PreAP Chemistry Chapter 4 Notes
Section 4.1 The Development of a New Atomic Model
Previously, Rutherford reshaped our thought of the atom by showing the protons were located in the _____________________ of the atom,
but he could not model for us where the electrons were, other than outside the nucleus somewhere. Fortunately, studies into the properties of
light and the effects of light on matter soon gave clues to where electrons actually are.
Light is a small part of all the radiation (something that spreads from a source) called _____________________ radiation. Electromagnetic
radiation is _____________________ in the form of waves (of electric and magnetic fields). Electromagnetic radiation includes radio waves,
microwaves, infrared, visible light, X-rays, and Gamma rays. All these together are considered the Electromagnetic
_____________________
As all the forms of electromagnetic radiation are waves, they all have similar properties.

All electromagnetic radiation travels at the ________________ ___ _____________ (c), 299,792,458 m/s (3 x 108) in a vacuum

The _____________________ is the top of the waves, the _____________________ is the bottom of the waves, and the
_____________________ is a measurement from the rest or zero line to a crest or trough

The _____________________ (λ – lambda) is the distance between successive crests/troughs and is measured in
_____________________ (often nm = 10-9 m)

The _____________________ (ν – nu) is the number of waves that pass a point in one second and is measured in
_____________________ (per second – can be written as s-1) or Hz (Hertz)
c
The speed of a wave is directly proportional to the wavelength and the frequency; c = λν is the formula.
λ
Example: A certain violet light has a wavelength of 413 nm. What is the frequency of the light?
ν
Unfortunately, thinking of light as waves led to a problem. It was noticed that if light strikes a metal, then sometimes it could cause
_____________________ to be emitted (leave the atoms entirely – like in a solar panel); called the _____________________ effect. If light
was a wave, then all amounts of light energy should cause this to happen, but this was not the case. It always took some
_____________________ amount of energy to get the electrons to be emitted. This lead Max Planck to theorize that light must carry energy
in basic minimum amounts that he called _____________________ Like a delivery person cannot correctly deliver half a box, the electrons
in atoms cannot gain a fraction of a quantum of energy (it has to be in _____________________ numbers). He proposed that this energy was
directly proportional to the frequency of the electromagnetic radiation and a constant, now called Planck’s constant.
E=hν
where
E
E = energy, measured in Joules (J)
h = Planck’s constant, 6.626 x 10-34 Js
ν = frequency, in 1/s
Example: What is the energy content of one quantum of the light with a wavelength of 413 nm?
h
ν
In 1905 Einstein used Planck’s work to propose that electromagnetic radiation had a dual ______________-_____________________ nature.
As a particle, electromagnetic radiation carries a quantum of energy of energy, has no mass, and is called a _____________________. So to
get an electron to emit from a metal, it must be struck with a photon having quantum energy big enough, or nothing will happen. Each metal
requires different quantum energy, thus each metal can be identified by the frequency of light needed to emit electron.
This idea was expanded upon to develop an idea of where the _____________________ were in an atom. It was found that low pressure
_____________________ could be trapped in a tube and electrified, and would then glow a color particular to the gas inside. Furthermore
this light could be passed into a prism, and instead of getting the entire spectrum (rainbow) of colors, only certain wavelengths of light would
be seen as small bars of color, called a ________________-_____________________ spectrum.
This would indicated that the electrons in an atom were only absorbing _____________________ amounts of energy from the electricity,
causing the electrons to move from their _____________________ state (normal position close to the nucleus) to an
_____________________ state (higher energy position further away from the nucleus). The electrons do not stay in the excited state for long
and fall back to their ground state, losing the energy _____________________ to what they gained.
Niels Bohr used this to develop a model of the atom where the electrons could only be in certain, specific _____________________ level (n)
orbits around the nucleus. Just as you cannot go up half a rung on a ladder, the electron could not go up a partial energy level. The electrons
gained or lost enough energy to move a _____________________ number amount of energy levels (n) away from or closer to the nucleus, or
it did not move. He calculated the amount of energy needed for an electron of hydrogen to move between each energy level (n) (which was
not constant) and his calculations _____________________ with experimental results.
The _____________________ series of hydrogen spectral lines refer to the four lines seen in the visible light region (the four colored bars).
If the electron was excited to energy level (n) 6, 5, 4, or 3 and fell to energy level (n) 2, the resulting energy given off would have a frequency
in the _____________________ region of electromagnetic radiation. (One line for dropping from 6 to 2, one for 5 to 2, one for 4 to 2, and
one for 3 to 2).
However, there are other possibilities. If the electrons drop from n=6, 5, or 4 to n=3, then the energy given off is not big enough to be seen as
it is in the _____________________ region. These three lines in the infrared region are referred to as the _____________________ series.
If the electrons drop to n=1, then the five lines given off are too high in energy to be seen, as they are in the _____________________ region.
These lines are referred to as the _____________________ series.
Model of Atom Review:
1.
Thomson’s Plum Pudding Model – the atom is a ball of evenly spread _____________________ stuff with random
_____________________ particles (electrons).
2.
Rutherford’s Nuclear Model – the atom has a central _____________________ containing the positive particles (protons) with the
electrons outside.
3.
Bohr’s Orbital Model – The electrons circle the nucleus in specific energy _____________________ like the planets orbit the sun.
Unfortunately this only works for atoms with _____________________ electron…
4.
Quantum Mechanical Model – electrons are found in specific _____________________ around the nucleus, but the exact location
of the electrons inside the regions _____________________ be determined
Section 4.2 The Quantum Model of the Atom
The quantum mechanical model of the atom is built on the ideas and calculations of several scientists.

