U
Element Symbol
(D)
238
92
3
1H
(T)
U
2.3
2.3
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A molecule is an aggregate of two or more atoms in a definite
arrangement held together by chemical bonds
Do You Understand Isotopes?
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How many protons, neutrons, and electrons are in
C?
H2
6 protons, 8 (14 - 6) neutrons, 6 electrons
H2O
NH3
CH4
A diatomic molecule contains only two atoms
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6
How many protons, neutrons, and electrons are in
C?
6 protons, 5 (11 - 6) neutrons, 6 electrons
H2, N2, O2, Br2, HCl, CO
A polyatomic molecule contains more than two atoms
O3, H2O, NH3, CH4
2.3
An ion is an atom, or group of atoms, that has a net positive or
negative charge.
A monatomic ion contains only one atom
cation – ion with a positive charge
If a neutral atom loses one or more electrons
it becomes a cation.
Na
11 protons
11 electrons
2.5
Na+, Cl-, Ca2+, O2-, Al3+, N3-
11 protons
10 electrons
Na+
A polyatomic ion contains more than one atom
anion – ion with a negative charge
If a neutral atom gains one or more electrons
it becomes an anion.
Cl
17 protons
17 electrons
OH-, CN-, NH4+, NO3-
17 protons
18 electrons
Cl-
2.5
Do You Understand Ions?
How many protons and electrons are in
2.5
The black body radiation
"Blackbody radiation" or "cavity
radiation" refers to an object or
system which absorbs all radiation
incident upon it and re-radiates
energy which is characteristic of
this radiating system only, not
dependent upon the type of
radiation which is incident upon it.
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3+ ?
13 Al
13 protons, 10 (13 – 3) electrons
How many protons and electrons are in
78 Se 2- ?
34
The black body radiation curve
(Fig1) shows that the black body
does radiate energy at every
wavelength.
34 protons, 36 (34 + 2) electrons
2.5
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The Photoelectric Effect
The black body radiation
•  it was observed that many metals emit electrons when a
Wien's displacement law
light shines on their surface
 this is called the Photoelectric Effect
•  classic wave theory attributed this effect to the light
energy being transferred to the electron
•  according to this theory, if the wavelength of light is
Stefan–Boltzmann law
made shorter, or the light waves intensity made
brighter, more electrons should be ejected
 remember: the energy of a wave is directly proportional to its
amplitude and its frequency
 if a dim light was used there would be a lag time before
electrons were emitted
 to give the electrons time to absorb enough energy
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The Photoelectric Effect
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The Photoelectric Effect
The Problem
•  in experiments with the photoelectric effect, it
was observed that there was a maximum
wavelength for electrons to be emitted
 called the threshold frequency
 regardless of the intensity
•  it was also observed that high frequency light
with a dim source caused electron emission
without any lag time
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Einstein’s Explanation
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Ejected Electrons
•  Einstein proposed that the light energy was
delivered to the atoms in packets, called quanta
or photons
•  the energy of a photon of light was directly
proportional to its frequency
 inversely proportional to it wavelength
 the proportionality constant is called Planck’s
Constant, (h) and has the value 6.626 x 10-34 J·s
Tro, Chemistry: A Molecular Approach
Tro, Chemistry: A Molecular Approach
•  1 photon at the threshold frequency has just
enough energy for an electron to escape the atom
 binding energy, φ
•  for higher frequencies, the electron absorbs more
energy than is necessary to escape
•  this excess energy becomes kinetic energy of the
ejected electron
Kinetic Energy = Ephoton – Ebinding
KE = hν - φ
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Spectra
•  when atoms or molecules absorb energy, that energy is
Emission Spectra
often released as light energy
 fireworks, neon lights, etc.
•  when that light is passed through a prism, a pattern is
seen that is unique to that type of atom or molecule –
the pattern is called an emission spectrum
 non-continuous
 can be used to identify the material
 flame tests
•  Rydberg analyzed the spectrum of hydrogen and found
that it could be described with an equation that
involved an inverse square of integers
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Tro, Chemistry: A Molecular Approach
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Examples of Spectra
Exciting Gas Atoms to Emit Light
with Electrical Energy
Oxygen spectrum
Hg
He
H
Neon spectrum
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Identifying Elements with
Flame Tests
Na
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K
Li
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Emission vs. Absorption Spectra
Spectra of Mercury
Ba
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Bohr Model of H Atoms
Bohr’s Model
•  Neils Bohr proposed that the electrons could only have
very specific amounts of energy
 fixed amounts = quantized
•  the electrons traveled in orbits that were a fixed
distance from the nucleus
 stationary states
 therefore the energy of the electron was proportional the
distance the orbital was from the nucleus
•  electrons emitted radiation when they “jumped” from
an orbit with higher energy down to an orbit with lower
energy
 the distance between the orbits determined the energy of the
photon of light produced
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Electron Diffraction
Wave Behavior of Electrons
•  de Broglie proposed that particles could have wave-like
character
•  because it is so small, the wave character of electrons is
significant
•  electron beams shot at slits show an interference
pattern
 the electron interferes with its own wave
•  de Broglie predicted that the wavelength of a particle
was inversely proportional to its momentum
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Uncertainty Principle
Complimentary Properties
•  when you try to observe the wave nature of the
•  Heisenberg stated that the product of the uncertainties
electron, you cannot observe its particle nature –
and visa versa
in both the position and speed of a particle was
inversely proportional to its mass
 x = position, Δx = uncertainty in position
 v = velocity, Δv = uncertainty in velocity
 m = mass
 wave nature = interference pattern
 particle nature = position, which slit it is passing
through
•  the means that the more accurately you know the
•  the wave and particle nature of nature of the
position of a small particle, like an electron, the less
you know about its speed
electron are complimentary properties
 as you know more about one you know less about
the other
Tro, Chemistry: A Molecular Approach
Tro, Chemistry: A Molecular Approach
however, electrons actually
if electrons behave like
present an interference
particles, there should
pattern, demonstrating the
only be like
two waves
bright spots
behave
on the target
 and visa-versa
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The Franck-Hertz experiment
The Franck-Hertz experiment
•  The electrons were accelerated by an electric field (Figure 4.) between the
cathode (C) and the grid (G) in a tube filled with Hg vapour. After
passing through the grid, an opposite field slowed down the electrons and
prevented them from reaching the anode (A), unless they gained enough
kinetic energy in the previous acceleration. The electrons reaching the
anode formed a measurable current, I.
