Chapter 9
Electrons in Atoms
and the
Periodic Table
2
Waves
•Amplitude-- the height of the wave.
•Wavelength (λ) ---distance from one crest to
another
3
Frequency of a Wave
• Frequency (n) is the number of complete cycles
per given time.
Units are hertz (Hz), or cycles/s = s-1.
1 Hz = 1 s-1
• Higher Frequency– many cycles per second
• Lower Frequency--- less cycles per second
4
Low Frequency Wave
High Frequency Wave
l
l
l
5
The Nature of Light
• Light --- can be defined as a form of energy
and a form of electromagnetic waves.
• Electromagnetic waves
Has an Electrical and magnetic component
interacting perpendicularly
6
Electromagnetic Waves
• The properties of a wave are determined by:
wave speed
height (amplitude)
wavelength
frequency.
• All electromagnetic waves move through
space at the same, constant speed.
3.00 x 108 m/s = The speed of light, c.
7 Tro's "Introductory Chemistry",
Chapter 9
Electromagnetic Spectrum
8 Tro's "Introductory Chemistry",
Chapter 9
Types of Electromagnetic Radiation
• Classified by the Wavelength Radiowaves = l > 0.01 m.
Low frequency and energy.
Microwaves = 10-4m < l < 10-2 m.
Infrared (IR) = 8 x 10-7 < l < 10-5 m.
Visible = 4 x 10-7 < l < 8 x 10-7 m.
ROYGBIV.
Ultraviolet (UV) = 10-8 < l < 4 x 10-7 m.
X-rays = 10-10 < l < 10-8 m.
Gamma rays = l < 10-10.
High frequency and energy.
9
The Electromagnetic Spectrum • Light passed through a prism is separated into all its
colors. This is called a continuous spectrum.
• The color of the light is determined by its wavelength.
10 Tro's "Introductory Chemistry",
Chapter 9
Color
• The color of light is determined by its wavelength.
• White light is a mixture of all the colors of visible light.
RedOrangeYellowGreenBlueViolet.
• When an object absorbs some of the wavelengths of
white light while reflecting others, it appears colored.
The observed color is predominantly the colors reflected.
11 Tro's "Introductory Chemistry",
Chapter 9
Light’s Relationship to Matter
• Atoms can acquire extra energy, but
they must eventually release it.
• When atoms emit energy, it usually is
released in the form of light.
• However, atoms don’t emit all colors,
only very specific wavelengths.
Because energy is quantized
12 Tro's "Introductory Chemistry",
Chapter 9
Emission Spectrum
13 Tro's "Introductory Chemistry",
Chapter 9
Spectra
14
The Bohr Model of the Atom:
Electron Orbits • In the Bohr model, electrons travel in orbits
around the nucleus.
• Bohr’s major idea was that the energy of the atom was it is quantized.
The integer, n, is called a
quantum number.
n is directly proportional to energy
15 Tro's "Introductory Chemistry",
Chapter 9
The Bohr Model of the Atom:
Ground and Excited States
Ground state: The least energy state an atom can exist in.
Excited State: When an atom absorbs energy.
When the atom gains energy, the electron leaps to a higher energy orbit. The atom is less stable in an excited state and so it will release the extra energy to return to the ground state. Energy is released in the form of light.
16 Tro's "Introductory Chemistry",
Chapter 9
The Bohr Model of the Atom
17 Tro's "Introductory Chemistry",
Chapter 9
The Quantum-Mechanical Model
of the Atom
• Bohr model accurately predicts the spectrum of
hydrogen.
• However, it fails when applied to multi-electron
atoms.
• The quantum-mechanical model was proposed.
• According to this model it is possible to locate
regions in an atom where there is a higher
probability of finding an electron
• The region is called an orbital
18 Tro's "Introductory Chemistry",
Chapter 9
The Quantum-Mechanical Model:
Quantum Numbers
• In Schrödinger’s wave equation,
there are 3 integers, called
quantum numbers, that
quantize the energy.
