Atoms Part 3: Atoms And Heat

Incandescent light bulbs, stars, and hot briquettes on the barbecue, like the ones below, all glow brightly. The atoms in them emit light.

(Jihei;Wikipedia)

When atoms are very close together, they interact with each other, and when they do, they can sometimes emit light. In this case, the colour of the light they emit depends on their average energy. Some light bulbs, especially the old incandescent ones, as well as briquettes and stars glow brightly because they are hot.

From A Line Spectrum To A Continuous Spectrum

Hydrogen, the simplest atom there is, can emit dozens of different photon wavelengths when it is excited. That's a lot of lines in hydrogen's electromagnetic (EM) emission spectrum, as we saw in Atoms Part 2, but it is nowhere near a continuous EM spectrum. The visible emission spectrum of hydrogen is even simpler - it contains just four lines:

The sample of hydrogen that scientists use to obtain these spectra is very pure. A sample mixture of excited and ionized hydrogen gas containing both atoms and molecules would make the emission spectrum more complex. It would more vaguely resemble a full continuous spectrum like the one you see through a prism held up in sunlight. If we think about what the emission spectra of mixtures of atoms might be, say bound up in molecules, we start to wonder if we're going to approach a continuous emission spectrum, where all possible lines (wavelengths of emission photons) show up, and we do get closer to one. But even emission spectra of molecules resolve into distinct spectral lines, when you use a very good spectroscope. The Sun's emission spectrum is a continuous spectrum, which means there are no, even tiny, white gaps between spectral lines. The Sun's visible emission spectrum is a what you see when you put a sunbeam through a prism, or a rainbow:










Something else must be going on here, besides atoms and possibly a few molecules, emitting their emission spectra.








(D-Kuru;Wikipedia)

The Sun's Continuous Emission Spectrum

With what we've learned about emission spectra, you might expect to see something like a combination of hydrogen and perhaps helium emission line spectra superimposed on each other when you look at sunlight through a prism. And why shouldn't you? Our Sun is made of mostly hydrogen (71% of its mass; 91% of its atoms) and a little helium (27% of its mass; 8.7% of its atoms). But you don't. You see a continuous spectrum of visible colour rather than spectral lines. You would see a continuous spectrum even through a good spectroscope.

In this case, atoms emit light as they did before. But the spectrum they emit is different . . .

Black Body Radiation

If you looked at our Sun from outer space it would look like a glowing white sphere. Our atmosphere makes it look more yellow and it's a very common misconception that the Sun looks somewhat yellow from space:

(NASA image from International Space Station)

There is something very interesting about the Sun's white glow. It doesn't matter what the Sun is made of! It doesn't depend on the kinds of atoms involved. It could be a solid ball of iron and, if that iron ball were the same temperature as our Sun, it would glow the same white colour. Any other star at the same temperature as our Sun glows the same white colour, no matter what its composition is.

How can this possibly be true?

At the turn of the century, scientists didn't know much about electron orbitals in atoms. They didn't understand emission line spectra. They were focused instead on the colours of glowing hot objects. For example, Max Planck, looking at hot metals, wondered why any kind of metal, when it's heated, goes through the same sequence of glowing colours. It gets hot and begins to glow dim red. Then the red gets brighter and begins to shift to yellow, then to white, and finally to bluish-white. Why?



He realized that a hot object gives off light, or more precisely, EM radiation, known as thermal radiation. He also realized that as an object grows hotter, the intensity of the EM radiation increases as the wavelength decreases. The light it emits gets brighter as it gets hotter, and it shifts from longer wavelength red to shorter wavelength blue. Here is the relationship, based on experimental data, in graph form, below right:

The wavelength of the EM emission from the hot object is shown in micrometres (millionths of a metre) along the bottom of the graph. The intensity of the EM emission is plotted along the y-axis. At 300 K (around 27°C), the object emits only a few infrared photons. Notice how the blue line peaks at around 10 µm? We sense this wavelength of radiation as feeling warm to the touch. At around 1000 K (725°C; green line) the object is hot enough that it starts to glow dull red. At around 6000 K (red line; that's about how hot the surface of the Sun is) the object looks bright white. See how the red line peaks right at the visible spectrum? If you look at sunlight through a spectrum you see a nice continuous rainbow. This spectrum is called the Sun's black body spectrum. So what's a black body?

Black Body

Our Sun is a good, but not perfect, example of a black body. A perfect black body, in physics, is an object that absorbs all incoming EM radiation but doesn't reflect any back. A black charcoal briquette (below right) comes close to a perfect black body:

The briquette is opaque, unreflective (almost) and black. It absorbs all the (at least visible) photons that strike it. Its visible absorption spectrum is continuous. Its emission spectrum, on the other hand, would be black if it were perfectly cold. But nothing in the universe is perfectly cold. It emits at least a few infrared photons. We sense these photons as warmth. Infrared photons are invisible to us, so a perfect black body at room temperature would look perfectly black like the charcoal. If we heat the charcoal it will get hot. It will radiate more infrared photons.

                                            (Vladsinger;Wikipedia)

As we heat it further it will start to glow dull red at around 1000 K, then yellow, white and eventually blue. Heated further still, over 10,000 K, it will radiate ultraviolet (UV) photons, but it will continue to radiate some blue photons too so it will still look blue to us. It won't turn invisible. And it will continue to radiate lots of infrared photons as well. In fact, it will radiate a whole continuous series of photon wavelengths, some of higher and lower wavelengths, but most will be UV photons at this temperature.

It seems really odd that the wavelength of these emission photons doesn't depend on the charcoal being made of carbon atoms. It only depends on how fast the carbon atoms are oscillating. More precisely, it depends on how fast charges are oscillating among the atoms. When thermal energy (heat) is applied, atoms oscillate faster and faster. The electrons in them are along for the ride. They oscillate too.

How Atoms Emit Black Body (Thermal) Radiation

Let's stop here a moment and examine the electron. The nuts and bolts of black body radiation really focuses on what the electrons are doing. Just as in excited atoms, it's the movement of electrons that creates light.

The electron is a quantum wave function. According to quantum mechanics, each electron is defined by four quantum numbers. We already know one number. It describes the energy of the electron. Other numbers describe its spin (electrons come in two spins - up and down), its angular momentum (this defines the shapes of its orbitals) and its magnetic moment. This last number tells us that electrons act like tiny quantum magnets. Electrons also have a negative charge. As the electrons oscillate along with their atoms, they set up tiny oscillating perpendicular electric and magnetic fields. These fields propagate, or move, as electromagnetic (EM) waves:



How fast an atom (and electrons) vibrates determines the oscillation (frequency) of the EM radiation it emits.

