Electromagnetic spectrum Visible light: all that can be seen by
the human eye *** Blue flames have a shorter wavelength than yellow
flames ***Blue flames are higher in energy and hotter! The rest
cant be seen, but is still considered light
Slide 4
Graphing Direct and Inverse Proportions Direct Proportion:
Inverse Proportion: Y X Y X
Slide 5
Frequency and Wavelength
Slide 6
Relationships Between Energy, Wavelength and Frequency of
Waves: Energy and Frequency: Wavelength and Frequency: Energy and
Wavelength: E E Direct Inverse
Slide 7
Problems: 1. If the wavelength of a wave doubles, and its
energy and speed dont change, what will happen to its frequency? 2
The frequency () will be cut in half. 2. If the frequency of a wave
triples, what happens to the energy needed for the wave? 3 3E You
will need THREE times the energy 3. A student modifies the
wavelength of a wave by changing the energy. If the energy is cut
in half, what must have happened to the wavelength? E 2
Slide 8
Light has two components: Wavelength: in (m)Frequency: in (1/s)
= = Energy: in (J) E = h or c = speed of light (3 x 10 8 m/s) h =
Plancks constant (6.6 x 10 -34 Js) 1 x 10 9 nm = 1 m Or 1nm = 1.0
10 -9 m E E
Slide 9
Energy WS problems 1. Find the energy of a wave if its
frequency is 2.2 x 10 16 1/sec E = h E = (6.6 x 10 -34 Js) x (2.2 x
10 16 1/sec) E = 1.45 x 10 -17 J
Slide 10
6. Find the wavelength of light if its energy is 2.8 x 10 -19
J. What color light would you see? Use E = h to find : 2.8 x 10 -19
J = (6.6 x 10 -34 Js) x = 4.24 x 10 14 1/s Then solve for
wavelength: = 7.07 x 10 -7 m = 707 nm Red color = =
Slide 11
#6 (using one equation) E= =
Slide 12
Review: Neil Bohr Confined electrons to energy levels Quantum
leaps: Energy (light) being released when e- jump from excited
state to ground state.
Slide 13
Atomic emission spectrum: Each element has a unique
line-emission spectrum Emission spectrum of Hydrogen and Iron:
Slide 14
Drawing Bohrs atom: Energy levels# of e- 12 28 318 432
Slide 15
Present Day Model Electrons are located in an orbital Orbitals-
regions around a nucleus that correspond to specific energy levels
Orbitals are also called electron clouds Dont know the exact
position of the electron- creates fuzzy image Electron cloud
Slide 16
Quantum Numbers Quantum Mechanical Model - present day model of
the atom Electrons do NOT orbit nucleus in circular pattern How can
we keep track of the electrons? Four Quantum numbers - defines
region in which electrons can be found
Slide 17
The fixed energies an electron can have are called energy
levels. electrons energy & distance from the nucleus {analogy
is floors in apartment building} 1. Energy Level + n= 1, 1 st
energy level, 1 st floor n= 2, 2 nd energy level, 2 nd floor Think
of the atom as an apartment building with each floor representing
an energy level
Slide 23
4. Spin indicates the orientation of an electrons magnetic
field relative to an outside magnetic field A single orbital can
hold a maximum of 2 electrons, which must have opposite spins
{analogy is roommates in a room}
Slide 24
S: 1 orbital Max e- = 2 P:3 diff. orbitals Max e- = 6 D:5 diff.
orbitals Max e- = 10 f: 7 diff. orbitals Max e- = 14
Slide 25
Aufbau principle states that electrons fill orbitals that have
the lowest energy first (e.g. 1s before 2s). Overlapping
Orbitals
Slide 26
Electron Configuration arrangement of the electrons around the
nucleus Based on the quantum model of the atom 1 H = 1s 1 # of
electrons Sublevel 1 st energy level n = 1 1H1H Orbital
Notation
Slide 27
Hunds Rule states that orbitals of equal energy are each
occupied by one electron before any orbital is occupied by a second
electron, and all electrons in singly occupied orbitals must have
the same spin.
Slide 28
Pauli exclusion principle Maximum 2 electrons can occupy a
single orbital and they must have opposite spins ( or ). No 2 atoms
in the same atom can have the same 4 quantum numbers Electrons in
orbitals can be represented by arrows in boxes
Slide 29
Give the full electron configuration for... Li Na K 1s 2 2s 1
1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 The electrons in
the outer most shell are called the valence electrons. Electrons in
the inner shells are core electrons.
