The Periodic Table of Elements V: _Periodicity

Beyond atomic mass

Dmitri Mendeleev, the Russian chemist whose work is discussed extensively in our Periodic Table II module knew a bit about fluorine when organizing his table in 1864. He did not know about atomic structure, like protons, neutrons, and electrons – and “atomic number” wasn't even part of the vocabulary of chemists at the time. What Mendeleev did know is how fluorine behaves when combined with other elements and its high reactivity. Chemists describe how an element tends to react with other elements using the term “chemical properties.”

Chemists of Mendeleev's time also understood the concept of atomic mass. Mendeleev tried organizing his elements according to mass, but it didn’t quite work. For example, chlorine is much heavier than fluorine but shares similar chemical properties. On the other hand, several elements with masses closer to fluorine’s do not share its properties. That’s when Mendeleev combined the atomic mass approach with chemical properties.

Mendeleev certainly wasn't the only chemist of the mid-1800s working to categorize the elements. John Newlands, a British chemist, also tried grouping elements by chemical properties and increasing atomic mass. And Newlands noticed an odd pattern. When ordered by increasing atomic mass, it almost seemed that chemical properties repeated every eighth element.

“Preposterous!” Newland's colleagues said. “You may as well organize them by alphabetical order,” they joked. It didn't help Newland’s case that his pattern seemed to fall apart after the element calcium (Ca).

Still, Newland was onto something. Today, our modern periodic table shows the gradual change in chemical properties that reset at the eighth element. We now know that this pattern stems from how electrons fill up electron shells, as discussed in our Periodic Table III: Electron Configuration module. Each row of the periodic table – called a “period” – shows this chemical property reset. That’s why we use the word periodicity to describe these changes across the table.

We know that elements above one another – like chlorine and fluorine – share similar properties. And elements near one another in a period, are similar as well. Using this, we can broadly group the elements into a few different classifications.

Metals vs. nonmetals

Let’s take two of the most dangerously reactive elements from the periodic table: cesium and fluorine. You’ll never stumble upon either in its pure form. At least, not in nature you won’t. Cesium and fluorine react so readily that we only unearth them after they’ve formed chemical compounds. But scientists have mixed them in the lab – in only in tiny quantities, though. Together, they burst into white fire. It is one of the most vigorous reactions possible between pure elements.

Think of cesium and fluorine as chemical opposites. They represent the most reactive of two major classes of elements: metals and nonmetals. These two major categories of elements are introduced in our module Periodic Table IV: Chemical Families. But here’s a quick summary.

Atoms work to achieve a stable noble-gas electron configuration through the give, take, or sharing of electrons. Refer to Periodic Table III: Electron Configuration for more about such configurations.

To chemists, “metals” are elements that tend to achieve stable electron configurations by giving away electrons, while “nonmetals” are elements that tend to achieve stable electron configurations by taking electrons from other atoms.

Think of metals and nonmetals as a spectrum. Each element lies somewhere between the most metallic (cesium) to the most non-metallic (fluorine).

Tracking metallic character across the periodic table

Take a peek at a periodic table (Figure 1). Metal elements sit on the left side, nonmetal elements on the right. Chemists use the term “metallic character” to describe where an element falls on the metal-to-nonmetal spectrum (Figure 3).

From right-to-left across the periodic table, metallic character increases. In other words, from left-to-right, metallic character decreases. That means elements become more nonmetal-like as we move to the right. Using a more technical chemistry definition, this means elements progressively become more likely to pull electrons from other elements. But wait, this is important. The term “metallic character” sometimes confuses students, especially when chemists use the term “metallic character” to describe nonmetal elements. But it’s simple, really. All elements have metallic character. Elements with high metallic character are called metals. “Nonmetals” are just elements low in metallic character.

Metallic character trends up and down the table as well. When scrolling down the periodic table, metallic character increases. Chemically speaking, this means elements further down the periodic table lose their electrons more easily than those near the top.

Cesium is a silvery-white metal that is liquid at room temperature with a consistency like melted candle wax. Find cesium on the periodic table’s bottom-left. Only francium (Fr) – a lab-made element that doesn’t occur in nature – sits lower. That position marks cesium as the naturally occurring element highest in metallic character.

The most elements lowest in metallic character occur toward the top-right, with the exception of the noble gases. The noble gases neither give nor take electrons due to their stable electron configurations.

