Quantum Theory
I. Light
A. light acts as both a particle & a wave
B. So what? Electrons also have a dual nature
C. Production of color
1. When light “hits” matter/atoms the
electrons (which
have very little mass &
thus very little inertia –
resistance to a
change in motion) in the matter
are
forced to vibrate. They are the only parts
that have a small enough mass to be
able to
vibrate as fast as light.
a. An electron holds the photon of light
that hits it
briefly trying to see if it can
absorb it
permanently.
b. As it holds the photon it jumps
(vibrates) farther
away from the
nucleus to an energy level with
greater E & is said to be in an “excited state”.
c. When it loses (re-emits) the photon it
sinks back
down to its original energy
level aka its ground
state.
http://www.worsleyschool.net/science/files/electr
onphoton/interaction.html
d. Different photons have different
frequencies
(vibrations/sec) &
different atoms have different
frequencies.
e. The photon’s E = the difference
between the 2
energy levels.
f. If an electron can be caused to resonate
by the
photon it can keep the photon
& won’t re-emit
it. It absorbs that
frequency of light.
1) If it resonates with the photon it
jumps even
farther away from the
nucleus & hits other
atoms
causing them to vibrate. When this
happens heat is produced.
{So, Light E is converted to Kinetic E which is
converted to Thermal E (heat) plants use this
process. An e- in a chlorphyll molecule is
excited, it jumps to a higher E level & in the
process is used to split a water molecule. This
starts a series of chemical reactions & ends in
the production of food (glucose).}
g. If an electron holds the photon for a
SHORT
period of time but can’t
resonate it re-emits the
photon. The
frequency of this re-emitted photon
determines what color we see. We use the
phenomenon & put chemical dyes in
different
things to make
pretty
colors!
{If you think of the electrons as being on a
staircase they have a set distance for each
step. If the photon can give them the right
amount of E they can get up to the next
step & stay (the e- resonates). If not they
jump up attempting to get to the next step
but ultimately fall back down (releasing the
photon as color).}
h. If something absorbs & keeps all the
photons we
see a lack of reflected color
& call this black.
i. If something absorbs NONE of the
photons we see
all the colors
combined...we see light without
color or
white light.
D. quantum theory was developed to explain
this dual
nature
1. a quantum is the minimum amt of E that
can be lost or
gained by an atom...each type
of atom will be
slightly different
2. a photon is a particle of electromagnetic
radiation (
ROY G BIV) w/zero mass &
carrying a quantum
of E
a. each type of photon carries a different
quantum
of E
3. Heisenberg Uncertainty Principle – can’t
determine
both the position & velocity of eat the same
time...you try to find the
location & the process
knocks the e- out of
position
4. There are 4 quantum numbers
a. principal q# - tells main E-level
occupied by ep.101
Energy level 1 (1st orbit) can only hold
a
maximum of 2 electrons.
Energy level 2 (2nd orbit) level can
hold a
maximum of 8
electrons.
Energy level 3 (3rd orbit) can hold up
to a
maximum of 18.
Energy level 4 can hold a maximum of
32
electrons.
Rather than memorizing the number of
electrons each energy level (orbit) can hold, use
2n2 to find the number of electrons possible for
each energy level (n = the energy level (orbit)
number)
b. angular momentum q# - tells shape of
orbital
p.102!!!
sp d f
c. magnetic q# - tells orientation of orbital
around
nucleus (in 3D) p.102!!!
s has only 1 orientation/sublevel &
can hold
2 ep has 3 dif. sublevels & can hold 6
e(2 in each orientation)
d has 5 dif. sublevels & can hold 10
e(2 in each one)
f has 7 dif. sublevels & can hold 14
e(2 in each)
d. spin q# - tells 2 spin states of e(like clockwise or counterclockwise)
↑↓
Pauli exclusion principle – no 2 e- in the same
atom can have the same set of 4 q#’s
spdf charts
*The electrons in an atom are basically lazy
and will fill a position in the atom which is of
lowest energy, beginning with Energy level
orbit 1, s sublevel. –Aufbau Principle
*When E-level 1 is full, Energy level 2 will
begin to fill, first with 2 e- in the s orbital
and then with 6 e- in the p orbital.
*Then the electrons will begin to fill the next
highest-energy positions, the 3 "s", then the 3
"p".
*Knowing this allows us to write what is
called the electron configuration or a
description of where the electrons in that
atom are actually located. ex: 1s2 2s2 2p6
3s2......
