Video Transcript
In this video, we will learn how to
describe the trends and properties across periods of the periodic table.
The periodic table is an amazing
display of information. Simply by listing the elements in
order of their number of protons and starting new rows at the right time, we can
arrange all of the elements into a visual representation that reveals important
information and trends. In chemistry, periodicity is the
idea that there are trends and patterns that we can find in the periodic table. For the simplest trends, we will
see whether a certain characteristic increases or decreases as we move in two
directions in the periodic table. We can look at the trend down the
group, and we can also look at the trend across the period. In this video, we will take a look
at the periodicity of four different characteristics: atomic radius, ionic radius,
melting point, and electrical conductivity.
First, let’s take a look at atomic
radius. If we consider the structure of an
atom with a central nucleus and an outer ring of electrons, the atomic radius is the
distance between the nucleus of the atom and its outer electrons, the radius of the
atom. However, it’s hard to directly
locate the outermost electron for this measurement, so chemists measure the atomic
radius a different way. They take two atoms bonded to one
another and measure the distance between two nuclei to find a distance equal to two
atomic radii. For example, the distance between
adjacent nuclei in a chlorine gas molecule is 202 picometers. Since we consider the two molecules
to be adjacent with no space between them, the atomic radius of chlorine is half of
that value or 101 picometers.
So, what are the periodic trends
for atomic radius? The first trend is that atomic
radius increases as we go down a group. To investigate this trend, let’s
take a look at the electron shells of two elements in the same group. Sodium is lower than lithium in
group one. As a result, it has more
electrons. And those electrons fill more
electron shells. The end result is that sodium has a
larger atomic radius than lithium. And in fact, in any group on the
periodic table, as we move down the group, we will increase in atomic radius. The trend in atomic radius in the
other direction is that as we move from left to right within a period, the atomic
radius decreases. To investigate this trend, let’s
take a look at two elements in the same period.
Nitrogen is to the right of carbon
in period two. Therefore, it has more protons. More protons means more positive
charge, which means a stronger pull on the outer electrons of the atom. The extra inward pull on the
electrons reduces the atomic radius. So, nitrogen has a smaller atomic
radius than carbon. And as we move to the right in a
period on the periodic table, the atomic radius gets smaller and smaller. The second pattern sometimes seems
counterintuitive. At first glance, you might think
that adding protons and electrons could increase the atomic radius.
As a metaphor to help us understand
the second trend, we can think of the atom as a concert with a popular proton boy
band on stage with an audience of adoring electron fans. The concert has a certain radius,
with fans standing in neat rows. Adding more fans to an existing row
doesn’t change the radius of the concert. But if another boy band member
emerges on stage, the adoring electron fans might move in closer to get a better
look, taking up less space as a result. It’s only when we add fans to a new
row or add electrons to a new electron shell that the radius of the concert
increases. With this metaphor in mind, we can
better understand why adding more protons and electrons increases the attractive
forces between them and decreases the size of the atom.
We have talked about atomic radius
or the radius of neutral atoms. And we can also talk about ionic
radius or the radius of charged ions. The periodic trends for ionic
radius are the same as the periodic trends for atomic radius. Ionic radius increases down a group
and decreases across a period. The trends are the same for both
positive and negative ions. If we look at our previous example
with lithium and sodium but with ions instead of atoms, we can see that a sodium ion
still has more electron shells and thus still has a larger atomic radius than a
lithium ion.
For the periodic trend, we can look
at two ions from the same period of the periodic table, such as an oxygen ion and a
fluoride ion. Whether these particles are ions or
atoms, the one with more protons will have a stronger pull on the electrons and thus
a smaller atomic or ionic radius. We can extend this pattern a bit
further if we look at a group of ions or atoms that are isoelectronic. Two or more atoms or ions are said
to be isoelectronic if they have the same number and arrangement of electrons. This group of atoms and ions here
is isoelectronic. They all have 10 electrons,
completely filling the first two electron shells. The main difference between these
ions and atoms is the number of protons.
Nitrogen on the left has seven
protons and the protons increase by one up to magnesium on the right, which has 12
protons. The same rule from earlier still
applies. More protons means more pull on the
outer electrons, which means a smaller atomic radius. So, as we look at the isoelectronic
atoms and ions with more and more protons, the atomic radius decreases. So the second pattern holds for
ions within a period, as well as isoelectronic ions across several periods.
Another characteristic we can
investigate is melting point. Melting a substance requires adding
enough energy to overcome the attractive forces holding the solid structure
together. So, a high melting point is an
indication that the substance contains strong attractive forces holding its
particles together. This relationship is apparent when
we compare metals and nonmetals. Metals are held together by strong
metallic bonds that take a lot of energy to break. As a result, in general, metals
have high melting points. Nonmetals, on the other hand, are
held together by the much weaker London dispersion forces. As a result, in general, nonmetals
have lower melting points than metals.
This pattern matches what we
observe in real life. At room temperature, the vast
majority of metals are solid, meaning that their melting point is higher than room
temperature. The vast majority of nonmetals, on
the other hand, including oxygen gas and nitrogen gas, are gaseous at room
temperature. For these substances, their melting
points as well as their boiling points are below room temperature. It’s worth noting that this is a
general trend. There are some specific exceptional
values that may deviate from this pattern.