A real object contains atoms that are oscillating at a variety of frequencies. They have various amounts of kinetic energy. The temperature of the object reflects the average kinetic energy of its atoms. That average determines the general frequency of EM radiation it emits, what colour it glows when it's hot.

The emission spectrum of black body EM radiation reflects a whole variety of frequencies, or wavelengths, of photons, with a majority vibrating at some particular wavelength. The peak wavelength is the average energy of the atoms in the object. The overall average of thermal energies is what the thermometer reads. There are no forbidden photon wavelengths in black body radiation. Thermal radiation, and the black body spectrum of that radiation, depends only on the temperature of the object, not on its elemental make up.

Thermal radiation is yet another way that the electrons in atoms can emit light. The mechanism behind thermal EM radiation is subtly different from atomic EM emission, but in both cases, energized atoms release excess energy in the form of photons. When atoms are excited, the excess energy is localized in their electrons. When atoms gain thermal energy, the excess energy is distributed across the entire atom in the form of kinetic energy. It moves faster. Both mechanisms can and often do happen at the same time in one atom. Very hot materials are full of atoms that are both excited and thermally energized.

Why a Black Body is Assumed to Be a Solid Body

When we look at sunlight through a prism, we see a black body spectrum rather than an atomic emission spectrum. This has to do with how close the sun's atoms are. Thermal energy is atoms with kinetic energy. Thermal energy is what makes atoms jostle around amongst each other when they are close together in a solid. They can't go anywhere because they are chemically bound together, so they oscillate or vibrate instead. Pure elements in solid form melt when they get hot enough. If we take a ball of iron and heat it, it's going to glow red but its also going to melt. Molten iron glows the same red, then white, then blue. Some iron, if hot enough, will also vaporize into gaseous iron. In gases, atoms can move more freely. In this case, atoms, with their shells of electrons, act like tiny beach balls that bounce off each other. The electron shells have a spring-like quality - they deform a little like a spring, transferring thermal energy from atom to atom when they collide. A gas tends to expand when it is heated. The atoms in a (non-compressed) gas don't have much of a chance to set up oscillations. That is why scientists tend to focus on oscillations that set up thermal radiation in solid materials where atoms are forced to stay close together and vibrate.

Hmm, the Sun isn't solid is it? It is in a fourth kind of physical state called plasma. Plasma is a state that is physically similar to a gas but the Sun is so massive that its atoms, even though they contain a great deal of thermal energy, are forced to stay close together, close enough to set up thermal oscillations, because they are pushed down from the outside in by gravity. The Sun is surrounded by a cloud of gas, called a corona, made visible during a solar eclipse, below:

(Luc Viatour / www.Lucnix.be;Wikipedia)

Like some planets, it has an atmosphere. Most of this gas consists of ionized and excited atoms. Gravity pushes atoms into the Sun but the energy from its nuclear fusion blows atoms away from it too. If we blocked out the hot Sun itself and looked at the spectrum of the gas around the Sun, we would see atomic emission line spectra of those gases.







How Do Hot Objects Cool Down?

In a hot object, energy is transferred from atom to atom and it is also radiated away as thermal radiation (heat and sometimes, light). Hot objects eventually cool down as their thermal energy is transferred to other atoms, such as those in the air or in other nearby objects, and as kinetic energy of atoms is converted into, and carried off as, EM radiation.  An atom all by itself can have kinetic energy, and that energy is reflected in the atom's velocity. But it isn't accurate to say an individual atom has thermal energy. The atom can't interact with other atoms, so it can't release its kinetic energy as thermal radiation. It can only release excess energy by emitting photons from its excited electrons. High up in the atmosphere where the thermosphere lies, atoms are very few and far between, but those that are present are continually bombarded by fast moving electrons and other particles streaming from the Sun and accelerated along Earth's magnetic field lines. These collisions excite the atmospheric atoms and transfer a lot of kinetic energy to them. That energy makes the thermosphere incredibly hot, up to 2000 C (calculated based on average velocity). But, if you placed a thermometer up there, it would read well below freezing because so few atoms would collide with the mercury inside it and transfer their kinetic (thermal) energy to it. Thermal energy is the energy of multiple atoms close enough to each other to conduct and transfer kinetic energy.

A Closer Look at the Sun's Spectrum

There are no perfect black bodies in real life, with one exception - black holes. They absorb all EM radiation. Their black body radiation is better described as Hawking radiation, as temperature depends on the mass of the black hole, and it is the subject my article on black holes.

Our Sun, however, is a pretty good approximation of a black body but it also contains excited hydrogen atoms, within its chromosphere and photosphere, two layers that are cool enough for these atoms to exist in excited, but not ionized, atomic form. Deeper inside the Sun, hydrogen atoms are fully ionized into seething plasma made of protons and free electrons. Excited hydrogen atoms are emitting line spectra that signal their presence but the Sun is a very hot object. Its intense black body radiation (its thermal radiation) overwhelms its atomic emission spectra. Discrete spectral emission lines can be carefully picked out from the Sun's black body emission spectrum, however, representing excited hydrogen atoms, hydrogen ions, a few helium atoms and ions, etc. The radiation from the Sun is a very complex mixture of not only excited atom and ion emissions but of intense thermal radiation as well as intense EM radiation from nuclear fusion as well as other processes. Its radiation covers the whole visible spectrum and much of the entire EM spectrum as well - as gamma rays, X-rays, UV rays, radio waves, etc., all emanate from it.  The Sun is a mess of radiation, reflecting an overall surface temperature (an average energy) of around 6000K. That is why it looks white.

The Sun's Line Emission and Absorption Spectra

Scientists have some tricks to get around the overwhelming black body radiation from the Sun in order to see its hydrogen emission spectrum. They can block out the Sun itself and examine its corona through a spectroscope. Or, they can look through a telescope with a hydrogen-alpha filter. This optical filter allows only a narrow range of light around the H-alpha wavelength (the bright red line emission in the Balmer emission spectrum). The Sun seen through this filter, you guessed it, looks like a giant red ball:

(Eyrian;en.Wikipedia)

Scientists can do one even better by looking very carefully at the emission spectrum of the Sun. By doing so, they can make out hydrogen's absorption spectrum within the Sun's emission spectrum. They can make out many fine black lines in it, called Fraunhofer lines:

You need a very good prism to see them. Joseph von Fraunhofer himself was shocked to see these dark lines through a set of new improved prisms he had just made. He spent many years studying and measuring these lines. They show up exactly where the hydrogen gas emission lines do and they represent hydrogen's absorption spectrum embedded in the Sun's black body spectrum. It means that near the surface of the Sun, between it and our spectroscope, a modern version of Fraunhofer's prism, hydrogen atoms are cool enough to absorb EM radiation. They block out emission lines in the black body spectrum. If you look carefully at the spectrum above, however, you will notice that there are many more lines here than those that match the visual absorption spectrum for hydrogen. The Sun also contains some helium, a little bit of carbon and even very tiny amounts of iron and other heavy atoms. These elements also contribute Fraunhofer lines. If you read Atoms Part 1, you might be asking yourself how these heavy elements got into the Sun. It's a small star and it doesn't make them. 1.5% of the dust cloud that condensed into the Sun billions of years ago contained heavy elements. Much of that cloud was dust leftover after large ancient stars blew up at the end of their lifespans.

