Download Modern Atomic Theory

Modern Atomic Theory
CP Chemistry ~ Unit 4
Chapter 11 (p. 323)
CP Chemistry Unit 7
3.1.10 B Describe concepts of models as a way
to predict and understand science and
technology.
3.1.10 C Apply patterns as repeated processes
or recurring elements in science and technology.
3.4.10 A Explain concepts about the structure
and properties of matter.
3.4.10 B Analyze energy sources and transfers
of heat.
Directory
I.] Atomic Models, Energy, & Light (EMR)
– Atomic Models … History
– Energy, Waves, & Light
II.]Excited Atoms, Hydrogen, & Bohr
III.] Wave-Mechanical Model
Atomic Models … History
Match each name below
with the correct term(s)
on the right!
Greeks (400 BC)
Boyle (1627-1691)
Dalton (1766-1844)
JJ Thomson (1890’s)
Lord Kelvin (1824-1907)
Rutherford (1910)
Chadwick (1932)
Plum Pudding model
“atomos”
Neutron
Gold Foil Experiment
Indivisible spheres
CRT
Elements can’t be broken
down
Proton
Electron
Nuclear atom
Ia.] History of Atomic Theory
Unanswered questions abound …
1. What are the electrons doing?
2. What keeps the electrons from being drawn
into the nucleus?
Ia.] History of Atomic Theory
Greeks … smallest particle of matter is
“atomos”
Ia.] History of Atomic Theory
John Dalton (1766 – 1844)
Elements are made of tiny particles called
atoms
All atoms of an element are identical
Atoms of an element are different from those of
other elements
Atoms combine with other atoms to form
compounds
Atoms are indivisible, neither created nor
destroyed
Ia.] History of Atomic Theory
JJ Thomson (1856 – 1940)
– Discovered electrons (CRT)
– Electrons have a negative charge
Ia.] History of Atomic Theory
Lord Kelvin (William Thomson) (1824-1907)
– “Plum Pudding” model of an atom (positive
cloud, pudding, with negative electrons,
plums, in it.
Ia.] History of Atomic Theory
5. Ernest Rutherford (1871-1937)
– Gold Foil Experiment … used alpha particles
“Like shooting a gun at a piece of paper and
having the bullet bounce back”
– “Nuclear atom” … atom with a nucleus
– Nucleus … small, dense, and positive
– Found protons (1919) in nucleus, positive
Ia.] History of Atomic Theory
6. James Chadwick (1891-1971)
– Discovers the neutron, mass approximately
equal to the proton, no charge (neutral)
7. 1920’s atom …
– Nucleus with protons and neutrons
– Electrons outside of nucleus
Ib.] Energy, Waves, & Light
Characteristics of Waves …
Wavelength (λ)
Frequency (ν)
Speed (c)
Amplitude
Crest
Trough
Ib.] Energy, Waves, & Light
Types of waves …
– Transverse … displacement is perpendicular
to the direction of propagation
– Longitudinal … displacement is parallel to the
direction of propagation
Ib.] Energy, Waves, & Light
Properties wave demonstrate …
– Reflection
– Refraction
– Interference
– Diffraction
Ib.] Energy, Waves, & Light
What do the following have in common?
– Microwaves
– X-Rays
– ROY G BIV
– Radio / TV waves
– Tanning beds
– Cell phones
Ib.] Energy, Waves, & Light
Definition: EMR (electromagnetic radiation)
radiant energy that exhibits wave-like
behavior and travels through space at
the speed of light in a vacuum.
Composition:
- Electric component
- Magnetic component
Ib.] Energy, Waves, & Light
Formula …
C = f (λ)
E = hf (energy and frequency are directly
proportional
E = hc / λ (energy and wavelength are inversely
proportional
Ib.] Energy, Waves, & Light
EM Spectrum
Ib.] Energy, Waves, & Light
What is light?
Answer …
a wave and a particle (photon)
(wave – particle duality)
Important:
1. Different color = different wavelength
2. Different wavelength = different energy
Ib.] Energy, Waves, & Light
So …
If we study the
electromagnetic spectra
of elements we will get
clues into atomic
structure!
Flame Test: Strontium
II.] Excited Atoms, Hydrogen &
Bohr
When atoms are heated
…
1. They absorb energy =>
electrons jump to a
higher energy “excited
state” (unstable)
2. When excited electrons
fall back to lower levels,
they release energy
(visible light).
II.] Excited Atoms, Hydrogen &
Bohr
The energy released
corresponds to the
available energy
levels in that atom.
Only certain energies
are released
(quantized energy).
II.] Excited Atoms, Hydrogen &
Bohr
If you look at a flame test through a
spectroscope, what you see are lines of
color that correspond to the energy states
of that atom!
II.] Excited Atoms, Hydrogen &
Bohr
Short Response …
What is the significance of the observed
lines in the spectroscope when viewing an
“excited” atom?
II.] Excited Atoms, Hydrogen &
Bohr
Niels Bohr (1885 –
1962)
– Solar system model of
the atom
– Circular orbits of
electrons correspond
to energy levels
– Helpful model …
unable to explain
atoms other than
hydrogen
III.] Wave Mechanical Model (11.6)
Also known as the Quantum Mechanical
Model (energy is “quantized;” a quantity)
Proposed by Louis Victor de Broglie
(1892-1987) and Erwin Schrödinger
(1887-1961)
Schrödinger’s Equation 
III.] Wave Mechanical Model (11.6)
Mechanics is a study of forces and
motions
– Newtonian Mechanics … describes the
behavior of visible objects traveling at
ordinary velocities.
