1. Chapter 9 Electrons in Atoms and the Periodic Table
Helium gas
He has 2 protons and
if neutral how many
electrons?
2006, Prentice Hall

2. CHAPTER OUTLINE Waves Electromagnetic Radiation Dual Nature of Light
Bohr Model of Atom
Quantum Mechanical Model of Atom
Electron Configuration
Electron Configuration & Periodic Table
Abbreviated Electron Configuration
Periodic Properties

3. Blimps
blimps float because they are filled with a gas that is less dense than the surrounding air
early blimps used the gas hydrogen, however hydrogen’s flammability lead to the Hindenburg disaster
blimps now use helium gas, which is not flammable

4. Why is hydrogen gas diatomic but helium gas not?
Because hydrogen is reactive and helium is inert.
What makes hydrogen reactive?
Recall that when elements are arranged in order of
increasing atomic number (# of protons), certain sets
of properties periodically recur
-All group I elements have similar reactivity
-All noble gases are inert
-Alkali metals- +1; metals
-Alkaline earth metals- +2; metals
-Halogens- -1; nonmetals
-Noble gases- 0; nonmetals
We need a theory/model that would explain why these occur:

5. Classical View of the Universe
since the time of the ancient Greeks, the stuff of the physical universe has been classified as either matter or energy
we define matter as the stuff of the universe that has mass and volume
therefore energy is the stuff of the universe that does not have mass and volume
we know that matter is ultimately composed of particles, and the properties of particles determine the properties we observe
energy therefore should not be composed of particles, in fact the thing that all energy has in common is that it travels in waves

6. Light: Electromagnetic Radiation
light is a form of energy
light is one type of energy called electromagnetic radiation with electric and magnetic field components and travels in waves

7. Electromagnetic Waves
velocity = c = speed of light
its constant! = x 108 m/s (m•sec-1) in vacuum
all types of light energy travel at the same speed
amplitude = A = measure of the intensity of the wave, “brightness”
height of the wave
wavelength = l = distance between crests
generally measured in nanometers (1 nm = 10-9 m)
same distance for troughs or nodes
determines color
frequency = n = how many peaks pass a point in a second
generally measured in Hertz (Hz),
1 Hz = 1 wave/sec = 1 sec-1

8. Electromagnetic Waves
Wavelength (λ) is the distance between any 2 successive crests or troughs.
Amplitude
Nodes

9. inversely proportional
WAVES
Frequency (nu,) is the number of waves produced per unit time.
Wavelength and frequency are inversely proportional.
Speed tells how fast waves travel through space.
inversely proportional
As wavelength of a wave increases
its frequency decreases
10.1

10. ELECTROMAGNETIC RADIATION
Energy travels through space as electromagnetic radiation. This radiation takes many forms, such as sunlight, microwaves, radio waves, etc.
In a vacuum, all electromagnetic waves travel at the speed of light (3.00 x 108 m/s), and differ from each other in their frequency and wavelength.

11. ELECTROMAGNETIC RADIATION
These waves range from -rays (short λ, high f) to radio waves (long λ, low f).
The classification of electromagnetic waves according to their frequency is called electromagnetic spectrum.
Long wavelength Low frequency
Short wavelength High frequency

12. ELECTROMAGNETIC RADIATION
Visible light is a small part of the EM spectrum
Infrared waves have longer λ but lower  than visible light
X-rays have longer λ but lower  than -rays
10.2

13. Particles of Light
Albert Einstein and other scientists in the early 20th century showed that wave properties do not completely explain electromagnetic radiation (EM) and showed that EM was composed of particle-like properties called photons
photons are particles of light energy
each wavelength of light has photons that have a different amount of energy
the longer the wavelength, the lower the energy of the photons

14. DUAL NATURE OF LIGHT
Scientists also have much evidence that light beams act as a stream of tiny particles, called photons.
Red light has longer wavelength and less energy than blue light
A photon of red light
A photon of blue light

15. DUAL NATURE OF LIGHT
Scientists, therefore, use both the wave and particle models for explaining light. This is referred to as the wave-particle nature of light.
Scientists also discovered that when atoms are energized at high temperatures or by high voltage, they can radiate light. Neon lights are an example of this property of atoms.

