1. Ch. 5 Electrons in Atoms

2. The Development of Atomic Models
Ernest Rutherford’s existing model
Ask students to recall what Rutherford’s Gold Foil experiment proved about the atom. (the atom consists of protons and neutrons that make up the nucleus, with electrons surrounding it) After discovering the atomic nucleus, Rutherford used existing ideas about the atom and proposed an atomic model in which the electrns move around the nucleus, like planets move around the sun.

3. Models of the Atom > The Rutherford Model
5.1
Models of the Atom > The Rutherford Model
What was inadequate about Rutherford’s atomic model?
Rutherford’s atomic model could not explain the chemical properties of elements.
Rutherford’s atomic model could not explain why objects change color when heated.
Observing that the as the horseshoe is heated, the color changes from black to red, then yellow, and then white. This observation could be explained only if the atoms in the iron gave off light in specific amounts of energy. Explaining what leads to the chemical properties of elements requires a model that better describes the behavior of electrons within atoms.

4. The Development of Atomic Models
5.1
The Development of Atomic Models
The timeline shoes the development of atomic models from 1803 to 1911.
These illustrations show how the atomic model has changed as scientists learned more about the atom’s structure.

5. The Development of Atomic Models
5.1
The Development of Atomic Models
The timeline shows the development of atomic models from 1913 to 1932.
These illustrations show how the atomic model has changed as scientists learned more about the atom’s structure.

6. Models of the Atom > The Bohr Model
5.1
Models of the Atom > The Bohr Model
Niels Bohr (1913) changed Rutherford’s model to include how the energy of an atom changes when it absorbs or emits light.
Bohr proposed a model of the atom showing a dense nucleus with electrons found in definite paths called orbitals surrounding the nucleus.

7. Models of the Atom > The Bohr Model
Each possible electron orbit in Bohr’s model has a fixed energy.
The fixed energies an electron can have are called energy levels.
A quantum of energy is the amount of energy required to move an electron from one energy level to another energy level.

8. Models of the Atom > The Bohr Model
5.1
Models of the Atom > The Bohr Model
Like the rungs of the strange ladder, the energy levels in an atom are not equally spaced.
The higher the energy level occupied by an electron, the less energy it takes to move from that energy level to the next higher energy level.
These ladder steps are somewhat like energy levels. In an ordinary ladder, the rungs are equally spaced. The energy levels in atoms are unequally spaced, like the rungs in this ladder. The lowest rung of the ladder corresponds to the lowest energy level. The higher energy levels are closer together and father from the nucleus.
A person can climb up or down a ladder by going from rung to rung. Similarly, an electron can jump from one energy level to another. A person on a ladder can not stand in between the rungs. And an atom can not be between energy levels.
To move from one energy level to another, an electron must gain or lose just the right amount of energy. The amount of energy an electron gains or loses in an atom is not always the same. It takes less energy to climb from one rung to the next near the top of the ladder because they are closer together.

9. Models of the Atom > The Bohr Model
Bohr identified 7 different orbits, each with a maximum number of electrons possible
K – L – M – N – O – P – Q

10. Filling Bohr Models To fill in a Bohr Model:
Identify the number of protons and neutrons in the nucleus of an element.
Fill the electrons in the energy levels starting at the level closest to the nucleus.
Each level can only hold the designated number of electrons.
2 – 8 – 18 – 32
Energy level:
(Your Periodic Table has the Bohr configuration)

11. Let’s Try Na O

13. Models of the Atom > The Bohr Model
Bohr’s model was supported by emission spectra of hydrogen, but does not work for larger atoms

14. Quantum Mechanical Model
Models of the Atom > The Quantum Mechanical Model
Quantum Mechanical Model
Proposed by Erwin Schrödinger (1926)
Electrons are still present in discrete energy levels
Instead of orbits, the electrons are in clouds of probability around nucleus
The energy levels are divided into sublevels called “orbitals”
Determines the allowed energies an e- can have & how likely it is to find the e- in various locations around the nucleus.
The Rutherford planetary model and the Bohr model of the atom are based on describing paths of moving electrons.

15. E- location is based on probability
Models of the Atom > The Quantum Mechanical Model
E- location is based on probability
What is the probability of reaching in the bag and pulling out a blue marble?
What is the probability of reaching in the bag and pulling out an orange marble? 3 4
= 33%
= 25% 12 12

16. Models of the Atom > The Quantum Mechanical Model
Probability of finding an e- within a certain volume of space surrounding the nucleus can be represented by a fuzzy cloud.
The cloud is more dense where the probability of finding an e- is high.
The cloud is less dense where the probability of finding an e- is low.
Compare the electron moving around the nucleus to that of the motion of a rotating propeller blade. Propeller blade as has same probability of being anywhere in the blurry region it produces in the picture, but you cannot tell its precise location at any instant.

