Download Periodic Table of Elements II

Periodic Table of Elements II
Chasity F. Jeffrey, B.S. Education
Ck12 Science
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AUTHORS
Chasity F.
Education
Ck12 Science
Jeffrey,
B.S.
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Complete terms can be found at http://www.ck12.org/terms.
Printed: October 10, 2014
iii
Contents
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Contents
1
Orbitals
1
2
Valence Electrons
5
3
Electron Configurations
7
4
Hund’s Rule and Orbital Filling Diagrams
10
5
Aufbau Principle
14
6
Pauli Exclusion Principle
17
7
Electron Affinity
19
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C ONCEPT
Concept 1. Orbitals
1
Orbitals
• Draw the shapes of s, p, d, and f orbitals.
• Relate the four quantum numbers for an electron to a specific orbital.
How is it that so many planes are able to fly without running into each other?
The flight path of a commercial airliner is carefully regulated by the Federal Aviation Administration. Each airplane
must maintain a distance of five miles from another plane flying at the same altitude and 2,000 feet above and below
another aircraft (1,000 feet if the altitude is less than 29,000 feet). So, each aircraft only has certain positions it is
allowed to maintain while it flies. As we explore quantum mechanics, we see that electrons have similar restrictions
on their locations.
Orbitals
We can apply our knowledge of quantum numbers to describe the arrangement of electrons for a given atom. We do
this with something called electron configurations. They are effectively a map of the electrons for a given atom.
We look at the four quantum numbers for a given electron and then assign that electron to a specific orbital.
s Orbitals
For any value of n, a value of l = 0 places that electron in an s orbital. This orbital is spherical in shape:
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p Orbitals
From Table 1.1 we see that we can have three possible orbitals when l = 1. These are designated as p orbitals and
have dumbbell shapes. Each of the p orbitals has a different orientation in three-dimensional space.
d Orbitals
When l = 2, ml values can be -2, -1, 0, +1, +2 for a total of five d orbitals. Note that all five of the orbitals have
specific three-dimensional orientations.
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Concept 1. Orbitals
f Orbitals
The most complex set of orbitals are the f orbitals. When l = 3, ml values can be -3, -2, -1, 0, +1, +2, +3 for a total
of seven different orbital shapes. Again, note the specific orientations of the different f orbitals.
TABLE 1.1: Electron Arrangement Within Energy Levels
Principal
Quantum
Number (n)
Allowable Sublevels
Number of Orbitals per Sublevel
1
2
s
s
p
s
p
d
s
p
d
f
1
1
3
1
3
5
1
3
5
7
3
4
Number
of
Orbitals
per
Principal
Energy Level
1
4
9
16
Number
Electrons
Sublevel
of
per
2
2
6
2
6
10
2
6
10
14
Number
of
Electrons
per Principal
Energy Level
2
8
18
32
Summary
• There are four different classes of electron orbitals.
• These orbitals are determined by the value of the angular momentum quantum number l.
Practice
Questions
Use the link below to answer the following questions:
http://www.chem4kids.com/files/atom_orbital.html
1. What is a shell?
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2. What do the letters K-Q stand for?
3. How many electrons does the K shell hold?
4. What is the maximum number of electrons any shell can hold?
Review
Questions
1.
2.
3.
4.
•
•
•
•
•
What is an electron configuration?
How many electrons are in the n = 1 orbital?
What is the total number of electrons in a p orbital?
How many electrons does it take to completely fill a d orbital?
electron configuration: A map of the electrons for a given atom.
d orbital: Five orbitals characterized by l = 2.
f orbitals: The most complex set of orbitals with seven orbital shapes for l = 3.
p orbitals: Three dumbbell shaped orbitals for l = 1.
s orbitals: Spherical orbitals where l = 0.
References
1. User:Baseball Bugs/Wikimedia Commons. http://commons.wikimedia.org/wiki/File:NWA_Plane_over_La
ke_Harriet_020623.JPG .