16. Light’s Relationship to MatterHe
Atoms can acquire extra energy, but they must eventually release it
When atoms emit energy, it always is released in the form of light
However, atoms don’t emit all colors, only very specific wavelengths
in fact, the spectrum of wavelengths can be used to identify the element Hg

17. When an atom absorbs energy it reemits it as light
emission spectrum of Hydrogen
White Light Source
Emits at every wavelength
(all colors) also called a
continuous spectrum
H produces its own unique and
distinctive emission spectrum

18. Spectra

19. The Bohr Model of the Atom
The Nuclear Model of the atom does not explain how the atom can gain or lose energy
Neils Bohr developed a model of the atom to explain the how the structure of the atom changes when it undergoes energy transitions
Bohr’s major idea was that the energy of the atom was quantized, and that the amount of energy in the atom was related to the electron’s position in the atom
quantized means that the atom could only have very specific amounts of energy
1885 to1962
Nobel Prize in Physics in 1922

20. The Bohr Model of the Atom Electron Orbits
the Bohr Model, electrons travel in orbits around the nucleus
more like shells than planet orbits
the farther the electron is from the nucleus the more energy it has
Rank the electrons from
highest to lowest energy e- e- e- e- e-

21. The Bohr Model of the Atom Orbits and Energy
each orbit has a specific amount of energy
the energy of each orbit is characterized by an integer - the larger the integer, the more energy an electron in that orbit has and the farther it is from the nucleus
the integer (whole #s), n, is called a quantum number
Bohr orbits are like steps in a ladder. It is possible to be on one
step or another, but it is impossible to be between steps.

22. BOHR MODEL OF ATOM
In this model, the electrons could only occupy particular energy levels, and could “jump” to higher levels by absorbing energy.
Neils Bohr, a Danish physicist, studied the hydrogen atom extensively, and developed a model for the atom that was able to explain the line spectrum.
Bohr’s model of the atom consisted of electrons orbiting the nucleus at different distances from the nucleus, called energy levels.

23. BOHR MODEL OF ATOM
The lowest energy level is called ground state, and the higher energy levels are called excited states.
When electrons absorb energy through heating or electricity, they move to higher energy levels and become excited.
energy

24. BOHR MODEL OF ATOM
When excited electrons return to the ground state, energy is emitted as a photon of light is released.
The color (wavelength) of the light emitted is determined by the difference in energy between the two states (excited and ground).
Lower energy transition give off red light
Higher energy transition give off blue light

25. The Bohr Model of the Atom Ground and Excited States
The lowest amount of energy hydrogen’s one electron can have corresponds to being in the n = 1 orbit –this is the ground state
when the atom gains energy, the electron leaps to a higher energy orbit –this is the excited state
the atom is less stable in an excited state, and so it will release the extra energy to return to the ground state
either all at once or in several steps

26. The Bohr Model of the Atom Hydrogen Spectrum
every hydrogen atom has identical orbits, so every hydrogen atom can undergo the same energy transitions
however, since the distances between the orbits in an atom are not all the same, no two leaps in an atom will have the same energy
the closer the orbits are in energy, the lower the energy of the photon emitted
lower energy photon = longer wavelength
therefore we get an emission spectrum that has a lot of lines that are unique to hydrogen

27. The Bohr Model of the Atom Hydrogen Spectrum
Which e- has longer wavelength and lower
energy (red, violet or blue-green)?

28. The Bohr Model of the Atom Success and Failure
the mathematics of the Bohr Model very accurately predicts the spectrum of hydrogen
however its mathematics fails when applied to multi-electron atoms
it cannot account for electron-electron interactions
a better theory was needed

29. QUANTUM MECHANICAL MODEL OF ATOM
In 1926 Erwin Shrödinger created a mathematical model that showed electrons as both particles and waves. This model was called the quantum mechanical model.
This model predicted electrons to be located in a probability region called orbitals.
An orbital is defined as a region around the nucleus where there is a high probability of finding an electron.
1887 to 1961
Nobel Prize in Physics in 1933

30. Orbits vs. Orbitals Pathways vs. Probability
Orbital – acts as a wave and particle
thus could end up anywhere in a
probability map
Orbit – acts as a particle and
follows a well-defined path

31. QUANTUM MECHANICAL MODEL OF ATOM
Based on this model, there are discrete principal energy levels within the atom.
Principal energy levels are designated by n.
The electrons in an atom can exist in any principal energy level.
As n increases, the energy of the electron increases

32. QUANTUM MECHANICAL MODEL OF ATOM
Each principal energy level is subdivided into sublevels.
The sublevels are designated by the letters s, p, d and f.
As n increases, the number of sublevels increases.
10.7, 10.8

