Metals 101-1 What are Metals

Metals 101 Home

Next: The Structure of Metals

Video Transcript

Metals are all around us, but did you ever take time to really understand what makes a metal a metal, and what makes metals different from the other materials?

Here is a dictionary definition of what metals are:

1 : any of various opaque, fusible, ductile, and typically lustrous substances that are good conductors of electricity and heat, form cations by loss of electrons, and yield basic oxides and hydroxides; especially : one that is a chemical element as distinguished from an alloy

Let’s start near the end of this definition and work backwards.

“ a chemical element” A good place to start looking at metals is on the periodic table of the elements.  By just glancing around this table you may recognize some of the more common metals.  Here is Iron, copper, aluminum, silver and gold.

But what you may not have seen before is that most of the elements on this this table are actually metals.   Here, I’ll remove the elements that are not metals from the table. Ok.  Now you can see that there are many, many metals on the periodic table.  You will also notice that the metals are all grouped together on the left side of the  table. This is because of how the atoms that make up the metals are put together.   If we take a closer look at this we can get a better understanding of what makes a metal a metal and why they have the properties they do.

Now, you are probably familiar with the fact that inside each atom, there is a nucleus.  Inside the nucleus there are protons and neutrons.  It is the number of protons in the nucleus that determines what the element is.  For instance, this atom has three protons in its nucleus. That means that this element is lithium.  The number of neutrons  in the nucleus of an atom is usually pretty close to the number of protons in a atom, but it doesn’t have to be.  The number of neutrons in an atom determines the isotope of a particular element.  Lithium usually has three or four neutrons in its nucleus.  We call these the lithium-6 [pause] and lithium-7 [pause] isotopes.  You see the number comes from adding the number of protons and neutrons in the nucleus.

Now if we look at this simplified version of an atom, we can see that there are electrons flying around the nucleus.  Remember that electrons have a negative charge, and in a neutral atom the number of electrons equals the number of protons.  BUT it is possible to have a different number of protons and electrons. When the number of protons and electrons is not matched, the atom becomes electrically charged, and we refer to it as an ion.  If an ion has more electrons than protons, it is negatively charged and is called an anion.  If it has more protons that electrons, it is positively charged and becomes a cation.  We can see from our definition that metals tend to give up electrons and become cations.

Now let’s look at the shells around the nucleus that the electrons swim around in.   There are rules about how many electrons can go in each shell. We are going to simplify these rules just a little.  The inside shell can hold two electrons, and for the most part, the outside shell can hold up to eight. If there are shells in-between, they can hold different numbers, but we are mostly interested in the outside shells right now.  An important fact in chemistry is that atoms really like to have eight electrons in their outer shell if they can.

Let’s build a few atoms and see how the periodic table is arranged according to this pattern.

What do we need to build our first atom?  We need a proton. Here we go. One proton.  That means we are dealing with hydrogen. We can give it a neutron or even two neutrons if we like, these are the isotopes of hydrogen called protium, deuterium and tritium.  Now, to make it electrically neutral, we can give it an electron. So here we only have one shell. It is the inside shell so remember it can only hold up to two electrons. Let’s add a proton and see what happens.  With two protons we don’t have hydrogen anymore, we now have helium. Helium is a well-balanced element. If we give it two neutrons and two electrons, it will be quite content not to react with any other element. It’s outer shell has the two electrons to fill it up, and it’s charge is balanced.  Since the first shell is filled, the first row of the periodic table is filled.

Now things get interesting when we add another proton to create lithium.  For lithium to be electrically balanced, it needs three electrons. But if we try to give lithium three electrons they won’t all fit in the two-electron inner shell, and the outside shell begins filling up. This shell can hold eight electrons but there is just one electron sitting there.  This makes lithium feel a little self-conscious. Atoms want to look well-balanced on the outside. So what lithium would really like to do would be to get rid of the single electron in its outer shell. Sometimes an atom would rather carry a charge than to have one electron in their outer shell.   The elements in this whole column all work this way. They are called alkali metals. Alkali metals react easily because this single outer electron is very unstable. Alkali metals are famous for their violent reactions with water. Just look them up on youtube and you’ll see what I mean.

Now, let’s add more protons and electrons to see the shells fill up.   Here’s Beryllium, carbon, nitrogen, oxygen, fluorine, and finally Neon.  Do you see how we have filled up all the available spaces in the shells? Neon, like helium has a completely filled outer shell.  In fact, it is true of all of the elements in this column. The elements here are called the noble gases. Because their outer shell is full, they tend not to react with anything.  This is a useful property. For example, Argon and Helium are often used to prevent reactions from taking place during welding.

As we keep adding protons and filling shells with electrons, we can see that a pattern emerges.  As the shells fill up, the properties of the elements seem to repeat. Here we are back to the first column, the alkali metals.  In the second column, here, the outer shell has two electrons. The metals in this column are called “alkaline earth” metals, and while they don’t react as violently as the alkali metals, they do react easily.  That is because they have two electrons in their outer shell. They would rather lose these electrons than be electrically balanced.