Scientists can look at stars and tells us how hot they are as well as what they are made of. If a star looks bluish-white, they know its temperature is around 8000K. If they look at it through a spectroscope, they will see a continuous spectrum, with a brightened blue region, and various absorption lines too, like the Fraunhofer lines. These lines are the absorption lines of atoms near the surface of the star; they tell scientists what the star's surface is made of.

Here is a black body applet to play with and get a feel for the concept:

Temperature versus Wavelength of Black Body Radiation

From Black Body Radiation To The Modern Quantum Mechanical Model Of The Atom

Planck understood black body radiation as a classical effect of heat applied to an object. He knew that when an object absorbs enough energy (heat) its atoms begin to vibrate. The relationship between the wavelength and the intensity of light emission from hot objects suggested to Planck that atoms vibrate like little harmonic oscillators, but how?

The Accidental Beginning of Quantum Mechanics

Before he could get very far, Planck saw a problem with his theory and it was a big one. He was working with classical electromagnetic theory to understand how atoms, when heated, act like harmonic oscillators. This sounds a little familiar so far doesn't it? The problem is that when he used Rayleigh-Jeans law, a classical law that's part of this theory, to calculate the relationship between the intensity and the wavelength of the objects' radiation, the solution didn't match his experimental results. According to this law, the number of electromagnetic vibrational modes is proportional to the square of the frequency. Each mode has the same energy so as the number of modes increases, energy increases. As the modes increase, wavelengths get shorter, and intensity approaches infinity:

This implies that a black body (at a high enough steady temperature) emits EM radiation of infinite intensity. Our Sun, at around 5000 K, would have grown infinitely bright. Naturally, it hasn't. Planck had to fix the problem and he did so by suggesting that EM radiation did not follow classical laws. Instead, EM radiation could only be emitted in discrete packets of energy that are directly proportional to the wavelength. Applying this adjustment brought the theoretical solution into line with the experimental results. For Planck it was a desperate shot in the dark, and at first he didn't know why it worked. He effectively reduced the number of electromagnetic modes possible. What at first seemed like an arbitrary fix, became an object of fascination to scientists at the time. Several physicists (Einstein, Heisenberg, Born, Bohr, de Broglie, Dirac, etc.), seeing a giant new possibility open up, took this work and expanded on it to create the modern quantum mechanical model of the atom, a giant breakthrough in our understanding of how atoms work. The energy of a photon is directly proportional to its wavelength. The intensity of the light from a hot object depends only on the number of photons coming from it, not on the energy of the photons. Intensity is delivered in discrete packets or quanta in other words. The reduced modes of electromagnetic vibration that Planck came up with are what we now call photons.

Conclusion

Over this article and the previous one, we took a close look at the emission and absorption spectra of hydrogen, as well the black body spectrum of an object almost entirely composed of hydrogen, in order to see how electron movement inside atoms works. It all boils down to two simple concepts:

First, electrons move away from the nucleus into higher energy orbitals when they absorb energy, and they release that energy by emitting light when they return to their lowest energy state.

Second, when an object gets hot, its atoms bounce against each other and oscillate, and the electrons in them vibrate. The electromagnetic vibration of electrons is what light is made of - a propagating electromagnetic oscillation.

The emission of light from countless sources around us owes itself to how electrons behave in atoms. By examining how atoms emit light, we followed in the footsteps of many physicists who devoted their lives to understand how the atom works: the modern quantum mechanical model of the atom.

Next we'll look at how atoms bond together in Atoms Part 4 - Atoms and Chemistry.

Atoms Part 4A: Atoms and Chemistry – Atomic Orbitals and Bonding

In Atoms Part 2 and Part 3, we explored how the electrons in atoms emit photons of electromagnetic radiation by comparing different kinds of spectra from hydrogen atoms and from the Sun, essentially a huge ball of hydrogen.

We're not done with electrons in atoms - not only do they change energy states by moving into different orbitals, they interact with electrons in other nearby atoms. In a nutshell, this is what chemistry is.

Chemistry is All About Bonds

You may have heard of two basic kinds of bonds: ionic and covalent.  We will explore and compare them in the next article. They both achieve the same goal: atoms bond with each other because they want to find the lowest possible energy state. Chemical bonding helps atoms achieve this goal. In order to understand bonding, we first need to look at the atoms themselves and how their electrons are configured, and that is what we're about to do here.

ORBITALS AND BONDING

Orbitals: Their Shapes and Their Electrons

In atoms, electrons occupy orbitals. An orbital is a cloud-like area around the atom's nucleus where an electron is likely to be found. To give you an idea of what these shapes look like, the first five atomic orbital shapes are shown below.

From left to right they are the 1s, 2s, 2p_x, 2p_yand 2p_z orbitals. The 2p_x, 2p_y and 2p_z are actually suborbitals, subtypes of the 2p orbital. All this will become clearer as we go on.

Orbitals play a crucial role in how atoms bond with each other. There is an orbital rule every atom must obey - only two electrons can occupy the same orbital, and only if their spins are opposite. So, the 1s and 2s orbitals can fit two electrons in them, and the 2p_x, 2p_y and 2p_zorbitals, likewise, each can only fit in two electrons.

This rule is based on the Pauli exclusion principle, which says that no two electrons in an atom can share all four of the same quantum numbers. Quantum numbers are part of the quantum mechanical model that describes the atom. These numbers describe the energy, angular momentum, magnetic moment and spin of electrons. The shape of each electron orbital depends on the electron's angular momentum. The energy of the electron determines which orbital it occupies. The electron's magnetic moment is behind all magnetic phenomena. That leaves spin, and electrons can choose one of two possible spins. This means that two, and only two, electrons can share the same orbital (or suborbital) because each electron can have one of two possible quantum spins. All the other quantum numbers of these two electrons will be the same.

