1. Arrangement of Electrons in Atoms
Chapter 4
Arrangement of Electrons in Atoms
The emission of light is fundamentally related to the behavior of electrons.

2. The Concept of Energy Energy – the ability (capacity) to do work
Units of Energy: Calorie (Cal), calorie (cal), Joule (J)
Work is done when an object moves some distance in response to a force (pushing or pulling).
work = force x distance

3. PE – “energy of position”
Types of Energy
kinetic energy – the energy that objects have because they are moving
KE – “energy of motion”
potential energy – the energy that is available for doing work at some later time
PE – “energy of position”

4. Kinetic vs. Potential
Potential energy (in fact any type of energy) can be converted into other forms of energy. For example if an object is dropped then it speeds up (gains kinetic energy) as at falls (loses potential energy).

5. Light is a Form of Energy
During the 1900s, it was discovered that
light has certain particle-like properties.
electrons (particles) exhibit some wave-like properties.
This became known as the dual wave-particle nature of light.
If we study waves, we can know more about the behavior of electrons.

6. Thomas Young's Double Slit Experiment

8. Particles act like waves
One does not observe the wave-like quality of everyday objects because the associated wavelengths of people-sized objects are exceedingly small.

9. Measurable Properties of Light
Wavelength (l) – the distance between corresponding points on adjacent waves
Units: meters, nanometers, etc.
1m = 1x109nm

10. Parts of a Wave

11. Measurable Properties of Light
Frequency (u)- number of waves that pass a given point in a specific time (usually one second)
Unit: Hertz (Hz)

12. Radio Stations FM – frequency in megahertz (MHz) example: 90.5 FM
= x 106 Hz
= 90.5 x 106 1/s
AM – frequency in kilohertz (kHz)
example: 1020 AM
= 1020 kHz
= x 103 Hz
= x 103 1/s

13. Wave Description of Light
Electromagnetic Radiation- form of energy that exhibits wavelike behavior as it travels through space.
Electromagnetic Spectrum- all the forms of electromagnetic radiation.

17. Types of Electromagnetic Radiation
Gamma rays
X-rays
Ultraviolet
Visible
Infrared
Microwaves
Radio waves
High energy, E
High frequency, u
Violet
Indigo
Blue
Green
Yellow
Orange
Red
High wavelength, l

18. c is from the Latin word celeritas meaning "swiftness"
What is the constant, c ?
Speed of Light (c)- all forms of electromagnetic radiation move at a speed of 3.00 X 108 m/s in a vacuum.
c is from the Latin word celeritas meaning "swiftness"

19. Relationship between u and l
c = lu
c = speed of light in m/s
l = wavelength in m
u = frequency in 1/s

20. Relationship between u and E
In order to get more cycles per second (u), the wave needs more energy (E).
A direct proportion is show as
Changing it into an equality with a constant, it becomes

21. What is the constant, h? E = energy in J h = constant in J.s
u = frequency in 1/s or Hz

22. Combining the Equations you get…l u =
E = h u

23. Check for Understanding
13. What would the energy be for light with a frequency of 5.68 x /s?
14. What is the energy of light that has a wavelength of 7.89 x m?

24. Check for Understanding

25. Continuous Spectrum
Continuous Spectrum- a spectrum in which all (¥) wavelengths within a given range are included.
Example: Visible spectrum- ROY G. BIV
The range of wavelengths in the visible spectrum is approximately 400nm (blue region) to 700nm (red region).

26. Continuous Spectrum

27. Absorption vs. Emission
Absorption – when energy is “taken in” by electrons
Emission – when energy is “given off” by electrons

28. Energy States of an Electron
Ground State- the lowest energy state of an atom. (stable)
Excited state- state in which an atom’s potential energy is increased from that of the ground state. (unstable).

29. Exciting Atoms
e- can be excited by passing electric current through them or by heating them (giving them E).
When an e- gets excited it becomes unstable and wants to return to its ground state. In order to do this the atom releases specific amounts of energy (quanta) in the form of electromagnetic radiation.

30. Spectroscopy
Spectroscope – used to separate light into a spectrum by wavelength so it can be examined.
Line-Emission Spectrum – produced when an electron jumps from a higher energy level to a lower energy level.
Acts as an atomic fingerprint.

31. Discovery of the Line-Emission Spectrum
Hydrogen atoms were excited by passing a high-voltage current through hydrogen gas causing the gas to glow a lavender color.
When viewed with a spectroscope (diffraction grating or prism) the lavender light separated into four narrow lines of different color.

32. Hydrogen Line Spectrum
Each of the lines seen in the hydrogen spectrum is a result of light at a different wavelength.
Since light of a particular wavelength has a definite frequency and a definite energy, the lines of the hydrogen spectrum must be a result of the emission of photons with specific energies.

33. Hydrogen Line Spectrum

34. The Bohr Model

35. Bohr Model of the Atom
The fact that hydrogen atoms only released photons of specific frequencies indicated that differences between the atom’s energy levels were fixed.
The puzzle behind the hydrogen-atom spectra was solved in 1913 by Niels Bohr.

36. Bohr Model of the Atom
The Bohr model of the atom explains the emission of photons by the hydrogen atom.
Electrons can circle the nucleus in allowed paths or orbits.
Orbits are also referred to as energy levels.
When an electron is in one of these orbits it has a fixed energy.

37. Bohr Model of the Atom
The orbits are separated from one another by empty space where the electrons cannot exist.
The energy of the electrons becomes higher as they get farther away from the nucleus.
Electrons can move from one energy level to the next by gaining or losing a finite amount of energy.

38. Bohr Model of the Atom K L M

39. Emission
Emission- when an electron falls to a lower energy level, and a photon is emitted.
The photon’s energy is equal to the difference between the higher energy orbit and the lower energy orbit.

40. Absorption
Absorption- when the minimum required energy is added to an atom causing the electron to move from a lower energy level to a higher energy level.

41. Shortcomings in the Bohr Model
Bohr’s model of the atom only worked for hydrogen.
Bohr’s model did not explain the chemical behavior of atoms.

42. Louis De Broglie
Evidence from the photoelectric effect and hydrogen’s line emission spectrum proved that light had both particle and wavelike properties.
De Broglie stated that electrons could behave a dual wave-particle nature as well.