*The e-configuration of an atom’s highest
occupied E-level determines the atom’s
chemical properties!
1s
2s 2p
3s 3p
3d
4s 4p
4d
4f
5s 5p
5d
5f
5g
6s 6p
6d
6f
6g
6h
7s 7p
7d
7f
7g
7h
7I
8s 8p
8d
8f
8g
8h
8I
8j
last orbital being filled w/electrons
1s – turquoise
s – outlined in purple
p – peach
d – yellow
f – blue
II. Periodic Table & Trends
A. History
1. early 1800’s scientists tried to
classify
known elements
a. Dobreiner (German chemist
1817) came
up with the idea of
elements being grouped by
properties or trends.
2. Newlands (English chemist
1863) organized elements by
atomic masses & came up with
law of octets – repetition of
similar properties every 8
elements
3. 1st created by Mendeleev – (Russian
Chemist 1869)
put elements in order by atomic
masses
found they naturally fell
in order by
properties (almost
the periodic law)
a. said their properties were
caused by
their at. masses &
that they were
periodic
b. left several blank spaces
where he
thought
undiscovered elements
should go ( they were discovered
w/in 15 years)
c. there were a few
discrepancies
4. In the early 1900’s Moseley (English
scientist 1913) rearranged
them by atomic
number (# of
protons) & they fit much
better.
a. periodic law – properties of
the
elements are periodic
functions of
their at. #’s!
B. Periodic Table – arrangement of
elements in
order of at# so that
elements with similar
properties fall
in the same column--- properties
show periodicity – regular, repeating
patterns
C. Structure
1. split into 3 major types of elements
a. metals – ductile, malleable,
lustrous,
good conductors
of heat & electricity—
left side
of table
b. nonmetals – opposite of metals –
right
side of table
c. metalloids – has char. of both
metals &
nons – found along
stairstep B, Si, As,
Ge, Sb, Te
only
2. periods –
a. horizontal rows of elements
arranged
by at #
b. elements in the same periods
have the
same highest
occupied E-level
3. groups- vertical columns of elements
with
similar char.
a. Noble gases gp 18– elements
with a
stable number of
outer E-level e1). outer E-level e- are called
valence electrons!
a) these are the ones that are
involved in chem
reactions
b) noble gases have 8
valence e- (Helium has 2)
octet rule!—8 e- in outer
E-level is stable
2). they are very un-reactive
3). He – Rn
b. Lanthanides & Actinides –
elements with VERY similar
properties… they’re
put
separately down at the bottom just
for space-saving purposes
c. Halogens – gp 17 –
7 valence emost reactive nonmetals
react w/metals to form salts
d. alkali metals – gp 1—
1 valence esilvery
soft
extremely reactive
reacts strongly w/water & air
so
must be stored in
kerosene
not found pure in nature
e. alkaline earth metals – gp 2—
2 valence eharder, denser, stronger than
gp 1
metals
less reactive than group 1 but
still not found pure in
nature
f. transition metals—gps 3-12—
strong metals
less reactive than gps 1 & 2
many found in pure state
D. Periodic table trends
1. ATOMIC RADIUS - the size of
atomic
radius cannot be
measured exactly because
an atom
does not have a sharply-defined
boundary.
--so, it’s ½ the distance from the
nucleus of 1 atom
that’s bonded to an
identical
atom
a. size of an atom INCREASES as
you
because:
move DOWN a group
1) the atoms have more and
more orbits
of electrons
(more layers)
2) shielding is
increasing down the
column which means the outermost
electrons are not held as tightly
and
"roam" out to a bigger
diameter
a) shielding – layers of ebetween
the nucleus &
the outer E-level elessen the pull of the nucleus
b. size DECREASES as you move
across
1) more p+ in nucleus = greater
pull but
no increase in Elevels
2. ELECTRON AFFINITY – amt of E
absorbed or given off when
a neutral atom
acquires an e- (most
atoms release E & is
shown as a
negative #)
a. trends not as regular
b. decreases down a group
(becomes
more +)
c. generally increases as you go
across a
period (becomes more
neg.)
3. IONIZATION ENERGY – E
required to
remove 1 e- from a
neutral atom of an
element
a. shows how easy it is to remove
an eb. IE decreases down a group
c. IE increases as you move across
a period
4. ELECTRONEGATIVITY – ability of an
atom in
a compound to attract ea. smaller atoms with less shielding
tend to
easily attract eb. decreases down a group
c. increases across a period