Atomic Structure,
Quantum #’s, e- configurations
- Periodic
Table
- Created in 1869 by
Mendeleev
- Organized so that
columns (groups) have similar properties
Electromagnetic (EM) Spectrum
- Radiant energy (light,
rays, waves) travel in waves
- Parts of a wave
- wavelength (l) distance from peak
to peak (units = m)
- frequency (n) - how many peaks
pass per second (units =1/s = s-1 = Hz = Hertz)
- energy (E) - higher
frequency means more energy (units = Joules = J = kg m2 / s2)
- All the frequencies put
together is the EM Spectrum (important)
- Also see Fig 5.2
- All of these waves are
radiation. Radiation is mostly good - sun, TV, cell phones,
microwaves, medical applications. Some is bad is exposed too much -
X rays, gamma rays. Click
here and go to Activities then EM spectrum activity.
- Speed of light (c) =
3.00 x 108 meters per second (m/s). Which travels
faster? UV light or microwaves? Well guess what all EM radiation travels
at the same speed - light speed.
- As l increases, E decreases
and n
decreases.
Longer waves have less energy and lower frequency.
- c = ln (important) \
- Atomic
Spectra (no
math)
- The sun, old style
bulbs give off white light - all visible light together. If put white
light through a prism it splits into all the colors of the rainbow and is
continuous.
- If we heat up pure
elements or add electricity to them they emit different colors. We use
this principle in neon signs. Take glass tubes shaped however you
want and fill with different gases. Attach to electricity and the
"excited" gases emit specific colors of light. Many gases
are used besides Neon. Fluorescent bulbs use mercury gas!
Electrical energy is converted into light energy.
- If you put the light
from these heated or electrified elements through a prism you can see the
exact wavelengths involved - just a few lines show up - not all of them like a rainbow since the light is not white.
We call this line spectra. Examples:
- More are at the bottom
of this review. Line spectra are unique for each element. Can be used to
identify unknown gases. Note there may be some wavelengths in the IR or
UV regions - we just can't see them.
- Examples of burning
metals in flame tests are at the end of this review. Click
here go to Activities then flame test movie.
- Later we'll learn WHY
we get distinct or discrete lines for elements instead of a continuum.
- Plank
- Planck (~1900)
discovered E = hn where h =
6.626 x 10-34 Js (Joules x seconds) (important)
- often used to solve
problems with c = ln
- What radio station
(MHz) uses a 3.05 m carrier wave? What is the energy of this
wave? Well plug in: 3.00 x 108 m/s
= (3.05m)n
so solve for n
= 9.84 x 107 s-1 or Hz. Now change to MHz. 9.84 x 107Hz ( 1 MHz / 106 Hz) = 98.4 MHz. The E = (6.626 x 10-34 Js)(9.84 x 107 s-1) = 6.52 x 10-26
J.
- We see that radiation
(light, waves, rays) are wavelike. Well radiation is also particle
like. Einstein (1905) finally explained the photoelectric effect
(when you hit metals with radiation above a certain minimum energy it
causes electrons to be ejected off the surface of the metal.)
Einstein said this is because light is a beam composed of what he called photon
particles. Particles have mass, does light? Well yes since
black holes suck in light by gravity. Light and all the other EM
radiation is both wave and particle like. (important)
- EM
radiation exists in very small discrete units called photons. Just
like matter exists in very small units called atoms.
- Quantized means exists
in discrete units like steps, stairs, shelves, and now light and EM
radiation. We say a photon is a quantum (discrete unit) of light
energy. (important)
- de Broglie
- In 1924 de
Broglie derived this equation l = h /
mc. If something other than EM radiation is traveling then
light speed just becomes velocity: l = h /
mv. Because of h being in Joules ( J = kg m2 / s2), the mass must be in what units? yep, kg
- The mass of
an electron is 9.11 x 10-31 kg and the velocity is 2.2 x 106
m/s. Check yourself that you get l = 3.3 x 10-10
meters. You mean an electron has a wavelength? Yes, because an
electron is both wave like and particle like.