Fluorine: electron gobbler supreme

In 1907, roughly a century before McKay's fluorine-related studies in Colorado Springs, Henri Moissan accepted the 1907 Nobel Prize in Chemistry. Two months later, the 54-year-old French chemist died. “Acute appendicitis” states his death certificate. But many scientists, including Henri himself, attribute his early death to damage caused by his 20-year pursuit to isolate fluorine in its purest form, F₂, or fluorine gas. (When chemists say “fluorine,” they refer to F₂).

Nonmetal elements readily take electrons. Fluorine – positioned nearly top-right – earns its position as the most aggressive electron-gobbler of them all.

That “savage beast among the elements” is how Peter Klason, Nobel prize awards president, described fluorine in 1906. For good reason, too. Fluorine rips away electrons from almost all other elements. Even highly unreactive substances – such as the mineral asbestos, which is used as a flame-retardant because it is so unreactive – can ignite upon contact with fluorine.

Moissan wasn’t the first fluorine-focused chemist to suffer a suspicious or unnatural death. No other element has left such a body trail of chemists across the world trying to isolate the element’s pure form. Belgian chemist Paulin Louyet and French chemist Jerome Nickels died from inhaling the dangerous gas. English chemist George Gore blew up his lab and nearly himself.

A tale of two metals

Periodicity refers to how chemical properties trend when moving up and down the periodic table. Knowing these trends allows us to compare two elements. Let’s consider an example of two metals shown in Figure 2: sodium (Na) and cesium (Cs).

Sodium and cesium both have high metallic character, as they both reside in column 1. Cesium sits lower than sodium, and therefore exhibits higher metallic character. Or to put it another way, sodium is the less metallic of these two metals.

Now consider an example of two nonmetals shown in Figure 3: fluorine (F) and chlorine (Cl).

Both F and Cl appear in column 17 of the table, so we know they have low metallic character and tend to steal electrons. But which steals electrons better? Due to its higher periodic placement, the answer is fluorine.

Metallic character goes a long way toward predicting the reactivity of elements. However, why do elements differ in such a way? Electron configuration explains in part why different atoms react as a metal or nonmetal. For instance, fluorine – being only one electron short of a noble gas electron configuration – takes electrons more readily than sodium, which has one electron too many. But electron configuration alone doesn’t explain why fluorine reacts more readily than others in its column. Fluorine (F) and iodine (I) are both nonmetals, but fluorine’s violent reactivity greatly exceeds that of iodine.

Despite metallic character being a helpful description of an element’s reactivity, it isn’t a fundamental property. In other words, metallic character arises from the combined effects of other properties, three of them. Those three are atomic radii, ionization energy, and electron affinity.

Atomic radius: how can we know an atom’s size?

Let's start with atomic radius, or atomic size, which is the distance from the center of an atom to its outer edge.

If we can't see an atom, how can we possibly measure its size? X-ray crystallography is a tool that allows us to do this. In the early 19th century, the German physicist Max von Laue was experimenting with a curious form of radiation called X-rays. Von Laue and colleagues noticed that X-rays passed through solids, including crystals. The resulting diffraction pattern (the bending of waves when they meet an obstruction) differed depending on the type of crystal that they were studying.

A father-and-son team of British physicists, William Henry Bragg and William Lawrence Bragg, later showed that the pattern reflected the distance between elements bonded together in the crystal. By measuring this distance in crystals in which similar atoms were bonded together, they were able to determine the actual size of atoms. The Braggs published their measurements of atomic size, or “atomic radii,” in 1920.

Findings from atomic radii studies revealed an unexpected trend. Atomic radius decreases when moving from left to right across a period, which is a row on the periodic table.

For example, let’s compare fluorine (F) and lithium (Li). Which do you think is the larger atom? Consider what you know about the differences in these two elements (Figure 4), then venture a guess.

Fluorine, with an atomic number of 9, has greater atomic mass. Fluorine contains more of every type of subatomic particle (protons, neutrons, and electrons) than lithium. For many people, it seems natural to assume fluorine is the bigger atom with the widest atomic radius. But surprisingly, it’s not.

