Transcript Electrons in Atoms

Electrons in Atoms
Chapter 5
In this chapter…
• Scientists pursue an
understanding of how electrons
are arranged within atoms
• Electron arrangement plays a role
in chemical behavior
• Chemists found Rutherford’s nuclear
model to be lacking because it did not
begin to account for the differences in
chemical behavior among various
elements
• Early 1900’s- scientists observed that
certain elements emit visible light when
heated in a flame chemical behavior
Wave Nature of Light
• Electromagnetic Radiationform of energy that exhibits
wavelike behavior as it travels
through space
Vocabulary to know..
• Wavelength- shortest distance
between equivalent points on a
continuous wave
• Symbol- λ (lambda)
• Unit- meters, centimeters, or
nanometers (1 nm= 1x10-9m)
• Frequency- the number of waves that
pass a give point per second
• Symbol- ν (nu)
• Unit- Hertz (SI Unit)= (1/s)= (s-1)
• Amplitude- the wave’s height from
the origin to a crest, or from the
origin to a trough
How are they related?
•
ALL electromagnetic waves,
including visible light,
travel at a speed of
c= 3.00x108 m/s
(MEMORIZE)
C= λν
• Speed of light= wavelength x frequency
Electromagnetic
Spectrum
• Aka EM Spectrum
• Encompasses all forms of
electromagnetic radiation
• The only differences in the types
of radiation being their
wavelengths and frequencies
ROYGBIV
Calculations 
• Microwaves are used to transmit
information. What is the
wavelength of a microwave having
a frequency of 3.44 x109 Hz?
• Know: C= λν
• C= 3.00x108 m/s
• ν = 3.44 x109 Hz
• λ = ???
 
λ = c/ ν
8
3.00 x10 m/s
9
3.44 x10 s
-1
λ= 8.72 x10-2 m
DON’T FORGET YOUR SIG FIG RULES!!!
Particle Nature of Light
• Quantum Concept
• Explained why colors of heated
matter correspond to different
frequencies and wavelengths
• Max Plank- “matter can gain or
lose only in small, specific
amounts called quanta”
• Quantum- the minimum amount
of energy that can be gained or
lost by an atom
• Energy of a quantum is related
to the frequency of the emitted
radiation by the equation
•
•
•
•
Equantum= hv
E= energy
h = Plank’s Constant (6.626x10-34J)
v= frequency
Joule (J)= SI unit for energy
Photoelectric Effect
• Electrons, called
photoelectrons, are emitted
from a metal’s surface when
light of a certain frequency
shines on the surface
• Photon- a particle of EM radiation
with no mass that carries a
quantum of energy
Ephoton= hv
Atomic Emission
Spectra
• Set of frequencies of
the electromagnetic
waves emitted by
atoms of the element
• Example- The light of neon
sign is produced by passing
electricity through a tube
filled with neon gas. Neon
atoms release energy by
emitting light.
An atomic emission spectrum is characteristic of
the element being examined and can be used to
identify that element
Section 5.2
Bohr Model of the Atom
• Proposed that the hydrogen atom has
only certain allowable energy states
• Ground State- lowest allowable energy
state of an atom
• Bohr’s model worked well to
explain Hydrogen- however it
did not explain other elements
• Substantial evidence indicates
that electrons do not move
around the nucleus in circular
orbits
• The de Broglie equation predicts
that all moving particles have wave
like characteristics
λ= h/mv
• The Heisenburg uncertainty
principal - states that it is
fundamentally impossible to know
precisely both the velocity and
position of a particle at the same
time
The Quantum Mechanical
Model (QMM)
• 1926- Austrian physicist Erwin
Schrodinger used the results of
Rutherford and Bohr to devise and
solve a mathematical equation
describing the behavior of the
electron in a hydrogen atom
• Unlike the Bohr model, the quantum
mechanical model does not involve an
exact path the electron takes around
the nucleus
• The quantum mechanical model
determines the allowed energies an
electron can have and how likely
(probability) it is to find the electron in
various locations around the nucleus
• The cloud is more dense where the
probability of finding an electron is
high
• Atomic Orbital- a 3D region
around the nucleus describing
the electron’s probable location
Atomic Orbitals
• Principal Quantum Number (n)indicates the relative sizes of
energies of atomic orbitals
• Principal Energy Levels- the
major energy levels of an atom
• Energy Sublevels- the energy
levels contained with the
principal energy level
Hydrogen’s First 4 Principal Energy Levels
Principal
Quantum
Number (n)
Sublevels
Number of
orbital's related
to sublevel
Max Number of
Electrons
1
s
1
2
2
s
p
1
3
8
3
s
p
d
1
3
5
18
4
s
p
d
f
1
3
5
7
32
Electron Arrangement in
Atoms
• Electrons and the nucleus
interact to make the most
stable arrangement possible.
• Electron Configurations- the
ways in which electrons are
arranged in various orbitals
around the nuclei of atoms
3 Rules for electron
configurations of atoms
1. Aufbau Principle : Electrons
occupy the orbitals of lowest
energy first
2. Pauli Exclusion Principle: an
electron orbital may describe at
most two electrons
• To occupy the same orbital, two
electrons must have opposite spins
(↓or ↑)
3. Hund’s Rule- electrons occupy
orbitals of the same energy in a
way that makes the number of
electrons with the same spin
direction as large as possible.
Noble Gas Configuration
• Method of representing electron
configurations of noble gases using
bracketed symbols.
• Neon= [Ne]
• Also used to shorten electron
configurations
• Sodium: #11- instead of 1s22s22p63s1
can be shortened to [Ne] 3s1
Exceptions to predicted
configurations
• Chromium- [Ar] 4s13d5
• Copper - [Ar] 4s13d10
• Illustrates the increased
stability of half-filled and filled
sets of s and d orbital's
Valence Electrons (V.E.)
• Electrons in the atom’s
outermost orbital's
• Determine the chemical
properties of an element
• V.E. are used in forming
chemical bonds
Electron Dot Structures
• Consists of the element’s symbol
and inner-level electrons surrounded
by dots representing the atom’s
valence electrons
• V.E. are placed one at a time on the
four sides of the symbol and then
paired up until all are used
• Examples
 Homework 
• Page 146
• 33-35, 38-39, 43-46, 49, 52,
57-61, 64-69, 72-73, 78-82,
86-89