Download Chapter 4 - Tolland High School

Chapter 4
Spectroscopy &
Arrangement of Electrons(e-)
Properties of Light
• Electromagnetic Radiation- form of
energy w/ wavelike properties as it
travels through space
• Electromagnetic Spectrum- classifies as
electromagnetic radiation based on
wavelength(λ) and frequency(ν)
The Electromagnetic Spectrum
Visible Light
Violet
Blue
Green Yellow Orange
400nm 450nm 500nm 550nm 600nm 700nm
Gamma Rays X-Rays UV-Rays
v
Red
1019 Hz
1017 Hz
1015 Hz
Infrared Microwave Radio
1014 Hz
109 Hz
Long-Wave
108-106 Hz 105 Hz
Short Wavelength
Long Wavelength
High Frequency
Low Frequency
• Frequency(ν)- # of waves that pass a given
point in a specific time (usually 1s)
– Measured in waves per second (1/s) or Hz
• Wavelength(λ)- the distance between
corresponding points on a wave
(ex: peak to peak)
– Measured in meters (m)
Speed of Light(c)
• Speed of Light = wavelength x frequency
c = λν
• Speed of Light is a Constant:
c = 3.0 x 108 m/s
• Frequency and Wavelength are Inversely Proportional
:
v
Speed of Light Sample
Calculations
• What is the wavelength of EM Radiation
that has a frequency of 1014Hz?
Speed of Light Sample
Calculations
• What is the frequency of microwaves with a
wavelength of 0.01m?
Dual Wave-Particle Nature of Light
• Light has both the properties of waves and
of particles
– Wave Properties: light can be bent as it passes
through objects
– Particle Properties: photons have mass and
exert force on other objects
Photoelectric Effect
• When light particles (photons) collide with a
metal, the photons knock electrons (e-) loose
• These electrons move toward the positive terminal
creating an electric current (electricity)
Light as Particles
• Photons- particles of light carrying a
quantum of energy
(Max Planck)
• A Quantum of Energy- the minimum quantity of
energy that can be lost or gained by an atom
E = hν
E = energy of a photon (J)
h = Planck’s constant (6.626 x 10-34Js)
ν = frequency (Hz)
Energy of a Photon Sample
Calculations
• What is the energy of violet light with a
frequency of 7.50 x 1014 s-1?
Energy of a Photon Sample
Calculations
• What is the frequency of UV light that has
an energy of 2.39 x 10-18J?
Sample Calculations
• What is the energy of EM radiation with a
1.0 x 10-6m wavelength?
Emission Spectra
• Emission spectra are fingerprints of an atom
• Every atom gives off different colors from
the visible light spectrum when they
release absorbed energy
Energy States of Atoms/Electrons
• Ground State- lowest energy level of an
electron within an atom
• Excited State- a higher energy level within
an atom that an electron may exist in
– Energy must be absorbed for an electron to go
from ground to excited state
– Energy is given off as visible light when an
atom returns to ground state
– Every atom gives off a unique spectrum based
on the movement of it’s electrons
Energy of a Photon
• Ephoton = Efinal - Einitial
E2
Excited State
E2 – E1 = Ephoton = hν
E1
Ground State
Bohr Model of the Hydrogen Atom
• Bohr’s model indicates that as atoms
absorb energy their electrons move to
higher energy levels
• When the absorbed energy is given off as
visible light the electrons return to their
ground state
Bohr Model
Drawing Bohr Diagrams
• Shows the total number of electrons
• The energy levels fill before an e- can be put
in the next energy level
– 1st shell = 2 e– 2nd shell = 8 e– 3rd shell = 8 e-
Practice Drawing Bohr
Diagrams
• Draw a Bohr Diagram for:
–
–
–
–
–
C
H
O
Ar
F
Quantum Model of the Atom
• The Bohr model was more accurate than
previous models but was only completely
accurate for Hydrogen, other elements did
not behave exactly as Bohr predicted
• The Quantum model was later developed
based on work of many scientists including
Schrodinger, Heisenberg, & Einstein
Quantum Model of the Atom
Electrons as Waves
• Louis deBroglie proved that electrons had
wave properties by showing that electrons
could produce interference patterns like
sound and light waves
• Passing electrons through a crystal also
caused the stream of electrons to bend like
light waves do
Interference Patterns
