Download Periodic Trends - Harrison High School

Periodic Trends
• Introduction
• If you were to look carefully at many of the properties
of the elements, you would notice something besides
the similarity of the properties within the groups. You
would notice that many of these properties change in a
fairly regular fashion that is dependent on the position
of the element in the periodic table. And of course,
that is what you will do next. As you compare elements
from left to right across the periodic table, you will
notice a trend or regular change in a number of
properties. The same thing happens if you go up and
down on the periodic table and compare the
properties of the elements.
Atomic Mass
250
200
150
Period 7
100
Period 6
period 5
50
Period 4
Period 3
2
3
4
5
Period 7
Period 6
6
period 5
7
Period 4
Period 3
8
Period 2
Period 1
10
Period 2
Period 1
• The first of these properties is the atomic size. You know that each
atom has a nucleus inside and electrons zooming around outside
the nucleus. It should seem reasonable that the size of an atom
depends on how far away its outermost (valence) electrons are
from the nucleus. If they are very close to the nucleus, the atom
will be very small. If they are far away, the atom will be quite a bit
larger. So the atomic size is determined by how much space the
electrons take up.
• Measuring the size of atoms is, in some ways, like measuring the
size of cotton balls or automobile tires. The value you get depends
on the conditions under which they are measured. A "free" cotton
ball has a different size than when it is in the package. The radius of
the tire is different when measured to the top of the tire than when
measured to the bottom of the tire resting on the ground. Different
values for the sizes of atoms are obtained depending on both the
method used and the conditions in which the atoms find itself - free
or bonded to other atoms. The following table gives a variety of
values collected from a variety of sources. Whichever set of values
you choose to use, note the trends.
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Let's make some comparisons in a family and in a period. In a family--like from hydrogen to lithium
to sodium on down--the atomic size increases. As you go down a group, the size increases. As you
go across a period, as from lithium to neon, notice that the size decreases. You need to remember
(or memorize) those trends.
Now let's talk about why that's the case and relate it back to the various factors presented earlier.
Remember that the nuclear charge and the shielding electrons combine to make the effective
nuclear charge. That is a very important factor when you are comparing elements in a period. As
you go across a period, the nuclear charge increases and the number of energy levels stays the
same. Consequently, the number of shielding electrons stays the same and the effective nuclear
charge increases. As the effective nuclear charge increases, it pulls the electrons in closer and
closer to the nucleus. So as you go across a period, the increase in the nuclear charge causes a
decrease in the atomic size because the electrons in the valence energy level are pulled in closer
and closer.
Now let's make comparison within a family such as hydrogen down to francium (Fr). It is true that
the nuclear charge is increasing, but so is the number of shielding electrons. The number of
shielding electrons increases by the same amount that the nuclear charge increases. So the
effective nuclear charge felt by the valence electrons stays the same. There is no increase in the
effective nuclear charge but there is an increase in the number of energy levels that are being used.
Consequently, the electrons in the valence energy level are further and further away from the
nucleus because they are in higher energy levels. Consequently, the important factor in a vertical
comparison on the periodic table is the number of energy levels that are being used because the
increase in the number of shielding electrons cancels out the increase in the nuclear charge.
To summarize, as you go across a period, the increase in the nuclear charge is the most important
factor because the number of energy levels stays the same. As you go down a group, the increase in
shielding electrons more or less cancels out the increase in nuclear charge, leaving the increase in
the number of energy levels as the most important factor. This is true not only for atomic size but
for other properties as well.
Atomic Radius
Atomic radius => distance from the center of an atom's nucleus to its outer
most electron
• Atomic Radii
• 1) As you move down a group, atomic radius increases.
• WHY? - The number of energy levels increases as you move down a
group as the number of electrons increases. Each subsequent
energy level is further from the nucleus than the last. Therefore,
the atomic radius increases as the group and energy levels
increase.
• 2) As you move across a period, atomic radius decreases.
• WHY? - As you go across a period, electrons are added to the same
energy level. At the same time, protons are being added to the
nucleus. The concentration of more protons in the nucleus creates
a "higher effective nuclear charge." In other words, there is a
stronger force of attraction pulling the electrons closer to the
nucleus resulting in a smaller atomic radius.