In fact, losing or giving away electrons is what metals do so well.  As we move across the table, we pass through the transition metals. It’s easy to find familiar metals in this group of elements.  Here is Titanium, Chromium, Iron, nickel and copper. Look at what happens as we add protons to move across the table. See how the middle shells are beginning to fill up.  But the inside shells still has two and the outside almost always has less than eight. You can see that the metals have low numbers of electrons in their outer shells. This means they can easily give up or “donate” electrons.

So that’s three groups of metals so far, the Alkali metals, alkaline earth metals, and transition metals.  There is one more group, here the post-transition metals. These are sometimes called the poor metals or “other metals.”  Here too, you will see some familiar metals. Here is aluminum, tin and lead.

Now to contrast, let’s look at some of the nonmetals.  Here is choline. The outer shell of chlorine is almost completely full, with seven electrons in its outer shell.  Chlorine would love to get hold of an an extra electron and become positive than to have this almost-filled outer shell.  The nonmetals tend to be takers of electrons instead of givers like the metals are.

So as a general rule, metals like to donate electrons and nonmetals tend to take electrons. Let’s quickly look at how this makes atoms stick together to form molecules, because that will give us some more insight into what metals are and why they have the properties they do.

You probably remember that there are different types of chemical bonds.  Ionic, Covalent, and Metallic. Let’s review those and then look in particular at the metallic bonds.

Ionic bonds

Let’s look at sodium and chlorine for a minute. See how they are at different ends of the periodic table?  Sodium has one electron in its outer shell. Remember that the metals like to give electrons? Sodium would very easily give this electron up to have a nice, neat outer shell with eight electrons.   And here is Chlorine. It has seven electrons in its outer shell. If it could just get one more it would become stable like a noble gas. So when sodium and chlorine are close to each other, and are given just a little push of energy, the electron leaves the sodium for good and becomes part of the chlorine atom. This gives the sodium a positive charge, and the chlorine a negative charge.  Do you remember the name of an electrically charged atom? They are called ions and ions of different charges attract each other. So the sodium and chlorine stay attached to each other by this newly formed ionic bond.  Notice that with this arrangement, we can say these particular electrons belongs to the chlorine atom, and these belong to the sodium atom.

Covalent bonds

Let’s take a look at an element that has a half-filled outer shell.  This is carbon. It has four electrons in a shell that could ideally hold eight.  Carbon is not eager to give up electrons, and its not likely to just pull in four electrons just to have a filled outer shell.  What it does do nicely is to work out a sort of time share agreement with elements like hydrogen. You see if hydrogen and carbon are encouraged just a little, they can work out a sharing plan between themselves.  A sort of custody agreement with the electrons. You see, four hydrogen atoms and one carbon atom can share electrons with each other. Hydrogen can have two electrons in its outer shell and it will be stable, carbon can have eight to be stable.  They can share electrons to make this happen. This electron is shared between this hydrogen atom and the carbon atom. If carbon can do this with enough other atoms, it can look like it has eight electrons in its outer or “valence” shell. This sharing type of bond is called a covalent bond, and it you can see it being used in all kinds of neat molecules.  For example, this paraffin molecule.

Metallic bonds

When metals are around other metals, a really neat thing happens.  The metal atoms all share their outer electrons freely. Since they hold their outer electrons very loosely the electrons are free to move around this group of metals atoms.  The atoms are positively charged ions or cations because they have all given up their outer electrons. This field (or “sea”) of electrons attracts the metal cations and holds them together because of their opposite charges.  This is a metallic bond. It is a strong bond that only occurs to hold metals to other metals. In fact, it is this unique bonding phenomenon that explains a lot of the properties of metals.

Let’s look back at our definition:

any of various opaque, fusible, ductile, and typically lustrous substances that are good conductors of electricity and heat, form cations by loss of electrons, and yield basic oxides and hydroxides; especially : one that is a chemical element as distinguished from an alloy

When light strikes a metal, it hits the freely mobile electrons and bounces right off.  This explains why light won’t go through the metal, and in fact this bouncing back of light explains the shininess or lustrous property of metals.

Metals are also ductile and malleable.  This means they can be bent or hammered without breaking.  The metallic bonds allow the atoms that make up the material to slide past each other and still be bonded.  Ionic bonds and covalent bonds do not allow this type of behavior.

It’s probably easy to see at this point why metals are such good conductors of electricity.  Electrical current is the flowing of electrons. The valence electrons in metals are free-flowing.  They are not attached to any atom in particular. This free movement also allows heat to be conducted through the metal.

When you look at the periodic table, you might notice that some metals seem to be missing. For instance, there is no steel or brass on the table.    That is because those metals are not pure, elemental metals. They are what is known as alloys. Alloys are metals that are combined with other elements (either metals or nonmetals).  For example, brass is an alloy, it is a combination of copper and zinc. Steel too is an alloy, it is a combination of iron and carbon.

Now, we will deal with oxidation in a later video, but based on what you know about metals and their electrons, you will probably have a pretty good head start on understanding what happens during oxidation.  Just look up a video on oxidation and reduction reactions. I’ll link to one in the description below.

So in review, metals are the elements on the left side of the periodic table.  They tend to donate electrons from their outer shell. Metals form metallic bonds with each other and this special type of bond is responsible for many of the properties that characterize metals.  Alloys are not pure, elemental metals, they are a combination of a metal with other elements.