17. Atomic Orbitals
An atomic orbital is often thought of as a region of space in which there is a high probability of finding an electron.
Each energy sublevel corresponds to an orbital of a different shape, which describes where the electron is likely to be found.

18. Atomic Orbitals
Different atomic orbitals are denoted by letters. The s orbitals are spherical, and p orbitals are dumbbell-shaped.

19. Atomic Orbitals
Four of the five d orbitals have the same shape but different orientations in space.
The d orbitals are illustrated here. Four of the five d orbitals have the same shape but different orientations in space. Interpreting Diagrams How are the orientations of the dxy and dx2 – y2 orbitals similar? How are they different?

20. Atomic Orbitals
The numbers and kinds of atomic orbitals depend on the energy sublevel.

21. Atomic Orbitals
Observe the characteristics of atomic orbitals
The number of electrons allowed in each of the first four energy levels are shown here.
Each orbital can have a max of 2 e-’s
Click link when done covering slide.

22. Bohr vs. Quantum
Electron configuration for Rubidium

24. Were You Paying Attention?
1. Rutherford's planetary model of the atom could not explain
any properties of elements.
the chemical properties of elements.
the distribution of mass in an atom.
the distribution of positive and negative charges in an atom.

25. Were You Paying Attention?
2. Bohr's model of the atom proposed that electrons are found
embedded in a sphere of positive charge.
in fixed positions surrounding the nucleus.
in circular orbits at fixed distances from the nucleus.
orbiting the nucleus in a single fixed circular path.

26. Were You Paying Attention?
3. What is the lowest-numbered principal energy level in which p orbitals are found? 1 2 3 4

27. Bell – Ringer Date: Think about the stability of these arrangements
Does this scene look natural to you?
The Thinker is a bronze sculpture on marble pedestal by Auguste Rodin was his first cast in 1902, now located in a museum in Paris.
Rock arrangements like these are natural, but rare because they are unstable. Unstable arrangements, whether grains of sand in a sandcastle or the rock formation shown here, tend to become more stable by losing energy.
If this rock were to tumble over, it would end up at a lower height. It would have less energy than before, but its position would be more stable.
You will learn that energy and stability play an important role in determining how electrons are configured in an atom.

28. Electron Configurations
5.2
Electron Configurations
Electron configurations – the ways in which electrons are arranged in various orbitals around the nuclei of atoms
Three rules tell you how to find the electron configurations of atoms:
Aufbau Principle
Pauli Exclusion Principle
Hund’s Rule

29. Electron Configurations
5.2
Aufbau Principle
According to the aufbau principle, electrons occupy the orbitals of lowest energy first. In the aufbau diagram below, each box represents an atomic orbital.
This aufbau diagram shows the energy levels of the various atomic orbitals. Orbitals of greater energy are higher on the diagram. Using Tables Which is of higher energy, a 4d or a 5s orbital?

30. Electron Configuration
Diagonal Rule:
Lowest
Highest

31. Electron Configurations
5.2
Electron Configurations
Pauli Exclusion Principle
States that an atomic orbital may describe at most two electrons.
To occupy the same orbital, two electrons must have opposite spins; that is, the electron spins must be paired.

32. Electron Configurations
5.2
Hund’s Rule
When filling sublevels other than s, electrons are placed in individual orbitals before they are paired up.
“…up, up, up… down, down, down…”
Electrons fill like people do on a bus. You would never sit right next to someone you did not know if there are free seats available, unless of course all the seats are taken then you must pair up.
So when working with the p sublevel, electrons fill like this....up, up, up...down, down, down...take a look

33. Electron Configurations
5.2
Orbital Filling Diagram

34. Write the electron configuration for phosphorous.
An element found in matches. Atomic #15 1s2 2s2 sp6 3s2 3p3

35. Write the electron configuration for carbon.
Atomic #6 1s2 2s2 2p2

36. Write the electron configuration for argon.
Atomic # s2 2s2 sp6 3s2 3p6

37. Write the electron configuration for nickel.
Atomic # s2 2s2 2p6 3s2 3p6 3d8 4s2

38. Write the electron configuration for boron
Write the electron configuration for boron. How many unpaired electrons does each atom have?
Atomic # s2 2s2 2p1 1 unpaired electron

39. Write the electron configuration for silicon
Write the electron configuration for silicon. How many unpaired electrons does each atom have?
Atomic # s2 2s2 2p6 3s2 3p2 2 unpaired electrons