33. QUANTUM MECHANICAL MODEL OF ATOM
The number of orbitals within the sublevels vary with their type.
Within the sublevels, the electrons are located in orbitals. The orbitals are also designated by the letters s, p, d and f.
s sublevel
= 1 orbital
= 2 electrons
p sublevel
= 3 orbitals
= 6 electrons
d sublevel
= 5 orbitals
= 10 electrons
f sublevel
= 7 orbitals
= 14 electrons
An orbital can hold a maximum of 2 electrons

34. How does the 1s Subshell Differ from the 2s Subshell

35. Probability Maps & Orbital Shape s Orbitals

36. Probability Maps & Orbital Shape p Orbitals

37. Probability Maps & Orbital Shape d Orbitals

38. ELECTRON CONFIGURATION
The distribution of electrons into the various energy shells and subshells in an atom’s ground state is called its electron configuration
The electrons occupy the orbitals from the lowest energy level to the highest level (Aufbau Principal).
The energy of the orbitals on any level are in the following order: s < p < d < f.
Each orbital on a sublevel must be occupied by a single electron before a second electron enters (Hund’s Rule).

39. ELECTRON CONFIGURATION
Electron configurations can be written as:
Number of electrons in orbitals
2 p6
Principal energy level
Type of orbital

40. ELECTRON CONFIGURATION
Another notation, called the orbital notation is shown below:
Electrons in orbital with opposing spins
Principal energy level
1 s
Type of orbital

41. Filling an Orbital with Electrons
each orbital may have a maximum of 2 electrons with opposite spins
Pauli Exclusion Principle
electrons spin on an axis
generating their own magnetic field
when two electrons are in the same orbital, they must have opposite spins
so there magnetic fields will cancel

42. ELECTRON CONFIGURATION↑ 1s H
1s1
Hydrogen has 1 electron. It will occupy the orbital of lowest energy which is the 1s. ↑ 1s ↓ He
1s2
Helium has two electrons. Both helium electrons occupy the 1s orbital with opposite spins.

43. ELECTRON CONFIGURATION↑ ↓ 1s ↑ ↓ 2s ↑ 2p B
1s22s22p1
Boron has the first p electron. The three 2p orbitals have the same energy. It does not matter which orbital fills first. ↑ ↓ 1s ↑ ↓ 2s ↑ 2p ↓ ↑ C
1s22s22p2
The second p electron of carbon enters a different p orbital than the first p due to Hund’s Rule.

44. ELECTRON CONFIGURATION↑ ↓ 1s ↑ ↓ 2s ↑ 2p ↓ ↑ ↓ ↑ ↓ Ne
1s22s22p6
The last p electron for neon pairs up with the last lone electron and completely fills the 2nd energy level. ↑ 3s Na
1s22s22p6
3s1
Sodium has 11 electron. The first 10 will occupy the orbitals of energy levels 1 and 2.
core electrons
valence electron

45. ELECTRON CONFIGURATION
As electrons occupy the 3rd energy level and higher, some anomalies occur in the order of the energy of the orbitals.
Knowledge of these anomalies is important in order to determine the correct electron configuration for the atoms.

46. ELECTRON CONFIG. & PERIODIC TABLE
10.15

47. ELECTRON CONFIG. & PERIODIC TABLE
The horizontal rows in the periodic table are called periods. The period number corresponds to the number of energy levels that are occupied in that atom.
The vertical columns in the periodic table are called groups or families. For the main-group elements, the group number corresponds to the number of electrons in the outermost filled energy level (valence electrons).

48. ELECTRON CONFIG. & PERIODIC TABLE
One energy level
4 energy levels
3 energy levels

49. ELECTRON CONFIG. & PERIODIC TABLE
3 valence electrons
1 valence electron
5 valence electrons

50. ELECTRON CONFIG. & PERIODIC TABLE
Note that elements in the same group have similar electron configurations.
The valence electrons configuration for the elements in periods 1-3 are shown below.
10.15

51. Arrangement of orbitals in the periodic table
ELECTRON CONFIG. & PERIODIC TABLE
Arrangement of orbitals in the periodic table
10.16

52. d orbital numbers are 1 less than the period number
ELECTRON CONFIG. & PERIODIC TABLE
d orbital numbers are 1 less than the period number
10.16

53. f orbital numbers are 2 less than the period number
ELECTRON CONFIG. & PERIODIC TABLE
f orbital numbers are 2 less than the period number
10.16

54. ABBREVIATED ELECTRON CONFIG.
When writing electron configurations for larger atoms, an abbreviated configuration is used.
In writing this configuration, the non-valence (core) electrons are summarized by writing the symbol of the noble gas prior to the element in brackets followed by configuration of the valence electrons.

