By Jake Grodsky and Sarine Hagopian
Image From: http://www.eoearth.org/files/115601_115700/115629/350px-Spectrum.jpg
1. An electron in the atom absorbs energy from heat,
electricity, radiation, etc.
2. That electron moves to an orbital at a higher
energy level
3. Later, the excited electron returns to a lower
energy level
4. Excess energy lost by electron is released as light
or other electromagnetic radiation
Since each element has its orbitals at slightly
different energies, each spectrum has a unique
finger print.
http://www.youtube.com/watch?v=QI50GBUJ48s
•According to Niels Bohr’s theory: electrons can
only exist in certain possible energy levels.
•Energy of an electron is proportional to its
distance from the nucleus
Image From: http://hyperphysics.phy-astr.gsu.edu/hbase/imgmod/bohr1.gif
Image From: http://reich-chemistry.wikispaces.com/file/view/rutherford_bohr_model.gif/103784023/rutherford_bohr_model.gif
•When light is shone on metal, electrons are emitted
from the metal
•The effect can be used to switch a light signal into an
electric current
Bright light
current
KEejected e-=Ephoton-Ethreshold
Dim light
0
fthreshold
•For a long time it was believed that light was
solely a wave
•Both light and electrons have a dual nature
•They exhibit characteristics of both waves and
particles
•The photoelectric effect proves that light has a
particle nature as well
•The wave properties of electrons are shown
through the DeBroglie Hypothesis
Constant:
h = 6.626 × 10-34 J•s
λ=
velocity
wavelength of a particle
Constant: me = 9.11 × 10-31 kg
•The wave mechanical model is the most recent model of the atom
•Improvements were made on Bohr’s model, specifically dealing with electrons
•Electrons are treated as waves instead of particles- electron has more in common with
light, tv, radio waves, microwaves, and x-rays than it does with protons and neutrons
•Orbitals are the regions in atoms which are most likely to have electrons in them
•The model is more statistical than visual
•This model includes energy levels which are numbered 1-7(closest to farthest) which
indicates how far a given electron is from the nucleus
•The energy level can be viewed in the same way as Bohr’s model viewed the shell
•Lower energy levels are
always filled first
•Ions have less electrons
than the neutral parent
atom
• this means that
electron configurations
of ions look like those
of other neutral
elements (even from a
different atom)
Image From: http://www.mikeblaber.org/oldwine/chm1045/notes/Struct/EPeriod/IMG00011.GIF
•Configuration can be abbreviated
•The last Noble Gas element symbol is
put in brackets and remainder of the
electron configuration is written out
Example:
Electron Configuration of Zinc:
1s22s22p3s23p64s2 3d10
Abbreviated Electron
Configuration of Zinc: [Ar] 4s2 3d10
Image From: http://www.mpcfaculty.net/mark_bishop/abbreviated_electron_configuration_help.htm
Key:
n= principal quantum number
l= angular momentum number [0- (n-1]
ml= magnetic quantum number [-l – l]
•The final quantum
number is the ms
number. This is ± ½,
depending on the spin
of the electron
Image From: http://library.thinkquest.org/19662/images/eng/pages/improved-bohr-2.jpg
•Electron configurations can be expressed as orbital diagrams as pictured below by
visualizing each individual electron and its corresponding spin, as well as the orbital and
energy level that it is a part of.
•To form an orbital diagram:
1. Determine the electron configuration of the atom and the total
amount of electrons.
2. Following Hund’s Rule, begin to fill orbitals from lowest energy level to
highest, remembering the Pauli Exclusion Principal and having an
upward and downward arrow in each orbital, representing the positive
and negative electron spin (responsible for the +½ and -½ values of ms.
Oxygen
1s
2s
2px 2py 2pz
Electron Configuration: 1s22s22p4
Total number of e-: 8
What would the quantum number be for this e-?
2, 1, -1, -½
Atomic Radius: Size of an atom which is influenced by the
volume of the e- orbitals (clouds)
Increases
Decreases
Why does atomic radius increase as you go
down a group?
• more energy levels so the new levels are
“blocked” and therefore not as tightly pulled to
the center
Why does atomic radius decrease as you go across a period?