40. Predicting electron configurations using the periodic table.

42. Electron Configuration Practice
Li Quantum 1s2 2s1
Ne Quantum 1s2 2s2 2p6
Se Quantum 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p4 K Mn Br

43. Exceptions to the Aufbau Principle
Half-filled sublevels are not as stable as filled sublevels, but they are more stable than other configurations.
Examples:
Cr 1s2 2s2 2p6 3s2 3p6 3d4 4s2
Cu 1s2 2s2 2p6 3s2 3p6 3d9 4s2
1s2 2s2 2p6 3s2 3p6 3d5 4s1
1s2 2s2 2p6 3s2 3p6 3d10 4s1
WRONG
RIGHT

44. Visualize the exceptions Cr. 1s2 2s2 2p6 3s2 3p6 3d4 4s2
Visualize the exceptions Cr 1s2 2s2 2p6 3s2 3p6 3d4 4s2 1s2 2s2 2p6 3s2 3p6 3d5 4s1
WRONG
RIGHT

45. Visualize the exceptions Cu. 1s2 2s2 2p6 3s2 3p6 3d9 4s2
Visualize the exceptions Cu 1s2 2s2 2p6 3s2 3p6 3d9 4s2 1s2 2s2 2p6 3s2 3p6 3d10 4s1
WRONG
RIGHT

46. Were You Paying Attention?
1. Identify the element that corresponds to the following electron configuration: 1s22s22p5. F Cl Ne O

47. Were You Paying Attention?
2. Write the electron configuration for the atom N.
1s22s22p5
1s22s22p3
1s22s1p2
1s22s22p1

48. Were You Paying Attention?
The electron configurations for some elements differ from those predicted by the aufbau principle because the
the lowest energy level is completely filled.
none of the energy levels are completely filled.
half-filled sublevels are less stable than filled energy levels.
half-filled sublevels are more stable than some other arrangements. `

49. The Quantum Mechanical Model > Light
5.3
The amplitude of a wave is the wave’s height from zero to the crest.
The wavelength, represented by  (the Greek letter lambda), is the distance between the crests.
The frequency, represented by  (the Greek letter nu), is the number of wave cycles to pass a given point per unit of time.
The SI unit of cycles per second is called a hertz (Hz).
Each complete wave cycle starts at zero, increases to its highest value, passes through zero to reach its lowest value, and returns to zero again.

50. The Quantum Mechanical Model > Light
5.3
The Quantum Mechanical Model > Light
Wavelength and frequency are inversely proportional to each other
As wavelength increases, frequency decreases

51. Light
5.3
The electromagnetic spectrum consists of radiation over a broad band of wavelengths.
Radiation - the process in which energy is emitted as particles or waves.
Light exists in waves
The electromagnetic spectrum consists of radiation over a broad band of wavelengths. The visible light portion is very small. It is in the 10-7m wavelength range and 1015 Hz (s-1) frequency range. Interpreting Diagrams What types of nonvisible radiation have wavelengths close to those of red light? To those of blue light? (infrared, ultraviolet)
What color in the visible spectrum has the longest wavelength? (red)

52. Electromagnetic Radiation
Energy that travels in waves
We organize it by wavelength: a a e b e
b _________________
c _________________
d _________________
Crest – trough - wave length a c d

53. The Quantum Mechanical Model > Light
5.3
Electromagnetic radiation includes radio waves, microwaves, infrared waves, visible light, ultraviolet waves, X-rays, and gamma rays.

54. The Quantum Mechanical Model > Light
5.3
Sunlight consists of light with a continuous range of wavelengths and frequencies.
When sunlight passes through a prism, the different frequencies separate into a spectrum of colors.
In the visible spectrum, red light has the longest wavelength and the lowest frequency.
The color of light for each frequency found in sunlight depends on the frequency. All colors are present in white light. When sunlight passes through a prism, the different frequencies separate into a spectrum of colors; rainbow.

55. 5.3
Atomic Spectra
A prism separates light into the colors it contains. When white light passes through a prism, it produces a rainbow of colors.
When light from a helium lamp passes through a prism, discrete lines are produced.
A prism separates light into the colors it contains. Light from a helium lamp produces discrete lines. Identifying Which color has the highest frequency?

56. 5.3
Atomic Spectra
The frequencies of light emitted by an element separate into discrete lines to give the atomic emission spectrum of the element.
Mercury
Nitrogen
No two elements have the same emission spectrum. a) Mercury vapor lamps produce a blue glow. b) Nitrogen gas gives off a yellowish-orange light.

