Download Chemistry - Unit 6 What do you need to know?? This chapter is on

Chemistry - Unit 6
What do you need to know??
This chapter is on Modern Atomic Theory and it is, in some ways, a history lesson in the development
of the body of knowledge surrounding the atom.
We've learned earlier, about
 Democritus - 465 BC - came up with notion of atoms which he described as an elemental
particle
 Sir Robert Boyle - latter half of 1600s - came up with gas laws (P1V1 = P2V2), and introduced
the idea that smaller particles of elements join together as molecules
 Lavoisier - later half of 1700s - discovered nitrogen and is generally onsidered Father of
chemistry
 Dalton, John - (1766-1844)- The first great chemical theorist; proposed atomic theory (1807);
stated law of partial pressure of gases. His ideas led to laws of multiple proportions, constant
composition and conservation of mass.
 Dalton's theory had four main concepts:

1. All matter is composed of indivisible particles called atoms. Bernoulli, Dalton, and
others pictured atoms as tiny billiard-ball-like particles in various states of motion. While
this concept is useful to help us understand atoms, it is not correct as we will see in later
modules on atomic theory linked to at the bottom of this module.
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2. All atoms of a given element are identical; atoms of different elements have different
properties. Dalton’s theory suggested that every single atom of an element such as oxygen is
identical to every other oxygen atom; furthermore, atoms of different elements, such as
oxygen and mercury, are different from each other. Dalton characterized elements according
to their atomic weight; however, when isotopes of elements were discovered in the late
1800s this concept changed.

3. Chemical reactions involve the combination of atoms, not the destruction of atoms.
Atoms are indestructible and unchangeable, so compounds, such as water and mercury calx,
are formed when one atom chemically combines with other atoms. This was an extremely
advanced concept for its time; while Dalton’s theory implied that atoms bonded together, it
would be more than 100 years before scientists began to explain the concept of chemical
bonding.
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4. When elements react to form compounds, they react in defined, whole-number ratios.
The experiments that Dalton and others performed showed that reactions are not random
events; they proceed according to precise and well-defined formulas.

Avogadro - right around 1800 - said that an equal volume of gas at STP contains the same
number of molecules -- 6.02 x 1023 for 22.4 liters of gas.
 Mendeleev - latter half of 1800s - discovered periodicity of elements and created the periodic
table of elements
 Thomson, Sir J.J. - 1896 - Research on cathode rays resulted in proof of existence of electrons
and said that atoms were not indivisible (as Dalton had tought) but were comprised of smaller
particles
Rutherford, Sir Ernest - early 1900s - Famous for his gold foil experiment that demonstrated that
atoms were comprised of a dense core and the electrons were tiny and surrounded the core. Discovered
half-life of radioactive elements. We also learned about Charles (Charles' Law) a Guy-Lussac (you
guessed it, Guy-Lussac's Law), but I'm leaving them out right now as we focus on the development of