Despite having fewer subatomic particles (like electrons and neutrons), lithium atoms are bigger than fluorine atoms. And it’s not even a close call. Scientists measure atomic radii using the unit picometers (pm). As you can probably guess, a picometer is tiny. There are 1,000,000,000,000 (one trillion) picometers in a meter. Lithium's atomic radius is 145 pm – almost three times larger than fluorine atoms, which measure only 50 pm. How could this be? After all, fluorine has nine electrons and protons compared to lithium's three.

Let's learn how fluorine, with 3 times higher mass than lithium, can also be only a third of lithium’s size.

A periodic paradox

Fluorine lies to the right of lithium on the periodic table. As you move down the periodic table, atomic radii tend to increase. And that makes sense, because atomic mass increases further down the periodic table. However, as you move from left to right, atomic radii tend to decrease (Figure 5) – even as atomic mass increases.

Be careful. It’s not always the case that higher atomic mass means smaller size. To understand this paradox, we first must understand what gives an atom its size. (Hint: It may not be what you think!)

Protons and neutrons make up 99.9% of an atom's mass. They’re also the largest of the three subatomic particles. Electrons are about 2,000 times smaller and less massive than protons or neutrons. But surprisingly, protons and neutrons take up almost none of the atom’s volume. That’s because most of the volume an atom takes up is empty. In other words, atoms are mostly empty space. Imagine a single atom but one the size of a baseball field. How big might its nucleus be? The answer may surprise you. A typical nucleus would be about the size of a green pea!

That mostly empty space around the nucleus is where the tiny electrons reside, as Ernest Rutherford discovered with his iconic 1911 gold foil experiment. Read more about Rutherford’s work in our Atomic Theory I module. Rutherford’s work gave rise to new atomic model, one in which a teeny-tiny nucleus of protons and neutrons packs 99.9% of the atom’s mass. Outside that tiny nucleus lay the vast, almost empty electron fields.

In other words, electrons may be smaller than protons, and less massive, but their orbits still define an atom’s outer bounds. The distance electrons reside from the nucleus determines an atom's radius.

Any force that pulls electrons closer to the nucleus also shrinks the atom. Conversely, any force that pushes electrons further apart enlarges the atom.

An electron tug-of-war

At all times, electrons are experiencing both push and pull forces. The “pull” comes from the nucleus. Positively charged protons attract electrons, pulling them inward and shrinking the atomic radius. The more protons an atom has, the stronger that attraction (called the “Coulomb force”). Thus, atoms with more protons tend to be smaller in size.

At the same time, electrons repel one another. (That’s the “push” force.) And the closer electrons come to the nucleus, the closer electrons draw to one another and the stronger those repulsions grow. It may help to think of each atom’s electrons as locked in a never-ending struggle – a tug-of-war – between attraction and repulsion forces.

Pull forces increase gradually and consistently across a period. A higher atomic number means stronger attraction for electrons, which increase gradually from left to right across a period.

But periodicity trends reveal that push forces do not change gradually. Electron repulsion forces remain pretty steady across a period. Then it leaps up at the start of the next period.

For elements of the same period, such as lithium and fluorine, that means that both experience the same electron repulsion push forces but different pull forces. Fluorine’s nine protons exert a stronger attractive pull on its electrons than lithium’s three protons. The result: Fluorine draws its more electrons in closer than lithium. Therefore, fluorine is the smaller atom.

Once we move down a period, electron repulsion takes a big jump. Why does electron repulsion remain consistent across a period, but skyrocket at the start of a new one? The answer relates to atomic structure, especially the positioning of electrons around a nucleus.

Recall from our Periodic Table I module that individual electrons reside in electron shells. Those electron shells “fill up” with electrons in a predictable pattern. This pattern is described in detail in our Periodic Table III module. But in summary, the first electron shell fills up first, maxing out at two electrons. Overflow must go to the second electron shell, which fits eight electrons, as does the third electron shell.

On the periodic table, each period corresponds to the filling-up of a specific electron shell, so all elements of the same period are working on filling the same outermost electron shell. This shell is so important that it gets its own name: the valence shell.

When an inner electron shell is full, it blocks some of the nucleus’ attractive force from outermost electron shells. This effect is called “electron shielding.” It doesn’t shield electrons entirely. If it did, those electrons wouldn’t remain in the atom; they’d leave. But electron shielding does weaken the nucleus’s hold enough so that the electrons behind the shield distance themselves out a bit.