• The principal quantum
number, n, specifies the main
energy level for the orbital.
19
Quantum Numbers, Continued
• The quantum-mechnical model uses parameters known
as quantum numbers to describe the structure of the
atom.
The principal quantum number designates the energy level.
• Each energy level is made of orbitals
• The quantum number that designates the orbital is often
given a letter.
s, p, d, f.
The number of orbital types = the principal quantum number.
20
Shells and Subshells
21
Orbital Types and Shapes:
s-orbital
s-orbital has a spherical shape
-has only one subshell
22
p-orbital has a dumbell shape
-has only three subshells
Orbital Types and Shapes:
p-orbital
23
Orbital Types and Shapes:
d-orbital
-has only five subshells
Electron Configurations
• Electron configuration-- distribution of electrons
in an atom
• Each energy shell and subshell has a maximum
number of electrons it can hold.
24
Orbital(or shell) Maximum number of
electrons
s 2
p 6
d 10
f 14
25
Filling an Orbital with Electrons
• Each orbital may have a maximum of 2 electrons.
Pauli Exclusion principle.
• When filling orbitals that have the same energy,
place one electron in each before completing pairs.
Hund’s rule.
• When filling, put electrons in lower energy orbitals first. e.g. a 1s orbital will be filled before the 2s
Aufbau principle.
• When two electrons are in the same orbital, they must have opposite spins.
26
Filling Orbitals According to Energy
Levels
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d
7s
1s orbital is the
lowest, so filling
should start there,
then 2s, 2p, 3s, 3p ….
Example:
Kr = 36 electrons = 1s22s22p63s23p64s23d104p6
27
Orbital Diagrams
• Orbital is represented as a square and the
electrons in that orbital as arrows.
Orbital with
1 electron
Unoccupied
orbital
Orbital with
2 electrons
28
Electron Configurations
29 Tro's "Introductory Chemistry",
Chapter 9
1. Determine the atomic number of the element
from the periodic table.
This gives the number of protons and electrons in
the atom.
Mg Z = 12, so Mg has 12 protons and 12 electrons.
Example—Write the Ground State
Orbital Diagram and Electron
Configuration of Magnesium.
30 Tro's "Introductory Chemistry",
Chapter 9
2. Draw 9 boxes to represent the first 3 energy
levels s and p orbitals.
1s 2s 2p 3s 3p
Example—Write the Ground State
Orbital Diagram and Electron
Configuration of Magnesium,
Continued.
31
3. Add one electron to each box in a set, then
pair the electrons before going to the next set
until you use all the electrons.
• When paired, put in opposite arrows.
1s 2s 2p 3s 3p
Example—Write the Ground State
Orbital Diagram and Electron
Configuration of Magnesium,
Continued.
32 Tro's "Introductory Chemistry",
Chapter 9
Example—Write the Ground State
Orbital Diagram and Electron
Configuration of Magnesium,
Continued. 4. Use the diagram to write the electron
configuration.
Write the number of electrons in each set as a
superscript next to the name of the orbital set.
1s22s22p63s2
1s 2s 2p 3s 3p
33
Example—Write the Full Ground State Electron
Configuration of Sc3+.
Sc Z = 21, therefore 21 e−
therefore Sc3+ has 18 e−
s subshell holds 2 e−
p subshell holds 6 e−
d subshell holds 10 e−
f subshell holds 14 e−
Therefore the electron configuration is 1s22s22p63s23p6
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d
7s
34
Practice—Write the Full Ground State Orbital
Diagram and Electron Configuration of Potassium.
35
Practice—Write the Full Ground State Orbital
Diagram and Electron Configuration of F−.
36
Valence Electrons
• The electrons in all the subshells with the highest principal energy shells are called the valence electrons.
• Electrons in lower energy shells are called core electrons.
