Transcript Chapter 6 - Olmsted - Seton Hall University Pirate Server

Chapter 6
Characteristics of Atoms
Department of Chemistry
and Biochemistry
Seton Hall University
Characteristics of Atoms
• Atoms posses mass
– most of this mass is in the nucleus
• Atoms contain positive nuclei
• Atoms contain electrons
• Atoms occupy volume
– electrons repel each other, so no other
atom can penetrate the volume
occupied by an atom
• Atoms have various properties
– arises from differing numbers of
protons and electrons
• Atoms attract one another
– they condense into liquids and solids
• Atoms can combine with one
another
2
Wave aspects of Light
• Most useful tool for studying
the structure of atoms is
electromagnetic radiation
• Light is one form of that
radiation
• Light is characterized by the
following properties:
– frequency, , nu
– wavelength, , lambda
– amplitude
3
Electric and magnetic
field components of
plane polarized light
• Light travels in z-direction
• Electric and magnetic fields travel at
90° to each other at speed of light in
particular medium
• c (= 3 × 1010 cm s-1) in a vacuum
4
Connections between
wavelength and
frequency
• c = 3108 m/s in a vacuum
• make sure the units all agree!
  c
5
Characterization of
Radiation
λ, υ, υ or energy
hc
hc
E
or λ 
λ
E
erg sec
h  6.626 10
molecule
1
c(cm
sec
)
-1
υ(sec ) 
λ(cm)
1
υ
λ(cm)
 27
hc
ΔE(erg molecule )  hυ 
 hc υ
λ
6
1
Wavelength and Energy
Units
• Wavelength
– 1 cm = 108 Å = 107 nm = 104  =107
m (millimicrons)
– N.B. 1 nm = 1 m (old unit)
• Energy
– 1 cm-1 = 2.858 cal mol-1 of particles
= 1.986  1016 erg molecule-1
= 1.24  10-4 eV molecule-1
– E (kcal mol-1)  (Å)
= 2.858  105
– E(kJ mol-1) = 1.19  105/(nm)
297 nm = 400 kJ
7
The photoelectric effect
• A beam of light impacts on a
metal surface and causes the
release of electrons (the
photoelectron) if certain
conditions are satisfied
• Conditions
– light must have a frequency above
the threshold, o
– number of photoelectrons
increases with light intensity, but
not the kinetic energy
8
Explanation of the
photoelectric effect
• Ephoton = hphoton
• h = Planck’s constant = 6.626 
10-34 J s
• Applying the Law of the
Conservation of Energy
– energy of the photon is absorbed
by the metal surface and is
transferred to the photoelectron
– the minimum frequency is the
binding energy of the electron
– the remaining energy shows up as
the kinetic energy of the electron
9
Photoelectric effect
• Electron kinetic energy = Photon
energy - Binding energy
• Ekinetic(electron) = h - ho
• Comments
– if frequency is too low, the photo
energy is insufficient to overcome the
binding energy of the electron
– energy in excess of the binding energy
shows up as the kinetic energy of the
electron
– increasing the intensity of the light
increases the number of photons
impacting on the metal
10
Particle properties of
light
• Light has a dual nature of acting
like a wave and acting like a
particle
• The photoelectric effect
confirmed that light occurs as
little packets of energy
• Light is still diffracted like a
wave, has wavelength and
frequency
11
Light and atoms
• When matter absorbs photons of
light, the energy of the photon is
transferred to the matter
• In the case of atoms, the absorption
process yields information about the
atom
• Absorption of a photon transforms
the atom to a higher energy state
• All higher energy states are referred
to as excited states
• The most stable state is the ground
state
12
Absorption and
Emission
• White light (light containing all
energies of light) is passed through a
sample
• Sample absorbs some of the light
• Light that passes through the sample
is dispersed by a prism or other
wavelength selecting device
• Photodetector records the intensity
of the light passing through the
sample, which is then interpreted as
absorption of light
13
Beer’s Law
log 10
•
•
•
•
•
I0
 A   lc
I
Io = Intensity of incident light
I = Intensity of transmitted light
 = molar extinction coefficient
l = path length of cell
c = concentration of sample
14
UV Spectral
Nomenclature