You can quickly find out how many electrons any atom has by looking at the periodic table:

The number in each square is the atomic number. That's the number of electrons (and protons) in each atom. Hydrogen [H] has only one electron while chlorine [Cl] has 17 electrons, for example.

Hydrogen just has one (ground state) orbital, 1s¹. The superscript "1" tells you the number of electrons occupying the orbital, in this case just one. Lithium [Li] has three electrons so it fills up the 1s orbital (this orbital can fit two electrons) and sets one electron in the next highest energy 2s shell: 1s²2s¹. Phosphorus [P] with 15 electrons is written as 1s²2s²2p⁶3s²3p³. Notice that the 2p shell allows six electrons in it, while the 3s orbital allows 2 electrons. Remember that the 2p orbital is actually made up of three suborbitals (take a quick look at the previous orbital diagram, at the three suborbitals to the right). Each of these three suborbitals holds two electrons, one of each spin. The suborbitals keep the electron pairs separate. A way to visualize the p orbital is shown below. Each p suborbital (a separate colour) has two lobes and it lines up along one of three axes:

This is what a typical p orbital looks like, whether it is 2p, 3p, 4p and so on. As the number in front increases, the p orbital's energy increases, so a 3p orbital has more energy than a 2p orbital, for example. The overall shape stays the same. The atomic nucleus is in the middle where the lobes intersect.

Now let's compare the p orbitals of two atoms. Argon [Ar], a noble gas atom, has 18 electrons (1s²2s²2p⁶3s²3p⁶). Its outermost 3p orbital is full; it has 6 electrons in it. Another atom, phosphorus has 15 electrons (1s²2s²2p⁶3s²3p³). It has a 3p orbital that's only half filled, with three electrons. Argon's filled orbitals make it very stable. The more stable an atoms is, the lower its potential energy is. This allows argon to exist in a very low energy state and it likes to stay that way. It is a very nonreactive, or inert, atom. The half-filled 3p orbital in phosphorus, on the other hand, makes its electron configuration a bit less stable. This atom is reactive; it can adopt a lower energy argon-like stability if it attracts three electrons to itself. It can do this by forming bonds. This is how atomic orbitals make chemistry possible.

The d orbital is higher energy yet. It can hold up to ten electrons. It holds a maximum of two (opposite spin) electrons in each of five suborbitals. This orbital looks a bit like a fancy three-dimensional daisy, shown bottom middle, below:

(User:Sven;Wikipedia)

Electron Orbitals Versus Electron Shells

You have probably heard of electron shells and valence shells before. For example, argon, a noble gas, has a full valence shell of eight electrons.

Talking about orbitals and shells can get a bit confusing, so let's compare the two. If you look at a typical Bohr diagram of argon, shown below, you will see rings, or shells as they are called, representing specific energies at which electrons may be found.

(commons:User:Pumbaa (original work by commons:User:Greg Robson); Wikipedia)

The maximum number of electrons per shell depends on the shell. The first or closest shell can hold two electrons. The second one can hold eight, the third can hold 18, the fourth can hold 32, and so on. Inner shells generally must be filled before outer ones. The outermost shell of electrons is called the valence shell. Notice that argon has eight electrons in its third shell, which is its valence shell. At first glance it looks like argon's valence shell is less than half full. It has eight out of 18 possible electrons in it. Shells consist of subshells. Subshells and orbitals are the same thing. So, a p subshell is a p orbital (and remember the p orbital can be further divided into three suborbitals). The number in front of the subshell indicates what the electron energy is. A 3p subshell belongs to shell n = 3 in a Bohr diagram. Argon has eight n = 3 shell electrons. The three p suborbitals account for six electrons in total, so where do the other two come from? From the 3s orbital - it also has n = 3 energy, and being an s orbital, it contains two electrons when filled.

Argon has eight electrons in its valence shell (n = 3), which consists of two subshells or orbitals: 3s and 3p. The 3s (two electrons) and 3p (six electrons) subshells make up the n = 3 electron shell (eight electrons). As I said, the n = 3 shell can hold even more electrons – 18 in total. There is another orbital available to electrons with n = 3 energy: the d orbital. The 3d orbital can hold up to ten electrons in it. Take another quick look at the daisy-shaped d orbital diagram above. Argon has no electrons in this orbital. Its 3s and 3p orbitals, however, are full and that makes it stable.

Transition Metals: An Exception to the Orbital-Filling Scheme

There are a few exceptions to the general orbital filling scheme. One exception you might come across is the transition metal group. Many gems such as rubies and sapphires, owe their brilliant colours to the uniquely spaced electron orbitals of transition metal atoms inside them. It is also why transition metals can form a variety of different oxidation states. For example, Iron [Fe] can form both Fe²⁺ and Fe³⁺ ions when it reacts with oxygen to form iron (II) oxide and iron (III) oxide. Vanadium [V] has even more oxidation states – four, each with its own attractive colour in solution.

Nickel [Ni] is a transition metal atom with ten more electrons than argon. It does not just go ahead and fill up the 3d orbital, however. It fills up the 4s orbital first.

(commons:User:Pumbaa (original work by commons:User:Greg Robson); Wikipedia)

In transition metals, the inner 3d orbital generally has more energy than the valence shell 4s orbital.

Noble Gas Orbital Writing Shortcut

A shortcut to writing out long orbital notations is to use a noble gas core. Noble gases (group 18 on the periodic table) have completely filled orbitals. Our example argon is one of them. A shortcut for phosphorus, for example, uses the next smallest noble gas core. Phosphorus has 15 electrons, five of which are valence electrons, confined in two more orbitals of electrons than neon has. Neon [Ne], with 10 electrons, is written as 1s²2s²2p⁶, so for phosphorus we can write [Ne]3s²3p³. Nickel, above, can be written as [Ar]4s²3d⁸.

Halogens (group 17 on the periodic table) are just the opposite of the noble gases. They have an outer shell of seven electrons rather than the stable eight of the noble gases. These elements want just one more electron very badly and that makes them highly reactive. They gain this electron by reacting with other elements.

Fluorine [F], shown as a Bohr diagram below, is a great example.

(Pumbaa;Wikipedia)

It's the most reactive element there is. It will attack even inert materials like glass, and it will even form compounds with the heavier noble gases, which are generally nonreactive. Once fluorine bonds with something, that bond is so strong that almost nothing can break it. This makes Teflon so perfect as a non-stick non-reactive coating. It's made of carbon bonded with fluorine.

Now that we understand what makes some atoms more stable than other ones, we can explore how atoms increase their overall stability by bonding with each other, in Atoms Part 4B – Ionic and Covalent Bonds (coming soon).