Heisenburg Uncertainty Principle
• This principle states that it is impossible to
determine simultaneously both the position
and velocity of an electron
• This theory led to the concept of the
electron cloud
Quantum Theory
• This theory describes mathematically the
wave properties of electrons based on
probability
– Electrons are not in set energy levels
– Electrons are in 3D orbits around the nucleus
called orbitals
Orbitals
• Orbitals are 3D regions around the nucleus
of an atom that indicate the probable
location of an electron
Quantum #’s
• Quantum #’s are used to specify the
properties and location of electrons in
orbitals around the nucleus
• There are 4 quantum #’s, each is
represented by a letter : n, l, m, & s
Principle Quantum # (n)
• The principle quantum number indicates the
main energy level of an electron and it’s
distance from the nucleus
• n=1 : e- close to nucleus
• n=7 : e- further from nucleus
* As the n value increases so does the energy
of the e-, and the distance from the nucleus
Angular Momentum Quantum #(l)
• The angular momentum quantum # indicates the
shape of an orbital (s, p, d, f)
• Each l value has a corresponding shape
l = 0 : s-shape, sphere orbital around nucleus
- an s-orbital can hold up to 2 electrons
l = 1 : p-shape, 2 lobes on either side of nucleus
- an atom can have 3 p-orbitals, one in each plane (x,
y, z)
- p-orbitals can hold up to
6 electrons
Angular Momentum Quantum #(l)
• l = 2 : d-shape, 4 lobes “clover” around
nucleus
– An atom can have up to 5 d-orbitals
– d-orbitals can hold up to 10 electrons
Angular Momentum Quantum #(l)
• l = 3 : f-shape
– f-orbitals can hold up to 14 electrons
Magnetic Quantum # (m)
• The magnetic quantum # indicates the
orientation of an orbital around the nucleus
– Ex. Would indicate if your p orbital is on the x,
y, or z plane
Spin Quantum # (s)
• The spin quantum number indicates the
direction that an electron is spinning, either
clockwise or counterclockwise
• s = +1/2 : clockwise spin
• s = -1/2 : counterclockwise spin
* 2 electrons in the same orbital must have
opposite spins
Quantum #’s
n – principle quantum #
distance from nucleus / main energy level
l – angular momentum quantum #
shape of orbital
m – magnetic quantum #
orientation of orbitals around nucleus
s – spin quantum #
direction of e- spin around nucleus
Pauli Exclusion Principle
• The Pauli Exclusion principle states that no
2 electrons in the same atom can have the
same combination of 4 quantum #’s
• This means that no 2 electrons could be in
the same place at the same time
Orbital Notation
• Orbital notation uses boxes and arrows to
indicate electrons in orbital by energy level
• Aufbau Principle- an electron will occupy
the lowest possible energy level that can
hold it
• Hund’s Rule- orbitals of equal energy will
each receive one electron before they
receive a second electron
e
Configurations
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d
7s 7p
Order of Atomic Sublevels
• Orbital Notation:
• P:
• F:
* Each p-orbital gets 1 ebefore it gets a 2nd e-
Examples
• N:
• Si :
• Fe :
• Mg :
Electron Configuration Notation
• 1s22s22p63s23p64s23d104p65s24d105p66s2…
Principle Quantum #
or Energy Level
Angular Momentum Quantum #
or Orbital Shape
• Ex: Carbon = 6 electrons
• C = 1s22s22p2
• Ex: Sodium = 11 electrons
• Na = 1s22s22p63s1
# of Electrons in Orbital
Examples
• N:
• Si :
• Fe :
• Mg :
Noble Gas Notation (Shortcut)
• To eliminate repetitive electron configuration for
elements with large #’s of electrons the symbol of
a Nobel Gas can be substituted for a portion of the
electron configuration
• Ex: K = 1s22s22p63s23p64s1
[Ar]4s1
• Ex: Zn = 1s22s22p63s23p64s23d10
[Ar]4s23d10
Examples
• N:
• Si :
• Fe :
• Mg :
END OF CHAPTER 4
NOTES !!!
Flame Test Lab
Compound
NaCl
Na(NO3)
Sr(NO3)2
Ca(NO3)2
Color Description
Unknown
A
B
C
Ba(NO3)2
K(NO3)
D
Cu(NO3)2
E
Cu(SO4)
F
Li(NO3)
Color
Compound
Description
Calculations
Compound
Flame
Color
λ (nm)
ν (Hz)
Sr(NO3)2
CuSO4
NaCl
General Equations:
Speed of Light: c
c = 3.0 x 108 m/s
=λν
Energy: E
=hν
h = 6.626 x 10-34 Js
E (J)