Video Transcript
Metallic Bonding
The buildings we live in, the
vehicles we ride in, the device we’re watching this video on, they all harnessed the
characteristics of metals in particular ways. But why do metals behave the way
they do? Well, it has to do with the
arrangement of their atoms and the electrons from those atoms; namely, that their
valence electrons float freely in a sea of electrons. In this video, we will learn how to
describe metallic bonding and the effect it has on the physical and chemical
properties of metals.
Which elements are metals? Let’s take a look at the periodic
table. We can divide the table into three
sections: metals, nonmetals, and metalloids. Metals are the elements that tend
to form positively charged ions. We find them in the left-hand
portion of the periodic table. Nonmetals, which typically form
negatively charged ions, are found on the right-hand side of the periodic table,
while metalloids are in between the two other groups. In their elemental form, metals are
shiny. They conduct heat and
electricity. They are malleable, which means
they can be formed into a specific shape, and they are ductile, meaning they can be
pulled into a wire. Nonmetals tend to be the
opposite. They are dull, poor conductors of
heat and electricity, brittle, and nonductile.
Metalloids have some metallic
properties and some nonmetallic properties. Metals include some familiar names,
such as copper and iron, and they also include substances we may not think of as
everyday metals, such as sodium and calcium. You may think to yourself, “Sodium
is a metal! But table salt, sodium chloride,
has sodium in it, and it’s a white crystal, not a shiny metal.” Well, to answer this question, we
need to look at the different types of substances and compounds that can be formed
by bonds between metal and nonmetal elements.
Before we look at metallic bonding,
let’s recall a couple different types of commonly seen bonds and compare metallic
bonding to these types. Ionic bonds occur between metal and
nonmetal elements. An ionic bond typically occurs when
a metal atom donates one or more electrons and a nonmetal atom receives one or more
electrons. The result is two ions of opposite
charge with a strong electrostatic attraction. While ionic compounds like sodium
chloride may include metal elements, they’re not themselves metallic substances. Covalent bonds occur between atoms
of nonmetal elements. In covalent bonds, the electrons
are shared between the atoms rather than being donated completely to one side or the
other. Simple examples of substances that
contain covalent bonds include water and carbon dioxide.
Our substances of interest, metals,
are made up of atoms of one or more metallic elements. Single-element examples include the
elemental forms of iron, gold, and sodium. Metals that combine different types
of metallic atoms include brass, a combination of copper and zinc, and bronze, a
combination of copper and tin. We call the metallic substances
made of two or more metal elements alloys.
The electrons and metallic
substances flow freely in what we call a sea of delocalized electrons. Instead of the electrons being
bound to a specific atom or shared between two atoms, each atom’s valence electrons
move around freely in the space surrounding the atom. Let’s take a more detailed look at
this arrangement of electrons. The particles in metallic bonds are
arranged like this. The positive ions form a lattice,
while the electrons flow freely in the remaining space. While the valence electrons move
around freely, the other parts of the metal atoms, the positive ions, arrange
themselves in repeated layers. The layers of different metals will
have different geometric arrangements, like the square patterns of gallium or the
hexagonal patterns of sodium that we see here.
The specific arrangement isn’t
always important. But we should know that the
positive metal ions made up of the positively charged nuclei and their inner
electron shells arrange themselves in predictable repeated lattices. Within these lattices, the valence
electrons or the electrons in the outermost electron shells move about freely. We typically describe this
arrangement as a sea of delocalized electrons. Delocalized, meaning that the
electrons don’t tend to belong to any one particular atom but are instead shared
across the network.
In Ionic bonds, the electrons are
gripped by the negative ions. In covalent bonds, each electron is
shared between two particular atoms not shared by the whole network like in metallic
bonding. These free-flowing electrons have
interesting effects on the metal as a whole. As a reminder, in ionic bonds, the
electrons are gripped by particular ions, while in covalent bonds, the electrons are
held in place between particular pairs of atoms. So, the delocalized sea of
electrons in metals gives the network of atoms flexibility and its constituent
particles a lot of freedom of motion. Interestingly, the microscale
freedom of motion of the particles translates into macroscale physical
properties. Let’s take a look at how the
arrangement of particles in metallic bonding leads to some of the physical
properties of metallic substances.
First, metals are malleable. They can easily be formed into
shape, like a bendy sheet of aluminum foil, rather than snap, like a brittle ionic
compound such as a caped-together salt crystal. When a force is applied to a metal,
the shape of the lattice can change, while the particles are still held together by
an electrostatic attraction between the positive ions and the negative
electrons. Thanks to the delocalized sea of
electrons, the positive ions in the metal have more freedom of motion than in a
brittle ionic compound. Metals are also ductile, meaning
they can be pulled into a wire for similar reasons. When a tensile force or a pulling
force is applied, the lattice can rearrange laterally to form the shape of a
wire. Thanks again to the freedom of
motion provided by the sea of delocalized electrons.
Metals conduct electricity as
well. Electricity is the flow of charged
particles. So when a metal is connected to a
circuit, the free electrons will be drawn to the positive terminal. The flow of electrons through the
metal, made easier by their loose arrangement, is what we call electricity. When there are no free charged
particles, like in solid sodium chloride, electricity will not flow through the
substance. But we can create free charged
particles in the form of ions by dissolving the salt to create salt water, which
will indeed conduct electricity.