Atoms Part 4B: Atoms and Chemistry - Ionic and Covalent Bonds

All atoms want to achieve a lowest possible energy state. They can accomplish this by increasing the stability of their electron configuration. Noble gases have the most stable electron configurations because their electron orbitals are full. An atom can achieve a noble gas-like configuration by sharing, donating or accepting electrons from other atoms. This concept leads to the octet rule in chemistry.

Octet Rule

The octet rule is a useful rule of thumb in chemistry, especially when we visualize chemical bonds using Lewis dot diagrams. Atoms, particularly those with atomic numbers less than 20, may combine in such a way that they attain eight electrons in their valence shells. It's a very stable electron configuration like that of a noble gas. An example is carbon dioxide, shown below in a Lewis dot diagram. Carbon [C] atoms have four valence electrons (1s² 2s² 2p²) and oxygen [O] atoms have six valence electrons (1s² 2s² 2p⁴). Remember that the inner 1s orbital electrons are not valence electrons. Carbon forms a stable bond by sharing valence electrons with two oxygen atoms, with each bond contributing two electrons to carbon's valence shell, making it a stable octet. Each oxygen atom also forms its own octet:

The shared electrons count in both the oxygen octets and the carbon octet. The total energy of the carbon dioxide molecule is lower than the energy of the three separate atoms involved. Ionic and covalent bonds, two main kinds of bonding, usually follow the octet rule.

BONDING: IONIC VERSUS COVALENT

The first thing to keep in mind is that ionic and covalent bonds aren't mutually exclusive. There are many examples of bonds that fall in between ionic and covalent, forming a spectrum of bonding with covalent at one end and ionic at the other. Pure ionic bonding, in fact, doesn't even exist. There is always at least some covalent nature, some electron sharing in other words, to it. Some elements can form either kind of bond. Hydrogen [H], for example, forms bonds ranging from strongly ionic in nature to strongly covalent. We'll define these bonds as we go on.

Electronegativity: How Atoms Choose Which Kind of Bond

The type of bonding that happens depends on the electronegativity of each atom. Electronegativity is a chemical property that describes the tendency of an atom to attract electrons to itself. It depends on the number of electrons in an atom's valence or outermost shell as well as how far those valence electrons are from the atom's nucleus. This brief two-minute video describes the difference between ionic and covalent bonding:



The screen shot from Wikipedia below shows how electronegativity trends in the periodic table of elements:

The higher the number, the more electronegative the atom is. Sodium [Na] has a relatively low number, 0.93. It tends to form positive ions (Na⁺), for example, whereas fluorine [F] has the highest number of all, 3.98. It will only form a negative ion when it bonds and it really wants to attract an electron to itself.

The difference in electronegativity between two atoms determines whether they will form an ionic bond or a covalent bond. A larger difference tends toward an ionic bond and a smaller difference or no difference tends toward a covalent bond.

Sodium Chloride: A Typical Ionic Bond

Sodium [Na] and chlorine [Cl] are two atoms with a large difference in electronegativity, 3.16 - 0.93. They react, creating a strong ionic bond to form sodium chloride or table salt.

Sodium, an alkali metal belonging to group 1 on the periodic table, has one more electron than a noble gas, meaning that it has a lone valence electron, while chlorine, a halogen belonging to group 17, is one electron short of a noble gas. Ionic bonding for them is the perfect solution. As you might have guessed, the chlorine and the sodium atoms have a lower energy total state when they bond. That is what drives the bonding reaction. It has to be energetically favourable in order to happen. The reaction Na⁺ (g) + Cl^- (g) → NaCl (s) releases 787 kJ/mol of energy. Likewise, it takes 787 kJ/mol of energy to convert sodium chloride back into its two gas components. Interestingly, it actually takes a little energy to remove an electron from the sodium atom but the accepting of the electron by the chlorine atom releases energy and the attraction of the two ions together lowers the total energy of the system. This energy savings overcomes the smaller amount of energy that must be put into the reaction. For ionic solids such as sodium chloride, the bond energy (787 kJ/mol) is often called lattice energy. It's the total electrostatic potential energy of the lattice formed when these atoms bond with each other. So what is lattice structure all about?

Bonding Determines The Shape Of A Molecule

In ionically bonded materials, many of which are solids arranged in a three-dimensional lattice like sodium chloride, atoms tend to pack themselves in as close as they can to each other. This represents the lowest possible overall energy of the system. Larger chlorine ions pack in as tightly as possible with smaller sodium ions to create a tight regular lattice like this:

The atoms in an ionic lattice are in ion form, with positive cations such as sodium ions and negative anions such as chlorine ions, reflecting the fact that they exchange electrons when they bond.

Covalent bonding is a little different from ionic bonding. It happens when pairs of electrons are shared among atoms, rather than exchanged. It operates on the same general principles as ionic bonding but it achieves its favourable lower energy state a little bit differently.

Covalently bonded atoms attain stable (lowest energy) electron configurations by sharing electrons, rather than through the attraction of oppositely charged ions. Possible molecular shapes, therefore, can be much more variable than they are for ionic molecules. Organic chemistry (the study of carbon-based compounds) and stereochemistry (the science of molecular shapes) focus on covalent bonding. Consider the complex shape of a protein molecule, for example, shown below. It is an organic molecule. The peptide bond is an example of a covalent bond and it is highlighted in the box:

(Chemistry-grad-student; Wikipedia)

Covalent Bonding: Another Way to Lower Energy

Covalent bonding is possible because electrons can pair up. They do this for the same basic reason that ionic bonds attempt to fill and stabilize orbitals - to find a lower-energy, more stable state.  Like ionic bonding, electron pairing happens in the valence shell of each atom. Electrons in the same orbital must have opposite spins, but they're both negatively charged too, and they tend to repel each other. This repulsion is overcome by the lower energy state that pairing up can offer. In the diagram below, called a molecular orbital (MO) diagram, electron spins are shown as red arrows. Two unpaired electrons, with the same spins, represent a considerably higher energy state than two paired (opposite spin) electrons, shown below.

This pair can be a lone pair, not involved in chemical bonding, or a bonding pair. Covalent bonds can be created with single or multiple electron pairs, in the latter case forming double and triple bonds. Double bonds do not change the overall shape of the molecule but they are stronger than single bonds. Triple bonds are stronger yet. Remember that covalent bonding is simply a more even sharing of electrons between two atoms than ionic bonding. It is favoured when the electronegativity of the two atoms is similar or identical (as when two atoms of the same element bond together).

