Transcript Chapter 13 Electrons in Atoms

Chapter 13
Electrons in Atoms
Atomic Models
• Democritus
400 BC
1. A Greek philosopher described matter
more than 2400 years ago
2. His theory: Matter could be divided
into smaller pieces only so far
3. Named the smallest particle of
matter “atomos”
4. His theory was ignored for a long
time.
 Aristotle
1. More popular and respected
2. His theory: 4 elements
fire, air, earth and water
• Dalton’s Model
1800’s
1. Chemist
2. Performed a number of experiments
that lead to the acceptance of the idea
of atoms
3. His theory:
-All elements are composed of
atoms. Atoms are indivisible and
indestructible particles
- Atoms of the same element are
exactly alike
- Atoms of different elements are
different
- Compounds are formed by the joining
of atoms of two or more elements
(This became of the foundations for modern
chemistry)
• Thomson’s Plum Pudding Model
1897
1. A scientist
2. Provided the first hint that an atom
has smaller particles.
3. His theory:
Atoms were made from a positively
charged substance with negatively
charged electrons scattered about
4. His experiment:
- He passed an electrical current
through a gas.
- When the current passed through
the gas, it gave off rays of
negatively charged particles
- Discovered there were smaller
particles in the atom
- Called the negatively charged
“corpuscles” now known as electrons
- Since the gas was known to be
neutral, he reasoned that there must
be positively charged particles in the
atom
- He could never find them
 Ernest Rutherford
1908
1. Gold Foil Experiment
- He fired a stream of tiny positively
charged particles at a thin sheet of
gold foil
- most of the + charged particles
passed right through the gold
atoms in the sheet of gold foil
- Some of the + particles bounced back
from the gold sheet (positive repels
positive)
- The experiment explained that the gold
atoms in the sheet were mostly open
space.
- He concluded that an atom had a
small, dense, positively charged center
that repelled the positively charged
particles
- He called the center of the atoms the
nucleus
- He reasoned that all of the atom’s +
particles were contained in the
nucleus. The negatively charged
particles were scattered outside the
nucleus.
• Niels Bohr
1913
1. His theory: proposed that the
electrons were in a specific energy
level
2. Electrons move in definite orbits
around the nucleus, much like
planets circle the sun. These orbits
or energy levels are located at certain
distances from the nucleus.
• The Quantum Mechanical Model
(Schrodinger)
Modern atomic theory
describes the electronic
structure of the atom as
the probability of finding
electrons within certain regions of space.
- It is impossible to determine the exact
location of an electron. The probability
of where the electron is located is
based on the energy the electron has.
Electron Cloud:
Depending on their energy they are locked
into a certain area in the cloud.
Electrons with the lowest energy are found
in the energy level closest to the nucleus.
• Electrons with the highest energy are
found in the outermost energy levels,
further from the nucleus.
Atomic Orbitals
- regions where an e- resides 95% of the
time
4 shapes
letter
shape
max eorbitals
s
sphere
2
1
p
dumbell
6
3
d
4-leaf clover 10
5
f
dragonfly
14
7
 Atomic Orbitals diagrams
s orbital
p orbitals
d orbitals
Energy and Electrons
 In nature, changes generally proceeds
toward the lowest possible energy level.
 High energy systems are unstable and
lose energy to become more stable.
 Electrons are arranged with lowest
possible energy level (electron
configurations)
Aufbau Principle
 Electrons enter orbitals of lowest energy
first.
 Follow the path
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d
6p 7s 5f 6d 7p
Examples:
Show the electron configuration for
1. sodium atomic number is 11
e- = 11
1s2 2s2 2p6 3s1
2. phosphorus atomic number =15
1s2 2s2 2p6 3s2 3p3
Short-hand configurations
Rule: Place the [ preceding noble gas]
then the rest of the electron
configuration
Examples: Br e- = 35
 Hund’s Rule
There must be one electron in each
orbital of a sublevel before
doubling occurs
One arrow equals one electron
So: s has 2 electrons, one orbit
s _____
p ___ ____ ____
d
f
Exceptions:
s1 exception
one electron leaves the “s” and
goes to the “d”
Nb, Cr, Mo, Tc, Ru,Rh, Cu, Ag,
Au, Pt
Four Quantum Numbers
1. Principal Quantum number (n)
The maximum distance an electron’s
orbital is from the nucleus.
n = 1, 2, 3,…..