modern atomic theory.
As you may have already guessed, Electrons are the thing! They absolutely hold the key to
understanding why substances behave as they do. So this whole lesson is on how we've learned what
we've learned about electrons and how the chemists have calculated their function(s)!
From here on out in this study guide, I'm going to lead you through parsing the absolute essentials from
the interesting-but-not-necessary-to-memorize bits of this Unit 6!
This sentence is the link between electrons and the first sub-section:
We understand electronic structure (number of electrons, where they are found in
the atom and the energies they possess) from the study of light emitted or light
absorbed by substances.
This is where electromagnetic radiation comes in. This radiation:
 carries energy through space
 includes visible light along with everything from gamma rays to radio waves with x-rays,
microwaves, infrared etc. etc. included also!
8
 ALL move through a vacuum at 3.00 x 10 m/sec (aka "Speed of light"!)
 are characterized by wave-like behavior
Draw a series of waves:
Label the wavelength and the amplitude.
Now, draw a palm tree somewhere in the middle of your series of waves. Can you calculate the
velocity of the wave as they pass by your palm tree? You could draw yourself in the palm tree
(remember Harold and the Purple Crayon??) and count how many peaks pass by per second. Or you
could use a formula. Since we know that all electromagnetic radiation moves at the same speed, we
can figure out that the wavelength and the frequency (how many wave peaks passed by your plam tree
in a second).
Velocity (speed of light (c) or 3.00 x 108 m/sec) = wavelength (λ) x frequency (ν)
c = (λ) (ν)
Can you rewrite this formula in terms of calculating wavelength and in terms of calculating frequency?
~ Sodium vapor lamps are used for public lighting. They give off a yellowish light with a wavelength
of 589 nm (nm stands for nanometers and is equal to 10-9 meters). What is the frequency of this
radiation?
The color of light depends upon the wavelengths of which it is comprised. The observed color of an
object is due to the wavelengths that were not absorbed by the object. It is also dependent upon the
color of light in which the object is observed. The color of light depends on the wavelengths of which
it is composed. In the above example, the yellowish color of sodium vapor lamps is due to the light
emitted by sodium - of a particular wavelength.
If you look at a red rose under light emitted by a sodium vapor lamp versus in the sunlight, it will
appear to be different colors - this is due to the fact that the object can only reflect wavelengths of light
that are being emitted by the light source!

An absorption spectrum will show only those wavelengths of light which are reflected by an
object and not absorbed.

An emission spectrum will show those wavelengths of light emitted by a light source.

Different elements have different emission spectrums. That's why you can tell what gas is
responsible for what color in the Northern Lights. Those electrons in each specific elemental
gas emit a unique spectra -- and we see it as a unique color!
Here's a clarification of a term Thinkwell uses - intensity of light - This means the most intense
wavelength of the spectrum obtained. In a band of rainbow colors that represent the emission spectrum
of gases, those colors that are most prominent are of greater light intensity. And that explains why a
poker stuck into a fire first becomes "red-hot", then "white-hot". (Hmmm... You want further
explanation?)
 The pokers are at different temperatures but made of the same elemental materials.
 The pokers emit different intensities and wavelengths of electromagnetic radiation in the visible
spectrum at different temperatures.
The fact that these alleged pokers do not emit intensities predicted by classical physics is what
Thinkwell introduced as the "ultraviolet catastrophe". The catastrophe was this: Here these scientists
were thinking they were getting a handle on how light/electrons/atoms worked and then their elegant
prediction failed in the real world. Catastrophe! If you were one of those scientists.
But this was important to the development of the modern atomic theory because it led to the idea that
energy is quantized. That is, it is not a continuously emitted stream but rather it has a lower limit of
indivisibility - a point at which it exists as a 'chunk' that cannot be broken down further. This was a
revolutionary finding!




Max Planck is the scientist credited with this idea of energy quantization in 1900.
The smallest packet of light is called a quantum or photon.
The energy (E) of a quantum is related to its frequency (n) by Planck's constant (h) 6.63 x 10 -34
Joule seconds
Electromagnetic energy is always emitted or absorbed in whole number multiples of (h*n).
So, wonder of wonders, a formula was derived to allow scientists to calculate the energy of a quantum.
It is this:
Energy = Planck's Constant (h) * frequency (ν)
E = (6.63 x 10-34) (ν)
This formula allows you to calculate the value of the energy in one quantum. To calculate more than
one, the right side of the equation is multiplied by (n) or the number of quanta to become:
E = (6.63 x 10-34) (ν) (n)
~ Calculate the amount of energy (one quantum) that an object can absorb from yellow light with a
wavelength of 589 nm. (Remember, if you know wavelength and the speed of light, you can calculate
frequency.) Also remember, the joule is the SI unit for measuring energy and = kg • m2
s2
Watch your units as you do this calculation.
The point of this formula is to explain why the intensities of light emitted do not follow the predicted
model using classic physics. Planck's constant makes the difference! And demonstrates that energy is
quantized.
The Photoelectric Effect
Once the chemists had this packet business figured out, they went on to wonder why, when they shined
light on metals, their ammeter (a gadget that measures current or amps) showed an increased current.