55. ABBREVIATED ELECTRON CONFIG.
Previous noble gas K
Z = 19
1s22s22p63s23p6
4s1
valence electron
core electrons
[Ar]
4s1

56. ABBREVIATED ELECTRON CONFIG.
Z = 35
1s22s22p63s23p6
4s2
3d10
4p5
valence electrons
core electrons
[Ar]
4s23d104p5

57. Electron Configuration of As from the Periodic Table1 2 3 4 5 6 7 2A 3A 4A 5A 6A 7A
3d10 Ar As
4s2
4p3
As = [Ar]4s23d104p3
As has 5 valence electrons

58. TRENDS IN PERIODIC PROPERTIES
The electron configuration of atoms are an important factor in the physical and chemical properties of the elements.
Some of these properties include: atomic size, ionization energy and metallic character.
These properties are commonly known as periodic properties and increase or decrease across a period or group, and are repeated in each successive period or group.

59. ATOMIC SIZE
The size of the atom is determined by its atomic radius, which is the distance of the valence electron from the nucleus.
For each group of the representative elements, the atomic size increases going down the group, because the valence electrons from each energy level are further from the nucleus.

60. ATOMIC SIZE

61. Be (4p+ & 4e-) Mg (12p+ & 12e-) Ca (20p+ & 20e-) Group IIA Be Mg Ca 4

62. ATOMIC SIZE
The atomic radius of the representative elements are affected by the number of protons in the nucleus (nuclear charge).
For elements going across a period, the atomic size decreases because the increased nuclear charge of each atom pulls the electrons closer to the nucleus, making it smaller.

63. ATOMIC SIZE

64. Li (3p+ & 3e-) Be (4p+ & 4e-) B (5p+ & 5e-) C (6p+ & 6e-)
Period 2
2e-
1e-
3 p+
2e-
4 p+
2e-
3e-
5 p+
Li (3p+ & 3e-)
Be (4p+ & 4e-)
B (5p+ & 5e-)
6 p+
2e-
4e-
8 p+
2e-
6e-
10 p+
2e-
8e-
C (6p+ & 6e-)
O (8p+ & 8e-)
Ne (10p+ & 10e-)

65. IONIZATION ENERGY Na (g) + IE Na+ + e-
The ionization energy is the energy required to remove a valence electron from the atom in a gaseous state.
When an electron is removed from an atom, a cation (+ ion) with a 1+ charge is formed.
Na (g) + IE Na+ + e-

66. IONIZATION
ENERGY
The ionization energy decreases going down a group, because less energy is required to remove an electron from the outer shell since it is further from the nucleus.
Larger atom
Less IE

67. IONIZATION
ENERGY
Going across a period, the ionization energy increases because the increased nuclear charge of the atom holds the valence electrons more tightly and therefore it is more difficult to remove.

68. IONIZATION
ENERGY
In general, the ionization energy is low for metals and high for non-metals.
Review of ionization energies of elements in periods 2-4 indicate some anomalies to the general increasing trend.

69. IONIZATION ENERGY Be 1s2 2s2 N 1s2 2s2 2p3 B 1s2 2s2 2p1 O 1s2 2s2 2p4
These anomalies are caused by more stable electron configurations of the atoms in groups 2 (complete “s” sublevel) and group 5 (half-filled “p” sublevels) that cause an increase in their ionization energy compared to the next element.
More stable (1/2 filled)
Higher IE
More stable
Higher IE Be
1s2 2s2 N
1s2 2s2 2p3 B
1s2 2s2 2p1 O
1s2 2s2 2p4

70. METALLIC
CHARACTER
Metallic character is the ability of an atom to lose electrons easily.
This character is more prevalent in the elements on the left side of the periodic table (metals), and decreases going across a period and increases for elements going down a group.

71. Most metallic elements Least metallic elements
CHARACTER
Most metallic elements
Least metallic elements

72. Larger due to less nuclear charge
Example 1:
Select the element in each pair with the larger atomic radius: or K Br
Larger due to less nuclear charge

73. Highest IE due to most nuclear charge
Example 2:
Indicate the element in each set that has the higher ionization energy and explain your choice: or F N C
Highest IE due to most nuclear charge

74. THE END