2. CK-12 Foundation - Christopher Auyeung. .
3. CK-12 Foundation - Christopher Auyeung. .
4. CK-12 Foundation - Christopher Auyeung. .
5. CK-12 Foundation - Richard Parsons. .
4
www.ck12.org
Concept 2. Valence Electrons
C ONCEPT
2
Valence Electrons
• Define valence electron.
• Be able to indicate valence electrons when given the electron configuration for an atom.
What makes a particular element very reactive and another element non-reactive?
A chemical reaction involves either electron removal, electron addition, or electron sharing. The path a specific
element will take depends on where the electrons are in the atom and how many there are.
TABLE 2.1: Electron Configurations of Second-Period Elements
Element Name
Lithium
Beryllium
Boron
Carbon
Nitrogen
Oxygen
Fluorine
Neon
Symbol
Li
Be
B
C
N
O
F
Ne
Atomic Number
3
4
5
6
7
8
9
10
Electron Configuration
1s2 2s1
1s2 2s2
1s2 2s2 2p1
1s2 2s2 2p2
1s2 2s2 2p3
1s2 2s2 2p4
1s2 2s2 2p5
1s2 2s2 2p6
In the study of chemical reactivity, we will find that the electrons in the outermost principal energy level are very
important and so they are given a special name. Valence electrons are the electrons in the highest occupied principal
energy level of an atom. In the second period elements listed above, the two electrons in the 1s sublevel are called
inner-shell electrons and are not involved directly in the element’s reactivity or in the formation of compounds.
Lithium has a single electron in the second principal energy level and so we say that lithium has one valence electron.
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Beryllium has two valence electrons. How many valence electrons does boron have? You must recognize that the
second principal energy level consists of both the 2s and the 2p sublevels and so the answer is three. In fact, the
number of valence electrons goes up by one for each step across a period until the last element is reached. Neon,
with its configuration ending in s2 p6 , has eight valence electrons.
Summary
• Valence electrons are the outer-shell electrons of an atom.
• Valence electrons determine the reactivity of an atom.
Practice
Use the link below to answer questions about valence electrons:
http://www.sciencegeek.net/Chemistry/taters/Unit3ValenceElectrons.htm
Review
Questions
1.
2.
3.
4.
5.
Define valence electron.
Define inner shell electron.
How many valence electrons are there in fluorine?
What are the 2s electrons in nitrogen?
How many inner shell electrons are there in beryllium?
• inner-shell electrons: Those electrons that are not in the outer shell and are not involved in the reactivity of
the element.
• valence electrons: The electrons in the highest occupied principal energy level of an atom.
References
1. User:Chemicalinterest/Wikipedia. http://commons.wikimedia.org/wiki/File:Cobalt_carbonate.JPG .
6
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C ONCEPT
Concept 3. Electron Configurations
3
Electron Configurations
• Use electron configuration notation to indicate the electron configuration of an atom.
How big is a file?
If you keep your papers in manila folders, you can pick up a folder and see how much it weighs. If you want to
know how many different papers (articles, bank records, or whatever else you keep in a folder), you have to take
everything out and count. A computer directory, on the other hand, tells you exactly how much you have in each
file. We can get the same information on atoms. If we use an orbital filling diagram, we have to count arrows. When
we look at electron configuration data, we simply add up the numbers.
Electron Configurations
Electron configuration notation eliminates the boxes and arrows of orbital filling diagrams. Each occupied sublevel
designation is written followed by a superscript that is the number of electrons in that sublevel. For example, the
hydrogen configuration is 1s1 , while the helium configuration is 1s2 . Multiple occupied sublevels are written one
after another. The electron configuration of lithium is 1s2 2s1 . The sum of the superscripts in an electron configuration
is equal to the number of electrons in that atom, which is in turn equal to its atomic number.
Sample Problem: Orbital Filling Diagrams and Electron Configurations
Draw the orbital filling diagram for carbon and write its electron configuration.
Step 1: List the known quantities and plan the problem.