Light shining on metals causes the surface to emit electrons. (Remember your physics when
you learned that current is merely a flow of electrons? Hence the ammeter reading!)
For each metal, there is a minimum frequency of light below which no electrons are emitted,
regardless of the intensity of light. This is called the threshold frequency, or v0
The higher the light's frequency above this minimum value, the greater the kinetic energy of the

released electrons.
Since different metals have different atomic structures, the quanta of light needed to overcome
the attractive forces within the atom differs for each element. It is known as the work function
of the metal or Φ.
It was Einstein, in 1905, who deduced the basis of the photoelectric effect.
I like an example that Dr. Blaber, at Florida State University uses:
The Photoelectric effect as a carnival game:
"A popular carnival game is where you are given a giant mallet and have to hit a pad on the
ground. This sends a small metal slug shooting up a vertical track and, if you hit hard enough,
it will hit a bell at the top. This is like the photoelectric effect - if the electron will be releaed
form the atom if it absorbs a photon with enough energy. Imagine that you have a 10 year old
child and Arold Schwartzenegger. The child is too weak to hit the pad hard enough to ring the
bell. It doesn't matter if you have an army of 10 year olds (intensity) lined up to take their turn none of them will ever hit it hard enough to ring the bell. However, Arnold (being Arnold) will
have no problem ringing the bell.
Thus, if a light shining on a metal does not have enough photons with the necessary energy to
cause an electron to be ejected, then it does not matter how bright the light is. The key thing is
to increase the energy of the individual photons, and this is achieved by increasing the
frequency (i.e. decreasing the wavelength)."
We would, therefore, expect to see more electrons from atoms to be ejected, with high kinetic energy,
the smaller the wavelength (and greater the frequency). And we do. X -rays are far more problematic
to body tissue for example, than are radio waves! You never have to listen to a radio wearing a lead
apron!
Which brings us to 1913. That's eight years after Einstein worked out the business about metals and
light and electrons.
The Bohr Model
Basically, what Bohr said, in the most simplistic language, is that if the energy of electrons is
quantized, then the distance from the nucleus is also quantized. His model includes these points:
 He assumed that electrons were orbiting the nucleus much like planets orbit the sun.
 From classical physics, a charge traveling in a circular path should lose energy by emitting
electromagnetic radiation.
 If the "orbiting" electrons lose energy, they should end up spiraling into the nucleus (which they
do not). Therefore, once again, classical physics does not apply!
 So Bohr explored the notion of quantized energy that Planck first talked about and and
proposed:
 that only orbits of certain radii, corresponding to defined energies, are 'permitted'
An electron orbiting in one of these "permitted" orbits:
 Has a defined energy state
 will not radiate energy

