Download CHAPTER 5

CHAPTER 5
THE ELECTRON
Newtonian Physics
• Describes normal particles
moving at normal speed and the
laws that pertain to them
• Example: car moving at 55 mph
QUANTUM PHYSICS
• Describes extremely small
particles moving at speeds
close to the speed of light
• Example: electron moving in an
atom
BOHR MODEL
• Said e- can be found at certain
energy levels (like a ladder)
• e- can move up and down that
ladder by absorbing quantums of
energy
Increasing
energy
nucleus
QUANTUM- Max Planck
• Quantum= amount of energy an
e- needs to absorb to move from
lower to higher energy level
• Photon = quanta of radiant energy
(little packet of light energy)
ELECTROMAGNETIC
ENERGY
• A radiant energy- types include
visible light, radio, infrared, UV
and X-Ray
• Travels at the speed of light;
has variations in electric and
magnetic fields
• If we plot the strength of these
variations vs. time, we get waves
l
• Wavelength (l)= the physical
distance between crests
• FREQUENCY (f )= the # of wave
breaks (crests) that occur in a
unit of time
• Measured in Hertz (Hz)- 1
peak/second
ELECTROMAGNETIC
SPECTRUM
• MOVIE
Electromagnetic Spectrum
Visible Light700 nm (red) – 400 nm (violet)
Low energy
high energy
SPECTROSCOPY
• The study of the interaction
between matter or substances
and radiant energy
• Radiant energy = continuous
excited energy
Explanation of Atomic
Spectra
• Spectrum = occurs when you
look at the light given off thru a
diffraction grating (separates
light into its diff. wavelengthsvisible light)
• Uses: can ID unknown
gases; determine
composition of the stars
How it works:
• Atoms can lose or gain energy
• An atom gets excited but can’t
hold onto that energy so it loses
the energy in the form of light
Excited electron- absorbed
a quantum of energy- moved
to higher energy level
e-
Electron can’t hold onto the
Energy- falls back down to
ground state- release
photon (packet of light
energy)
e-
Energy levels
Ground state (lowest energy level an electron
Occupies)
Types of Spectrums
Chemistry of Fireworks video
Fireworks
CONTINUOUS SPECTRUM
• White light, gives you an
uninterrupted spectrum of light
• Get the colors of the rainbow
blended together (ROYGBIV)
EMISSION SPECTRUM
• Created by the light emitted when
the e- drop from their excited state
to lower energy states
• Get a series of lines in different colors
on a blacked out background
• Each line corresponds to a particular l
of light emitted by an atom
Emission spectrum of hydrogen (bright line spectrum)
ABSORPTION SPECTRUM
• Spectrum created by determining
the energy absorbed by the atom as
it moves up to higher energy states
• See a rainbow background w/dark lines
in the spot where the l of light have
been absorbed
• These lines correspond exactly to the
colored lines of the emission spectrum
Absorption spectrum of hydrogen (dark-line spectrum)
PHOTOELECTRIC EFFECT
• Albert Einstein (1921)
• Found that when light shines on
a metal surface at a very
specific frequency, the metal
emits electrons
• Uses: automatic doors;
solar calculator
• Proved light has both wavelike
and particle like properties
WAVE PARTICLE
DUALITY OF NATURE
• Louis de Broglie
• Used Einstein and Planck’s
formulas
• Described the behavior of the
electron in orbit
• Derived this equation:
h
h = Planck’s constant
l = mv
6.626 x 10-34 J·s
V = velocity
M = mass
l = wavelength
• It predicts that all moving particles
have wave characteristics (even a
moving car- just that the l is so
small we can’t see it)
• WAVE PARTICLE DUALITY OF
NATURE = all particles have
properties of waves and waves
have properties of particles
HEISENBERG
UNCERTAINTY PRINCIPLE
(1927)
• Found that it is impossible to
make any measurement on an
object without disturbing the
object
• Electrons were detected by
their interactions with photons
(packets of light that have the
same energy as an electron)
Uncertainty cont.
• Interaction between the photon
and electron alters the position
and the velocity of the electron
• HEISENBERG UNCERTAINTY
PRINCIPLE= it is impossible to
know precisely both the location and
the velocity of a particle at the same
time
photon
Change of
Position and v
e-
collision
photon
e-
Schrödinger Wave
Equation (1926)
• Erwin Schrödinger developed a wave
equation that laid the foundation for
the quantum theory and quantum
mechanical view of the atom
• QUANTUM THEORY= describes the
mathematical wave properties of
electrons and other small particles
• He used an equation to predict the
most probable location of an
electron around a nucleus (found
that electron did not travel in neat
paths)
ATOMIC ORBITAL
• A 3D region around the nucleus
that indicated the probable
location of electron (the
electron cloud)
H- 53 pm from
nucleus (90%
probablity)
QUANTUM NUMBERS
• From Schrodinger’s equation we get
a set of variables that specify the
properties of the electron in orbit
• Each electron can be described by a
unique set of 4 quantum numbers
describing the electron’s behavior
• Principle, angular momentum, Magnetic
and spin
Video- Electron Behavior
PRINCIPLE QUANTUM
NUMBER (n)
• Indicates the main energy level
occupied by the electron
• As the energy level (n)
increases the electron’s energy
increases (more energy further
from the nucleus)
• Can have more than one
electron in the energy level
(n) cont.
• Maximum number of electron in each
level represented with the following
formula:
• 2n2
where n = energy level
n=1
2 (1)2 = 2en= 3
2 (3)2 = 18e-
n=2
2(2)2 = 8en=4
2 (4)2 = 32 e-
n = 5-7 same as n = 4; only 32 e-
ANGULAR MOMENTUM
QUANTUM NUMBER (L)
• The orbitals or clouds in each
energy level occur in different
shapes
• This quantum number indicates
the shape of the orbital or the
name of the sublevel
(L) cont.