Known
• atomic number of carbon, Z = 6
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FIGURE 3.1
Orbital filling diagram for carbon.
Use the order of fill diagram to draw an orbital filling diagram with a total of six electrons. Follow Hund’s rule.
Write the electron configuration.
Step 2: Construct Diagram
Electron configuration 1s2 2s2 2p2
Step 3: Think about your result.
Following the 2s sublevel is the 2p, and p sublevels always consist of three orbitals. All three orbitals need to be
drawn even if one or more is unoccupied. According to Hund’s rule, the sixth electron enters the second of those p
orbitals and with the same spin as the fifth electron.
Second Period Elements
Periods refer to the horizontal rows of the periodic table. Looking at a periodic table you will see that the first period
contains only the elements hydrogen and helium. This is because the first principal energy level consists of only
the s sublevel and so only two electrons are required in order to fill the entire principal energy level. Each time a
new principal energy level begins, as with the third element lithium, a new period is started on the periodic table.
As one moves across the second period, electrons are successively added. With beryllium (Z = 4), the 2s sublevel
is complete and the 2p sublevel begins with boron (Z = 5). Since there are three 2p orbitals and each orbital holds
two electrons, the 2p sublevel is filled after six elements. The Table 3.1 shows the electron configurations of the
elements in the second period.
TABLE 3.1: Electron Configurations of Second-Period Elements
Element Name
Lithium
Beryllium
Boron
Carbon
Nitrogen
Oxygen
Fluorine
Neon
Symbol
Li
Be
B
C
N
O
F
Ne
Atomic Number
3
4
5
6
7
8
9
10
Electron Configuration
1s2 2s1
1s2 2s2
1s2 2s2 2p1
1s2 2s2 2p2
1s2 2s2 2p3
1s2 2s2 2p4
1s2 2s2 2p5
1s2 2s2 2p6
Summary
• Electron configuration notation simplifies the indication of where electrons are located in a specific atom.
• Superscripts are used to indicate the number of electrons in a given sublevel.
Practice
Use the link below to practice solving electron configuration problems.
http://www.sciencegeek.net/Chemistry/taters/Unit2ElectronNotations.htm
8Review
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Concept 3. Electron Configurations
2. How do we know how many electrons are in each sublevel?
3. An atom has the electron configuration of 1s2 2s2 2p5 . How many electrons are in that atom?
4. Which element has the electron configuration of 1s2 2s2 2p6 3s2 ?
• electron configuration notation: Each occupied sublevel designation is written followed by a superscript
that is the number of electrons in that sublevel.
References
1. David R. Tribble (Wikimedia: Loadmaster). http://commons.wikimedia.org/wiki/File:DirectoryListing1.svg
.
2. CK-12 Foundation - Joy Sheng. .
9
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C ONCEPT
4
Hund’s Rule and Orbital
Filling Diagrams
• State Hund’s rule.
• Apply Hund’s rule to the filling of orbitals.
• Use orbital filling diagrams to describe the locations of electrons in an atom.
Have you ever wondered what those load limit signs mean on a bridge?
The sign above says that nothing over five tons is allowed because it will do damage to the structure. There are limits
to the amount of weight that a bridge can support, there are limits to the number of people that can safely occupy a
room, and there are limits to what can go into an electron orbital.
Hund’s Rule
The last of the three rules for constructing electron arrangements requires electrons to be placed one at a time in a set
of orbitals within the same sublevel. This minimizes the natural repulsive forces that one electron has for another.
Hund’s rule states that orbitals of equal energy are each occupied by one electron before any orbital is occupied by
a second electron and that each of the single electrons must have the same spin. The Figure 4.1 shows how a set of
three p orbitals is filled with one, two, three, and four electrons.
Orbital Filling Diagrams
An orbital filling diagram is the more visual way to represent the arrangement of all the electrons in a particular
atom. In an orbital filling diagram, the individual orbitals are shown as circles (or squares) and orbitals within a
sublevel are drawn next to each other horizontally. Each sublevel is labeled by its principal energy level and sublevel.
