Transcript Slide 1
Chapter 7: Periodicity and Atomic Structure
Renee Y. Becker Valencia Community College CHM 1045 1
Light and Electromagnetic Spectrum • Several types of electromagnetic radiation make up the electromagnetic spectrum 2
Light and Electromagnetic Spectrum
Frequency,
:
The number of wave peaks that pass a given point per unit time (1/s)
Wavelength,
:
The distance from one wave peak to the next (nm or m)
Amplitude:
Height of wave Wavelength x Frequency = Speed (m) x (s -1 ) = c (m/s) 3
Light and Electromagnetic Spectrum 4
The Planck Equation E = h E = hc / h = Planck’s constant, 6.626 x 10 -34 J s 1 J = 1 kg m 2 /s 2 5
Example1: Light and Electromagnetic Spectrum • The red light in a laser pointer comes from a diode laser that has a wavelength of about 630 nm. What is the frequency of the light?
c
= 3 x 10 8 m/s 6
Atomic Spectra •
Atomic spectra:
Result from excited atoms emitting light.
•
Line spectra:
Result from electron transitions between specific energy levels.
•
Blackbody radiation
is the visible glow that solid objects emit when heated.
7
Atomic Spectra •
Max Planck (1858–1947):
proposed the energy is only emitted in discrete packets called quanta.
The amount of energy depends on the frequency: E = energy = wavelength = frequency c = speed of light h = planck’s constant
E
=
h
=
hc
h
=
6.626
10
-
34
J
s
8
Atomic Spectra
Albert Einstein (1879 –1955):
Used the idea of
quanta
to explain the
photoelectric effect.
• He proposed that light behaves as a stream of particles called
photons
• A photon’s energy must exceed a minimum threshold for electrons to be ejected.
• Energy of a photon depends only on the frequency.
E = h 9
Atomic Spectra 10
Example 2: Atomic Spectra • For red light with a wavelength of about 630 nm, what is the energy of a single photon and one mole of photons?
11
Wave –Particle Duality •
Louis de Broglie (1892 –1987):
Suggested waves can behave as particles and particles can behave as waves. This is called wave – particle duality.
m = mass in kg p = momentum (mc) or (mv) For Light :
h
=
mc
=
h p
For a Particle :
h
=
mv
=
h p
12
Example 3: Wave –Particle Duality • How fast must an electron be moving if it has a de Broglie wavelength of 550 nm?
m
e = 9.109 x 10 –31 kg 13
Quantum Mechanics •
Niels Bohr (1885 –1962):
Described atom as electrons circling around a nucleus and concluded that electrons have specific energy levels.
•
Erwin Schrödinger (1887–1961):
Proposed quantum mechanical model of atom, which focuses on wavelike properties of electrons.
14
Quantum Mechanics •
Werner Heisenberg (1901 –1976):
Showed that it is impossible to know (or measure) precisely both the position and velocity (or the momentum) at the same time.
• The simple act of “seeing” an electron would change its energy and therefore its position.
15
Quantum Mechanics •
Erwin Schrödinger (1887–1961):
Developed a compromise which calculates both the energy of an electron and the probability of finding an electron at any point in the molecule. • This is accomplished by solving the Schrödinger equation, resulting in the wave function 16
Quantum Numbers • Wave functions describe the behavior of electrons.
• Each wave function contains four variables called quantum numbers: • Principal Quantum Number (
n
) • Angular-Momentum Quantum Number (
l
) • Magnetic Quantum Number (
m l
) • Spin Quantum Number (
m s
) 17
Quantum Numbers •
Principal Quantum Number (n):
Defines the size and energy level of the orbital.
n
= 1, 2, 3, – As
n
increases, the electrons get farther from the nucleus.
– As
n
increases, the electrons’ energy increases.
– Each value of
n
is generally called a
shell
.
18
Quantum Numbers •
Angular-Momentum Quantum Number (l):
Defines the three-dimensional shape of the orbital.
• For an orbital of principal quantum number
n
, the value of
l
can have an integer value from 0 to
n
– 1.
• This gives the subshell notation:
l =
0
= s
orbital
l =
3
= f
orbital
l =
1
= p
orbital
l =
2
= d
orbital
l =
4
= g
orbital 19
Quantum Numbers •
Magnetic Quantum Number (m
l
):
the spatial orientation of the orbital.
Defines • For orbital of angular-momentum quantum number, from
–l l
to , the value of
+l
.
m l
has integer values • This gives a spatial orientation of:
l = 0 giving m l = 0 l = 1 giving m l = –1, 0, +1 l = 2 giving m l on…...
= –2, –1, 0, 1, 2, and so
20
Quantum Numbers •
Magnetic Quantum Number (m
l
):
–l
to
+l
S orbital 0 P orbital -1 0 1 D orbital -2 -1 0 1 2 F orbital -3 -2 -1 0 1 2 3 21
•
Spin Quantum Number:
m s • The
Pauli Exclusion Principle
states that no two electrons can have the same four quantum numbers.
Quantum Numbers 22
Quantum Numbers 23
Example 4: Quantum Numbers • Why can’t an electron have the following quantum numbers?
(a)
n
= 2,
l
= 2,
m l
= 1 (b)
n
= 3,
l
= 0,
m l
= 3 (c)
n
= 5,
l
= –2,
m l
= 1 24
• Example 5: Quantum Numbers Give orbital notations for electrons with the following quantum numbers:
(a)n
= 2,
l
= 1 (b)
n
= 4,
l
= 3 (c)
n
= 3,
l
= 2 25
Electron Radial Distribution •
s Orbital Shapes:
Holds 2 electrons 26
Electron Radial Distribution •
p Orbital Shapes: Holds 6 electrons, degenerate
27
Electron Radial Distribution •
d and f Orbital Shapes:
d holds 10 electrons and f holds 14 electrons, degenerate 28
Effective Nuclear Charge • Electron shielding leads to energy differences among orbitals within a shell.
• Net nuclear charge felt by an electron is called the
effective nuclear charge
(
Z
eff ).
•
Z
eff
is lower than actual nuclear charge.
•
Z
eff increases toward nucleus
ns > np > nd > nf
29
Effective Nuclear Charge 30