Download Ch 6 notes 6.1 to 6.4

Chapter 6
Electronic Structure
of Atoms
Electronic
Structure
of Atoms
• The electronic structure of an atom
refers to its number of electrons, how
these electrons are distributed around
the nucleus, and to their energies.
• Much of our understanding of the
electronic structure of atoms has come
from the analysis of light either
absorbed or emitted by substances.
Electronic
Structure
of Atoms
6.1 The Wave Nature of Light
• To understand the electronic structure of
atoms, one must understand the nature of
electromagnetic radiation (EMR).
• Because EMR carries energy through space,
it is also known as radiant energy.
• There are many forms of EMR, to include
visible light, radio waves, infrared waves, Xrays, etc. (See fig 6.4).
• All EMR consists of photons, the smallest
increments of radiant energy.
Electronic
Structure
of Atoms
• Different forms of EMR share characteristics:
 All have wavelike characteristics (much like waves of
water).
 The distance between corresponding points on adjacent
waves is the wavelength ().
 The amplitude – or, maximum extent of oscillation of the
wave -- is related to the intensity of the radiation
Electronic
Structure
of Atoms
 The number of waves passing a given point per unit of
time is the frequency ().
 For waves traveling at the same velocity, the longer the
wavelength, the smaller the frequency.
Electronic
Structure
of Atoms
Electromagnetic Radiation Spectrum
 All EMR moves through travels at the same velocity:
the speed of light (c), 3.00  108 m/s.
 Therefore,
c = 
Electronic
Structure
of Atoms
• As you can see from the electromagnetic
spectrum, there are many forms of EMR.
• The difference in the forms are due to their
different wavelengths, which are expressed in
units of length.
• Wavelengths vary from10-11 m to 103 m.
• Frequency is expressed in cycles per second,
also called a hertz.
 Units of frequency are usually given simply as “per
second,” denoted as s-1 or /s. As in 820 kilohertz
(kHz), written as 820,000 s-1 or 820,000/s
• See Table 6.1 for types of radiation and
associated wavelength.
Electronic
Structure
of Atoms
Wavelength and frequency problems
Electronic
Structure
of Atoms
6.2 Quantized Energy and Photons
• The wave nature of light
does not explain how an
object can glow when its
temperature increases.
• Max Planck explained it by
assuming that energy
comes in packets called
quanta.
Electronic
Structure
of Atoms
• Einstein used this assumption
to explain the photoelectric
effect.
 When photons of sufficiently high
energy strike a metal surface,
electrons are emitted from the metal.
 Electrons are not emitted unless
photons exceed a certain minimal
energy.
 E.g., light with a frequency of 4.60 x
1014 s-1 or greater will cause cesium
atoms to emit electrons, but light of
lower frequency has no effect.
• He concluded that energy is
proportional to frequency:
E = h
where h is Planck’s constant,
6.626  10−34 J-s.
Electronic
Structure
of Atoms
• Therefore, if one knows the
wavelength of light, one
can calculate the energy in
one photon, or packet, of
that light:
c = 
E = h
Electronic
Structure
of Atoms
Energy of photon problems
Electronic
Structure
of Atoms
6.3 Line Spectra and the Bohr Model
Another mystery
involved the emission
spectra observed from
energy emitted by atoms
and molecules.
Electronic
Structure
of Atoms
• One does not observe a
continuous spectrum, as
one gets from a white
light source.
• Only a line spectrum of
discrete wavelengths is
observed.
Electronic
Structure
of Atoms
•
Niels Bohr adopted Planck’s assumption and
explained these phenomena in this way:
1. Electrons in an atom can only occupy certain orbits
(corresponding to certain energies).
2. Electrons in permitted orbits have specific, “allowed”
energies; these energies are not radiated from the
atom.
3. Energy is only absorbed or emitted in such a way as
to move an electron from one “allowed” energy state
to another; the energy is defined by
E = h
Electronic
Structure
of Atoms
• The energy absorbed or emitted
from the process of electron
promotion or demotion can be
calculated by the equation:
E = −RH (
1
1
nf2
ni2
)
where RH is the Rydberg
constant, 2.18  10−18 J, and ni
and nf are the initial and final
energy levels of the electron.
Electronic
Structure
of Atoms
Hydrogen emission spectra
Electronic
Structure
of Atoms
Energy states of hydrogen atom problems
Electronic
Structure
of Atoms
• Limitations of the Bohr Model
 The Bohr model explains the line spectrum of the
hydrogen atom, but not (accurately) the spectra of
other atoms.
 Also, the Bohr model assumes the electron behaves
as a particle.
 Electrons also have wave-like properties.
• However, Bohr model is important because:
 It shows electrons as existing in only certain discrete
energy levels, which are described by quantum
numbers.
 Energy is involved in moving an electron from one
level to another.
Electronic
Structure
of Atoms
6.4 The Wave Nature of Matter
• In the years after development of the Bohr model,
the dual nature of light became known: EMR
(i.e., light) can exhibit both particle-like (photon)
character as well as wave-like character.
• Louis de Broglie (in 1924) extended this idea to
electrons, proposing a relationship between the
wavelength of an electron (or any other particle),
its mass, and velocity:
h
 = mv
 h = Planck’s constant
Electronic
Structure
of Atoms
The Uncertainty Principle
• Heisenberg showed that the more precisely the
momentum of a particle is known, the less
precisely is its position known:
(x)(mv) 
h
4
• In many cases, our uncertainty of the
whereabouts of an electron is greater than the
size of the atom itself!
• In sum, you cannot accurately know both an
electron’s position and momentum at the sameElectronic
Structure
time.
of Atoms
Matter Waves
Electronic
Structure
of Atoms