Only full electron shells have this shielding ability – it’s almost all or nothing. Electrons within the same shell do not shield one another. That is why electron shielding stays mostly the same across a period.

Atoms shrink as you move from the left to the right because electron shielding remains about the same across a period, but pull forces from the nucleus consistently increase. And that cinches these shells in tighter, making a more compact atom. Take note that this means that electron shells are not a single fixed size. Two atoms can have the same number of electron shells, but those shells may be spaced differently.

As you move down the periodic table, atomic radii tend to increase. This makes sense because more electron shells mean increased electron shielding between those shells. That causes electron shells to space apart and increase the size of the atom.

Since atomic radii increase moving down and to the left, the largest elements must occur at the bottom-left. Cesium (Cs), with an atomic radius of 265 pm, boasts the widest radius of the naturally occurring elements (Figure 6).

Keep in mind we’re discussing general periodicity trends, not absolutes. Exceptions exist: the noble gases and transition metals, for instance. Transition metals are notable in their straying from many periodic table property trends. This unpredictability stems from their valence shell electrons occupying a sub-shell known as the “d orbital,” discussed in our Periodic Table III module.

A dental health mystery

Not everyone who lived in Colorado Springs in the early 1900s developed the “Colorado Brown Stain.” Dentist Frederick McKay soon found a connection between those who did: They all had lived in Colorado Springs when their permanent teeth emerged and calcified. People who moved to the region after that critical age period did not develop the stains. This suggested that the stains were due to something mineralizing into the developing teeth.

McKay (Figure 7) suspected that “something” was in the water. But it would take him the next 30 years to figure out what. In 1923, the residents of another town, Oakley, Idaho, reached out to McKay for help with the strange brown stains that appeared on their children’s teeth. The residents pointed out to McKay that the addition of a new water pipeline had marked the emergence of those stains. That path would lead McKay to fluorine and its enamel-strengthening abilities.

To avoid confusion, it’s fluoride, not fluorine, that appears in modern toothpaste. Fluoride is the negative ion, or anion, of fluorine. Too much fluoride causes discoloration of teeth in addition to the benefit of strengthened enamel. Later studies, however, would reveal that lower amounts of fluoride still granted the benefits while avoiding the staining side effect.

McKay's findings arrived at a time when dental decay in the United States was surging. But what had changed to cause such an increase? The answer is our diet. In particular, the rise and spread of the dental health supervillain processed sugar.

A pattern emerges if we track sugar throughout human history: Increased sugar availability leads to increased tooth decay. A research team from Norway followed one such change among the Alaskan Inuit: Before 1920, the Alaskan Inland Inuit of the Anaktuvuk Pass, Alaska (Figure 8), lived the same 2,000-year-old lifestyle as their ancestors. They survived mainly by fishing and hunting the caribou, native birds, and sea mammals of the Alaskan coast. The Norwegian team evaluated the community's diet and dental health starting in 1955. And their data show that most of the residents’ calories came from fat and protein. They ate almost no sugar. And despite limited access to dental care, most residents had no cavities.

But everything began to change.

Starting in 1953, a post office and "white trader store" had moved into the village. The next few years saw rapid upticks in visitors and vacationers. Souvenir shops began to open, as did the availability of processed food.

When the Norwegian team followed up in 1965, they found that carbohydrate consumption (such as sugar) had spiked by 50%. The native diet now comprised only 20% of the average resident's diet. The rest consisted of the refined, processed foods they purchased from the new local stores.

This caused a big change in the Alaskan Inuits’ teeth. Only a few years earlier, 74.5% of the residents had no cavities at all. But now, that number had dropped to zero. Everyone showed some degree of tooth decay.

The problem was sugar. Sugar stimulates the growth of acid-oozing bacteria inhabiting our mouths. Those acids strip calcium minerals from enamel, leaving teeth microscopically porous and vulnerable to even faster breakdown and decay. Our teeth need something that can cling to enamel's calcium and minerals, even in the face of enamel-leeching acids. And that something is fluoride.