- I throw a
16 pound bowling ball down an alley at 7.55 m/s. What is the
wavelength of the ball? Try it first. Answer below.
- 16 lbs (1
kg / 2.20 lbs) = 7.3 kg bowling ball. l = (6.626 x 10-34
kg m2 s/ s2) / (7.3kg)(7.55m/s)
= 1.2 x 10-35 m Does this really mean anything?
Not much for such a large object. It is more important with tiny
objects.
- Quantum
Mechanics and the Heisenberg Uncertainty Principle
- Heisenberg (1927) said
it is impossible to know both where an electron is AND its velocity at
the same time. Kind of like taking a picture of a moving car - you can't
tell how fast it is going.
- no math
- Quantum
Numbers
- Solving quantum
mechanics we get solutions that tell us where electrons are most likely
to exist = areas of probable electron density = wave functions and we call
these shapes of electron density orbitals. See electrons can't be
just anywhere, they can exist only in certain areas or levels.
Orbitals are quantized like steps. We call these areas shells.
- Quantum Numbers (only
cover first 2, not all 4) (important)
- Principle Quantum
Number (n) = shell number. n = 1, 2, 3, 4, 5... and represents the distance from the nucleus.
- Second quantum number
(l) = subshell number. Shells are divided into one or more
subshells. l=0 for s subshell. l=1 for p subshell. l=2
for d subshell. l=3 for f subshell. Rule:
n>l .
So on the first shell where n=1, l must =0. On the second shell
where n=2, l can be =0 or 1. On the third shell where n=3, l
can be =0, 1 or 2. And so forth.
- Examples: Give
the first two quantum numbers for the following subshells:
- 2nd shell, s
subshell-------n=2 and l=0
- 7th shell f
subshell--------n=7 and l=3
- 5th shell p
subshell----------n=5 and l=1
- 4th shell d
subshell----------n=4 and l=2
- What subshell is this?
- n=3 and l=0
-----------3s
- n=1 and l=1-----------no
such thing as 1p
- n=4 and l=3-----------4f
- n=7 and l=2-----------7d
- Subshells contain
orbitals. Each orbital can contain a maximum of two electrons. (important)
- s subshell has one
orbital so a max of 2 electrons
- p subshell has 3
orbitals so a max of 6 electrons
- d subshell has 5
orbitals so a max of 10 electrons
- f subshell has 7
orbitals so a max of 14 electrons
- Questions
go to 151 Help and click on quantum numbers (important)
- Shapes
of Orbitals
- s subshell has 1 orbital
(l=0) and it is spherical. The 1s is a sphere. The 2s
is a sphere within a sphere separated by a node (area of 0% electron
density). The 3s is three spheres separated by 2 nodes. And
so forth. Fig 5.9 and 5.10
- p subshell has 3
orbitals (l=1) and they are shaped like figure 8's or dumbbells.
Each follows an axis in space - x, y and z.
- d subshell has 5
orbitals (l=2) and four of them are clover like with 4 lobes, the
other is like a p orbital with a donut.
- f subshell has 7
orbitals (l=3) and most of them are eight lobed.
- Best pictures ever of
orbitals
- Back to line spectra introduced above. So why do
the elements when heated up emit only certain frequencies of light giving
us the distinct spectra instead of a continuous rainbow?
View a video here about emission lamps like neon signs. Because
atoms can only absorb wavelengths that make electrons jump from one shell
to another shell. Electrons can't exist between shells so only a few
energies are absorbed - the ones that exactly match the energy needed to
go between shells. (important)
-
- So an
absorption spectra is all the colors of the rainbow except those that
were absorbed - they are missing or black. An emission
spectra (pictured earlier) is when the atoms are given energy by heat or
electricity and then they emit the wavelengths of energy so just a few
lines show up. They are opposites of each other.