Electrons are indicated by arrows inside the circles. An arrow pointing upwards indicates one spin direction, while
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Concept 4. Hund’s Rule and Orbital Filling Diagrams
FIGURE 4.1
The 2p sublevel, for the elements boron
(Z = 5), carbon (Z = 6), nitrogen (Z = 7),
and oxygen (Z = 8). According to Hund’s
rule, as electrons are added to a set of orbitals of equal energy, one electron enters
each orbital before any orbital receives a
second electron.
a downward pointing arrow indicates the other direction. The orbital filling diagrams for hydrogen, helium, and
lithium are shown in Figure 4.2.
FIGURE 4.2
Orbital filling diagrams for hydrogen, helium, and lithium.
According to the Aufbau process, sublevels and orbitals are filled with electrons in order of increasing energy. Since
the s sublevel consists of just one orbital, the second electron simply pairs up with the first electron as in helium. The
next element is lithium and necessitates the use of the next available sublevel, the 2s.
The filling diagram for carbon is shown in the Figure 4.3. There are two 2p electrons for carbon and each occupies
its own 2p orbital.
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FIGURE 4.3
Orbital filling diagram for carbon.
Oxygen has four 2p electrons. After each 2p orbital has one electron in it, the fourth electron can be placed in the
first 2p orbital with a spin opposite that of the other electron in that orbital.
FIGURE 4.4
Orbital filling diagram for oxygen.
Summary
• Hund’s rule specifies the order of electron filling within a set of orbitals.
• Orbital filling diagrams are a way of indicating electron locations in orbitals.
Practice
Questions
Use the link below to carry out the following exercise:
https://www.caymanchem.com/app/template/chemAssistant,Tool.vm/itemid/4001
1. Select an atom from the list (you will probably want to do lower atomic numbers). Leave the number set a
zero.
2. Look up the atom on a periodic table and determine the number of electrons present.
3. Draw the orbital filling diagram for the atom.
4. Click on the “Calculate” button and compare your answer with the one provided.
Review
Questions
1. State Hund’s rule.
2. What is an orbital filling diagram?
3. Is the diagram in the Figure 4.5 correct? Explain your answer.
4. Is the diagram in the Figure 4.6 correct? Explain your answer.
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Concept 4. Hund’s Rule and Orbital Filling Diagrams
FIGURE 4.5
FIGURE 4.6
• Hund’s rule: Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a
second electron and that each of the single electrons must have the same spin.
• orbital filling diagram: A visual way to represent the arrangement of all the electrons in a particular atom.
References
1.
2.
3.
4.
5.
6.
7.
Laura Guerin. CK-12 Foundation .
CK-12 Foundation - Joy Sheng. .
CK-12 Foundation - Joy Sheng. .
CK-12 Foundation - Joy Sheng. .
CK-12 Foundation - Joy Sheng. .
CK-12 Foundation - Joy Sheng. .
CK-12 Foundation - Joy Sheng. .
13
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C ONCEPT
5
Aufbau Principle
• State the Aufbau principle.
• Use the Aufbau principle to determine the electron configuration of an atom.
How are buildings constructed?
Construction of a building begins at the bottom. The foundation is laid and the building goes up step by step. You
obviously cannot start with the roof since there is no place to hang it. The building goes from the lowest level to the
highest level in a systematic way.
Aufbau Principle
In order to create ground state electron configurations for any element, it is necessary to know the way in which the
atomic sublevels are organized in order of increasing energy. The Figure 5.1 shows the order of increasing energy
of the sublevels.
The lowest energy sublevel is always the 1s sublevel, which consists of one orbital. The single electron of the
hydrogen atom will occupy the 1s orbital when the atom is in its ground state. As we proceed with atoms with
multiple electrons, those electrons are added to the next lowest sublevel: 2s, 2p, 3s, and so on. The Aufbau principle
states that an electron occupies orbitals in order from lowest energy to highest. The Aufbau (German: “building up,
construction”) principle is sometimes referred to as the “building-up” principle. It is worth noting that in reality
atoms are not built by adding protons and electrons one at a time and that this method is merely an aid for us to
understand the end result.
