Download Chapter 5: Electrons in Atoms

Chapter 5: Electrons in Atoms
CHEMISTRY
5.1 Light and Quantized Energy
 Rutherford’s Nuclear Atomic Model was incomplete,
did not explain



How electrons were arranged
Why electrons were not pulled into the nucleus
The differences in chemical behavior
Wave Nature of Light
 In 1900s, observation
of light when elements were
heated were observed
 Analysis of light revealed an elements chemical
behavior is related to the arrangement of electrons
in its atoms
Electromagnetic Radiation
 Form of energy that exhibits wavelike behavior as it
travels through space
 Includes visible light which is only a small portion of
EMR
 Visible Spectrum – ROY G BIV
Electromagnetic Spectrum Chart
Wave
Wavelength
 Symbol: λ (Lambda)
 Shortest distance between equivalent points (Peak to
Peak)
 Units = m, cm, nm
 nm=nanometers (10-9)
Frequency
 Symbol: ν (nu)
 Number of waves that pass a given point per second
 Units: hertz (Hz or 1/sec)
 652 Hz=652 waves/sec=652/sec=652s-1
Waves (Continue)
 Wavelength and frequency have an inverse
relationship

As wavelength increases, frequency decreases
 Amplitude: Height of a Wave
 Speed and Amplitude are not affected by wavelength
Light
 All electromagnetic waves travel at 3.00 x 108 m/s in





a vacuum
This is known as the speed of light
Symbol is c
Speed of light is a product of wavelength and frequency
c = λν
All light travels at the same speed but with different
wavelength and frequency
Practice Problems
 A helium-neon laser emits light with a wavelength of
655 nm. What is the frequency of the light?
 What is the wavelength of x-rays having a frequency
of 4.80 x 1017 Hz?
 An FM radio station broadcasts at a frequency of
98.5 MHz. What is the wavelength of the station’s
broadcast signal?
Electromagnetic Spectrum
 Electromagnetic Spectrum: classifies light based on





wavelength and frequency
Encompasses all forms of electromagnetic radiation
Visible Spectrum: Roy G. Biv
Only difference in types of radiation is wavelengths
and frequencies
Energy increases with greater frequency
Violet light has shorter wavelength, greater
frequency than red light, therefore violet light has
more energy that red light.
Particle Nature of Light
 Wave model could not explain


Why heated objects emit light at given temperatures
Why some metals emit electrons when light shines on them
 Max Planck studied light emitted from heated objects,


Discovered that matter can gain or lose energy in small, specific amounts
called quanta
Quantum: the minimum amount of energy that can be lost or gained
by an atom

Proposed emitted light (from glowing objects) is quantized
 Equantum= energy
 h= Planck’s constant= 6.626 x 10-34 J.s
 J= Joules = Kg .m2/s2 (SI unit of energy)
 ν= frequency
Plank’s Equation
 Equantum=hν
 Planck’s Theory= for a given frequency matter can emit or
absorb energy only in whole number multiples ( 1hν, 2hν,
3hν, 4hν)
 As energy increases, frequency increases
 Equation explains why violet light has more energy
than red light
Photoelectric Effect
 Einstein reasoned that light acted both wave like and
particle like
 Reasoned that light acted like a stream of particles,
knocking electrons out of atoms
 Photoelectric Effect = electrons called photo
electrons are emitted from metals surface when light
of a certain frequency shines on surface
Photoelectric Effect
 Light of a certain minimum frequency ejects
electrons from the surface of metal
Planck and Einstein
 An energy increases, frequency increases
Practice Problems
 Calculate the energy of a gamma ray photon whose
frequency is 5.02 x 1020 Hz.
 What is the difference in energy between a photon of
violet light with a frequency of 6.8 x 1014 Hz and a
photon of red light with a frequency of 4.3 x 1014 Hz?
 Calculate the energy of a photon of ultraviolet light
that has a wavelength of 49.0 nm.
Atomic Emission Spectra
 Atomic Emission Spectra: when atoms of an element
in the gaseous state are excited by energy they emit
light
 The light can be broken down into a spectrum
consisting of discrete lines of specific frequencies, or
colors
 Atomic Emissions Spectra is unique for each element
 Example: Neon signs
5.2 Quantum Theory and the Atom
 Bohr Model of the Atom
 proposed quantum or planetary model of the atom
 Correctly predicted frequencies of the lines in H atomic
emission spectrum
 Related H atomic energy states to motion of e assigned Quantum numbers n to each orbit (n = energy level)
 Known as the principle quantum number
 n = 1, 2, 3, 4, 5, 6, 7 (numbers correspond to energy level as
well as quantum numbers)
Problems
 Bohr’s model only worked for hydrogen
 Electrons do not orbit in circular paths
Louis de Broglie
 Accounted for the fixed energy level of Bohr’s model
 Proposed/thought that if waves have particle-like
behavior then particles can have wave-like behavior.
 De Broglie Equation
 Predicts all moving particles have wave
characteristics
 Example: cars (have wavelengths to small to be seen
even with sensitive equipment)
Werner Heisenberg
 Heisenberg Uncertainty principle  it is
fundamentally impossible to know precisely both
the velocity and position of a particle at the same
time.
 We are certain about being uncertain
 Bumping into an electron while trying to determine
its position and movement transfers energy and
disrupts the electron
Erwin Schrödinger
 Treated hydrogen atoms electrons as a wave
 Applied well to other elements
 Led to quantum mechanical model of an atom
(atomic model in which electrons are treated as
waves)
 Wave equation was able to predict three-dimensional
regions of an e- probable location (ATOMIC
ORBITALS-electron cloud, fuzzy cloud)
Atomic Orbitals
 Principle Quantum numbers indicate the relative
sizes and energies of atomic orbitals
 The atom’s major energy levels are called principal
energy levels
 Principal energy levels contain sublevels
 The number of energy sublevels in a principal energy
level increases as the quantum number increases
Principal Quantum Number n
 Corresponds to energy level and contains 2n2







electrons
n = 1 contains 2e- [2(1)2]
n = 2 contains 8e- [2(2)2]
n = 3 contains 18e- [2(3)2]
n = 4 contains 32e- [2(4)2]
n = 5 contains 50e- [2(5)2]
n = 6 contains 72e- [2(6)2]
n = 7 contains 98e- [2(7)2
Energy Levels Contain Sub-Levels
 Known as Quantum number l
 s,p,d,f (Sublevels)
 s = 0, p = 1, d = 2, f = 3
 Number of orbitals in a sub-level = (2l +1)
 Number of electrons in a sub-level = 2(2l +1)
Electron Configuration
 Electron configuration-arrangement of electrons in
an atom, usually lowest amount of energy is more
stable.
 Three Rules or principles that define how electrons
can be arranged



The Aufbau Principle
The Pauli Exclusion Principle
Hund’s Rule
The Aufbau Principle
 States that each electron occupies the lowest energy
orbital available
Pauli Exclusion Principle
 States that a maximum of 2 electrons may occupy a
single orbital, but only if the electrons have opposite
spins
 ↑↓
Hund’s Rule
 States that a single electron with the same spin must
occupy each equal-energy orbital before additional
electrons with opposite spins can occupy the same
orbital
Orbital Diagrams and Electron Configuration
Notation
 Orbital Diagram Method
 Electron Configuration Notation
 Noble Gas Configuration
Valence Electrons
 Electrons in the outer most orbitals
 Determines the chemical properties of an element
 Generally the electrons in the highest energy level
 Elections involved in bonding
Electron-Dot Structures
 element symbol that represents atomic nucleus and
inner electrons surrounded by dots which represents
valence electrons