Carbon-Carbon Bonds are Typical Covalent Bonds

Let's examine a double bond between two carbon atoms. These bonds are often found in organic molecules. We'll take ethylene (C₂H₄) as our example to get an idea of how it works. It is drawn as a Lewis dot diagram below:

Carbon atoms have an orbital configuration of 1s²2s²2p². These atoms have four valence electrons, two 2s electrons and two 2p electrons, each one of which is available for bonding. Both the 2p and 2s orbitals are valence orbitals for carbon atoms. When carbon bonds with another atom, it has a fascinating ability to modify its orbitals to "fit together" with another atom's orbitals. The "fitting together" is actually the sharing of orbitals, which we'll explore next in Atoms Part 4C: Bond Hybridization.

Atoms Part 4C: Atoms and Chemistry - Bond Hybridization

In Atoms Part 4B, we learned that atoms form covalent bonds by sharing electrons. They do this by overlapping their electron orbitals. We took a good look at orbitals in Atoms Part 4A: 1s, 2s, 2p, 3s, and so on. Now we're ready for some pretty interesting bond chemistry - hybridization.

Remember that orbitals are described by the angular momentum of the electron. They are part of the wave function of the electron and they are solutions to Schrodinger's equations for quantum mechanics. By overlapping orbitals, electrons from two atoms are combining their wave functions to get a new shared wave function.

All the electrons involved in creating a bond must be in the same energy state and this often means that some energy must be put into the system so that an electron can move into a higher energy (excited) state to match that of its bonding companions. Ultimately, however, the creation of a covalent bond significantly lowers the energy of the system and that drives the bonding process. Carbon-carbon bonding is most often used as an example of hybridization but many molecules use it. As we will see, carbon hybridizes s and p orbitals to make bonds, but other atoms use other orbitals. For example, iron-iron (Fe-Fe) bonds are hybridized d orbitals.

First, we'll look at how carbon atoms form a double bond, in ethylene for example. Each carbon atom must first form hybrid sp² orbitals before it can bond. To bond with another carbon atom, a carbon atom modifies its 2s and 2p orbitals into three sp² hybrid orbitals and one p orbital. This is the hybridization process in a nutshell. It happens all at once but in order to understand it, we'll have to break it down into steps.

First, note that the "2" in "sp²" no longer refers to the number of electrons in the orbital. Nor does it tell us about the energy of the orbital. It now represents the type of orbital instead. An sp³ orbital, involved in a single carbon-carbon bond, or an sp orbital involved in a triple carbon-carbon bond, each has a different shape than an sp² orbital.

sp² orbitals are special hybrid orbitals. They are a mixture of p (dumbbell shape) and s (sphere shape) orbitals, so they form a geometry that's in between the shapes of the two orbitals. In each about-to-bond carbon atom, three sp² orbitals lie in a single plane with 120° angles between them. They look a bit like three lopsided dumbbells attached together at their middles, shown below left.

(Jfmelero;Wikipedia)

Remember from the previous article that carbon atoms have four valence electrons (1s² 2s² 2p²) for bonding. These three sp² hybrid orbitals are formed from three unpaired electrons by combining one 2s orbital electron with two unpaired 2p suborbital electrons (2p_x and 2p_y). The third 2p suborbital is also involved. It lies perpendicular to this plane (into and out of the page in the figure left where A is, and is not shown in the diagram) and we'll get to it in a moment.

Each carbon atom ultimately needs three unpaired electrons all in same-energy orbitals to make the three hybrid orbitals. A carbon atom doesn't have enough unpaired 2 p electrons to make a hybrid bond because it only has two of them in this orbital (see below left), so the atom promotes one 2s electron into the 2p_z suborbital, leaving the remaining 2s electron unpaired. The carbon atom is now in an excited state, below right:

Next, the lone electron orbitals hybridize into sp² hybrids, shown below. How far up the boxes are on the page indicates what the orbital energy is, so the energy of each of the hybrid orbitals is in between that of the original 2s and 2p orbitals:

When the carbon atoms approach each other to make a double bond, one sp² orbital from each atom overlaps with the other, forming what is called a sigma bond. There are different kinds of covalent bonds possible, based on orbital overlap, and the sigma bond is the strongest kind. Each of the other two sp² orbitals forms a hybrid orbital with a hydrogen atom electron orbital. Each carbon atom, in other words, sigma bonds with two hydrogen atoms.

Now to the third p suborbital: Two 2p_z suborbitals, one from each carbon atom, also overlap and they form a pi bond (perpendicular to the hybrid sp² orbitals). It too is a covalent bond but it's a bit weaker than the sigma bond. The sp² hybrid orbital overlap plus the 2p_z suborbital overlap together make a double C=C bond. The 30-second video below will help you visualize how this process works:



Carbon atoms can also form single bonds as well as triple bonds with each other. The double carbon bond is a bit shorter than a single C-C bond (an sp³ hybrid bond), another bond you will often see in organic molecules. The triple C≡C bond (an sp hybrid bond) is shorter still. A shorter bond is a stronger bond.

A single C-C bond is an overlap of one of four sp³ orbitals formed by hybridizing a 2s electron and three 2p electrons (again, one 2s suborbital is promoted to a 2p position). All four bonding electrons form sigma bonds this time. One sp³ orbital from each carbon atom overlaps to form a sigma bond between them. That leaves three sp³ orbitals that can sigma bond with other atoms. In the case of ethane, C₂H₆, the other sigma bonds attach three hydrogen atoms each. There's no pi bond in a C-C bond, so this bond is a bit weaker than the double bond.

A triple carbon bond is made with six carbon-carbon bonding electrons, instead of two (C-C) or four (C=C). In this case, a sigma bond and two pi bonds do the bonding, and it's a bit stronger (and shorter) than the double bond thanks to the extra pi bond involved. A 2s electron is promoted and hybridizes with only one p orbital electron, forming two sp hybrid bonds, shown below:

(Jfmelero;Wikipedia)

The other two p orbital electrons form two pi bonds. One sp bond from each carbon atom overlaps with the other one, forming a sigma bond. That leaves one hybrid sp orbital per each carbon atom free to overlap with another atom's electron orbital. In the case of ethyne, the other sp bond from each atom bonds a hydrogen to the molecule. While a single C-C bond can rotate about itself, the double and triple bonds cannot because p orbitals are bonded too and they have to stay aligned with each other.

Covalent bonds can be drawn simply as Lewis dot diagrams. Three Lewis dot diagrams below right show how single, double and triple carbon bonds are drawn, using ethane (C-C bond), ethylene (or ethene, C=C bond) and ethyne (or acetylene, C≡C bond) as examples:

These diagrams offer a simple two-dimensional picture of the actual shape of these three-dimensional molecules. In each case, electrons pair up in bonds to form a very stable octet of valence electrons around each carbon atom nucleus (remember that all the bond electrons count toward each carbon atom's octet because they are shared).

