Download Views on Atomic Stru..

Views on Atomic Structure
Classical View – electrons
and properties of electrons
Experiments with Light –
Quantum Theory
Quantum View – behavior of
electrons in atoms
1
Cathode Rays
Cathode rays are the carriers of
electric current from cathode to
anode inside a vacuumed tube
Cathode rays travel
in straight lines
2
Cathode Rays
Cause glass and
other materials to
fluoresce
Deflect in a
magnetic field
similarly to
negatively charged
particles
3
J. J. Thomson’s Experiment
Devised an experiment to find the
ratio of the cathode ray particle’s
mass (me) to the charge (e)
me /e = –5.686 × 10–12 kg C–1
4
The Electron
coined the term “electron”
Millikan
measured the
charge on an
electron - the
famous “oildrop”
experiment
5
Determined Electron Values
Robert Millikan then determined a value for the
charge
e = –1.602 × 10–19 C
From m/e and the charge, the mass of an electron
was determined to be
m = 9.109 × 10–31 kg/electron
6
J. J. Thomson – Atomic Model
Thomson proposed an atom
with a positively charged
sphere containing equally
spaced electrons inside
RAISIN BUN MODEL
7
Rutherford’s Model
Ernest Rutherford characterized alpha particles
through an experiment and discovered the positive
charge of an atom is concentrated in the center of an
atom, the nucleus
8
Rutherford’s Interpretation
9
Protons and Neutrons
From Rutherford’s experiments, he was able to
determine the amount of positive nuclear charge
The positive charge was carried by particles
called protons
Scientists introduced the atomic number,
which represents the number of protons in the
nucleus of an atom
James Chadwick discovered neutrons in the
nucleus, which have nearly the same mass as
protons and no charge
10
Mass Spectrometer
If a stream of positive ions
having equal velocities is
brought into a magnetic
field, the lightest ions are
deflected the most, making
a tighter circle
11
Wave Motion
Caused by a displacement in
a medium
Characterized by height of
crest (or depth of trough)
12
The Wave Nature of Light
Electromagnetic waves originate from the
movement of electric charges
The movement produces fluctuations in electric and
magnetic fields
13
Characterizing Waves
Electromagnetic radiation is characterized by its
wavelength, frequency, and amplitude
Wavelength (l) is the distance between any
two identical points in consecutive cycles
14
Characterizing Waves
Frequency of a wave is the number of cycles of
the wave that pass through a point in a unit of time
Amplitude of a wave is its height: the distance
from a line of no disturbance through the center of
the wave peak
15
The Electromagnetic
Spectrum
The electromagnetic spectrum is largely invisible to the
eye
16
The Electromagnetic Spectrum
• We can feel some radiation through other
senses (infrared)
• Sunburned skin is a sign of too much
ultraviolet radiation
• Materials vary in their ability to absorb or
transmit different wavelengths
– Our bodies absorb visible light, but transmit
most X rays
– Window glass transmits visible light, but
absorbs ultraviolet radiation
17
Continuous Spectra
White light
passed
through a
prism
produces a
spectrum –
colors in
continuous
form.