Small but fierce: how fluorine’s tiny size leads to huge reactivity

For nonmetals, a smaller atomic radius leads to higher reactivity. And fluorine is tiny. The reactivity of fluorine – and by extension, fluoride – proves one of few forces strong enough to combat the enamel-leaching acids brought on by sugar consumption. (Fluoride isn’t as small as its neutral counterpart, fluorine, but it’s still tiny for an anion.) Fluoride's high nuclear charge and small size grant it a powerful pull on any positively charged particles in the vicinity. When mineralized in enamel, fluoride snags onto the positively charged calcium ions that make up enamel. That strengthens enamel. And it’s all due to fluorine being the most reactive nonmetal of the periodic table.

But wait. What about the metals? Do atomic radii affect the reactivity of the metal elements? The answer is yes. But there’s one big difference. For metals, a higher atomic radius leads to higher reactivity, which is the opposite of what we see for nonmetals. Recall that by chemical definition, a “metal” element is an element that tends to lose valence electrons during chemical reactions. And valence electrons far away from the attractive forces of the nucleus are given up more easily. So metal elements become more reactive as atomic radius increases.

Ionization energy

Scientists today can measure how easily an element loses electrons. It’s called ionization energy (IE), and it underpins metal vs. non-metal character. Ionization energy is the energy required to remove an electron. (The gaining of an electron will be covered in the next section.)

To measure an element’s IE, chemists start with a sample of that element in vapor form. Then they blast the sample with high-velocity electrons, gradually increasing the voltage until the sample’s loosest bound electrons get kicked out of their electron shell, forming an ion. The chemist can tell an electron’s been booted by monitoring the sample's conductivity (its ability to carry an electric current). Once the sample becomes ionized, the conductivity spikes. Removing an electron leaves the atom more positively charged, forming a cation. So, it may help to think of ionization energy as "cationization energy." As for units, IE is described in kilojoules per mole (kJ/mole), which is the amount of energy it takes for all the atoms in one mole of that element to lose one electron each. The energy to remove the first electron is called 1st ionization energy, or 1st IE.

Chemists came up with the concept of IE years before it was actually measured. The 1925 Nobel Prize winning work by two German physicists, James Franck and Gustav Hertz, opened the door to the first IE measurements. The pair used the process described above on a sample of mercury (Hg). Their data supported the notion that a certain amount of “binding energy” held each electron to the nucleus. By blasting an atom with that binding energy, they found an electron could be dislodged. Take note that removing electrons always costs energy. Even metals, which by definition lose electrons easily, require energy input to lose an electron. Metals require less energy than nonmetals, though. Therefore, metals have lower IEs than nonmetals.

Ionization energy trends across the periodic table

Just like periodicity trends allow us to predict the atomic radii for elements, we can predict IE from periodicity trends too. IE increases from left to right across the period table (Figure 9).

Metals have the lowest IE values. Their electrons are the easiest to remove. The noble gases in column 18 show the highest IE values. Due to their stable electron arrangements, the noble gases hold tightly to their electrons, leading to the high IE values.

The same push/pull forces that influenced atomic radii also give rise to IE periodicity trends. Moving to the right across a period, the electron shielding force remains about the same, while attraction “pull” forces from the nucleus increase consistently. That stronger attraction for electrons makes it increasingly difficult to pluck away electrons. That’s why elements on the right side have a higher 1st IE than those on the left.

Once again, let’s make an example of sodium (Na) and fluorine (F), which both lie in the 3rd period. Sodium’s “first” IE would be described as follows:

Now, compare it with chlorine’s first IE:

We already know that nonmetals hold tighter to electrons than metals. IE is just a way to measure that difference. A neutral chlorine atom requires about 2.5X more energy to kick away an electron than does sodium.

Now let’s consider up-down periodicity trends for IE. When moving down the periodic table, IE decreases. Using what we already know, we can understand why. Moving down the periodic table means more electron shells, and therefore, more electron shielding. The same electron shielding force that relaxes the nucleus’s hold on electrons, allowing atomic radii to expand, also makes it easier to pluck away those valence electrons. That’s why IE decreases as you move down the table.

For example, potassium (K) sits just below sodium (Na) on the periodic table (Figure 10).

Potassium's first IE measures 419 kJ/mol, less than sodium's 496 kJ/mole. One step further down brings us to rubidium (Rb), which measures 403 kJ/mole. Based on these trends, the elements with the lowest first IE values should lie bottom left. And cesium – king of metals – measures just 375 kJ/mol.

To the second ionization energy, and beyond!

