1. Enriched Chemistry Chapter 4 – Arrangement of Electrons in Atoms
Section One – The Development of a New Atomic Model

2. 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 m/s through a vacuum and at slightly slower speeds through matter;

3. The Electromagnetic Spectrum
HIGH
ENERGY
LOW
ENERGY

4. 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

5. 
crest
origin
trough A

6. c =  c: speed of light (3.00  108 m/s) : wavelength (m, nm, etc.)
Frequency & wavelength are inversely proportional
c = 
c: speed of light (3.00  108 m/s)
: wavelength (m, nm, etc.)
: frequency (Hz)

7. Calculate the wavelength (in meters) of radiation a frequency of 5
Calculate the wavelength (in meters) of radiation a frequency of 5.00 x 1014 s¯1.

8. 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.

9. 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.

10. The energy of a photon is proportional to its frequency.
E = h
E: energy (J, joules) h: Planck’s constant (  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.

11. 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.

12. 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.

13. Hydrogen’s Line-Emission Spectrum

14. 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

15. Photon Emission and Absorption

16. 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.

17. 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 = hν, these frequencies corresponded to specific energies—the quantized energies of Bohr’s orbits.

18. 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.

19. Atomic Orbitals and Quantum Numbers

20. 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.

21. 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)

22. 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.

23. 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.

24. Shapes of s, p, and d Orbitals

25. Electrons Accommodated in Energy Levels and Sublevels

26. 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.

27. 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.