As seen in the Figure 5.1, the energies of the sublevels in different principal energy levels eventually begin to
overlap. After the 3p sublevel, it would seem logical that the 3d sublevel should be the next lowest in energy.
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Concept 5. Aufbau Principle
FIGURE 5.1
Electrons are added to atomic orbitals in
order from low energy (bottom of graph)
to high (top of graph) according to the Aufbau principle. Principal energy levels are
color coded, while sublevels are grouped
together and each circle represents an
orbital capable of holding two electrons.
However, the 4s sublevel is slightly lower in energy than the 3d sublevel and thus fills first. Following the filling of
the 3d sublevel is the 4p, then the 5s and the 4d. Note that the 4f sublevel does not fill until just after the 6s sublevel.
The Figure 5.2 is a useful and simple aid for keeping track of the order of fill of the atomic sublevels.
Summary
• The Aufbau principle gives the order of electron filling in an atom.
• It can be used to describe the locations and energy levels of every electron in a given atom.
Practice
Questions
Use the link below to answer the following questions:
http://ths.talawanda.net/~BrambleN/classroom/Chemistry/Notes/Section%202A/Exceptions&Shortcut.htm
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FIGURE 5.2
The Aufbau principle is illustrated in the
diagram by following each red arrow in
order from top to bottom: 1s, 2s, 2p, 3s,
etc.
1. Do the electrons in all atoms follow the Aufbau rule?
2. What happens to electrons in copper to make the atom more stable?
3. How does silver become more stable?
Review
Questions
1.
2.
3.
4.
What is the Aufbau principle?
Which orbital is filled after the 2p?
Which orbital is filled after 4s?
Which orbital is filled after 6s?
• Aufbau principle: An electron occupies orbitals in order from lowest energy to highest.
References
1. Gary Minnaert. http://commons.wikimedia.org/wiki/File:LACMA_BCAM02.jpg .
2. CK-12 Foundation - Christopher Auyeung. .
3. CK-12 Foundation - Christopher Auyeung. .
16
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C ONCEPT
Concept 6. Pauli Exclusion Principle
6
Pauli Exclusion Principle
• State the Pauli exclusion principle.
Can you name one thing that easily distinguishes you from the rest of the world?
And we’re not talking about DNA – that’s a little expensive to sequence. For many people, it is their email address.
My email address allows people all over the world to contact me. It does not belong to anyone else, but serves to
identify me. Electrons also have a unique set of identifiers in the quantum numbers that describe their location and
spin.
Pauli Exclusion Principle
When we look at the orbital possibilities for a given atom, we see that there are different arrangements of electrons
for each different type of atom. Since each electron must maintain its unique identity, we intuitively sense that the
four quantum numbers for any given electron must not match up exactly with the four quantum numbers for any
other electron in that atom.
For the hydrogen atom, there is no problem since there is only one electron in the H atom. However, when we
get to helium we see that the first three quantum numbers for the two electrons are the same: same energy level,
same spherical shape. What differentiates the two helium electrons is their spin. One of the electrons has a + 21 spin
while the other electron has a − 12 spin. So the two electrons in the 1s orbital are each unique and distinct from one
another because their spins are different. This observation leads to the Pauli exclusion principle, which states that
no two electrons in an atom can have the same set of four quantum numbers. The energy of the electron is specified
by the principal, angular momentum, and magnetic quantum numbers. If those three numbers are identical for two
electrons, the spin numbers must be different in order for the two electrons to be differentiated from one another.
The two values of the spin quantum number allow each orbital to hold two electrons. The figure below shows how
the electrons are indicated in a diagram.
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FIGURE 6.1
In an orbital filling diagram, a square represents an orbital, while arrows
represent electrons.
An arrow pointing upward represents one spin
direction, while an arrow pointing downward represents the other spin
direction.