Removing electrons gets harder each time. As a neutral atom becomes a cation, it holds tighter to the remaining electrons. So, second IEs are always higher than first IEs.

Consider the first IE for the element magnesium.

After magnesium’s first electron is removed, the remaining electrons are held more securely. That’s why magnesium’s second IE nearly doubles.

Magnesium’s third IE would be higher still, and so on.

Furthermore, certain steps from one IE to the next show gigantic leaps. This occurs when an ionization event results in an atom achieving a stable electron configuration, but then that same atom is forced to give up another electron.

For example, let’s return to sodium. Sodium has one valence electron, which occupies the third electron shell by itself. Losing that electron empties sodium’s valence shell. And that leaves sodium’s cation, Na+, with a stable electron configuration – the same as the noble gas neon (Ne).

But now, let’s consider removing sodium’s second electron, which is a non-valence electron or a core electron. Not only that, but remember sodium isn’t neutral anymore. Na⁺ is a cation, and that positive charge makes Na⁺ cling to its electrons even tighter. The result is this: Sodium’s first IE may only require 496 kJ/mole. But sodium's second IE is a tenfold leap!

Compare sodium's jump from its first IE to its second IE with magnesium's.

Magnesium's jump is a modest twofold increase. This is because magnesium has two valence electrons to lose before it reaches core electrons.

Electron affinity

Electron affinity (EA) and ionization energy are similar concepts, but in some ways they are opposites. Electron affinity is the energy change that occurs from adding an electron to an atom. It was first conceptualized in the early 1900s by American chemist Linus Carl Pauling in connection with his work to better understand the nature of chemical bonds.

Moving to the right across the periodic table, elements tend to accept electrons more readily than they lose them. And EA values reflect this trend by becoming more negative as you move across the table.

In contrast to EA, remember that IE refers to the energy input needed to remove an electron. Removing electrons always requires that energy be added, even for metals that lose them more easily than nonmetals. Because energy is added, not released, IE values are always positive values. Conversely, EA almost always involves release of energy as the atom accepts an electron. Because energy is usually lost, EA values are almost always negative.

In Figure 11, higher columns indicate more negative EA values. Generally, we see that EA values grow gradually more negative as you move to the right. Many exceptions balk this trend, however. The noble gases and a few other exceptions actually require energy input to accept electrons. Those unusual elements have positive EA values.

After exploring the more obvious periodicity trends for atomic radii and IE, the EA trend might appear more randomized. EA tends to increase when moving from left to right, but when moving down the periodic table, obvious trends vanish.

Fluorine and the other halogens show the most negative EA values, as you might expect. Take note of fluorine’s EA below.

Interestingly, fluorine’s EA is less negative than chlorine’s.

That’s unexpected. After all, we've talked a great deal about fluorine's reactivity and enamel-hardening capabilities. We’ve discussed how readily the tiny fluorine atom accepts electrons to become the anion, fluoride (F-). So shouldn’t fluorine’s EA value reflect that?

The reason for fluorine's less negative EA value derives from, again, fluorine's tiny atomic radius. Fluorine's atomic radius widens by accepting an electron like any other anion. But fluorine's radius changes less during the fluorine-to-fluoride transition than chlorine’s does during the chlorine-to-chloride transition. Therefore, fluorine's change in potential energy is slightly less than chlorine’s.

Applying the knowledge

McKay spent over thirty years working with the residents of Colorado Springs. In 1945, Grand Rapids, Michigan, became the first city to add fluoride to public tap water. Keeping the fluoride concentrations low still grants the benefits but helps avoid the undesirable “Colorado Brown Stain” effect. Over the years, other states followed Michigan’s example. Fluoride in tap water improves dental health. Today, regardless of socioeconomic class, millions of children benefit from fluoridation programs.

Fluoride’s chemical properties, such as atomic radii, ionization energy, and electron affinity, contribute to its enamel-strengthening power. And from fluorine’s placement on the periodic table, we can predict these properties today.

The periodic table allows us to predict the properties of elements. To use this chemistry cheat sheet, you just need a little knowledge of periodicity trends. Take any random element, and just from its placement on the periodic table, we can predict what kinds of bonds that element tends to form and how vigorously that element reacts. That’s what makes the periodic table one of chemistry’s most useful inventions.