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Notes # _____
Chemistry Chapter 5 Notes
Date: _______
5.1: Electrons
As mentioned in Chapter 4, the Bohr Model proposed that the electron cloud was broken
into regions called energy levels. Each energy level could contain a specific number of
electrons (2,8,8,18,18,32,32) each at a fixed distance from the nucleus, called a
________________. In the Bohr Model, electrons orbited about the nucleus much like
planets around the sun. In reality, the Bohr Model only really predicts the electron
behavior for the most simple atoms, hydrogen and helium. Further study of the energy
levels led to a surprising discovery, that each energy level is further split into slightly
different energies:
energy
___
___
___
___ ___ ___
___ ___ ___ ___ ___
(4th)
___ ___ ___ (3rd)
___ ___ ___ (2nd)
___ (1st)
The modern model of the atom, called the Orbital (or Quantum-Mechanical) Model, uses a
complex mathematical formula, the ____________________________ Equation. It describes
the electron as behaving more like waves than as particles and the electrons can exist in
discrete locations called __________________. __________________ are regions around the
nucleus that make the “electron cloud.” These regions indicate where an electron is most
likely found (so they are mathematical __________________). Imagine the electron cloud as
a large hotel--the orbitals can be thought of as individual rooms on different floors of that
hotel. The probability of an electron being found in any location produces a plot of a
variety of locations. Imagine a rapidly spinning propeller blade spinning in 3 dimensions:
According to the Schrodinger Equation, each orbital is described by a set of quantum
numbers. The first, called the ________________ quantum number (n), is an integer
(1,2,3,...) used to describe the energy level. The principal quantum number in turn
determines the number of _________________. So for n =1, there is ___ sublevel; for n = 2
there are ___ sublevels, and so on. Each sublevel is comprised of a certain number of
orbitals. There are 4 kinds of orbitals, always written in lowercase:
s,p,d,f
Each orbital has its own shape and number of electrons:
There is ______ s orbital
There are _____ p orbitals
There are _____ d orbitals
There are _____ f orbital
(really complex!)
Each individual orbital can contain at most ____ electrons so the total number of
electrons in any ________________ is equal to the number of orbitals in that sublevel times
2. So:
s sublevels contain _____ e- ( 1 x 2 )
p sublevels contain _____ e- (
)
d sublevels contain _____ e- (
)
f sublevels contain _____ e- (
)
So if n = 1 there can be at most _____ electrons:
if n = 2 there can be at most _____ electrons:
if n = 3 there can be at most _____ electrons:
if n = 4 there can be at most _____ electrons:
Note this pattern on the periodic table:
Notes # _____
Chemistry Chapter 5 Notes, Part II
Date: _______
5.2 Electron Configuration
When electrons are arranged in the electron cloud they do so in order to produce the
lowest energy, most _________ arrangement. The orbitals that make up the arrangement of
the electron cloud follow a very specific pattern seen in the periodic table:
Another way to look at this pattern is:
When you put together the pattern along with the number of electrons in a sublevel you
get the following, based upon the ______________ principle, that electrons occupy lowest
energy orbitals first There are two more “rules” to follow in order to determine the electron
configuration of an element:
(1)
_______________ __________________ Principle: any atomic orbital can contain at most
______ electrons but only if they have opposite __________. __________ is a quantum
mechanical property that can be described as either ____________________ or
________________________________.
_________ Rule: electrons will fill a sublevel such that electrons fill empty orbitals of
a sublevel before they partner.
Energy Increses ---->
(2)
A simple shorthand method for showing the electrons is called the _________________
________________________. The EC shows the location of every electron from lowest to
highest principal quantum number by sublevel and the number of electrons in each
sublevel by using this symbol:
#
n(s,p,d,f)
where n is the principal quantum number, and # is the number of electrons in that
sublevel. The sum of the superscripts will equal the _________________ for a neutral
element! For example, lithium has three electrons, so its EC will be ______________.