Let’s investigate melting point
more precisely by taking a closer look at one period of the periodic table. The third period of the periodic
table contains the elements sodium, magnesium, aluminum, silicon, phosphorus,
sulfur, chlorine, and argon. Starting with the elements to the
left, we have sodium with a melting point of 98 degrees Celsius, magnesium with a
melting point of 639 degrees Celsius, and aluminum with a melting point of 660
degrees Celsius. These three elements are metals
with relatively high melting points.
As we move to the right among these
three elements, we increase the number of free-floating electrons in the metallic
bond. Sodium has one free-floating
valence electron, magnesium has two free-floating valence electrons, and aluminum
has three free-floating valence electrons. These negatively charged particles
add extra attraction and therefore strength to the metallic bonds that make up the
metal substances. With stronger attractive forces
holding the substance together, the melting point of that substance is higher.
Silicon has the highest melting
point of any element in period three at 1410 degrees Celsius. This high melting point of silicon
is due to its structure, specifically its giant covalent structure. A giant covalent structure is held
together by many covalent bonds. Each silicon atom is covalently
bonded to four other silicon atoms in its immediate vicinity. The number and strength of the
bonds in the structure means that it requires lots of energy to break apart the
solid structure of the substance. As a result, the melting point of
silicon is very high.
Looking at the remaining elements
to the right, no clear pattern emerges. In general, we just need to
recognize that these are nonmetals and therefore they have lower melting points. These melting points are lower
because the particles within the substance are held together by the weaker London
dispersion forces. Although it is worth noting that
the nonmetal sulfur does have a higher melting point than the metal sodium. As a reminder, these are just
general trends. Other characteristics beyond
metallic character can influence the melting point.
Another characteristic relevant to
metals and nonmetals is electrical conductivity. Electrical conductivity is the
ability of a substance to carry an electric current since electricity is the flow of
charged particles. To know how conductive a substance
is, we need to ask how free are the electrons to move. On the left side of the periodic
table, we have metals. Metals are held together by
metallic bonds. Metallic bonds have delocalized
electrons. In other words, some of the
electrons float freely among the atoms of the substance. As a result, metals have high
electrical conductivity.
On the other side of the periodic
table, we have nonmetals. Nonmetals tend to form covalent
bonds, although some nonmetals exist as monatomic elements, in other words, as lone
atoms without any bonds. While the electrons of these bonds
do move around, they’re more confined to the space between atoms rather than flowing
freely about the substance. As a result, nonmetals have low
electrical conductivity.
In between the metals and nonmetals
on the periodic table are the semimetals. Semimetals sometimes behave like
metals and sometimes behave like nonmetals. So, semimetals have medium
electrical conductivity. By visually separating metals from
nonmetals, the periodic table helps us see this trend in electrical
conductivity. Now that we’ve learned about
periodicity, let’s do a practice problem to review.
Which of the following elements has
an atomic radius larger than that of aluminum? (A) Chlorine, (B) sodium, (C)
sulfur, or (D) silicon.
This question is asking us to
compare the atomic radii of different elements. The atomic radius of an element is
the distance from its nucleus to its outermost electrons. Scientists can measure the atomic
radius by taking half the distance between the two nuclei of adjacent atoms bonded
together. To answer this question, we need to
take a look at the periodic table, specifically at period three, which contains all
five of the elements mentioned in the question.
In order to answer this question,
we need to know how atomic radius changes across a period. Once we know this trend, we can
find which of the answer choices has the largest atomic radius. Let’s compare aluminum and silicon
to visualize this change. Since we don’t know what size they
are yet, we’ll start by drawing them as the same size. The first difference we can notice
is that when we move to the right in a period, the number of protons increases. Silicon has one more proton than
aluminum. With more protons in the nucleus,
there’s a stronger positive charge there and thus a stronger inward pull on the
electrons. As a result, there’s also a smaller
atomic radius.
So, the trend is clear. If more protons means a smaller
atomic radius, then the atomic radius decreases to the right and increases to the
left. If we wanna find the element with
an atomic radius larger than that of aluminum, we want to find the element to the
left of aluminum on the periodic table. In this case, that’s sodium. Sodium has fewer electrons than
aluminum and therefore a weaker inward pull on its electrons, which results in a
bigger atomic radius. Which of the following elements has
an atomic radius larger than that of aluminum? That’s answer (B) sodium.
Now that we’ve learned about
periodicity, let’s review the key points of the video. The word periodicity refers to the
trends and patterns found in the periodic table. One trend is that atomic radius
increases down a group and decreases across a period. These trends are due to the
layering of electron shells and the strength of the attractions between the nucleus
and the electrons, respectively.
Melting point is an indicator of
the strength of the attractive forces that hold the substance together. In general, metals held together by
strong metallic bonds have higher melting points than nonmetals, which are held
together by the weaker London dispersion forces. Another characteristic, electrical
conductivity, depends on the free movement of electrons. Metals are more conductive than
semimetals, which are more conductive than nonmetals.