- Electron
Spin
- Electrons spin clock or
counterclockwise creating tiny magnetic fields.
- Pauli exclusion principle (1925) states that the two electrons
in an orbital must have opposite spins. (important)
- Electron Configurations (important)
- The closer
to the nucleus the lower in energy (more stable). s subshells are closer to the nucleus than p subshells
so the s is more attracted to the protons in the nucleus and has lower
energy. Similar for d and f. E of s < E of p < E of d
< E of f.
- Ready for
filling electrons into their orbitals in the shells. Aufbau rules:
- Fill in
lower energy subshells first (we don't fill in by shell necessarily)
- Put 2
electrons per orbital and they must have opposite spin.
- Put 1
electron per orbital in subshells before pairing them up = Hund's
rule. For example you must put one electron in each of the 5 d
orbitals before pairing them up. Remember the orbitals in a
subshell are degenerate (equal in energy) This
means the 5 orbitals in a d subshell are all equal in energy to each
other.
- The actual
order we fill in is 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d
7p
- How can you
remember that? Well just follow the Periodic Table row by
row.
- Examples:
(superscript refers to how many electrons in that subshell, can't exceed
the max)
- H - 1s1
- F - 1s22s22p5
- P - 1s22s22p63s23p3
or [Ne]3s23p3
- I - 1s22s22p63s23p64s23d104p65s24d105p5
or [Kr]5s24d105p5
- You try U, Ag, O, Li,
and Fe. Don't look yet!
- U - 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s25f4 or
[Rn]7s25f4, Ag - 1s22s22p63s23p64s23d104p65s14d10
or [Kr]5s14d10, O - 1s22s22p4 or
[He]s22p4, Li - 1s22s1,
Fe - 1s22s22p63s23p64s23d6
- Practice Click
here go to Activities then electron
configuration activity and the movie.
- Another
great movie
Electron Configuration Trends
- All Alkali metals end
with s1
- All Alkaline Earth
metals end with s2
- All Halogens end with s2p5
- All Noble Gases end
with s2p6
- Valence electrons -
those in the outermost shell. (important)
- Electron
Configuration Anomalies
- When we did Ag above
did you think I made a typo? Well I did not! Some transition
metal configurations are not as expected. This is because it is
lowering in energy to have a half full or full d subshell. Atoms
would like to have 5 or 10 electrons in the d subshell if possible.
So they will steal one electron from the s to make this possible.
- Expect Mo to be [Kr] 5s24d4
but it is really [Kr] 5s14d5
- So that is why Ag is
not [Kr]5s24d9 but really [Kr]5s14d10
- What is the electron
configuration for Cu??? Expect [Ar]4s23d9
but it is really [Ar]4s13d10
- Discussed above
- Atomic
Radii
- Size increases f and i in the periodic table.
(important)
- It makes sense as you
down a group the atoms get bigger - we started filling in another shell.
But as you go across a row we are filling in the same shell and as more
electrons are added to that shell they are more attracted to the
increasing number of positive protons in the nucleus. So that attraction
pulls them a bit closer and the atoms get a tiny
bit smaller in general across a row. There is a big jump in size when we
start filling a new shell - Alkali metals are the biggest atoms.
- Zeff is the symbol for
effective nuclear charge which is the "feel" or attraction the
electrons feel for the protons in the nucleus. The electrons in the
1st shell feel more pull than the electrons in the 2nd shell etc.
This is because the inner electrons being negative actually repel the
outer electrons since they are also negative. So we say the inner
electrons "shield" the outer electrons from feeling the full
attraction of the protons. Zeff
increases g
and h
in the periodic table.
Enjoy and study real hard.
Totally Cool
Periodic Table with
all Emission
and Absorption Spectra
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Ba
flame test, Ca flame test
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K
flame test, Li flame test
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Na
flame test, Rb flame test
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