Nitrogen will be ___________________.
Practice Determine the EC for the following:
(a) oxygen: _________________________
(b) calcium: __________________________________
(c) iron: ____________________________
(d) bromine: __________________________________
Noble Gases are particularly stable so their EC are often used as shorthand for heavier
elements. For example, Ne is __________________, therefore every element in the next row of
the periodic table has the same first 10 electrons so the notation [Ne] is used to show
these 10 electrons. The EC for sodium could be written: _______________; calcium could be
written: _____________________.
There are a few exceptions to the Aufbau Principle. These occur when it is more stable to
have _______________ sublevels rather than partially filled ones. Although filled sublevels
are most stable, half-filled sublevels are usually more stable than partially filled ones. Two
major exceptions that you must know:
chromium “should be”: _____________________________
but “is”: ____________________________
copper “should be”: ______________________________
but “is”: _____________________________
There are other exceptions. Look at the 5th row of YOUR periodic table and see if you can
determine the elements that don’t follow the Aufbau Principle:
Unpaired Electrons
The most important information we can get from an element’s EC is the number of
__________________ electrons. These are electrons that occupy an orbital by themselves
that eventually can form a bond. Electrons that partner in an orbital are called orbital
__________. These can quickly be determined by subtracting the number of unpaired
electrons from the atomic number, then divide by two.
Practice:
Determine the number of unpaired electrons and orbitals pairs in the following atoms:
EC
(a) oxygen
(b) chlorine
(c) arsenic
(d) barium
(e) tellurium
Aufbau diagram
unpaired e-
orbital pairs
Notes # _____
Chemistry Chapter 5 Notes, Part III
Date: _______
5.3 Light
Waves are disturbances that transmit _______________ either through a substance, called a
________________, or through empty space ( a ________________). __________________________
waves are produced by disturbances in electric and magnetic fields (hence, EM).
____________________ often cause waves. Light is EM radiation that travels as a transverse
wave but is _______________ (given off) or _______________ (taken in) as a massless particle
called a _____________. This is called the “wave-particle duality” of light. James Maxwell
proposed the Electromagnetic theory that states an electromagnetic wave, such as light, is
a periodic disturbance involving electric and magnetic fields. The speed of light is affected
by its medium, decreasing when the medium gets ______________. Since light travels as an
EM wave it does not require a medium so it travels ____________ in a vacuum. The speed of
light, represented as ____, is equal to _____________ m/s ( ______________________m/s) or
about _______________ miles per second!! At this speed light would travel _____ times
around the Earth in one second; from the moon to the Earth in _____ seconds; and from
the Sun to Earth (93 million miles) in ______________________.
Wave Characteristics
The _____________ (
) is the time required for one full wave to pass a point--it is
measured in _____________. The _________________ (
is the reciprocal of the ______________ ( so
) is how often a wave passes by--it
= –—– ). So, if the period is 10.0 seconds
the frequency is ________ or __________. [ or ________ which is also _____________ (
Example 1:
_____________________________________________
since T = ______ f = ________ = _________
)]
Figure 5-10 (on p. 139) shows the range of visible light and the entire EM spectrum. Note:
_______ frequencies = _________ wavelengths and the ___________ the frequency the
___________ the energy. Visible light has wavelengths between ______ nm (
_______ nm (
).
Note 1 ________________ (
) = _______ m or _______ nm.
) and
Speed of Light
Recall that velocity equals distance ÷ time so:
Example 2:
What is the velocity of a 1.0 m wave if its period is 4.0 seconds?
Example 3:
What is the wavelength if a wave traveling at 15.0 m/s has a frequency of 5 Hz?
Production of Light
When an atom gains a specific energy, called a ______________, electrons in its electron
cloud will jump to higher energy levels (further from the nucleus) then immediately “fall”
back down to lower levels. Regardless of how the energy was originally gained by the
electron (heat, electricity, sound), the released energy is always light. Depending on the
amount of energy released different colors can be produced following the pattern:
ROY G BIV.