Tro, Chemistry: A Molecular Approach
•  The Hg atoms can not absorb any energy, just well defined values, namely
4.9 eV. This energy excites the ground state electron to the first excited state,
e.g. it equals exactly to the energy difference between the E1 and E2 states of
the Hg atom. Thus, this experiment gives an excellent proof of the validity of
Bohr’s theory.
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Tro, Chemistry: A Molecular Approach
orbital
n
Wave Function, ψ
l
3 quantum numbers
“address”
ml
magnetic
•  calculations show that the size, shape and
orientation in space of an orbital are determined
be three integer terms in the wave function
 added to quantize the energy of the electron
-l, …, l
orientation
angular momentum
•  these integers are called quantum numbers
0, 1, 2, …, (n - 1)
shape
 principal quantum number, n
 angular momentum quantum number, l
 magnetic quantum number, ml
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requires
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Principal
1, 2, 3, …
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Principal Quantum Number, n
size and energy
Principal Energy Levels in Hydrogen
•  characterizes the energy of the electron in a particular
orbital
 corresponds to Bohr’s energy level
•  n can be any integer ≥ 1
•  the larger the value of n, the more energy the orbital has
•  energies are defined as being negative
 an electron would have E = 0 when it just escapes the atom
•  the larger the value of n, the larger the orbital
•  as n gets larger, the amount of energy between orbitals
gets smaller
for an electron in H
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• 
ψ2
Probability & Radial Distribution
Functions
Probability Density Function
is the probability density
 the probability of finding an electron at a particular point in
space
 for s orbital maximum at the nucleus?
 decreases as you move away from the nucleus
•  the Radial Distribution function represents the total
probability at a certain distance from the nucleus
 maximum at most probable radius
•  nodes in the functions are where the probability drops to 0
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Radial Distribution Function
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The Shapes of Atomic Orbitals
•  the l quantum number primarily determines the
shape of the orbital
•  l can have integer values from 0 to (n – 1)
•  each value of l is called by a particular letter that
designates the shape of the orbital
 s orbitals are spherical
 p orbitals are like two balloons tied at the knots
 d orbitals are mainly like 4 balloons tied at the knot
 f orbitals are mainly like 8 balloons tied at the knot
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Tro, Chemistry: A Molecular Approach
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2s and 3s
l = 0, the s orbital
2s
n = 2,
l=0
•  each principal energy state
has 1 s orbital
3s
n = 3,
l=0
•  lowest energy orbital in a
principal energy state
•  spherical
•  number of nodes = (n – 1)
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p orbitals
l = 1, p orbitals
•  each principal energy state above n = 1 has 3 p orbitals
 ml = -1, 0, +1
•  each of the 3 orbitals point along a different axis
 px, py, pz
•  2nd lowest energy orbitals in a principal energy state
•  two-lobed
•  node at the nucleus, total of n nodes
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d orbitals
l = 2, d orbitals
•  each principal energy state above n = 2 has 5 d orbitals
 ml = -2, -1, 0, +1, +2
•  4 of the 5 orbitals are aligned in a different plane
 the fifth is aligned with the z axis, dz squared
 dxy, dyz, dxz, dx squared – y squared
•  3rd lowest energy orbitals in a principal energy state
•  mainly 4-lobed
 one is two-lobed with a toroid
•  planar nodes
 higher principal levels also have spherical nodes
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f orbitals
l = 3, f orbitals
•  each principal energy state above n = 3 has 7 f orbitals
 ml = -3, -2, -1, 0, +1, +2, +3
•  4th lowest energy orbitals in a principal energy state
•  mainly 8-lobed
 some 2-lobed with a toroid
•  planar nodes
 higher principal levels also have spherical nodes
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Stern-Gerlach experiment
Particles possess intrinsic angular momentum.
Spin angular momentum is quantized (it can only take on
discrete values)
For the completely filled shells,
subshell (4d10) the orbital magnetic
momentum is zero; for the 5s orbital
ML is also zero.
Hypothesis: the argent atom possesses
no magnetic momentum >> they move
in an inhomogeneous magnetic field
along the same path.
The electron possesses an intrinsic angular momentum and an
intrinsic magnetic momentum: spin
S — intrinsic angular momentum
Ms — intrinsic magnetic momentum
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