• Valence electrons determine the chemical and physical properties of elements
37 Tro's "Introductory Chemistry",
Chapter 9
Valence Electrons, Continued
Rb = 37 electrons = 1s22s22p63s23p64s23d104p65s1
• The highest principal energy shell of Rb that contains
electrons is the 5th, therefore, Rb has 1 valence
electron and 36 core electrons.
Kr = 36 electrons = 1s22s22p63s23p64s23d104p6
• The highest principal energy shell of Kr that contains
electrons is the 4th, therefore, Kr has 8 valence
electrons and 28 core electrons.
38
Practice—Determine the Number and Types
of Electrons in an Arsenic, As, Atom.
39
Electron Configuration and the
Periodic Table • The periodic table is divided into blocks—
s-, p-, d- and f-blocks.
• For the s- and p-blocks
the period number corresponds to the principal
energy level of the valence electrons.
• For the d-block
the period number corresponds to the principal
energy level minus 1 of the valence electrons.
40
s1
s2
d1 d2 d
3 d4 d5 d6 d7 d8 d9 d10
s2 p1 p2 p3 p4 p5
p6
f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14 f14d1
1
2
3
4
5
6
7
41
Electron Configuration from
the Periodic Table
P = [Ne]3s23p3
P has five valence electrons.
3p3
P
Ne
1
2
3
4
5
6
7
1A
2A 3A 4A 5A 6A 7A
8A
3s2
Use the periodic table to write
the electron configuration of P.
42
Electron Configuration from
the Periodic Table
As = [Ar]4s23d104p3
As has five valence electrons.
4s2
Ar 3d10
4p3
As
1
2
3
4
5
6
7
1A
2A 3A 4A 5A 6A 7A
8A
43
Transition Elements • For the d-block metals, the principal energy level is one
less than the period number
Zn
Z = 30, period 4, group 2B
[Ar] 4s23d10
Sn
Z = 50, period 5, group 4B
[Kr] 5s24d105p2
44
Practice—Use the Periodic Table to Write the Short
Electron Configuration of Na and Te
• Na (at. no. 11).
• Te (at. no. 52).
• Ca
• Rb
45 Tro's "Introductory Chemistry",
Chapter 9
Practice Cont’d—Use the Periodic Table to Write
the Short Electron Configuration of Ca and Rb
• Cl
• Ba
46 Tro's "Introductory Chemistry",
Chapter 9
Practice Cont’d—Use the Periodic Table to Write
the Short Electron Configuration of Cl and Ba
47
Periodic Trends in Properties of
Elements
• Effective nuclear charge
• Atomic size (atomic radius)
• Ionization energy
• Metallic character
Two important questions to be answered are:
• What happens when you go across a period (left to right)?
• What happens when you go down the group (top to bottom)?
48
Trends in Effective Nuclear Charge
• Effective nuclear charge--- The net positive charge experienced by an electron in a mutielectron atom
• Important Questions!!
What happens to the effective nuclear charge as:
- we go down a group?
-we go across a period?
49
4 p+
2e-
2e-
12 p+
2e-
8e-
2e-
Be (4p+ and 4e-)
going down a group
e.g., Group IIA
Mg (12p+ and 12e-)
Ca (20p+ and 20e-)
16 p+
2e-
8e-
2e-
8e-
Effective nuclear
charge decreases
down the group!!
50
Li (3p+ and 3e-)
What happens to effective nuclear charge as we move across
the period?
Be (4p+ and 4e-) B (5p+ and 5e-)
6 p+
2e-
4e-
C (6p+ and 6e-)
8 p+
2e-
6e-
O (8p+ and 8e-)
10 p+
2e-
8e-
Ne (10p+ and 10e-)
2e-
1e-
3 p+
2e-
2e-
4 p+
2e-
3e-
5 p+
Effective nuclear charge increases across the period!!
