1. Chapter 11
Modern Atomic Theory

2. Light Made up of electromagnetic radiation.
Waves of electric and magnetic fields are at right angles to each other.

3. Parts of a wave Wavelength l
Frequency = number of cycles in one second
Measured in hertz 1 hertz = 1 cycle/second

4. Frequency = n

5. The Nature of Waves

6. Kinds of EM waves There are many different l and n
Radio waves, microwaves, x rays and gamma rays are all examples.
Light is only the part our eyes can detect. G a m m a R a d i o R a y s w a v e s

7. Electromagenetic Spectrum

8. The speed of light In a vacuum is 2.998 x 108 m/s = c c = ln
Ex 1: What is the wavelength of light with a frequency 5.89 x 105 Hz?
Ex 2: What is the frequency of blue light with a wavelength of 484 nm?

9. In 1900
Matter and energy were seen as different from each other in fundamental ways.
Matter was particles.
Energy could come in waves, with any frequency.
Max Planck found that as the cooling of hot objects couldn’t be explained by viewing energy as a wave.

10. Energy is Quantized Planck found DE came in chunks with size hn
DE = nhn
where n is an integer (which we will ignore for our calculations at this point).
h is Planck’s constant=6.626 x J•s
These packets of hn are called quantum

11. Einstein is next
Said electromagnetic radiation is quantized in particles called photons.
Each photon has energy E= hn = hc/l

12. Fireworks in Washington D.C.

13. What are electrons doing in an atom?

14. Which is it? Is energy a wave like light, or a particle? Yes
Concept is called the Wave -Particle duality.
What about the other way, is matter a wave? Yes

15. Electrons move as waves
A wave has three characteristics: wavelength, frequency, and speed.
Wavelength-distance between wave peaks
Frequency-how many peaks pass a point per a certain amount of time
Speed-how fast a peak travels through water.

16. Spectrum The range of frequencies present in light.
White light has a continuous spectrum.
All the colors are possible.
A rainbow.

17. Hydrogen spectrum 656 nm 434 nm 410 nm 486 nm
Emission spectrum because these are the colors it gives off or emits.
Called a line spectrum.
There are just a few discrete lines showing
656 nm
434 nm
410 nm
486 nm

18. What this means
Only certain energies are allowed for the hydrogen atom.
Can only give off certain energies.
Use DE = hn = hc / l
Energy in the in the atom is quantized.

19. What will an electron do?
It has mass, so it is matter.
A particle can only go through one hole.
A wave through both holes.
An electron does go though both, and makes an interference pattern.
It behaves like a wave.
Other matter has wavelengths too short to notice.

20. The Bohr Ring Atom-1st quantum model

21. The Bohr Model Doesn’t work. Only works for hydrogen atoms.
Electrons don’t move in circles.
The quantization of energy is right, but not because they are circling like planets.

22. We are worried about the change
When the electron moves from one energy level to another.
DE = Efinal – Einitial

23. Electrons also travel as particles
n is the principal energy level
n = 1 is called the ground state and is the lowest energy state of an atom
Excited state-atom has a higher potential
When an excited atom returns to its ground state, it gives off energy in the form of a photon.
Photon-stream of tiny energy packets.

24. Electromagnetic Radiation Exhibits Wave Properties and Particulate Properties

25. (a)The Probability Distribution (b) The Probability Along a Line Drawn From the Nucleus Outward in Any Direction

26. Louis deBroglie
Louis deBroglie said electrons can be considered as waves that are fixed around a nucleus instead of orbits as Bohr stated.
This is known as the wave mechanical model.

27. Heisenburg
The Heisenburg Uncertainty Principle states that it is impossible to determine the position and velocity (speed) of an electron.

28. Quantum numbers
Angular momentum quantum number l , gives the shape of the orbital.
integer values from 0 to n-1
l = 0 is called s
l = 1 is called p
l =2 is called d
l =3 is called f
l =4 is called g

29. The Angular Momentum Quantum Numbers and Corresponding Letters Used to Designate Atomic Orbitals

30. Electron Shapes s (groups 1 and 2 including H and He)
holds up to 2 electrons
p (groups 13-18)
holds up to 6 electrons
d (groups 3-12)
holds up to 10 electrons
f (lanthanides and actinides)
holds up to 14 electrons

31. The Orbitals Being Filled for Elements in Various Parts of the Periodic Table

32. Representation of the 2p Orbitals (a) The Electron Probability Distribution for a 2p Oribtal (b) The Boundary Surface Representations of all Three 2p Orbitals

33. Representation of the 3d Orbitals (a) Electron Density Plots of Selected 3d Orbitals (b) The Boundary Surfaces of All of the 3d Orbitals

34. Quantum numbers Magnetic quantum number (m)
Tells direction in each shape.
Can have 2 values; either +1/2 or -1/2

35. A Picture of the Spinning Electron

36. The Periodic Table
Developed independently by German Julius Lothar Meyer and Russian Dmitri Mendeleev (1870”s).
Didn’t know much about atom.
Put in columns by similar properties and mass.
Predicted properties of missing elements.

37. Aufbau Principle Aufbau is German for building up.
As the protons are added one by one, the electrons fill up hydrogen-like orbitals.
Fill up in order of energy levels.

38. Increasing energy He with 2 electrons 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p3d 4d 5d 7p 6d 4f 5f
He with 2 electrons

39. Details
Valence electrons- the electrons in the outermost energy levels (not d).
Core electrons- the inner electrons.
Hund’s Rule- The lowest energy configuration for an atom is the one have the maximum number of unpaired electrons in the orbital.
C 1s2 2s2 2p2

40. Exceptions Cr =[Ar] 4s1 3d5 Cu=[Ar] 4s1 3d10 Half filled orbitals.
Scientists aren’t sure of why it happens

41. Electron Configurations for Potassium Through Krypton

42. Ionization Energy
Ionization Energy- The energy necessary to remove an electron from a gaseous atom.
Highest energy electron removed first.
It takes much more energy to remove a core electron than a valence electron because there is less shielding.

43. Trends in Ionization Energies (kj/mol) for the Representative Elements

44. Shielding
Electrons on the higher energy levels tend to be farther out.
Have to look through the other electrons to see the nucleus.
They are less effected by the nucleus.

45. Across a Period
Generally from left to right, IE increases because there is a greater nuclear charge with the same shielding.
As you go down a group IE decreases because electrons are farther away.

46. The Radius of an Atom (r) is Defined as Half the Distance Between the Nuclei in a Molecule Consisting of Identical Atoms

47. Atomic Radii (in Picometers) for Selected Atoms