15
UV and Visible
Spectroscopy
• Vacuum UV or soft X-rays
– 100 - 200 nm
– Quartz, O2 and CO2 absorb
strongly in this region
– N2 purge good down to 180 nm
• Quartz region
– 200 – 350 nm
– Source is D2 lamp
• Visible region
– 350 – 800 nm
– Source is tungsten lamp
16
Emission
• Sample is excited by light
• Excited sample emits the light
• Emitted light is wavelength
selected
• The light is detected by a
photodetector
• Plot of emission intensity vs
wavelength is generated
17
Quantization of
absorption and emission
• One of the three things that led
to quantum theory was that the
absorption and emission of light
occurred at discrete frequencies,
not continua
• Interpreted as the energy of the
photon must match the
difference in energy of two
energy levels in the atom or
molecule
18
Molecular process
• Absorption and emission of
visible and ultraviolet light
• Photon is annihilated upon
absorption, and the electrons in
the molecule are rearranged into
the excited state
• Emission results from the
conversion of excited electron
energy being converted to a
photon of light
• Ephoton = Eatom
19
Energy level diagrams
• Wiggle lines indicate radiative
processes
• Straight lines indicate nonradiative
processes
• Each energy level represents an
arrangement of electrons in the atom
h'
h
20
Properties of electrons
• Each electrons have the same
mass and charge
• Electrons behave like magnets
through a property called spin
(actually, magnets are magnets
because electrons have this
property)
• Electrons have wave properties
(diffract just like photons)
21
Heisenberg uncertainty
principle
• A particle has a particular location,
but a wave has no exact position
• The wave properties of electrons
cause them to spread out, hence the
position of the electron cannot be
precisely defined
• They are referred to as being
delocalized in a region of space
• Heisenberg proposed that the motion
and position of the particle-wave
cannot be precisely known at the
same time
22
Bound electrons and
quantization
• The properties of electrons bound to
a nucleus can only take on certain
specific values (most importantly,
energy)
• Absorption and emission spectra
provide experimental values for the
quantized energies of atomic
electrons
• Theory of quantum mechanics links
these data to the wave characteristics
of electrons bound to nuclei
23
The Schrödinger
Equation
• A second order partial differential
equation
• The solutions to such equations are
other equations
• These equations describe threedimensional waves called orbitals
• These solutions have indexes that are
integers (the solutions are quantized
naturally)
• These indexes are called quantum
numbers
24
Quantum numbers
• n - principle quantum number
– values of the positive integers
– n = 1,2,3,…
• l - azimuthal quantum number
– values correlate with the number
of preferred axes of a particular
orbital, indicating its shape
– l = 0,1,2,…(n - 1)
– value of l is often indicated by a
letter (s, p, d, f, for l = 0, 1, 2, 3)
25
Quantum number
• ml - magnetic quantum number
– directionality of orbital
– ml = 0, ±1, ±2, ±l
• ms - spin orientation quantum
number
– ms = ±½
• A complete description of an
atomic electron requires a set of
four unique quantum number
that meet the restrictions of
quantum mechanics
26
Shapes of atomic
orbitals
• Each atomic energy level can be
associated with a specific threedimensional atomic orbital
• Orbitals are maps of the probability
of the electron being in a particular
location around the nucleus
• While there are many
representations, the most important
to learn are the 90% probability
volumes (which I will draw for you)
27
Depictions of orbitals
• electron density plot - electron
density plotted against the
distance from the nucleus
• orbital density plots
• electron contour diagrams (90%
probability drawings)
• All are useful in helping us
visualize the orbital
28
Waves and nodes
29
A variety of radial
projections
30
Radial depictions
31
The p-orbitals
32
The d-orbitals
33
d-orbital radial
projection
34