18
The Continuous Spectrum
l ~ 650 nm
l ~ 575 nm
l ~ 500 nm
l ~ 480 nm
l ~ 450 nm
The different
colors of light
correspond
to different
wavelengths
and
frequencies
19
Line Spectra
Light passed
through a prism
from an
element
produces a
discontinuous
spectrum of
specific colors
20
Line Spectra
The pattern of lines emitted by excited atoms of
an element is unique
= atomic emission spectrum
21
Quantum Theory –
Black Body Radiation
Planck proposed that the vibrating atoms in a heated
solid could absorb or emit electromagnetic energy
only in discrete amounts
The smallest amount of energy, a quantum, is
given by:
E = hv
where h is Planck’s constant: = 6.626 × 10–34 J s
Planck’s quantum hypothesis states that energy can be
absorbed or emitted only as a quantum or as whole
multiples of a quantum
22
Quantum Theory –
Photoelectric Effect
Einstein considered electromagnetic energy to be
bundled into little packets called photons
Energy of photon is
E = hv
Photoelectric Effect
Movie
23
Bohr’s Hydrogen Atom
Niels Bohr found that the
electron energy (En) was
quantized, that is, that it can
have only certain specified
values
Each specified energy value is
called an energy level of the
atom
24
The Bohr Model
En = –B/n2 where B is a constant = 2.179 × 10–18 J
and n is an integer
The negative sign represents the forces of attraction
The energy is zero
when the electron is
located infinitely far
from the nucleus
25
Bohr Explains Line Spectra
Bohr’s equation is most useful in determining the
energy change (Elevel) that accompanies the leap
of an electron from one energy level to another
For the final and initial levels:
B
Ef  2
nf
and
B
Ei  2
ni
The energy difference between nf and ni is:
 B   B 
 1 1
E   2    2   B  2  2 
 nf   ni 
 ni nf 
26
Energy Levels and Spectral
Lines for Hydrogen
27
Ground States and Excited
States
Electrons in their lowest possible energy levels are
in the ground state
Electrons promoted to any level n > 1 are in an
excited state
Electrons are promoted by absorbing energy
e.g., electric discharge, heat, lasers (photons)
Electrons in an excited state eventually drop back
down to the ground state  “relaxation”
28
Electronic Transitions
Arrows represent transitions
between energy levels
Upward arrows (a) show energy
absorption, electrons move to
higher energy levels
Downward arrows (b)–(d)
represent energy release and
relaxation
The length of an arrow is inversely
proportional to photon wavelength
29
Electronic Transitions
The length of an arrow is inversely
proportional to photon wavelength
Shorter wavelengths,
higher energies
Longer wavelengths,
lower energies
30
De Broglie’s Equation
•Louis de Broglie speculated that matter can
behave as both particles and waves, just like light
•He proposed that a particle with a mass m
moving at a speed v will have a wave nature
consistent with a wavelength
h
l
mv
31
Wave Functions (y)
Quantum mechanics, or wave mechanics, is the
treatment of atomic structure through the wavelike
properties of the electron
Erwin Schrödinger developed an
equation to describe the hydrogen
atom
A wave function is a solution to the
Schrödinger equation and represents
an energy state of the atom
32
Interpretation of a
Wave Function
Wave mechanics provides a probability of where an
electron will be in certain regions of an atom
The Born interpretation:
The square of a wave function (y2) gives the
probability of finding an electron in a small
volume of space around the atom (orbital)
The interpretation leads to the idea
of a cloud of electron density rather
than a discrete location
33
The Uncertainty Principle
Werner Heisenberg’s
uncertainty principle states
that we can’t simultaneously
know exactly where a tiny
particle like an electron is
and exactly how it is moving
34
The Uncertainty Principle
In light of the uncertainty principle, Bohr’s model of
the hydrogen atom fails, in part, because it tells more
than we can know with certainty
Electron is spread out like a wave; the wave which
describes how the electron is distributed spacially is
called a wave function)
35
Quantum Numbers and
Atomic Orbitals
A wave function with a given set of these three
quantum numbers is called an atomic orbital
In quantum mechanics the atomic orbitals require
three integer quantum numbers to completely
describe the energy and the shape of the 3-D space
occupied by the electron (n, l, and ml)
36
Principal Quantum Number (n)
• Is independent of the other two quantum numbers
• Can only be a positive integer
• indicates the size of an orbital (distance from
the nucleus) and its electron energy
• n can be 1, 2, 3, 4, …
37
Orbital Angular Momentum Quantum
Number (l)
(aka Azimuthal quantum number)
• Determines the shape of the orbital: s, p, d, f which
corresponds to values of l = 0, 1, 2, 3
• Possible values: 0 to (n – 1); e.g., if n = 2, l can only be 0 or /1
• Each of these orbitals is a different region of space and a
different shape
•All the ‘l’ quantum values represent different subshells
•When n = 1, there is only 1 “l” value meaning there is only one
subshell in the first energy level; when n= 2; there are 2 values
for ‘l’ indicating two subshells in the second energy level
38
Magnetic Quantum Number
(ml)
Determines the orientation in space of the orbital;
so named because in a magnetic field, these
different orientations have different energies
Possible values: –l to +l;
e.g., if l = 2,
ml can be –2, –1, 0, 1, 2
The magnetic quantum number defines the number
of orbital in a shell. E.g. in the l = 0 subshell, there
is only one ml value, therefore there is only orbital in
this subshell; when l=1; there are 3 possible ml
values (-1, 0, +1) 3 orbitals in this subshell
39
Quantum Numbers Summary
Taken together the three quantum numbers specific
the orbital the electron occupies. Namely:
the energy of the orbital, the shape of the orbital, and
the orientation of the orbital
40
.