Metals also conduct heat. Heat is the transfer of energy due
to the movement and vibration of atoms. When a substance is heated, its
particles begin to vibrate. For metals, their positive ions are
more free to move back and forth, more readily transferring energy to nearby
particles, eventually transferring the energy to more and more particles all the way
across the substance. Think of a pan or a kettle on the
stove as a simple example of a metal quickly conducting lots of heat.
Metals also have high melting
points. The substance as a whole is held
together by the electrostatic attraction between the positive ions and the negative
valence electrons. Each ion is attracted to multiple
surrounding electrons, and each electron is attracted to multiple surrounding
ions. So, it requires a lot of energy to
separate a positive ion and its valence electrons from the many neighboring
particles of opposite charge that attract them. The more energy it takes to
separate the particles, the higher the melting point will be. We can also look at how different
characteristics of different metal atoms would affect the attractions in this
held-together network, thereby raising or lowering the boiling point.
One thing that can affect the
melting point of a metal is the number of valence electrons that each atom
contributes to the sea of delocalized electrons. For example, sodium atoms
contribute one electron to become Na+ ions, whereas magnesium atoms contribute two
electrons to become Mg2+ ions. So, how does this difference in the
number of valence electrons affect the melting point? Well, more valence electrons means
there are more particles for the positive ions to be attracted to. This increases the strength of the
metallic bond which leads to a higher melting point.
There is another pattern we can
look at as we move down a group in the periodic table. If we compare sodium to the element
directly below it on the periodic table, potassium, we can ask the question, “How
might the presence of more electron shells affect the melting point of the
metal?” It is worth noting that there are
many exceptions to this pattern. But atoms with more electron shells
have more separation between their nucleus and their valence electrons. In general, this means that atoms
further down in a group on the periodic table have weaker metallic bonds and lower
melting points. But again, there are many
exceptions to this pattern. So, don’t be surprised if you see
an element further down on the periodic table that ends up having a higher melting
point. As you can see, the micro and macro
properties of metals are intricately linked by the structure we call metallic
bonding.
Now that we’ve reviewed metallic
bonding and the physical properties of metals, let’s take a look at some practice
problems.
Which of the following diagrams
best represents metallic bonding?
The diagram that accurately
represents metallic bonding will be the one that has a lattice of positive ions with
a sea of delocalized electrons. Answer (B) is the correct answer,
as it represents the red positive ions in the form of a lattice as well as the sea
of delocalized electrons, the black dots floating in the space between.
Choice (C) has a lattice of
positive ions as well, but there are negative ions mixed into this lattice. Choice (C) depicts ionic
bonding. Choice (A) depicts hydrogen bonding
that occurs between the hydrogen atoms and the oxygen atoms of water molecules. The diagram in choice (D) depicts
network covalent bonding. In this specific case, it depicts
the repeated arrangement of carbon atoms to form diamond. So, the diagram that best
represents metallic bonding is choice (B).
Which of the following metals has
the strongest metallic bonding? (A) Sodium, (B) lithium, (C)
beryllium, (D) magnesium, or (E) aluminum.
This question is asking us to find
the strongest metallic bond. The strongest bond is the one that
has the strongest attractions between the particles. There are two key characteristics
that can affect the strength of the attractions from metal to metal. First, the more valence electrons
that a metal has, the stronger the bond. The extra electrons will add to the
overall attractive forces acting on the positive ions. The other relationship that could
come into play is that, in general, more electron shells means a weaker bond,
although there are many exceptions to this pattern.
The separation between the nucleus
and the valence electrons that the extra electron shells provide weakens the
attractive forces between the positively charged and negatively charged
particles. However, this second relationship
is not as strong as the relationship between valence electrons and bond
strength. So, let’s investigate that
relationship first by finding the number of valence electrons of each of these
metals. How do we find the number of
valence electrons? Well, we can take a look at the
periodic table.
As members of group one, sodium and
lithium have one electron in their outermost electron shell or one valence
electron. One column to the right, we find
beryllium and magnesium in group two with two valence electrons. Aluminum, the third element in its
row in group 13, has three valence electrons. Since it has the most valence
electrons, aluminum, choice (E), is the correct answer. In the case of aluminum, each
positive ion has more negative electrons around it, leading to more attractions and
a stronger bond.
In this problem, we did not end up
using the second relationship between the number of electron shells and the bond
strength. If there were multiple metals that
were tied for the highest number of valence electrons among the choices, we could
use this second relationship as a sort of tiebreaker. However, since there are many
exceptions to this pattern, it would be useful to confirm any assumptions about bond
strength or melting points gleaned from this pattern by looking them up. Of the choices, the metal with the
strongest metallic bonding is choice (E) aluminum.
Now that we’ve done some practice,
let’s review the key points of metallic bonding. Metal elements are the elements
that tend to form positive ions. They’re found on the left-hand side
of the periodic table. Substances are considered metallic
if they contain only atoms of metal elements. If a substance is made up of two or
more metallic elements, we call that substance an alloy. The structure of metallic bonding
involves a specific arrangement of the constituent parts. The positive ions form a lattice,
while the electrons flow free in a sea of delocalized electrons. This structure gives rise to many
physical properties of metals such as malleability, ductility, electrical
conductivity, heat conductivity, and high melting points.