The pattern of light produced is called the _____________ _____________ _______________ of
the element and might look like this for the element to the left:
Each element has its own unique fingerprint that
can be used to identify it. When all of the electrons
of an atom are in their lowest energy the atom is
said to be in its _______________ state. If a
quantum of energy is applied to the atom an
electron (or more) can jump to higher energy levels
and the atom is in an __________________ state.
When the electron falls back to the ground state the energy released can be found using
the equation:
Where E is the energy in J
h, is the Planck’s constant (6.626 x 10
–34
J·s)
–1
f is the frequency in hertz (s )
or, if one knows the wavelength the following equation can be used:
Where c is the speed of light
λ is the wavelength in meters
The atomic emission spectrum for hydrogen had been intensely studied. Since there is
only one electron in its spectrum we can look closely at its pattern. In the diagram below
one ground state and two excited states are shown. Each of the series has a name and a
corresponding energy in the EM spectrum:
Notes # _____
Chemistry Chapter 5 Notes, Part IV
Date: _______
5.3 Quantum Mechanics & Quantum Numbers
We now know that light can act as both a particle or a wave. The question we need to
consider is the opposite: can matter act as a wave also? That question was asked in the
1920’s by a French scientist ____________ ________________ and confirmed only a few years
later. This discovery eventually led to a remarkable device: the ________________
_____________________. Because the wavelengths of a moving electron are shorter than the
wavelength of light this microscope can “see” with millions of times more clarity (and
magnification). Since waves have a definite wavelength depending on their energy only
certain locations surrounding a nucleus can allow an electron:
De Broglie’s discovery led to a new branch of science: quantum mechanics. Unlike
______________________ mechanics that describes the behavior of larger bodies, quantum
mechanics describes the behavior of ______________________ particles and atoms in
wavelike terms. A side effect of quantum mechanics is that the behavior of an electron is
impossible to detect without affecting its energy. This is known as the ___________________
__________________________ Principle. To detect an electron one can imagine taking a
picture of the electron as it moves. To do so you would need to use a “flash” which, in
essence, would add energy to the electron. When the energy is absorbed by the electron, it
changes position, thus producing a different energy! Thus the more we attempt to
measure the position (or energy) of an electron, the more we affect it!
Quantum Numbers
Schrodinger’s Equation is based on very complex mathematics but at its base is a set of 4
simple numbers that are used to describe the energy and location of an electron:
n, is the primary quantum number an equal to the energy level. The lowest value for n is 1
and as it increases the electron moves further away from a nucleus. If the distance gets
great enough the electron leaves the atom producing a _________________ charged ion.
l, is the ________________ quantum number. Each value of n has values of l from 0 to (n-1)
So if n = 2, then l = ________. If n = 4, then l = __________. This is the quantum number
that determines the type of sublevel (s,p,d,f). l = 0 corresponds to an ‘s’ sublevel, l = 1
corresponds to a ‘p’ sublevel, l = 2 is a “d” sublevel, l = 3 is an “f” sublevel.
m, is the __________________ quantum number, due to the splitting of a sublevel into the
corresponding number of orbitals. The quantum number m can range from -l to +l in
whole number steps. Thus if l =1 then m can be ________________.
s, is the ___________ quantum number and is used to show the orientation of an electron
in an orbital. By convention the first electron is said to spin clockwise and is assigned a
value of +1/2; the second electron in an orbital must have counterclockwise spin and is
assigned a value of -1/2.
The combination of all four quantum numbers is unique to an electron in an atom--there
is a different quantum number for every electron in an atom. The symbol we used to show
this combination is:
| n, l, m, s >
For example: determine the quantum numbers for the following:
(a) the first electron in the second 4p orbital:
(b) the second electron in the sixth 5f orbital:
Draw the location of the following electrons:
(a) |3,2,+1,-1 /2 >
(b) |5,3,-2,+1 /2 >