• writing 3 quantum numbers to indicate
every possible orbital an electron can
occupy is cumbersome; instead do we do
the following:
• retain the numeric value of the principal
quantum number and we use a letter to
indicate the azimuthal quantum number:
• l = 0  s; l = 1 p; l = 2  d; l = 3  d
• When combined, they indicate an a
specific orbital e.g. 1s orbital; 2s orbital; 2p
orbital
41
Radial Distributions
Electrons are most likely to reside nearest the
nucleus because of electrostatic attraction
Probability of finding an electron
decreases as distance (radius) from the
nucleus increases
42
Electron Probabilities
and the 1s Orbital
The 1s orbital looks very much like a fuzzy ball,
that is, the orbital has spherical symmetry (the
probability of finding an electron is the same in
direction)
The electrons are more concentrated near the center
43
Electron Probabilities
and the 2s Orbital
The 2s orbital has two regions of high electron
probability, both being spherical
The region near the nucleus is separated from the
outer region by a spherical node - a spherical shell
in which the electron probability is zero
EOS
44
The Three p Orbitals
-There are 3 p orbital; each orbital is cylindrically
symmetrical with respect to rotation around one of the
3 axes, x, y, or z
Each ‘p’ orbital has two lobes of high probability
density separated by a node (region of zero
probability)
45
The Five d Orbitals
46
Electron Spin (ms)
The electron spin quantum number explains some
of the finer features of atomic emission spectra
The spin refers to a magnetic field induced by the
moving electric charge of the electron as it spins
Only possible values
= –1/2 to +1/2
EOS
47
The Stern-Gerlach Experiment
Interaction of the electron spin with the magnetic
field caused a splitting of the observed signal
EOS
48
Summary of Concepts
• Cathode rays are negatively charged
fundamental particles of matter, now called
electrons
• An electron bears one fundamental unit of
negative electric charge
• A nucleus of an atom consists of protons and
neutrons and contains practically all the mass of
an atom
• Mass spectrometry establishes atomic masses
and relative abundances of the isotopes of an
element
49
Summary of Concepts
• Electromagnetic radiation is an energy
transmission in the form of oscillating electric
and magnetic fields
• The oscillations produce waves that are
characterized by their frequencies (v),
wavelengths (l), and velocity (c)
• The complete span of possibilities for frequency
and wavelength is described as the
electromagnetic spectrum
50
Summary of Concepts
• Planck’s explanation of quantums gave us E
= hv
• The photoelectric effect is explained by
thinking of quanta of energy as concentrated
into particles of light called photons
• Wave functions require the assignment of
three quantum numbers: principal quantum
number, n, orbital angular momentum
quantum number, l, and magnetic quantum
number, ml.
• Wave functions with acceptable values of the
three quantum numbers are called atomic
orbitals
51
Summary of Concepts
• Orbitals describe regions in an atom that have a
high probability of containing an electron or a
high electronic charge density
• Shapes associated with orbitals depend on the
value of l. Thus, an s orbital (l = 0) is spherical
and a p orbital (l = 1) is dumbbell-shaped
• A fourth quantum number is also required to
characterize an electron in an orbital - the spin
quantum number, ms
52