• Only one p+ and one e- are added
•Increasing nuclear charge pulls outermost e-s closer and closer to the nucleus 
reduces atom size
•All additional e- go into same principle energy level so shielding is not an issue 
nucleus just gets stronger and squeezes everything closer
•Cations are smaller than their neutral parent atoms
•Cations have more protons than electrons (hence their positive charge)
•Protons more tightly pull the electrons towards the nucleus therefore reducing
the size of the atom
•Anions are larger than their neutral parent atoms
•Anions have more electrons than neutrons (hence their negative charge)
•Electrons are not as attracted to the nucleus therefore increasing atomic radius
Ionization Energy: amount of energy needed to remove an efrom an atom or ion
Electronegativity: a measure of the ability of an atom in a
chemical compound to attract/gain e-s
Increase
•As you go down a group, more orbitals
are added  valence e- are farther from
Decrease
nucleus so pull of p+s on the e-s is reduced
•As you go across a period, more protons
are added to the nucleus  valence
electrons are held more tightly
•Formed when e- pairs are shared amongst atoms
•Generally a metal and a nonmetal pair
Key Terms
•Lewis structure: representation of a covalently bonded molecule and its valence electrons
•Octet Rule: In a covalent molecule, each atom has eight electrons around it
•Lone Pair: pair of e-s not involved in bonding
•Bond pair: pair of e-s shared between two atoms
}
•Double bond: two pairs of e-s shared between atoms
•Triple bond: three pairs of
e -s
shared between atoms
•As the number
of shared e- pairs
goes up, bond
length goes
down
•Symmetrical arrangements are more likely than asymmetrical ones
•The less electronegative atom tends to be in the middle
•Subtract valence e-s from total electrons needed to complete octet/duet and
divide by two  this is the number of bonds that will need to be made
•If e- needed > e- remaining, add bonds
•If e- remaining > e- needed, add lone pairs to central atom
Image from:
https://vinstan.wikispaces.com/file/view/lewis_structure.gif/46694367/lewis_
structure.gif
Are any bonds polar?
No
Yes
Are polar bonds arranged
symmetrically?
Yes
Non-Polar Molecule
No
Polar Molecule
•Electronegativity difference of:
0.5 or less bond is nonpolar
greater than .5 bond is polar
greater than 1.7bond is ionic
Number of
Electron
Domains
Electron
Geometry
Bonding
Pairs
NonBonding
Pairs
Molecular
Geometry
Hybridization
Bond
Angle
2
Linear
2
0
Linear
sp
180º
3
Trigonal
Planar
3
0
Trigonal
Planar
sp2
120º
2
1
Bent
4
0
Tetrahedral
3
1
Trigonal
Pyramidal
sp3
109.5º
2
2
Bent
5
0
Trigonal
Bipyramidal
sp3d
120º
90º
sp3d2
90º
90º
Trigonal Pyramidal
Image from:
http://www.chem.hbnu.edu.cn/jysweb/whjys/
wangwd/jghxywkj/AX3E1.gif
4
Seesaw
Image from:
http://www.chem.hbnu.edu.cn/jysweb/whjys/
wangwd/jghxywkj/AX4E1.gif
5
6
Square Planar
Image from:
http://www.chem.hbnu.edu.cn/jysweb/whjys/
wangwd/jghxywkj/AX4E2.gif
Tetrahedral
Trigonal
Bipyramidal
Octahedral
4
1
Seesaw
3
2
T-Shaped
2
3
Linear
6
0
Octahedral
5
1
Square
Pyramidal
4
2
Square
Planar
• Helps decide which Lewis structure is most
reasonable
• It is the charge the atom would have if all the atoms
in the molecule had the same electronegativity
• To calculate: 1. Count all nonbonding electrons per
atom. 2. Count half of any bond. 3. Subtract valence
electrons by the number assigned to each atom.
Valence bond theory says
that electrons in a covalent
bond can be found in a
section that is the overlap
of the individual atomic
orbitals that are bonding
Image From: http://courses.chem.psu.edu/chem210/quantum/pictures/sigms.gif
Sigma bonds = formed by the overlap of two s orbitals, an s and a p
orbital, or two p orbitals
Pi bonds = formed by the overlap between two p orbitals oriented
perpendicularly to the internuclear axis
Ex) O=O has 1 sigma bond and 1 pi bond
N≡N has 1 sigma bond and 2 pi bonds
Where do Pi bonds come from?
•Only period 2 elements form pi bonds because they are small in size and
therefore form short bonds.
•Due to these short bonds when 2 of these atoms form a sigma bond, they are so
close together, that an additional energy level(orbital) overlap and a pi bond is
formed
½ [ (#bonding e-) - (#anti-bonding e-)]
• E = hf
• KE =
2
½mv
•c = 2.9979 ×
8
10
m/s
•h = 6.626 × 10-34 J•s
• c = fλ
•me = 9.11 × 10-31 kg
•λ=
•NA = 6.022 ×
23
10