• Types of Sublevels
•
•
•
•
s = sphere shape
p = peanut shape
d = double peanut
f = flower shape
•
•
•
•
n
n
n
n
• The value of L = the value of n
=
=
=
=
1 , has 1 sublevel (s)
2, has 2 sublevels ( s & p)
3, has 3 sublevels (s, p, &d)
4 has four sublevels (s, p, d, & f)
MAGNETIC QUANTUM
NUMBER (m)
• An orbital can have the same
shape but be found in different
orientations around the nucleus
• This quantum number
represents the number of
positions found in each sublevel
(m) cont.
•
•
•
•
s = 1 position
p = 3 positions
d = 5 positions
f = 7 positions
max
max
max
max
of
of
of
of
2e6e10 e14 e-
• Each position can contain a pair
of electrons
•
SHREK Example? Okay, um, ogres are like onions. (he holds out
his onion) DONKEY (sniffs the onion) They stink? SHREK Yes - No! DONKEY They make you cry? SHREK No! DONKEY You leave
them in the sun, they get all brown, start sproutin' little white
hairs. SHREK No! Layers! Onions have layers. Ogres have layers!
Onions have layers. You get it? We both have layers. (he heaves
a sigh and then walks off) DONKEY (trailing after Shrek) Oh, you
both have layers. Oh. {Sniffs} You know, not everybody likes
onions. Cake! Everybody loves cakes! Cakes have layers. SHREK
I don't care... what everyone likes. Ogres are not like cakes.
DONKEY You know what else everybody likes? Parfaits. Have
you ever met a person, you say, "Let's get some parfait," they
say, "Hell no, I don't like no parfait"? Parfaits are delicious.
SHREK No! You dense, irritating, miniature beast of burden!
Ogres are like onions! And of story. Bye-bye. See ya later.
DONKEY Parfaits may be the most delicious thing on the whole
damn planet. SHREK You know, I think I preferred your humming.
DONKEY Do you have a tissue or something? I'm making a mess.
Just the word parfait make me start slobbering.
SPIN QUANTUM NUMBER
(s)
• Like the earth, the electron in an
orbital can be thought of as spinning
on its axis
• This quantum number represents
that spinning action
• The electron can spin clockwise or
counterclockwise
• We represent the spin with up or
down arrows as the electron
1s
ELECTRON
CONFIGURATION
• DEF.=AN ARRANGEMENT OF
THE ELECTRONS IN AN ATOM
• RULES:
• 1. Aufbau principle- electrons
occupy the lowest energy level
that can receive it (fill the lower
level up first)
• 2. Pauli Exclusion Principle- no two
electrons in the same atom can have
the same set of quantum #’s
• 3. Hund’s Rule- orbitals of equal
energy are each occupied by 1
electron before any orbital is
occupied by a second electron
• all electrons in single occupied orbitals
must have the same spin
TYPES OF
CONFIGURATION
1. ORBITAL NOTATION
• Most detailed
• Uses all four quantum numbers
• Orbital – a space occupied by a
pair of electrons
• Example: aluminum; atomic #13
so 13 electrons
1s
2s
2p
3s
3p
B
2p
1s
2s
1s
2s
1s
2s
1s
2s
1s
2s
1s
2s
1s
2s
C
2p
N
2p
O
2p
F
2p
Cl
2p
Mn
3s
3p
3s
3p
2p
…
3d
4
DIAGONAL RULE
diagram)
(sublevel
• Overlapping of sublevels = when
you get to the larger energy
levels the d and f sublevels are
so close together they find
greater stability if they overlap
some of the lower orbitals
1s
2s 2p
3s 3p 3d
4s 4p 4d
5s 5p
6s
4f
5d 5f
6p 6d
7s 7p
8s
• 1s
2s
3s
4s
5s
6s
7s
2p
3p
4p
5p
6p
7p
3d
4d
5d
6d
4f
5f
• Diagonal rule = use to help
figure out when overlapping
occurs; follow arrows and list
the orbitals in the order listed
1s,2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s,
5f, 6d, 7p
Example: iron
electrons = 26
Example: iron
electrons = 26
1s
4s
2s
2p
3d
3s
3p
2. ELECTRON
CONFIGURATION
NOTATION
• Shorthand way of writing orbital
notation
• Replaces boxes with
superscripts
• Ex. Nitrogen
7e1s
2s
2p
Electron configuration =
3s
1s2 2s2 2p3
3p
• s orbital can hold 2 electrons
– (1 box)
• p orbital
can hold 6 electrons
– (3 boxes)
• d orbital
can hold 10 electrons
– (5 boxes)
• f orbital
can hold 14
electrons
– (7 boxes)
• Example- Tungsten = W = 74e-
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6
5s2 4d10 5p6 6s2 4f14 5d4
video
3. ELECTRON DOT
CONFIGURATION
• Valence electrons – the
electrons in the outermost
orbital; those in the s and p
sublevels of the highest
principle energy level (n)- they
determine the chemical
properties of the element
• Once the energy level is filled it is
stable (8 valence electrons)
• The valence electrons are used in
forming chemical bonds
• Electron dot configuration = method
of drawing dots around the symbol
of an element to represent the
valence electrons ( in highest s and
p sublevels)
Electron dot
configuration steps
• 1. Write symbol- the symbol represents
the inner stable energy levels
• 2. write the electron configuration of the
element
• 3. find the # of valence electrons by
looking at the s & p sublevels of the
greatest energy level
• 4. Place dots equal to the # of valence
electrons around the symbol (max. of 8
dots)
Example: neon
Ne
Lithium 1s22s1
Oxygen: 1s22s22p4