will not spiral into the nucleus
Bohr did a masterful job of explaining the behavior of the one electron in hydrogen. Unfortunately, it
does not work for other elements.
But his work was, and is, significant in the furthering of the current understanding of atomic theory.
The Heisenberg Uncertainty Principle
There were a couple of point here:
 One locates things by sight by detecting the photons that bounce off the object from a light
source. And the wavelength of light used to view an object must be smaller than the object.
However, the smaller the wavelength, the higher the frequency and the higher the energy of
each photon. At some point, this becomes destructive. (Remember, trying to locate a cat with
tennis balls??)
 The Heisenberg Uncertainty Principle says that the smaller an object the greater the relative
uncertainty in knowing both the position and the momentum of that object.
The point of this seemingly stray thought is that it applies to electrons. We cannot know where they are
in relation to the nucleus! And we cannot think of them as moving in well-defined orbits around the
nucleus, however nice that idea is.The formula that represents the principle is not something you need
to memorize for this year.
The year is 1924 and a scientist named de Broglie made a daring hypothesis:
If radiant energy could, under appropriate circumstances behave as though it were a stream of
particles, then could matter, under appropriate circumstances, exhibit wave-like properties?
For example, the electron in orbit around a nucleus. De Broglie suggested that the electron could be
thought of as a wave with a characteristic wavelength.
Here's his formula:
wavelength = Planck's constant
momentum
OR
λ = h/mass*velocity
To complete these calculations you need to know
-34
 Planck's constant - 6.63 x 10
kg * m2 ;
-10
 the mass of electron - 9.109 x 10
kg
And you must watch your units!!! ~ What is the characteristic wavelength of an electron with a
velocity of 5.97 x 106 m/s?
In 1927, Davisson and Germer proved that particles have wave properties just as de Broglie had
predicted when they passed electrons through a crystal and found an interference pattern. (Have you
ever sent two waves , one from each end, traveling the length of a stretched out Slinky? Sometimes the
waves multiply, sometimes they cancel each other out.)
Quantum Mechanics
Let's figure out what that term means - Quantum - having to do with those tiny packets of energy that
cannot be subdivided; Mechanics - how they move.
We give credit to a scientist named Schrödinger for developing an equation that incorporates both
wave-like and particle-like behaviors for electrons. He worked in the latter part of the 1920s.
Since we know from the Heisenberg Uncertainty Pricnicple that we cannot know for certain the exact
location of an electron, Schrödinger describes the probability of a location of an atom.
So what we have are probability distributions for a given energy level of the electron. Electrons still
have a set energy level associated with the energy of a photon, but the visualization of just where they
are in relation to the nucleus has changed. Thanks to Schrödinger!
We call these areas of probable electron inhabitance "orbitals" (very different from the Bohr's "orbits" ,
eh?)
Schrödinger devised an equation that included three variables and places electron probability in three
dimensions. His equation gives a series of wave functions that describe electron probability
distributions but we are NOT going to work these out ourselves! We are going to grasp some of the
implications.
Not surprisingly, the lowest energy level of the hydrogen atom calculates to be -2.18 x 1018 J - which is
the same thing that Bohr calculated. Remember that his model worked for hydrogen, or for one
electron, but then fell apart after that?
Schrödinger's calculated energy level also has a shape associated with it - that's the probability part.
The electrons are most likely somewhere in this defined region.
Bohr defined quantum levels, saying that each level corresponded with a certain energy. In our previous
calculations, n was used to denote the quantum level and could only be a positive integer.
Now, Schrödinger's model uses three quantum numbers: n, l and m1 to describe an orbital.
The principal quantum number 'n'
 has values of 1, 2, 3, etc. (positive integers)
 As n increases, the electron density is further away from the nucleus
 As n increases the electron has a higher energy and is less tightly bound to the nucleus
The second quantum number 'l'
 has integral (meaning they must be an integer) values from 0 to (n-1) for each value n
 instead of being listed as a numerical value, 'l' is referred to by a letter (s = 0, p = 1, d = 2, f = 3)
 defines the shape of the orbital
The third quantum number m1
 has integral values between 1 and -1 including 0
 describes the orientation of the orbital in space
For purposes of this class, you do not need to know which orbitals are which shapes. You just need to
know that we can no longer picture electrons in simple orbits around the nucleus. We still talk about
electron shells, those being the collection of orbitals with the same n value. Look at the chart below to
better put this all together. This chart outlines the description of the electron orbital with the principle
quantum number of 3.
n
(principal
quantum number)
l
(defines shape)
Subshell
designation
m1
(defines
orientation)
Number of orbitals
in subshell
3
0
3s
0
1
1
3p
-1, 0, 1
3
2
3d
-2, -1, 0, 1, 2
5
Of importance to you as we continue the study of chemistry is this:
 For each m1 quantum number within an electron shell, there are two electrons, each with an
opposite spin
 Only those electrons in the outermost shell are available for chemical interactions. They are
called valence electrons.
 Electrons can move from one shell to another (leaving the lower shell empty) with the addition
of energy. Moving back down emits energy.
We'll study next, how the electrons fill up the shells.
For now, you've learned ALOT about the history of the development of modern atomic theory.