Louis _____________________ suggested a way to show that a particle could have wave like behavior with the equation:
λ=
h
where
mv
h = Plamck’s Constant
m = mass of particle
v= velocity of particle

The _____________________ Uncertainty Principle states that it is impossible to know both the velocity and the location of an
electron at the same time. If the _____________________ was known, then there is no way to know where it has moved to, and if
the _____________________ is known, then there is no way to know where it was.

_____________________ developed wave-based equations that form the basis of the current Quantum theory, which
mathematically describe the probably location of electrons, often referred to as an electron cloud. The electrons clouds describe
areas around the nucleus with a 90% chance of finding the electron inside. Solving the equations has lead to
_____________________ ________________, which will be studied later.
The quantum mechanical model starts with a _____________________ Quantum Number (n), which is the basic energy level of an electron,
and often matches the _____________________ number. Possible values (currently) are 1-7.
Inside the principal quantum energy level are sublevels that correspond to different _____________________ shapes. The sublevels are
designated as s (sharp), p (principal), d (diffuse), and f (fundamental).
Inside the sublevels are orbitals, specific regions with a 90% probability of finding electrons.

s –orbitals are _____________________ shaped clouds around the nucleus

p -orbitals are _____________________ shaped clouds with the nucleus between the lobes

d and f are much more complex in shape
Each sublevel has room for a different amount of electrons, because an orbital can hold two electrons, then each sublevel has a different
amount of orbitals

s –sublevel can hold _____________________ electrons, so it has_____________________ orbital (shape)

p –sublevel can hold _____________________ electrons, so it has _____________________ orbitals (shapes)

d –sublevel can hold _____________________ electrons, so it has _____________________ orbitals (shapes)

f –sublevel can hold _____________________ electrons, so it has _____________________ orbitals (shapes)
To know the maximum amount of electrons that could be in any principal quantum level (and the number of elements that could be
represented) use the formula 2n2
if n=1, then _____________________ electrons will fit
if n=4, then _____________________ electrons will fit
Section 4.3 Electron Configurations
In order to show on paper where electrons are likely to be located in an atom, orbital filling diagrams and electron configurations are drawn
or written. When this is done, three rules must be followed:
1.
_____________________ principle – electrons fill lower energy levels first, thus 1 before 2 and s before p, etc.
a.
orbitals within a sublevel are _____________________ in energy (called degenerate)
b.
the principal energy levels often _____________________ making them seem a little out of order
c.
_____________________ are used to represent orbitals
d.
Another way of writing the aufbau principal diagram:
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d
7s 7p
2.
3.
________________Exclusion Principle – an orbital (box) can hold a maximum of two electrons (arrows)
a.
for two electrons to fit, they have to have _____________________ spins
b.
↑ for _____________________ electron in the orbital
c.
↑↓ for _____________________ electrons in the orbital (opposite spins)
________________Rule – when electrons occupy degenerate orbitals, one electron is placed into each orbital with parallel spins
before doubling up
____ ____ ____
NOT
____ ____ ____
3p
3p
http://intro.chem.okstate.edu/WorkshopFolder/Electronconfnew.html
Orbital Notation shows the arrows in the boxes to represent the electrons in an atom. To shorten this process, an electron configuration can
be written. It leaves out the information about the number of orbitals in each sublevel, so it will be expect you remember that information.
It has the general form ________________
n = principal quantum number (1-7…)
Δ = sublevel letter (s, p, d, or f)
° = number of e- in that orbital (1-14)
The sublevels can be listed in order of filling as from the _____________________ table, but for the notes, quizzes, and tests the sublevels
will be _____________________ together by principal quantum number.
Examples:
Ni
1s 2s 2p 3s 3p
3d
4s 4p
4d
4f 5s 5p
5d
5f
Sn
1s 2s 2p 3s 3p
3d
4s 4p
4d
4f 5s 5p
5d
5f
A few elements have electron configurations that _____________________ follow the normal pattern. These two _____________________
are represented by the first elements in their group:
By normal configuration:
Cr
1s 2s 2p 3s 3p
3d
4s 4p
4d
4f 5s 5p
5d
5f
Cu
1s 2s 2p 3s 3p
3d
4s 4p
4d
4f 5s
5d
5f
5p
By actual (seen in nature) configuration:
Cr
1s 2s 2p 3s 3p
3d
4s 4p
4d
4f 5s 5p
5d
5f
Cu
1s 2s 2p 3s 3p
3d
4s 4p
4d
4f 5s 5p
5d
5f
The other elements in those groups work in similar matters. Instead of writing _____________________ a half filled sublevel seems to be
more stable, so make it _____________________ Instead of writing a _____________________ a completely filled sublevel seems to be
more stable, so make it _____________________
If writing out the entire electron configuration is too much, we can use the previous (in the periodic table) _____________________ gas to
take the place of part of the electron configuration:
Examples:
Polonium:
1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p4
Xenon:
1s22s22p63s23p64s23d104p65s24d105p6
Polonium:
[Xe] 6s24f145d106p4
When the electron configuration is written for an element using the noble gas configuration the electrons written after the noble gas are the
ones that appear on the _____________________ of the atom. These electrons are called _____________________ electrons. When
elements bond to form compounds, it is these electrons that are involved. The _____________________ of valence electrons makes a big
difference in how the element will bond, so to make it easy to predict, we draw electron _____________________ diagrams.
A) In an electron dot diagram, we use the _____________________ of the element and dots to represent the number of valence
electrons.
B) Only s and p electrons with the _____________________ quantum number count for dot diagrams, even if there are d and f
electrons after the noble gas.
Examples:
Lithium
Li
Beryllium
Be
Boron
B
Carbon
C
Nitrogen
N
Oxygen
O
Fluorine
F
Neon
Ne