– Quantum Mechanics … describes the
behavior of extremely small particles at near
light velocities.
III.] Wave Mechanical Model (11.6)
Goal of mechanics related to atomic
structure … to describe the allowed
electron energy states.
Result … Quantum numbers (letters) tell
us four things about the electron …
– The distance of the electron from the nucleus
– The shape of the electron cloud (probability)
– The position of the electron in space
– The spin direction of the electron
III.] Wave Mechanical Model (11.6)
1. Electrons occupy orbitals (probability
based) not orbits!
Is the solar system model incorrect?
YES!
2. Orbitals give no information about …
– When an electron will be at a given point.
– The path of an electron.
III.] Wave Mechanical Model (11.6)
This diagram is a
representation of a
mathematical or
probability model of
an electrons location
around a nucleus.
What does this model imply?
What doesn’t it imply?
III.] Atomic Orbital Structures (11.7)
A. Principal Energy Levels (n = 1,2,3, …)
As level increases so does the energy.
(Roughly correspond to solar system rings.)
III.] Atomic Orbital Structures (11.7)
B. Sublevels … each Principal Energy level is
divided into 1 or more sublevels.
Example:
Level 1 … 1 sublevel “s” (“s” = sharp)
Level 2 … 2 sublevels s and “p” (“p” = principal)
Level 3 … 3 sublevels s, p, and “d” (“d” = diffuse)
Level 4 … 4 sublevels s, p, d, and “f” (“f” = fundamental)
III.] Atomic Orbital Structures (11.7)
C. Orbitals … each
sublevel contains
one or more
orbitals. Each
orbital is a
probability area for
electrons.
All “s” sublevels
contain 1 orbital
which is spherical
III.] Atomic Orbital Structures (11.7)
All “p” sublevels contain 3 orbitals
that are each dumbbell shaped.
III.] Atomic Orbital Structures (11.7)
All “d” sublevels contain 5 orbitals that
can only described as different.
All “f” sublevels contain 7 orbitals.
III.] Atomic Orbital Structures (11.7)
Pause …
Atoms contain _______________
Levels contain _______________
Sublevels contain ____________
Orbitals contain ___________
These electrons have opposite “spins,”
(designation … + or -, clockwise or
counterclockwise, up or down arrows)
III.] Atomic Orbital Structures (11.7)
Levels …
Sublevels …
Orbitals …
Electrons … “spin”
Pauli Exclusion Principal – “an atomic
orbital can hold a maximum of two
electrons and those electrons have
opposite spins.”
IV.] Wave Mechanical Model
(11.6)
Principle Quantum Number … LEVEL
– Corresponds to electron rings, shells, or orbits
in the Bohr model
– Represented by the letter n
– Primarily determines the electrons energy
(near nucleus = low energy)
– Always an integer greater than zero (n = 1, 2,
3, …)
– Within each level, the number of electrons =
2n2
IV.] Wave Mechanical Model
(11.6)
Angular Momentum Quantum Number …
SUBLEVEL
– Represented by the letter l
– Each level has a number of sublevels equal to
n
– The values for l= 0, 1, 2, 3, … n-1
– Within each level, the sublevels are named s,
p, d, f (sharp, principle, diffuse, fundamental)
– Increasing l only slightly increases the
electrons energy
Worksheet Diagram
n
l
m
s
IV.] Wave Mechanical Model
(11.6)
Angular Momentum Quantum Number … l
– n = 1 … l = 0 (s) spherical
– n = 2 … l = 0 & 1 (p) dumbbell shape
– n = 3 … l = 0, 1, & 2 (d) dumbbell w/ donut
– n = 4 … l = 0, 1, 2, & 3 (f) funny shape
IV.] Wave Mechanical Model
(11.6)
Magnetic Orbital Quantum Number …
ORBITAL
– Represented by the letter m
– Electrons will pair up, the space that a pair of
electrons occupies is an orbital
– m is always an integer; such that m = -l … 0
… +l
– Orbital has little effect on energy
Worksheet Diagram
n
l
m
s
IV.] Wave Mechanical Model
(11.6)
Number of Orbitals
– Sublevel 0 (s) has one orbital … [0]
– Sublevel 1 (p) has three orbitals …
[-1 ] [0] [+1]
– Sublevel 2 (d) has five orbitals …
[-2] [-1] [0] [+1] [+2]
– Sublevel 3 (f) has seven orbitals …
[-3] [-2] [-1] [0] [+1] [+2] [+3]
IV.] Wave Mechanical Model
(11.6)
Magnetic Spin Quantum Number … SPIN
– There are only two possible values … +1/2 or
– ½ (think of these as clockwise and
counterclockwise spins) (sometimes
represented as up and down arrows)
IV.] Wave Mechanical Model
(11.6)
Quantum Rules …
– Pauli Exclusion Principle: no two electrons
may have the same set of four quantum
numbers
– Hund’s Rule: when filling orbitals of equal
energy, fill so that as many electrons as
possible remain unpaired