2. Orbital quantum number (l)
The shape of an electron’s orbital
l = 0, 1, 2, 3, …. (n-1)
s p d f
3. Magnetic quantum number (m)
Shows how the electron’s orbital is
oriented in space
m = - l ….0….+l
s __
0
p __ __ __
-1 0 +1
d __ __ __ __ __
-2 -1 0 +1 +2
f __ __ __ __ __ __ __
-3 -2 -1 0 +1 +2 +3
4. Spin quantum number (s)
States in which direction the electron
spins
Uses the right hand rule from physics
s = +1/2
or -1/2
Pauli Exclusion Principle
No two electrons in the same atom can
have the same set of 4 quantum
numbers
Example: P
1s2 2s2 2p6
3s2
3p3
__ __ __ __ __ __ __ __ __
-1 0 +1
-1 0 +1
*
*
n=2
n= 3
l= 1
l=1
m = +1
m=0
s = -1/2
s= +1/2
State the element whose last electron has
the following quantum number’s
n= 5
l=2
m=0
s = +1/2
5d __ __ __ __ __
-2 -1 0 +1 +2
5d3 Ta
Energy levels
 Main areas where an electron could
be
 Closest to the nucleus has lowest
energy
 1 s, 2 s , 3 , 4
p
,5
,6
,7
Sublevels
The letters that stand for the shapes in
the different energy levels
 Example: 3s 3p 3d
Orbitals
determine how many electrons can be
held
 Ex. 3s ___ 3p ___ ___ ____
Other questions:
1. How many occupied energy levels are
in an atom of Ho (#67)
2. How many occupied sublevels?
3. How many half-filled orbitals?
Electron Dot Structure
 Keeps track of valence electrons
 valence electrons – outermost electrons
octet rule: has eight valence electrons,
stable
73
6
1
X
2
5
8 4
(right-left-top-bottom)
Example: Se (#34)
Behavior of electrons
Isaac Newton
1. Thought light as consisting of particles
2. Wave phenomenon
1700
Light and Atomic Spectra
 Light is wavelike
electromagnetic radiation: includes radio
waves, microwaves, infrared waves,
visible light, ultraviolet waves, xrays, and gamma rays
 waves travel in a vacuum at a speed of
3.0 x 1010cm/s or 3.0 x 108 m/s
wavelength: (λ) the distance between the
crest
amplitude: wave’s height from the origin to
crest
• Frequency: (ν) the number of wave cycles
to pass a given point
 c = λν
where c = speed of light
λ = wavelength
ν = frequency
units: c is m/s
λ is m
ν is hertz (Hz) = s-1
• Example: What is the wavelength of the
red light emitted by a barium lamp if the
frequency is 3.25 x 1014 s-1.
Atomic emission spectrum:
the relative intensity of each frequency
of electromagnetic radiation emitted by the
element’s atoms or the compound’s
molecule when they return to the ground
state
Einstein and Planck
Photoelectric effect
1. Reflected certain colors off a piece of
metal
2. Noticed electrons were released
3. Noticed not all colors did this
4. Found that certain color had specific
frequencies
5. Therefore light travels as particles
called photons
Planck’s constant
E = hv
h= Planck’s constant
= 6.626 x 10-34 Js
6. Light travels as both waves and
particles
Summary
1. Energy caused by electrons jumping
from high levels to low levels
2. Loss of energy is given to a photon
of light
3. Only specific colors, frequencies,
energies and jumps
de Broglie’s equation
predicts that all matter exhibits
wavelike motions
Classical mechanics vs. Quantum mechanics
1.Classical mechanics explains the motions
of objects larger than atoms. The object
gains or loses energy in any amount.
2.Quantum mechanics explains the motions
of subatomic particles and atom as waves.
These particles gain or lose energy in
packages called quanta.
Heisenberg Principle
1. Impossible to know both the position
and the path of the electron
2. Works better with smaller objects like
an atom than larger objects.