The arrangement and behavior of electrons bound to atoms gives those atoms all of their physical properties. Because the periodic table of the elements is organized according to those electronic properties, its structure can reveal many important trends in those physical properties. It turns out that at a glance you will be able to tell whether one atom has more or less of a given property than another, and that makes the periodic table an even more powerful tool.
These properties are:
*Don't worry if you don't understand some of these terms; we'll cover them below.
While the "size" of an atom is a bit of a soft target – atoms are "fuzzy" and difficult to measure, say in diameter or radius – it's also fair to say that atoms have relative size. That is, some atoms are larger than others, and atoms can be ranked by relative size.
While electrons don't take up much space (the radius of an electron is far smaller than the radius of a proton, which is also pretty small, though neither has been very accurately measured), they don't crowd together very well because their like negative charges keep them apart. More electrons means a bigger atom.
The number of protons and electrons in a neutral atom is equal, but because of the odd shapes of electron orbitals and other factors, increasing the number of protons in the nucleus effectively holds electrons more closely and more tightly to the nucleus. This effect is somewhat diminished in larger atoms, where outer electrons can be shielded from the nucleus by a screen of inner-shell electrons.
Finally, a full outer shell imparts a lot of energetic stability to an atom, which tends to allow its electron cloud to contract a bit, so as a shell is filled across a period of the periodic table, it tends to contract a little.
In general (but with a few exceptions) atomic size decreases from left to right and from the bottom to the top of the periodic table.
In this periodic table, the relative sizes (not absolute - there are no measurements) are shown. The sizes of the noble gas atoms (gray) can't really be compared to the other atoms because atomic sizes are generally averaged over many measurements of atoms bound in compounds with other kinds of atoms. The noble gases (with the exception of Xe) don't bind to other atoms, so their radii have to be measured another way.
Here are closeups of the size trends in the group 1A (alkaline earth) elements and the group 7A (halogens) elements.
These are very easy to understand. As we move down the periodic table within a group (same number of electrons in the outer shell), each row represents addition of a new shell of electrons. Those electrons take up space (mostly because they repel one-another), so more shells means a larger atom.
The exception to this rule can be seen in the middle of the periodic table, where d-subshells are being filled in metallic elements. D-orbitals are different in many respects in the chemistry of elements.
Here are parts of the second and third periods of the table. Notice that as we move from left to right across the period, atomic sizes diminish because of increased nuclear positive charge and completion of an octet. Remember that we have to throw out the noble gas sizes in this comparison; they can only be compared within their group, in which case they show the same group pattern as the figure above.
Electronegativity is not something you'll likely ever measure or use quantitatively, but it's an important concept to understand because it can explain many properties of many chemical compounds and reactions.
The concept is pretty simple: Electronegativity is a relative measure of the propensity (inclination or tendency) of an atom to attract to itself the electrons of another atom.
Electronegativity of an atom depends on two of its properties:
Gaining the 8th electron in the outer shell imparts an atom with a great deal of energetic stability, so there is a strong drive to gain that 8th electron. The stabilizing energy is lower for the 7th, 6th and lower electrons, so we say that atoms with 7, 6, 5, ... outer-shell electrons are less electronegative.
Electrons close to the nucleus are held more closely to it by the electrostatic (Coulomb) force than those farther from it. Therefore, attraction for the electrons of another atoms is greater for smaller atoms than for larger. In smaller atoms the nucleus is closer to the surface of the electron cloud and can exert more attractive force on the electrons of another atom. In larger atoms, the nucleus is both farther away and more "screened" by the intervening electrons, therefore larger atoms exert a smaller attractive force on the electrons of other atoms.
Atoms on the right side of the periodic table are more electronegative than those on the left. Atoms near the top of the table (smaller atoms) are more electronegative than those lower down.
It's worth studying some of the details of the periodic table above. First, consider hydrogen (H). the H-atom is a bare proton relatively unsheilded by its single 1s electron. It has a middle-of-the-road electronegativity, reflecting its high propensity for forming covalent bonds.
Copper (Cu) and silver (Ag) are exceptions to the rule of filling electrons by lowest-energy level first. The electron configuration of copper is
[Ar] 4s1, 3d10
You will recall that it is energetically favorable for copper to fill its 3d shell (10 electrons) with one of its 4s electrons. With a full d-shell, the propensity for copper to acquire another is low. The same is true of silver, with electron configuration
[Kr] 5s1, 4d10.
Copper and silver are lower-energy "excited states" of what we might have thought would have been the "ground state" of these atoms.
Gold (Au), however, is much larger, and its valence electrons do not enjoy such a large energy advantage upon rearrangement, so they do not. Therefore, gold has a tendency to take a tenth 4d electron from somewhere else to stabilize that shell.