57. Emission Spectra

58. The Quantum Mechanical Model > Light
5.3
The Quantum Mechanical Model > Light
The product of the frequency and wavelength always equals a constant (c), the speed of light.
Equation can be used to determine elements.
c = 2.99 x 108 m/s

59. 5.1 Calculating the Wavelength of Light
Calculate the wavelength of the yellow light emitted by the sodium lamp shown above if the frequency of the radiation is x 1014 Hz
5.1
Knowns
Unknowns
frequency v = 5.10 x 1014 Hz
= 5.10 x 1014 /s
wavelength λ = ? m
speed of light c = 2.99 x 108 m/s
Sodium vapor lamps produce a yellow glow.

60. 5.3
Atomic Spectra
When atoms absorb energy, electrons move into higher energy levels. These electrons then lose energy by emitting light when they return to lower energy levels.
Passing an electric current through a gas in a neon tube energizes the electrons of the atom of the gas, and causes them to emit light.

61. Electrons ABSORB energy to become EXCITED
When an electron gains energy, it jumps to a higher energy level.
Since it does not normally have this much energy, it starts to lose it right away and falls back to its original level.
The energy previously absorbed is now released as infrared, ultraviolet, or visible light.
Copper is absorbing energy from a Bunsen burner flame.

62. A few definitions:
Quanta – the amount of energy needed to move an electron from its present energy level to the next higher
Photons – a quanta of light. A wave/particle that travels at the speed of light. Has zero mass and zero electrical charge.
Ground State – all electrons are in their lowest possible states
Excited State – energy has been input to raise electrons to a higher level

63. Phil the photon is impulsive traveler
Phil the photon is impulsive traveler. His favorite vacation spot is the Canary Islands, in Spain, with more than 300 days a year of guaranteed sunshine. His eyes are typically blood shot due to constant travel, long layovers, and jet lag. His hotel preference depends solely on the presence of a microwave and a radio in his room. This much travel can be draining on a little particle, which would explain his zero mass and electric charge.

64. Electrons go to a higher level if they gain energy
That energy is released as light when the electrons drop back down
We can use this information to identify elements
Bullet 1: Ask students: “This is also known as?” (the excited state)
Bullet 2: Ask students: “When the electrons drop back down, they are returning to what?” (the ground state)

65. Ground vs. Excited Electron Configurations
Carbon
Atomic #6
Ground State
Excited State
C is in row 2 of P.T. and has 2 orbitals (energy levels)
Found in Group 4A so it has 4 valence electrons
1s2 2s1 2p3
1s2 2s2 2p2

66. Phosphorus (P): Ground State
What is the electron configuration?
P is in row 3 of P.T. and has 3 orbitals (energy levels)
P has 5 valence electrons (the # of electrons in the outermost shell)
1s2 2s2 2p6 3s2 3p3 is the P atom’s electron configuration

67. Phosphorus (P): Excited State
The ground state configuration changes.
But the total number of electrons stays the same.
Ground state  1s2 2s2 2p6 3s2 3p3
Total electrons = 15
Excited State  1s2 2s2 2p6 3s1 3p4

68. Some Practice!
Ground State Excited State Mg F

69. And Some More Practice!
Ground State Excited State S K

70. An Explanation of Atomic Spectra
5.3
An Explanation of Atomic Spectra
The light emitted by an electron moving from a higher to a lower energy level has a frequency directly proportional to the energy change of the electrons.
Quantum of energy E is related to the frequency v of the emitted light by the equation
E = h x v
where h is equal to x J·s
How are the frequencies of light an atom emits related to changes of electron energies?

71.  Each transition produces a line of a specific frequency in the spectrum
Lyman Series – transitions to the n=1 energy level
Balmer Series – transitions to the n=2 energy level
(visible light)
Paschen Series – transitions to the n=3 energy level

72. 5.3
Quantum Mechanics
The Heisenberg uncertainty principle states that it is impossible to know exactly both the velocity and the position of a particle at the same time.
This limitation is critical in dealing with small particles such as electrons.
This limitation does not matter for ordinary-sized object such as cars or airplanes.

73. Quantum Mechanics The Heisenberg Uncertainty Principle
The Heisenberg uncertainty principle states that it is impossible to know exactly both the velocity and the position of a particle at the same time.

74. 5.3 Section Quiz.
1. Calculate the frequency of a radar wave with a wavelength of 125 mm. Hz Hz Hz Hz

75. 5.3 Section Quiz.
2. The lines in the emission spectrum for an element are caused by
the movement of electrons from lower to higher energy levels.
the movement of electrons from higher to lower energy levels.
the electron configuration in the ground state.
the electron configuration of an atom.

76. 5.3 Section Quiz.
3. Spectral lines in a series become closer together as n increases because the
energy levels have similar values.
energy levels become farther apart.
atom is approaching ground state.
electrons are being emitted at a slower rate.