Hybrid bonds give us a peek into what is really going on with an atom's electron configuration when it covalently bonds with another atom. It gives us a feel for the mechanics of how atoms share electron orbitals. We'll look at two particular kinds of covalent bonds - polar bonds and metallic bonds, two different takes on electron orbital sharing - in Atoms Part 4D, next.

Atoms Part 4D: Atoms and Chemistry - Polar Covalent and Metallic Bonds

POLAR COVALENT BONDS: THE IN-BETWEEN BOND

In Atoms Part 4B, we learned that covalent bonding is favoured over ionic bonding when atoms are of similar electronegativity. Electrons are shared within electron pairs, forming a stable balance between attractive and repulsive forces. Electrons from both atoms are attracted to the two positively charged nuclei, while the electrons themselves, being negatively charged, experience a repulsive force between each other. Electrons push each other away like two south poles of a magnet. The reason they pair up is so that the most electrons possible can fit into the lowest possible energy orbitals, and by pairing up, electrons can share an orbital. This lowers the overall energy of the atom so it is favoured.

Covalent bonding doesn't necessarily mean, however, that the electron clouds involved are perfectly evenly dispersed across the molecule, a phenomenon called electron delocalization. When two or more atoms with different electronegativity bond covalently, the electrons may be shared unequally between them, creating areas with slightly different charges. These bonds are called polar covalent bonds. If the polarity is extreme, the bond created is an ionic bond. Polar covalent bonds, therefore, fall midway in the spectrum between non-polar covalent bonds and ionic bonds.

Water Has Strongly Polar Covalent Bonds

Polar bonds create tiny accumulations of charge. Atoms with high electronegativity, such as oxygen (3.44), pull on electrons strongly, so when oxygen bonds covalently with other atoms, such as hydrogen (2.20), it often forms strongly polar molecules. A water molecule is an example, shown below left as a Lewis dot diagram:

Notice the two lone pairs of electrons around the oxygen atom. The molecule appears bent because the two positively charged hydrogen atoms are flexed away from the negatively charged filled oxygen orbitals. Negative charge accumulates around the central oxygen atom (red), leaving the outer parts of the two hydrogen atoms with a net positive charge (blue) in the diagram below:

Polar molecules such as sugar, and ionic solids, such as sodium chloride, dissolve easily in (polar) water. Oils and waxes are nonpolar molecules, so they don't dissolve in water but they will dissolve in very nonpolar solvents like cyclohexane, which is a ring of six carbon atoms covalently bonded together. Charge is distributed very evenly in this molecule. "Like dissolves in like" is the common rule of thumb here. Cyclohexane, not surprisingly, won't dissolve in water. Instead it will form a separate distinct layer, just like oil does.

Water is a very fascinating polar molecule. I started off the "Our Universe" articles by exploring water. At very high temperatures and pressures, inside Jupiter for example, water transforms into a fully ionic state - a soup of positive hydrogen ions and negative oxygen ions. At even higher pressures, water will become a superionic solid. Oxygen crystalizes out of the soup and hydrogen ions float freely within the oxygen lattice.

One Molecule - Two Kinds of Bonds: Polyatomic Ions

In Atoms Part 4A, I mentioned phosphorus [P] as an example of a half-filled valence orbital. This makes phosphorus reactive - it wants to fill that orbital and find a more stable, lower-energy, argon-like electron configuration. One way it can do this is by reacting with four oxygen [O] atoms to form a complex ion, PO₄^3-. This ion is called phosphate. It is one of many different kinds of polyatomic ions, ions made of more than one kind of atom. The electronegativity of phosphorus (2.29) and oxygen (3.44) are close enough that they form covalent bonds, but the values are far enough apart that the covalent bonds are polar in nature. Negative charge accumulates around the oxygen atoms. A phosphate ion drawn as a Lewis structure is shown below left:

Lewis structures are often simplified using lines for bonds instead of drawing electron "dots." As we saw in the last article, a single line means a single bond (two dots). A double line means a double bond (four dots). This polyatomic ion (left) acts as a single unit that has an overall charge of -3. Like all covalently bonded molecules, phosphorus can lower its energy by sharing its five valence electrons with oxygen atoms. Do you notice something odd about the phosphorus atom here? It is breaking the octet rule. It has more electrons around itself than you'd expect - ten, rather than eight. Elements in the third period and below can accommodate more bonds than the rule allows, an exception to the octet rule.

Oxygen atoms (3.44) are more strongly electrostatic than phosphorus (2.19) (see the electrostatic periodic table far below), so each oxygen atom tries to draw two electrons to itself. If you look at the periodic table (below) you'll see that oxygen is two electrons short (eight) of the noble gas atom, neon (ten).

One oxygen atom uses two of its six valence electrons to form a double bond. The other four valence electrons of this oxygen form two lone pairs. The other three oxygen atoms can't do the same thing. Phosphorus only has five valence electrons available for bonding, so these oxygen atoms each attract one phosphorus electron into a single bond. In doing so, they are one electron short of an octet.

They fix this unstable situation by attracting an extra electron (which can be from water or another atom). In doing so, they each attain a single negative charge. It is energetically favourable for any atom to pair up its electrons. The three negative charges mean that the phosphate ion has a charge of -3.

In reality, phosphate doesn't arbitrarily choose one oxygen atom to double bond with and then stick with this particular atom, creating one shorter bond and three longer bonds. These bonds display resonance instead. All three bonds are the same length and strength - a hybrid between a double bond and a single bond.

The formation of a polyatomic ion with phosphorus is energetically favourable to oxygen atoms but it leaves three oxygen atoms one electron short of a stable full valence shell, a better but still energetically unfavourable situation. The complex ion has achieved its goal of saving energy by electron sharing but it could improve things even more. Three of the four oxygen atoms are ready to react . . .

A Brief Introduction To Acids And Bases

Many atoms will react with this polyatomic ion to form an ionic compound. Hydrogen [H] found in dissociated form in water is one example. Three hydrogen atoms will bond with phosphate's three oxygen atoms to create phosphoric acid, H₃PO₄. Phosphoric acid is an ionic compound containing a highly polar complex ion - phosphate. It is a crystalline powder at room temperature that dissolves easily in water, and when it does, just like table salt, it dissociates into two ions:

H₃PO₄ ↔ H⁺ + H₂PO₄-
H₂PO₄^- ↔ H⁺ + HPO₄^2-
HPO₄^2- ↔ H⁺ + PO₄^3-

Notice that I've drawn three reactions instead of just one. Phosphoric acid doesn't just dissociate once in water, it dissociates three times, each time releasing an H⁺ ion. H⁺ ions react with water to create H₃O⁺ ions. The concentration of H₃O⁺ ions, called hydronium ions, determines the pH of a solution. As the concentration of H₃O⁺ increases, the phosphate solution becomes more and more acidic, until eventually equilibrium is reached. This is basically how an acid works.