Slide 30
Electron Configuration Example: sulfur has sixteen electrons.
Its electron configuration is written as 1s 2 2s 2 2p 6 3s 2 3p 4
Shorthand Notation: It can also be written as follows by using the
previous noble gas:[Ne]3s 2 3p 4
Slide 31
Write the electron configuration for an atom whose atomic
number is 20. atomic number = number of protons = number of
electrons = 20 According to the aufbau principle, the order of
orbital filling is 1s,2s, 2p, 3s, 3p, 4s, 3d, 4p, and so on. Or
refer to the periodic table 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2
Abbreviated as follows: [Ar]4s 2
Slide 32
Slide 33
Try These 1. Try writing out the electron configuration for
Potassium. 2. Write the electron Configuration for an atom with an
atomic number 20. 3. Write an electron configuration for an atom of
an element whose atomic number is 8. 4. Write an electron
configuration for an atom of an element whose atomic number is
53.
Slide 34
Answers 1. K = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 2. Ca = [Ar]4s 2
3. O = 1s 2 2s 2 2p 4 4. I = [Kr] 5s 2 4d 10 5p 5
Slide 35
What is the correct electron configuration of a sulfur atom? A.
1s 2 2s 2 2p 4 3s 2 3p 6 B. 1s 2 2s 2 2p 6 3s 2 3p 3 C. 1s 2 2s 2
2p 6 3s 2 3p 4 D. 1s 2 2s 2 2p 6 3s 6 3p 2
Slide 36
Drawing Molecules using Lewis Electron-Dot Structures Lewis
structure: shows how the valence electrons are arranged among the
atoms in the molecule Valence electrons are represented by dots
Nuclei and electrons of the inner energy level (if any) are
represented by the symbol of the element
Slide 37
Slide 38
Exceptional Electron Configurations Cr 1s 2 2s 2 2p 6 3s 2 3p 6
3d 4 4s 2 Cu 1s 2 2s 2 2p 6 3s 2 3p 6 3d 9 4s 2 The correct
electron configurations are as follows: Cr 1s 2 2s 2 2p 6 3s 2 3p 6
3d 5 4s 1 Cu 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 1 These arrangements
give chromium a half-filled d sublevel and copper a filled d
sublevel. Some actual electron configurations differ from those
assigned using the aufbau principle because although half-filled
sublevels are not as stable as filled sublevels, they are more
stable than other configurations.
Slide 39
Copyright Pearson Education, Inc., or its affiliates. All
Rights Reserved. How are the quantum mechanical model and the Bohr
model alike? How are they different? Like the Bohr model, the
quantum mechanical model restricts the energy of electrons to
certain values. Unlike the Bohr model, the quantum mechanical model
does not specify an exact path the electron takes around the
nucleus.
Slide 40
Copyright Pearson Education, Inc., or its affiliates. All
Rights Reserved. Calculate the maximum number of electrons in the 5
th principal energy level (n = 5). The maximum number of electrons
that can occupy a principal energy level is given by the formula 2n
2. If n = 5, 2n 2 = 50.
Slide 41
How does Electron Configuration relate to the Periodic
Table?
Slide 42
Slide 43
2 6 10 14
Slide 44
Explain why the correct electron configuration of oxygen is 1s
2 2s 2 2p 4 and not 1s 2 2s 2 2p 3 3s 1. The 2p orbitals are lower
in energy than the 3s orbital, so they will be completely filled
before any electrons will be found in the 3s orbital.
Slide 45
Frequency and Wavlength Light has two components: wavelength (
read lambda ) and frequency ( read nu )
Slide 46
Slide 47
Slide 48
Energy Levels in Atoms The rungs on this ladder are somewhat
like the energy levels in Bohrs model of the atom. A person on a
ladder cannot stand between the rungs. Similarly, the electrons in
an atom cannot exist between energy levels.
Slide 49
Energy Levels in Atoms The rungs on this ladder are somewhat
like the energy levels in Bohrs model of the atom. The energy
levels in atoms are unequally spaced, like the rungs in this
unusual ladder. The higher energy levels are closer together.