51
Trends in Atomic Size (Atomic Radius)
• atomic radius increases down group
valence shell farther from nucleus
effective nuclear charge drops
• atomic radius decreases across period (left to right)
electrons added to same valence shell
effective nuclear charge increases
52
4 p+
2e-
2e-
12 p+
2e-
8e-
2e-
Be (4p+ and 4e-)
Group IIA
Mg (12p+ and 12e-)
Ca (20p+ and 20e-)
16 p+
2e-
8e-
2e-
8e-
53 Tro's "Introductory Chemistry",
Chapter 9
Li (3p+ and 3e-)
Period 2
Be (4p+ and 4e-) B (5p+ and 5e-)
6 p+
2e-
4e-
C (6p+ and 6e-)
8 p+
2e-
6e-
O (8p+ and 8e-)
10 p+
2e-
8e-
Ne (10p+ and 10e-)
2e-
1e-
3 p+
2e-
2e-
4 p+
2e-
3e-
5 p+
54
Trends in Atomic Size, Continued
55 Tro's "Introductory Chemistry",
Chapter 9
Example—Choose the
Larger Atom in Each Pair.
1. N or F
2. C or Ge
3. N or Al
56 Tro's "Introductory Chemistry",
Chapter 9
1. N or F
2. C or Ge
3. N or Al
1. N or F
2. C or Ge, Ge is further down
1. N or F
2. C or Ge
3. N or Al, Al is further down & left
1. N or F, N is further left
Example—Choose the
Larger Atom in Each Pair, Continued.
57
Ionization Energy
• Minimum energy needed to remove an electron
from an atom.
Gas state.
Endothermic process.
Valence electron easiest to remove.
M(g) + 1st IE M1+(g) + 1 e-
M+1(g) + 2nd IE M2+(g) + 1 e-
First ionization energy = energy to remove electron from
neutral atom; 2nd IE = energy to remove from +1 ion; etc.
58
General Trends in
Ionization Energy
• The larger the effective nuclear charge on the electron, the more energy it takes to remove it.
• The farther the electron is from the nucleus, the less energy it takes to remove it.
• First IE generally increases across the period.
effective nuclear charge increases
• First IE decreases down the group.
valence electron farther from nucleus
59 Tro's "Introductory Chemistry",
Chapter 9
Trends in Ionization Energy, Continued
60
1. Al or S
2. As or Sb, As is further up
1. Al or S
2. As or Sb
3. N or Si, N is further up
1. Al or S
2. As or Sb
3. N or Si
Example—Choose the Atom in Each Pair
with the Higher First Ionization Energy
1. Al or S, S is further right
61
Metallic Character • How well an element’s properties match the
general properties of a metal.
• Metals: Malleable and ductile as solids. shiny, lustrous, and reflect light. conduct heat and electricity. Form cations in solution. Lose electrons in reactions—oxidized.
62 Tro's "Introductory Chemistry",
Chapter 9
Metallic Character, Continued
• In general, metals are found on the left of
the periodic table and nonmetals on the
right.
• As you traverse left to right across the
period, the elements become less metallic.
• As you traverse down a column, the
elements become more metallic.
63 Tro's "Introductory Chemistry",
Chapter 9
Trends in Metallic Character
64 Tro's "Introductory Chemistry",
Chapter 9
1. Sn or Te
2. P or Sb, Sb is further down
1. Sn or Te
2. P or Sb
3. Ge or In, In is further down & left
1. Sn or Te
2. P or Sb
3. Ge or In
1. Sn or Te, Sn is further left
Example—Choose the
More Metallic Element in Each Pair
65
Recommended Study Problems Chapter 9
NB: Study problems are used to check the student’s understanding
of the lecture material. Students are EXPECTED TO BE ABLE
TO SOLVE ALL THE SUGGESTED STUDY PROBLEMS.
If you encounter any problems, please talk to your professor or seek
help at the HACC-Gettysburg learning center.
Questions from text book Chapter 9
5, 7, 13, 25, 27, 29, 33,37, 39, 43, 45, 49, 51, 55, 57, 61, 63, 65, 71,
77, 79, 85, 91, 95, 99, 101, 105, 107
ANSWERS
-The answers to the odd-numbered study problems are found at
the back of your textbook