Atomic Radius
Atomic radius => distance from the center of an atom's nucleus to its
outer most electron
Common Student Misconceptions
1. “The radius of an atom increases in a period with the addition of
electrons.”
From the data given on atomic radii, we see that with the addition of
electrons, the atomic radii decrease within a period. This is due to the
increase in the positive charge of the nucleus and its attraction for the
electrons in the atom. Although the number of electrons in the atom
increases as one goes across a period and although the electrons repel each
other potentially causing the atoms to grow in size, this effect is more than
opposed by the increase in the number of protons attracting these electrons.
Since these electrons are added to the same shell, the number of protons
attracting them causes the size of the atoms to decrease.
2. “All properties show periodic characteristics.”
There are deviations in all periods.
3. “Elements are arranged according to atomic weight.”
While this is true in general, the periodic properties are characteristic of
atomic numbers.
Practice with Comparing Atomic Size
• Now try your hand at answering the following
questions. Check your answers below and then
continue with the lesson.
• For each of the following sets of atoms, decide
which is larger, which is smaller, and why.
• a. Li, C, F
• b. Li, Na, K
• c. Ge, P, O
• d. C, N, Si
• e. Al, Cl, Br
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Answers for Comparing Atomic Sizes
Here are answers for the questions above.
a. Li, C, F
All are in the same period and thus have the same number of energy levels. Therefore, the important
factor is the nuclear charge. Li is the largest because it has the smallest nuclear charge and pulls the
electrons toward the nucleus less than the others. F is the smallest because it has the largest nuclear
charge and pulls the electrons toward the nucleus more than the others.
b. Li, Na, K
All are in the same group and thus have the same effective nuclear charge. Therefore, the important
factor is the number of energy levels. Li is the smallest because it uses the smallest number of electron
energy levels. K is the largest because it uses the largest number of electron energy levels.
c. Ge, P, O
All are in different groups and periods, therefore both factors must be taken into account.
Fortunately both factors reinforce one another. Ge is the largest because it uses the largest number of
energy levels and has the smallest effective nuclear charge. O is the smallest because it uses the
smallest number of energy levels and has the largest effective nuclear charge.
d. C, N, Si
Not all are in the same group and period, so, again, both factors must be taken into account. C and N
tie for using the smallest number of energy levels, but N has a higher effective nuclear charge.
Therefore, N is the smallest. C and Si tie for having the lowest effective nuclear charge, but Si uses more
energy levels. Therefore, Si is the largest.
e. Al, Cl, Br
Not all are in the same group and period, so, again, both factors must be taken into account. Cl is the
smallest because it has higher effective nuclear charge than Al and uses fewer energy levels than Br.
Which is largest is less straightforward. Al has a lower effective nuclear charge (by four), but Br uses
more energy levels (by one). Because the difference in effective nuclear charge is larger, it should be
the more important factor in this case, making Al the largest.
Al and Br can also be compared to one another indirectly by comparing both to Cl. Both Al and Br are
larger than Cl. Al is larger than Cl because it has lower effective nuclear charge (by four). Br is larger
than Cl because it uses more energy levels (by one). Because Al is larger than Cl by four "steps" and Br
is larger than Cl by only one "step", Al is likely the largest of the three.
Atomic Radius
Atomic radius => distance from the center of an atom's nucleus to its outer
most electron
Atomic Radius
Atomic Radius
Atomic Radius
4
3
2
1
1
2
3
4
5
6
7
8
0
Period 7
Period 6
Period 5
Period 4
Period 3
Period 2
Period 1
Atomic Radius
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ATOMIC RADIUS
Measures of atomic radius
Unlike a ball, an atom doesn't have a fixed radius. The radius of an atom
can only be found by measuring the distance between the nuclei of two
touching atoms, and then halving that distance.
As you can see from the diagrams, the same atom could be found to have
a different radius depending on what was around it.
The left hand diagram shows bonded atoms. The atoms are pulled closely
together and so the measured radius is less than if they are just touching.
This is what you would get if you had metal atoms in a metallic structure,
or atoms covalently bonded to each other. The type of atomic radius
being measured here is called the metallic
radius or the covalent
radius depending on the bonding.