If phosphate is added to a solution that is very acidic, around pH 2, almost all phosphate ions will be in the PO₄^3- form. If phosphate is added to pure water, around pH 7, you will get a mixture of phosphate ions, most of which will be HPO₄^2-. If you add phosphate to a very basic solution (pH 12) you will get mostly H₂PO^4- ions.

All hydrogen phosphate (ionic) compounds dissolve readily in water but the phosphate complex ion, being covalently bonded, does not. It stays intact.

Bonding Determines the Chemical and Physical Properties of Molecules

The strong electrostatic bonds in ionic compounds give these compounds typical chemical and physical properties. As we've seen with sodium chloride and phosphoric acid, ionic compounds tend to dissolve easily in water. They also conduct electricity very well in solution with water, as well as when they are in a molten state, but not in their solid state. Ionic compounds have very high melting and boiling points. They are almost always found as solids at room temperature, usually as tightly packed lattices or crystals. This tight packing means charged ions or electrons can't move easily and that is why they don't conduct electricity as solids.

The physical and chemical properties of covalent compounds tend to vary much more than they do for ionic compounds. These compounds may be solids, liquid or gases at room temperature. They are not usually electrically conductive. Polar covalent compounds are soluble in water but nonpolar ones are not. Examples of covalent compounds are water, carbon dioxide gas, and diamond.

This nine-minute video by Montana science teacher, Paul Andersen, reviews the covalent and ionic chemical bonds we've explored in a very user-friendly way:



METALLIC BONDING: A SPECIAL KIND OF COVALENT BOND

Metals make up a large part of the periodic table. You are probably familiar with metals like chromium [Cr], nickel [Ni], silver [Ag] and aluminum [Al], for example, but many other elements are also technically metals. In the periodic table below, every element left of the black line as well as the lanthanides and actinides are metals. This does not mean, however, that all of these elements form metallic bonds (I'll describe them in a moment) under normal conditions. Hydrogen, for example, is a gas (H₂) under normal pressure. It forms polar covalent bonds.

Metal reactivity tends to decrease as you move right on the periodic table shown below. Tin [Sn] is not very reactive but sodium [Na] is.

Metallic bonds, in which atoms of one type of element bond together, tend to be quite strong. That's why most metals have high melting and boiling points. Chromium's melting point is 1907°C, nickel's is 1455°C, silver's is 962°C, and aluminum's is 660°C.

Mercury [Hg], shown below, is a very interesting exception. It is a liquid at room temperature. It freezes into a solid below -38°C (that's its melting point) and it boils at 360°C.

(Materialscientist;Wikipedia)

Mercury has a unique electron shell configuration. It's a transition metal, but unlike most transition metals, which have partially filled d orbitals, mercury's electrons fill up all of its orbitals - [Xe] 4f¹⁴ 5d¹⁰ 6s² (we haven't covered the f orbital yet, it can contains a maximum of 14 electrons in it; notice that the d orbital with 10 electrons is full).

Mercury behaves more like a noble gas as a result of it s filled orbitals, so when its atoms bond with each other, they are unusually weak bonds. Mercury's valence electrons are not very reactive so its atoms do not share electrons readily with one another. Metallic bonding occurs but it is weak, so mercury melts far more easily than most other metals.

Metals (except mercury) form tightly packed lattices of atoms, much like ionic solids do, but these solids, unlike ionic solids, conduct electricity very well. Valence electrons are very mobile in these lattices because they are delocalized. This means that their orbitals extend over many adjacent atoms.

Sodium [Na], shown below, is a good example.

(Dnn87 at en.wikipedia)

It has the electron configuration 1s²2s²2p⁶3s¹. When two sodium atoms approach each other, the lone 3s orbital electrons pair up, just like how other covalent bonds form. The difference here is that each sodium atom is surrounded by eight other sodium atoms in a tightly packed lattice. The 3s orbital electron in the central atom overlaps with the 3s orbitals of all the surrounding atoms. Each orbital can only hold two electrons so this means there is a vast number of shared orbitals extending all over the sodium metal. There is no single metallic bond. They are all interconnected, so electrons in these orbitals can move freely all over the metal, within the shared orbitals.

Magnesium [Mg], shown below, has two 3s orbital electrons.

(Warut Roonguthal;Wikipedia)

Its delocalized electron "sea" has twice as the electron density of sodium. It has even stronger bonds and a higher melting point (650°C versus sodium's 98°C) as a result. It is also about twice as electrically conductive as sodium.

The delocalized electrons in transition metals have a greater density still, because both their 3d and 4s orbital electrons can take part. These metals are located in the light peachy pink central square in the periodic table above, called the d-block. The melting and boiling points of these metals tend to be very  high. My earlier examples of chromium, nickel and silver are transition metals but aluminum is not. Some of these metals are very electrically conductive but the trend based on electron density isn't always simple. Silver is the most electrically conductive metal of all, but iron [Fe] and platinum [Pt], also transition metals, are not very electrically conductive at all.

When metals melt, the metallic bond is still there - the electron orbital sharing still occurs - but the tight lattice structure breaks down and the bonds are weakened.

Under pressure, elements that are technically metals begin to show their metal character. The hydrogen atom I mentioned earlier is a gas in our atmosphere. It shows no metallic bonding at all. The two electrons are completely localized in their orbitals and the two hydrogen atoms bond covalently. But under extreme pressure, again scientists think this might happen deep inside Jupiter, hydrogen is a metal. The hydrogen atoms become so tightly packed together they begin to from a liquid metal and then with increasing pressure they form a solid lattice arrangement, and in this tight arrangement, electrons begin to delocalize. They are unbound to their proton nuclei and are free to move about in shared orbitals, making crystalline hydrogen metal very electrically conductive. (This could be the reason Jupiter has such an enormously powerful magnetosphere.)

Metals are strong, malleable, lustrous and opaque and they conduct electricity, all thanks to metallic bonding.

Science teachers Jonathan Bergmann and Aaron Sams explore metals and  how metallic bonding works in this six-minute video:



Next, we'll draw some brief conclusions about atoms, their electrons and their chemistry in Atoms Part 5E.