Wave and Particle Models of Light in Atoms

1. Chapter 5 Electrons in Atoms Wave and Particle Models of Light Frequency and Unique Characteristics of Atoms Bohr and Quantum Mechanical Model Orbital Notations

2. 5.1 Light and Quantized energy • Rutherford’s explanation of the atom was found to be incomplete. • Questions remained as to where the electrons in an atom were located and why the nucleus did not pull them into itself.

3. Wave Nature of Light • Visible light is part of a range of electromagnetic radiation, a form of wave energy that travels through empty space and is propagated in the form of alternating electric and magnetic fields. • All waves exhibit certain common characteristics.

4. Parts of a Wave

5. Frequency • All waves consist of a series of crests and troughs that travel away from their source at a velocity that is determined by the nature of the wave and the material through which the wave passes. • The rate of vibration of a wave is called the frequency and is defined as the number of waves that pass a given point per second. • The units of frequency is the hertz (Hz); one hertz equals one wave per second (s-1).

6. Wavelength • The velocity and the frequency of the wave determine the wavelength of the wave. • The equation that expresses this relationship is c = ln • In this equation , c equals the speed of light, 3.0 X 108 m/s, l equals wavelength in meters, and n is the frequency in hertz.

7. Practice • A helium-neon laser emits light with a wavelength of 633 nm. What is the frequency of this light? • What is the wavelength of X rays having a frequency of 4.80 x 1017Hz?

8. Particle Nature of Light • Not only does light behave as a wave, it also behaves as a particle. • Einstein’s explanation of the photoelectric effect helped display this quality. • The effect says that electrons are ejected from the surface of a polished metal plate when it is struck by light.

9. Photons • Einstein found that this could only happen if light behaved as particles. • These particles, or photons, of light at the high-frequency (or violet) end of the spectrum had greater energy and could therefore dislodge many more electrons. • He found that the energy of a photon of a certain frequency can be calculated by using the equation: Ephoton= h n where h (Planck’s constant) = 6.626 x 10-34 J s

10. Practice 3. 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 and a photon of red light with a frequency of 4.3 x 10 14 Hz?

11. Atomic Emission Spectra • When atoms of an element in the gaseous phase are excited by energy, they emit light. • This emitted light can be broken into a spectrum consisting of discrete lines of specific frequencies, or colors. • This pattern of frequencies is unique to each element and is known as the element’s atomic emission spectrum.

12. Bohr Model • According to Bohr, hydrogen’s single electron can only orbit at specific distances from the atom’s nucleus. • When close to the nucleus, it has low energy; when farthest away it has the highest energy possible. • Thus the electron can only occupy specified allowed orbits.

13. Quantums of Energy • Electrons that are excited by an input of energy only absorb the amount needed to jump to a higher energy orbit. • When it falls back to the lower level, it emits a quantum of energy equal to the difference in energy between the two orbits. • Because hydrogen emission spectra contained several frequencies, Bohr designated them using integers called quantum numbers.

14. Modern Atomic Model • Electrons occupy the space surrounding the nucleus and can exist in several discrete principal energy levels, each designated by one of the principal quantum numbers (n) that are integers 1, 2, 3, 4 and so on. [This number corresponds to the row number of the element]

15. Atomic Model (cont’d) • Electrons in successively higher principal energy levels have greater energy. • Each energy level consists of energy sublevels that have different energy values. These are designated by s, p, d, and f respectively. [The number of sublevels depends on the principal energy level number ]

16. Atomic Model (cont’d) • Each sublevel has orbitals, each of which van contain only 2 electrons. All of the orbitals in the same sublevel have the same energy. • Atomic orbitals are regions of space in which there is a high probability (90 %) of finding an electron. The electron can be anywhere in an orbital and there is a 10% chance they will be outside the orbital.

17. Practice • How many electrons can the second principal energy hold? How many electrons can the third principal energy level hold? Explain the difference in these numbers of electrons.

18. Electron Configurations

19. Electron Configurations • The number and arrangement of electrons around the nucleus of an atom determines its chemical properties. • Because of this, the electron arrangement, or electron configuration . • The electron configuration of an atom is written by stating the number of electrons in each energy sublevel. • The number of electrons in the sublevel is shown using a superscript.

20. Aufbau Principle • One rule governing electron configurations is aufbau principle which states that each successive electron occupies the lowest energy orbital available.

21. Orbital Filling • For elements in the third row (third energy level), once the s and p orbitals are filled you may expect the d orbitals to begin filling. • However, because the 4s sublevel is of lower energy than the 3d sublevel, 4s fills before 3d. • Remember that the configurations are written by increasing energy, not in numerical order. • The following diagram may be necessary to write configurations correctly.

22. Filling Diagram • Remember to follow from the tail of the arrow to the head of the arrow. • For example, francium.

23. Practice • Write the electron configuration of the following elements: • Cesium • Radium • Iridium • Holmium

24. Noble Gas Configuration • Since a new principal energy level always begins with the element immediately following on of the noble gases, we can use a noble gas configuration. • This uses the symbol of the previous noble gas to denote all of an atom’s inner-level electrons.

25. Practice • Write the noble gas configuration of the following elements: a. Fluorine b. Phosphorous c. Calcium d. Cobalt e. Selenium f. Technetium

26. Valence Electrons • When elements combine chemically, only the electrons in the highest principal energy level are involved. • These outermost electrons are called valence electrons, and they determine most of the chemical properties of an element. • Since bonding requires the valence electrons, the electron-dot structure is a useful tool.

27. Electron-dot Structures • The electron-dot structure contains the symbol of the element and dots around it. • A single dot is used to represent each valence electron. • One dot is placed on each of the four sides before any two dots can be paired together.

28. Practice • Write the electron-dot structure for the following elements. • Nitrogen • Aluminum • Neon • Strontium • Antimony • Iodine • Lead