Section One – The Development of a New Atomic
Model

The Wave Description of Light
■ Visible light is a kind of electromagnetic radiation;
■ Together, all forms electromagnetic radiation form the electromagnetic spectrum;
■ All forms e.r. move at a constant speed of 3.00 x 10 8 m/s through a vacuum and at slightly slower speeds through matter;

The Electromagnetic Spectrum
R
G
Y
E
N
E
I
H
G
H
E
R
E
N
G
Y
L
O
W

Describing Waves
■ Wavelength (  ) - length of one complete wave
■ Frequency (  ) - # of waves that pass a point during a certain time period
– hertz (Hz) = 1/s
■ Amplitude (A) - distance from the origin to the trough or crest

A
 crest origin trough

• Frequency & wavelength are inversely proportional c =
  c: speed of light (3.00

10 8 m/s)

: wavelength (m, nm, etc.)

: frequency (Hz)

Calculate the wavelength (in meters) of radiation a frequency of 5.00 x 10 14 s¯ 1 .

When certain frequencies of light strike a metal, electrons are emitted.
■ Photoelectric effect – refers to the emission of electrons from a metal when light shines on the metal.
– Scientists observed that for a specific metal, no electrons were emitted if the light’s frequency was below a certain minimum, regardless of the intensity.
– This puzzled scientists because it was not predicted by the wave theory of light.
https://www.youtube.com/watch?v=v-1zjdUTu0o

The Particle Description of Light
■ In 1900, German physicist Max Planck suggested that hot objects emit energy in small, specific packets called quanta.
– A quantum of energy is the minimum quantity of energy that can be lost or gained by an atom.

The energy of a photon is proportional to its frequency.
E = h

E: energy (J, joules) h: Planck’s constant (6.6262  10 -34 J·s)
 : frequency (Hz)
In 1905, Einstein introduced the idea of photons, which are particles of e.r. having zero mass and carrying a quantum of energy.

Electrons exist only in very specific energy states for every atom of each element.
■ Ground state – the lowest energy state of an atom.
■ Excited state – the atom has a higher potential energy than it has in its ground state.
– Basically, when an atom absorbs energy it moves to an excited state;
– When that atom returns to its ground state, it releases energy in the form of e.r.
– Neon signs are an example.

Hydrogen’s Line-Emission Spectrum
• Investigators passed electric current through a vacuum tube containing hydrogen gas at low pressure, they observed the emission of a characteristic pinkish glow.
• When a narrow beam of the emitted light was shined through a prism, it was separated into four specific colors of the visible spectrum.
• The four bands of light were part of what is known as hydrogen’s line-emission spectrum.
• Scientists had expected to observe the emission of a continuous range of frequencies of electromagnetic radiation, a continuous spectrum.

Hydrogen’s Line-Emission Spectrum

Bohr Model of the Hydrogen Atom
• Niels Bohr proposed a hydrogen-atom model that linked the atom’s electron to photon emission.
• According to the model, the electron can circle the nucleus only in allowed paths, or orbits.
• The energy of the electron is higher when the electron is in orbits that are successively farther from the nucleus.
• When an electron falls to a lower energy level, a photon is emitted, and the process is called emission.
• Energy must be added to an atom in order to move an electron from a lower energy level to a higher energy level. This process is called absorption.
• video

Photon Emission and Absorption

Section Two – The Quantum Model of the Atom
• In the same way that no two houses have the same address, no two electrons in an atom have the same set of four quantum numbers.
• In this section, you will learn how to use the quantum-number code to describe the properties and locations of electrons in atoms.

Electrons have wave-like properties.
• French scientist Louis de Broglie suggested in 1924 that electrons be considered waves confined to the space around an atomic nucleus.
• It followed that the electron waves could exist only at specific frequencies.
• According to the relationship E = h ν , these frequencies corresponded to specific energies—the quantized energies of Bohr’s orbits.

Heisenberg’s Uncertainty Principle
• German physicist Werner Heisenberg proposed that any attempt to locate a specific electron with a photon knocks the electron off its course.
• Electrons are detected by their interactions with photons.
• The Heisenberg uncertainty principle states that it is impossible to determine simultaneously both the position and velocity of an electron or any other particle.
http://www.teachertube.com/video/electrons-and-observer-heisenberg-copenhagen-131499

Atomic Orbitals and Quantum Numbers
■ https://www.youtube.com/watch?v=9E3QaRxqXZc

Atomic Orbitals and Quantum Numbers
• Quantum numbers specify the properties of atomic orbitals and the properties of electrons in orbitals.
• The principal quantum number, symbolized by n, indicates the main energy level occupied by the electron.

Quantum numbers (cont…)
■ The angular momentum quantum number, symbolized by l, indicates the shape of the orbital.
– We will learn about the s, p, d and f orbital shapes.
– N = 1 has one sublevel (s)
– N = 2 has two sublevels (s, p)
– N = 3 has three sublevels (s, p, d)
– N = 4 has four sublevels (s, p, d and f)

Atomic Orbitals and Quantum Numbers, continued
• The magnetic quantum number, symbolized by m, indicates the orientation of an orbital around the nucleus.
• The spin quantum number has only two possible values—(+1/2 , −1/2)—which indicate the two fundamental spin states of an electron in an orbital.

Let’s review two terms.
■ Orbital – a single allowed location for electrons capable of holding two electrons of opposite spin states.
■ Sublevel – includes all of the similarly shaped orbitals in an energy level.

Shapes of s, p, and d
Orbitals

Electrons Accommodated in Energy Levels and Sublevels

Electron Configurations
• The arrangement of electrons in an atom is known as the atom’s electron configuration.
• The lowest-energy arrangement of the electrons for each element is called the element’s ground-state electron configuration.

Rules Governing Electron
Configurations
■ According to the Aufbau principle, an electron occupies the lowest-energy orbital that can receive it.
■ According to the Pauli exclusion principle, no two electrons in the same atom can have the same set of four quantum numbers.