Slide 50
A prism separates light into the colors it contains. White
light produces a rainbow of colors. Light and Atomic Emission
Spectra Light bulb SlitPrism Screen
Slide 51
Light from a helium lamp produces discrete lines. Light and
Atomic Emission Spectra SlitPrism Screen Helium lamp Spectrum:
wavelengths of visible light that are separated when a beam of
light passes through a prism; range of wavelengths of
electromagnetic radiation
Slide 52
Hydrogens Line-Emission Spectrum Spectrum: wavelengths of
visible light that are separated when a beam of light passes
through a prism; range of wavelengths of electromagnetic
radiation
Slide 53
An Explanation of Atomic Spectra The three groups of lines in
the hydrogen spectrum correspond to the transition of electrons
from higher energy levels to lower energy levels.
Slide 54
s 1 orientaion Max e- = 2 p 3 orientaion 3 diff. orbitals Max
e- = 6 s orbitals are spherical p orbitals are dumbbell-shaped
Slide 55
f 7 orientaion 7 diff. orbitals Max e- = 14 D 5 orientaion 5
diff. orbitals Max e- = 10
Slide 56
Copyright Pearson Education, Inc., or its affiliates. All
Rights Reserved. The Heisenberg Uncertainty Principle The
Heisenberg uncertainty principle states that it is impossible to
know both the velocity and the position of a particle at the same
time. This limitation is critical when dealing with small particles
such as electrons. But it does not matter for ordinary-sized
objects such as cars or airplanes. Quantum Mechanics
Slide 57
Copyright Pearson Education, Inc., or its affiliates. All
Rights Reserved. To locate an electron, you might strike it with a
photon. The electron has such a small mass that striking it with a
photon affects its motion in a way that cannot be predicted
accurately. The very act of measuring the position of the electron
changes its velocity, making its velocity uncertain. Quantum
Mechanics Before collision: A photon strikes an electron during an
attempt to observe the electrons position. After collision: The
impact changes the electrons velocity, making it uncertain.
Slide 58
Amplitude: the height of a waves crest Wavelength (,lambda):
distance between adjacent crests of a wave Frequency (, nu): #of
wave cycles to pass a given point per unit of time.
Slide 59
Electromagnetic spectrum Speed of light: 3 10 8 m/s in a
vacuum. E = h
Slide 60
Frequency and Wavelength Light has two components: wavelength (
read lambda ) and frequency ( read nu ) Frequency and wavelength
are inversely proportional ( = 1/ ) E = h E E
Slide 61
Each element has a unique line-emission spectrum Emission
spectrum of Hydrogen and Iron: Atomic emission spectrum: the
pattern formed when light passes through a prism or diffraction
grating to separate it into the different frequencies of light it
contains
Slide 62
1 st level = 1 orbital, shape is spherical (s) 1s orbital-
(shape- spherical) (1s) 2 nd level = 4 orbitals 1 s orbital-
(shape- spherical) (2s) 3 p orbitals- (shape- dumbbell) (2p) 3 rd
level= 9 orbitals 1 s orbital- (shape- spherical) (3s) 3 p
orbitals- (shape- dumbbell) (3p) 5 d orbitals- shape- most are four
leaf clover (3d) 4 th level=16 orbital 1s orbital- (shape-
spherical) (4s) 3 p orbitals- (shape- dumbbell) (4p) 5 dorbitals-
shape- most are four leaf clover (4d) 7 f orbitals- shapes are
complicated (4f)
Slide 63
Atomic Orbitals Summary of Principal Energy Levels and
Sublevels Principal energy level Number of sublevels Type of
sublevel Maximum number of electrons n = 111s (1 orbital)2 n = 222s
(1 orbital), 2p (3 orbitals)8 n = 33 3s (1 orbital), 3p (3
orbitals), 3d (5 orbitals) 18 n = 44 4s (1 orbital), 4p (3
orbitals), 4d (5 orbitals), 4f (7 orbitals) 32 The numbers and
types of atomic orbitals depend on the principal energy level.
Slide 64
Do Now Ho does E, , change as the wave goes from A to B? A.
B.
Slide 65
Do Now #1 Draw a picture of a wave with high frequency and low
frequency. Label wavelength Circle the wave with high energy high
frequencylow frequency.
Slide 66
The visible red light has a wavelength of about 650 nm. At
sunrise and sunset, red or orange colors are present because the
wavelengths associated with these colors are less efficiently
scattered by the atmosphere than the shorter wavelength colors
(e.g., blue and purple). A large amount of blue and violet light
has been removed as a result of scattering and the longwave colors,
such as red and orange, are more readily seen.