Summary
• The Pauli exclusion principle specifies limits on how identical quantum numbers can be for two electrons in
the same atom.
Practice
Questions
Use the link below to answer the following questions:
http://www.nobelprize.org/nobel_prizes/physics/laureates/1945/pauli-bio.html
1.
2.
3.
4.
When was Pauli born?
How old was he when he received his doctorate?
What was he the first to recognize?
When did he win his Nobel Prize in Physics?
Review
Questions
1. What is the difference between the two helium electrons?
2. What does the Pauli exclusion principle state?
3. What does the two values for the spin quantum number allow?
• Pauli exclusion principle: No two electrons in an atom can have the same set of four quantum numbers.
References
1. Pixabay:Nemo. http://pixabay.com/en/mail-internet-icon-electronic-35636/ .
2. CK-12 Foundation. .
18
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C ONCEPT
Concept 7. Electron Affinity
7
Electron Affinity
• Define electron affinity.
• Describe trends in electron affinity in the periodic table.
Do you tend to overpack before going on trips?
Packing for a trip can be very challenging. What do you take with you? Where will you be going and what will
you need? We usually pack too much (like the suitcase above) and then find it hard to close the suitcase. When the
suitcase is over-full, there is stress on the system and forces pushing the suitcase open. When electrons are added
to an atom, the increased negative charge puts stress on the electrons already there, causing energy to be released.
When electrons are removed from an atom, that process requires energy to pull the electron away from the nucleus. Addition of an electron releases energy from the process.
Electron Affinity
In most cases, the formation of an anion by the addition of an electron to a neutral atom releases energy. This can be
shown for the chloride ion formation below:
Cl + e− → Cl− + energy
The energy change that occurs when a neutral atom gains an electron is called its electron affinity. When energy
is released in a chemical reaction or process, that energy is expressed as a negative number. The figure below
shows electron affinities in kJ/mole for the representative elements. Electron affinities are measured on atoms in the
gaseous state and are very difficult to measure accurately.
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FIGURE 7.1
Electron affinities (in kJ/mol) of the first
five periods of the representative elements.
Electron affinities are negative
numbers because energy is released.
The elements of the halogen group (Group 17) gain electrons most readily, as can be seen from their large negative
electron affinities. This means that more energy is released in the formation of a halide ion than for the anions of any
other elements. Considering electron configuration, it is easy to see why. The outer configuration of all halogens
is ns2 np5 . The addition of one more electron gives the halide ions the same electron configuration as a noble gas,
which we have seen is particularly stable.
Period and group trends for electron affinities are not nearly as regular as for ionization energy. In general, electron
affinities increase (become more negative) from left to right across a period and decrease (become less negative)
from top to bottom down a group. However, there are many exceptions, owing in part to inherent difficulties in
accurately measuring electron affinities.
Summary
•
•
•
•
Electron affinity is a measure of the energy released when an extra electron is added to an atom.
Electron affinities are measured in the gaseous state.
In general, electron affinities become more negative as we move from left to right on the periodic table.
In general, electron affinities become less negative from top to bottom of a group.
Practice
Questions
Use the link below to answer the following questions:
http://www.chemguide.co.uk/atoms/properties/eas.html
1. Which groups (using the old Roman numeral system) of elements are primarily involved with issues of electron
affinity?
2. What does a negative energy imply?
3. Why is the electron affinity value for fluorine less than that of chlorine?
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Concept 7. Electron Affinity
Review
Questions
1.
2.
3.
4.
Define “electron affinity.”
Does addition of an electron to a neutral atom require energy or release energy?
Describe the general trend for electron affinity values moving from left to right on the periodic table.
Describe the general trend for electron affinity values moving from top to bottom in a group on the periodic
table.
5. Why is more energy released in the formation of a halide ion than with other elements?
• electron affinity: The energy change that occurs when a neutral atom gains an electron.
References
1. Image copyright Africa Studio, 2014. Suitcase .
2. CK-12 Foundation - Christopher Auyeung. .
21