Ionic Radius
Ionic Radius Isoelectric Ions
Ionic Charges
Group Ionic Charges
Group Ionic Charges
Atomic Radius and Ionic Radius
Ionic radius
• When atoms gain or lose electrons, the atom becomes
an ion. When an atom gains an electron, it becomes a
negatively charged ion that we call an anion. Anions
are larger in size than their parent atoms because they
have one or more additional electrons, but without an
additional proton in the nucleus to help moderate the
size.
• When an atom loses an electron, it becomes a
positively charged ion called a cation. Cations are
smaller than their parent atoms because they have lost
electrons (sometimes the entire outermost energy
level) and the electrons that remain behind simply
don't take up as much room.
Ionic Radius
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Ionic Radii
1) Anions (negative ions) are larger than their respective atoms.
WHY?
Electron-electron repulsion forces them to spread further apart.
Electrons outnumber protons; the protons cannot pull the extra electrons
as tightly toward the nucleus.
• 2) Cations (positive ions) are smaller than their respective atoms.
• WHY?
• There is less electron-electron repulsion, so they can come closer
together.
Protons outnumber electrons; the protons can pull the fewer electrons
toward the nucleus more tightly.
If the electron that is lost is the only valence electron so that the electron
configuration of the cation is like that of a noble gas, then an entire energy
level is lost. In this case, the radius of the cation is much smaller than its
respective atom.
Practice Comparing Ionic Sizes
• Try your hand at the making the following comparisons
(also shown in exercise 9 in your workbook), based on your
understanding of ionic size comparisons and without
reference to the wall chart, except to check your answers.
Answers also follow on this page.
• For each of the following sets of atoms and ions, decide
which is the smallest and which is the largest.
• a. Na, Na+
• b. Cl, Cl• c. Na+, Cl• d. H+, H, H• e. Fe2+, Fe3+
Answers for Comparing Ionic Sizes
• Here are the answers to the questions above. For
each of the following sets of atoms and ions,
decide which is the smallest and which is the
largest.
• a. Na is largest, Na+ is smallest.
• b. Cl is smallest, Cl- is largest.
• c. Na+ is smallest, Cl- is largest.
• d. H+ is smallest, H, H- is largest.
• e. Fe2+ is largest, Fe3+ is smallest
Electronegativity
• Electronegativity—the strength with which an
atom pulls on the electrons it shares in a bond
(the attractions that hold molecules together)
is its electronegativity. The trend is the same
as the electron affinity trends.
Electronegativity
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ELECTRONEGATIVITY
This page explains what electronegativity is, and how and why it varies
around the Periodic Table. It looks at the way that electronegativity
differences affect bond type and explains what is meant by polar bonds
and polar molecules.
If you are interested in electronegativity in an organic chemistry context,
you will find a link at the bottom of this page.
What is electronegativity
Definition
Electronegativity is a measure of the tendency of an atom to attract a
bonding pair of electrons.
The Pauling scale is the most commonly used. Fluorine (the most
electronegative element) is assigned a value of 4.0, and values range
down to caesium and francium which are the least electronegative at 0.7.
Electronegativity Chart
Period 7
Period 6
Period 1
Period 2
4
Period 5
Period 4
Period 3
Period 3
3.5
3
Period 4
2.5
Period 5
2
1.5
Period 6
1
0.5
Period 7
0
1
2
3
4
5
6
7
8
Period 2
Period 1
• Trends in electronegativity across a period
• As you go across a period the electronegativity increases. The chart
shows
• electronegativities from sodium to chlorine - you have to ignore
argon. It
• doesn't have an electronegativity, because it doesn't form bonds.
Trends in electronegativity down a group
• As you go down a group, electronegativity decreases. (If it increases
up to
• fluorine, it must decrease as you go down.) The chart shows the
patterns
• of electronegativity in Groups 1 and 7.
Explaining the patterns in
electronegativity
• The attraction that a bonding pair of electrons
feels for a particular nucleus depends on:
• ● the number of protons in the nucleus;
• ● the distance from the nucleus;
• ● the amount of screening by inner electrons.
Why does electronegativity increase
across a period?
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Consider sodium at the beginning of period 3 and chlorine at the end
(ignoring the noble gas, argon). Think of sodium chloride as if it were
covalently bonded.
