General Chemistry/Periodicity and Electron Configurations
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[edit] Blocks of the Periodic Table
The Periodic Table does more than just list the elements. The word periodic means that in each row, or period, there is a pattern of characteristics in the elements. This is because the elements are listed according to their electron configuration. The Alkali metals and Alkaline earth metals have one and two valence electrons (electrons in the outer shell) respectively, and are thus very reactive. These elements are the s-block of the periodic table. The p-block, on the right, contains common non-metals such as chlorine and helium. The noble gases, in the column on the right, almost never react, since they have eight valence electrons. The halogens, directly to the left of the noble gases, readily react with metals. The s and p blocks make up the main-group elements, also known as representative elements. The d-block, which is the largest, consists of transition metals such as copper, iron, and gold. The f-block, on the bottom, contains rarer metals including uranium. Elements in the same Group or Family have the same configuration of valence electrons, making them behave in chemically similar ways.
[edit] Causes for Trends
There are certain phenomena that cause the periodic trends to occur. You must understand them before learning the trends.
[edit] Effective Nuclear Charge
The effective nuclear charge is the amount of positive charge acting on an electron. It is the number of protons in the nucleus minus the number of electrons in between the nucleus and the electron in question. Basically, the nucleus attracts an electron, but other electrons in lower shells repel it (opposites attract, likes repel).
[edit] Shielding Effect
The shielding (or screening) effect is similar to effective nuclear charge. The core electrons repel the valence electrons to some degree. The more electron shells there are (a new shell for each row in the periodic table), the greater the shielding effect is. Essentially, the core electrons shield the valence electrons from the positive charge of the nucleus.
[edit] Electron-Electron Repulsions
When two electrons are in the same shell, they will repel each other slightly. This effect is mostly canceled out due to the strong attraction to the nucleus, but it does cause electrons in the same shell to spread out a little bit. Lower shells experience this effect more because they are smaller and allow the electrons to interact more.
[edit] Coulomb's Law
Coulomb's law is an equation that determines the amount of force that two charged particles attract or repel each other. It is 
, where Q is the amount of charge (+1e for protons, -1e for electrons), r is the distance between them, and k is a constant. You can see that doubling the distance would quarter the force. Also, a large number of protons would attract an electron with much more force than just a few protons would.
[edit] Trends in the Periodic table
Most of the elements occur naturally on Earth. However, all elements beyond uranium (number 92) are called trans-uranium elements and never occur outside of a laboratory. Most of the elements occur as solids or gases at STP. STP is standard temperature and pressure, which is 0° C and 1 atmosphere of pressure. There are only two elements that occur as liquids at STP: mercury (Hg) and bromine (Br).
Bismuth (Bi) is the last stable element on the chart. All elements after bismuth are radioactive and decay into more stable elements. Some elements before bismuth are radioactive, however.
[edit] Atomic Radius
Leaving out the noble gases, atomic radius is larger on the left side of the periodic chart and gets smaller as you move rightward across the period. Additionally, as you move down the group, radius increases.
Atomic radius decreases along a period due to greater effective nuclear charge. Atomic radius increase down a group due to the shielding effect of the additional core electrons, and the presence of another electron shell.
[edit] Ionic Radius
For nonmetals, ions are bigger than atoms, as the ions have extra electrons. For metals, it is the opposite.
Extra electrons (negative ions, called anions) cause additional electron-electron repulsions, making them spread out farther. Less electrons (positive ions, called cations) cause fewer repulsions, bringing them in closer.
[edit] Ionization Energy
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Ionization energy is the energy required to strip an electron from the atom (when in the gas state).  |
Ionization energy is highest on the top right, lowest on the bottom left.
Ionization energy decreases going left across a period because there is a lower effective nuclear charge keeping the electrons attracted to the nucleus, so less energy is needed to pull one out. It decreases going down a group due to the shielding effect. Also, ionization energy decreases in both cases due to a larger atomic radius. Remember Coulomb's Law: as the distance between the nucleus and electrons increases, the force decreases at a quadratic rate.
[edit] Electron Affinity
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Electron affinity is the opposite of ionization energy. It is the energy released when an electron is added to an atom.  |
Electron affinity is highest in the upper left, lowest on the bottom right. However, electron affinity is actually negative for the noble gasses. They already have a complete valence shell, so there is no room in their orbitals for another electron. Adding an electron would require creating a whole new shell, which takes energy instead of releasing it. Several other elements have extremely low electron affinities because they are already in a stable configuration, and adding an electron would decrease stability.
Electron affinity occurs due to the same reasons as ionization energy.
[edit] Electronegativity
Electronegativity is how much an atom attracts electrons within a bond. It is measured on a scale with fluorine at 4.0 and francium at 0.7. Electronegativity decreases from upper left to lower right.
Electronegativity decreases because of atomic radius, shielding effect, and effective nuclear charge in the same manner that ionization energy decreases.
[edit] Metallic Character
Metallic elements are shiny, usually gray or silver colored, and good conductors of heat and electricity. They are malleable (can be hammered into thin sheets), and ductile (can be stretched into wires). Some metals, like sodium, are soft and can be cut with a knife. Others, like iron, are very hard. Non-metallic atoms are dull, usually colorful or colorless, and poor conductors. They are brittle when solid, and many are gases at STP. Metals give away their valence electrons when bonding, whereas non-metals take electrons.
The metals are towards the left and center of the periodic table—in the s-block, d-block, and f-block . Poor metals and metalloids (somewhat metal, somewhat non-metal) are in the lower left of the p-block. Non-metals are on the right of the table.
Metallic character increases from right to left and top to bottom. Non-metallic character is just the opposite. This is because of the other trends: ionization energy, electron affinity, and electronegativity.
 Sodium, very metallic |
 Sulfur, very non-metallic |