We can make a similar argument as we move down the 6B group from chromium (Cr) to tungsten (W). Recall that the electron configurations of these elements are also exceptions. They take an s electron from an s orbital in order to half-fill the d-orbital below it. For example, the lowest-energy electron configuration of Cr is
[Ar] 4s1, 3d5
Tungsten, on the other hand, is not such an exception, so a borrowed 5th 3d electron is energetically favorable, thus its relatively high electronegativity.
Finally, what's up with the electronegativities of krypton (Kr), xenon (Xe) and radon (Rn) ?
These electronegativities can't be scaled directly with those of, say, the halogens, because large noble gases don't react in the same ways to make similar compounds. They do react to form some compounds, however. For example, Xe and strongly electronegative compounds like fluorine (F) and oxygen (O) to form compounds like XeF, XeF6 XeO3F2
Below is a simple periodic table showing the gross trend – without any exceptions – of electronegativity. As you move up and to the right, from francium (Fr) to fluorine (F), electronegativity increases.
Electron affinity (EA) is another measure of the tendency of an atom to attract an electron from another atom, but it has a more precise definition than electronegativity.
EA is the the energy change that occurs when an electron e- is added to a neutral atom A in the gaseous state (i.e. not bonded to anything else):
A + e- → A-
EA is defined as the energy of A- minus the energy of A + e-, so if EA is negative, the atom A has a greater affinity for absorbing another electron than an element with a more positive EA.
When we look at the periodic table, the EA trend is very similar to the electronegativity trend. Here I've made the noble gas elements gray to indicate that, because of the full valence shell, they have little affinity for taking on another electron.
In contrast to EA, the ionization energy of an element is the energy it takes to remove the outermost valence-shell electron from the atom — that's the easiest electron to remove.
Strictly-speaking, we should say "first ionization energy" because there are also second, third ... and so on ionization energies, the energy required to remove the second electron, the third, and so on. We'll concentrate here on the first.
You might want to think back to our study of electrons and the photoelectric effect to help you understand ionization energy. For example,if we perform experiments in which we fire a laser at, say, a beam of gas-phase atoms, there is a threshold frequency (remember that frequency is directly proportional to energy) of light that will detach an electron. No lower energy will get the job done, and with higher frequency light, the electron will just speed away faster after it detaches from the atom.
As you might expect, the first ionization energy follows the electronegativity and electron affinity trend, except that the noble gases fit right in this time; they have the highest ionization energies. The tighter an electron is held, the more difficult it is to remove.
In the figure below, the ionization energies (in units of electron-volts, eV) are plotted against atomic number of the elements. Notice that ionization energy increases across each period of the table, culminating with the noble gases, which bind that last electron the most tightly. Notice that the alkali metals (Li, Na, K, Rb, Cs and Fr) are the easiest to ionize, and recall that these atoms mainly exist as +1 cations.
Notice also that there are some other "islands" of stability. Nitrogen (N) and phosphorus (P) have half-filled valence p-subshells, which appears to make them a bit more stable than their near-neighbors on the periodic table. Likewise, Zn, Cd and Hg all have full outer d-subshells.
In order to think about metallic character, we really have to ask what we mean by "metal." What is a metal?
Metals have some familiar properties. Most, but not all are shiny. Most, but not all are malleable (able to be shaped into flat sheets and other shapes) and/or ductile (able to be pulled into long filaments like wires). To varying degrees, metals carry electric current and transmit heat readily.
The outer-shell electrons of metals are either weakly bound (like the alkaline earth elements) or they occupy d-orbitals. D-orbitals are unique in that when many d-orbital containing metal elements bond together, the d-orbitals combine to create energy levels that allow electrons to roam freely from one atom to another. The best conductors, like the coinage metals, copper, silver and gold, have mostly-filled d-subshells. The periodic table below shows the metallic elements and the general at-a-glance trend. In general, as we move down and to the left, elements become more metallic.
| 1. |
Rank these elements in order of increasing atomic radius: C, Al, O, K
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| 2. |
Rank these elements in order of increasing electronegativity: O, S, Ne, Al
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| 3. |
Which atom of each pair has the largest atomic radius? (a) Al | B (b) Na | Al (c) Mg | Ca (d) O | F (e) Br | Cl (f) S | O
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| 4. |
One of the highly-possible products of fallout from nuclear explosions or nuclear power plants is 90Sr, a radioactive isotope of strontium (Sr). Hypothesize about why 90Sr might cause bone cancers, leukemias and other cancers in children exposed to such nuclear accidents.
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| 5. |
Which has the least attraction for electrons in a chemical bond? – O, Cl, F, N
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| 6. |
Why does fluorine (F) have a higher first ionization energy than iodine (I)?
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| 7. |
Of phosphorus (P) or iodine (I), which has the greater first ionization energy?
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| 8. |
Of sodium (Na) or rubidium (Rb), which has the greater first ionization energy?
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