Both sodium and chlorine have their bonding electrons in the 3-level. The
electron pair is screened from both nuclei by the 1s, 2s and 2p electrons,
but the chlorine nucleus has 6 more protons in it. It is no wonder the
electron pair gets dragged so far towards the chlorine that ions are
formed.
Electronegativity increases across a period because the number of charges
on the nucleus increases. That attracts the bonding pair of electrons more
strongly.
Electronegativity
Electronegativity
Ionization energy
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Now on to another property. It's called ionization energy. It can be defined as being the energy required to remove the
outermost electron from a gaseous atom. A "gaseous atom" means an atom that is all by itself, not hooked up to others in a
solid or a liquid. When enough energy is added to an atom the outermost electron can use that energy to pull away from the
nucleus completely (or be pulled, if you want to put it that way), leaving behind a positively charged ion. That is why it's
called ionization, one of the things formed in the process is an ion. The ionization energy is the exact quantity of energy that
it takes to remove the outermost electron from the atom.
In your lab work on atomic spectra you observed that a gas would conduct electricity and emit light when it was subjected
to a high voltage. When there is little or no voltage applied to the gas in the tubes, no light is emitted and the gas does not
conduct electricity. One method for measuring the ionization energy of a gas is to slowly increase the voltage applied to it
until it does conduct electricity and emit light. The voltage at which that occurs can be used to calculate the ionization
energy.
If the ionization energy is high, that means it takes a lot of energy to remove the outermost electron. If the ionization energy
is low, that means it takes only a small amount of energy to remove the outermost electron.
Let’s use your understanding of atomic structure to make some predictions. Think for a minute about how ionization energy
would be affected by three of the factors we were talking about earlier: (1) nuclear charge, (2) number of energy levels, and
(3) shielding.
As the nuclear charge increases, the attraction between the nucleus and the electrons increases and it requires more energy
to remove the outermost electron and that means there is a higher ionization energy. As you go across the periodic table,
nuclear charge is the most important consideration. So, going across the periodic table, there should be an increase in
ionization energy because of the increasing nuclear charge.
Going down the table, the effect of increased nuclear charge is balanced by the effect of increased shielding, and the
number of energy levels becomes the predominant factor. With more energy levels, the outermost electrons (the valence
electrons) are further from the nucleus and are not so strongly attracted to the nucleus. Thus the ionization energy of the
elements decreases as you go down the periodic table because it is easier to remove the electrons. Another way of looking
at that is that if you are trying to take something from the first energy level, you have to take it past the second, the third,
the fourth and so on, on the way out. But if something is already in the third or fourth energy level, it doesn't have to be
taken as far to get away from the nucleus. It is already part way removed from the nucleus.
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The periodic nature of ionization energy is emphasized in this diagram. With each new period the
ionization energy starts with a low value. Within each period you will notice that the pattern is
really kind of a zigzag pattern progressing up as you go across the periodic table. The zigs and zags
on that graph correspond to the sublevels in the energy levels. So far in this lesson we have
presumed that all the electrons in the second energy level are pretty much the same. Two factors
make that not completely true. One factor is that because s and p orbitals have different shapes,
the electrons in p orbitals have more energy and are further from the nucleus. The other factor is
that when electrons are paired up in an orbital, they repel one another somewhat. Those two
factors account for the zigzag nature of the increase in ionization energy. Nevertheless, as a general
trend, from left to right across the periodic table, ionization energy does increase. Also as you go
down the periodic table, the ionization energy does decrease for the reasons given.
Note that the trends in the periodic properties of atomic size and ionization energy are related.
Going across the periodic table from left to right, the electrons are more tightly held by the
nucleus, causing the atoms to be smaller and the ionization energy to be higher. As you go down
the periodic table, the electrons are further from the nucleus, causing the atoms to be larger and
the ionization energies to be lower.
Ionization Energy
Ionization Energy
Ionization energy
25
20
Period 7
Period 6
15
Period 5
Period 4
10
Period 3
Period 1
Period 2
Period 3
Period 4
5
0
1
Period 5
2
3
Period 6
4
5
6
Period 7
7
8
Period 2
Period 1
Practice with Comparing Ionization
Energies
• Please take some time now to do the following exercises
(also shown in example 7 in your workbook). When you
have done that, check your answers below and then
continue.
• For each of the following sets of atoms, decide which has
the highest and lowest ionization energies and why.
• a. Mg, Si, S
• b. Mg, Ca, Ba
• c. F, Cl, Br
• d. Ba, Cu, Ne
• e. Si, P, N
Answers to Comparing Ionization
Energies
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Here are answers to the exercises above.
a. Mg, Si, S
All are in the same period and use the same number of energy levels. Mg has the lowest I.E. because it
has the lowest effective nuclear charge. S has the highest I.E. because it has the highest effective nuclear
charge.
b. Mg, Ca, Ba
All are in the same group and have the same effective nuclear charge. Mg has the highest I.E. because it
uses the smallest number of energy levels. Ba has the lowest I.E. because it uses the largest number of
energy levels.
c. F, Cl, Br
All are in the same group and have the same effective nuclear charge. F has the highest I.E. because it
uses the smallest number of energy levels. Br has the lowest I.E. because it uses the largest number of
energy levels.
d. Ba, Cu, Ne
All are in different groups and periods, so both factors must be considered. Fortunately both factors
reinforce one another. Ba has the lowest I.E. because it has the lowest effective nuclear charge and uses
the highest number of energy levels. Ne has the highest I.E. because it has the highest effective nuclear
charge and uses the lowest number of energy levels.
e. Si, P, N
Si has the lowest I.E. because it has the lowest effective nuclear charge and is tied (with P) for using the
most energy levels. N has the highest I.E. because it uses the fewest energy levels and is tied (with P) for
having the highest effective nuclear charge.
Ionization energy and the Loss of
electrons
• Tendency to Lose Electrons
• Ionization energy measures how difficult it is
for atoms to lose electrons but quite often we
will want to talk about how easy it is for atoms
to lose electrons. A low ionization energy
means that it is easy for an atom to lose
electrons. A high ionization energy means that
it is hard for an atom to lose electrons.
Ionization energy and gaining
electrons
• Tendency to Gain Electrons
• Next let's consider the opposite of losing electrons (
ionization of atoms), that is the gaining of electrons.
Atoms can attract additional electrons if there is room
for them in the valence energy level. When an extra
electron moves into the valence shell, it can feel the
attraction exerted by the effective nuclear charge.
Because the effective nuclear charge is largest for the
elements on the right side of the periodic table, those
atoms provide the greatest attraction for electrons and
have the greatest tendency to gain electrons.
IE affect on Gaining Electrons
• Practice Comparing Tendencies to Gain Electrons
• For each of the following sets of atoms, decide which
has the least and which has the greatest tendency to
gain electrons and why. Check your answers below and
then continue with the lesson.
• a. Li, C, N
• b. C, O, Ne
• c. Si, P, O
• d. K, Mg, P
• e. S, F, He
Answers for Comparing Tendencies to
Gain Electrons
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Here are answers to the exercises above.
a. Li, C, N
Li has the least tendency to gain electrons because it has the lowest effective nuclear charge (and all use the
same number of energy levels). N has the greatest tendency to gain electrons because it has the highest effective
nuclear charge (and all use the same number of energy levels).
b. C, O, Ne
Ne has the lowest tendency to gain electrons because its outer energy level is full and there is no room for an
additional electron. O has the greatest tendency to gain electrons because it has a higher effective nuclear charge
than C (and both use the same number of energy levels).
c. Si, P, O
O has the greatest tendency to gain electrons because it has the highest effective nuclear charge and uses the
smallest number of energy levels. Si has the lowest tendency to gain electrons because it has the lowest effective
nuclear charge and is tied (with P) for using the most energy levels.
d. K, Mg, P
P has the greatest tendency to gain electrons because it has the highest effective nuclear charge and is tied
(with Mg) for using the smallest number of energy levels. Neither Mg nor K have much attraction for electrons,
but K has the lowest tendency to gain electrons because it has the lowest effective nuclear charge and uses the
most energy levels.
e. S, F, He
He has the lowest tendency to gain electrons because its outer energy level is full and there is no room for an
additional electron. F has the greatest tendency to gain electrons because it has a higher effective nuclear charge
and uses fewer energy levels than S.