>> Key Terms (indicated in blue within the text below):

antibonding orbital bonding orbital covalent bond diamagnetic double bond electron affinity electronegativity formal charge

ionic bond ionization energy Lewis structure lone pair molecular orbital theory nonpolar bond octet rule paramagnetic

pi () bond polar bond resonance structure sigma () bond single bond triple bond valence electron

In Chapter 3 the periodic trend associated with ionization energy was observed. Ionization energy, the energy required to remove an electron from a gaseous atom, increases from left to right and from bottom to top across the periodic table.

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A similar chemical property, electron affinity, also has a periodic trend. Electron affinity is the energy required to add an electron to a gaseous atom. However, not all atoms require energy when an electron is added. Some give off energy instead. To express that energy is given off, a negative value of energy is used. (Thus positive values of energy express that energy is required in a process.) One periodic trend of electron affinity is that it is easier to add an electron (more negative or less positive energies) to atoms further to the right until the noble gases. Noble gases do not have an affinity for an electron. With this trend the column of noble gases actually fits best on the leftmost side of the periodic table. The other trend is for electron affinity to increase (more negative, less positive) from bottom to top. The reason for both these trends is the change in effective nuclear charge, Z_eff. From left to right across the table, more protons are added but the n quantum number, distance, remains the same. Thus Z_eff increases. From top to bottom the n quantum number and the distance of the outermost electron from the nucleus increase. Not only does the distance decrease Z_eff, but the presence of more electrons between the nucleus and the outermost electron shields the outermost electron from the nuclear charge.

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Both ionization energy and electron affinity refer to gaseous atoms. A more useful property of atoms is electronegativity. Electronegativity () is the ability of an atom to attract electrons to itself when combined with other atoms. The periodic trend of electronegativity is to increase from left to right (except for the noble gases) and from bottom to top. Thus fluorine is the most electronegative of all atoms. It is often useful to know the top four electronegative atoms. These are, in order, fluorine, oxygen, nitrogen, and chlorine. Hydrogen is a special, but important, case. If the periodic table were totally organized strictly by trends in electronegativity, it would fit between carbon and boron.

Electronegativity is important in the way that atoms combine, or bond. If the electronegativity difference is large ( > 2.5), one atom will actually give an electron to the other so that both atoms can obtain a noble gas arrangement of electrons. The atom that gained an electron also gained a negative charge to become an anion, and the other (the atom that lost the electron) becomes a cation. The positive charge and negative charges are attracted to each other to form an ionic bond. Because these charges attract generally (cations attract any negative charge, not just the atom that gave up the electron), compounds with ionic bonds usually come in groups of very large and nonspecific numbers of ions. However, the net charge must be zero, so their formulas are expressed as the simplest ratio of ions that result in a zero charge. Large electronegativity differences are normally found between metals and nonmetals. Ions involved in ionic bonding are not true molecules, since the number of atoms involved is large and nonspecific.

On the other hand, atoms with similar electronegativities combine by sharing their electrons. Sharing of electrons normally occurs between nonmetals and is called a covalent bond. In covalent bonds, both atoms get to count the shared electrons toward completing the noble gas configuration. There are eight s and p electrons in a noble gas configuration. Thus atoms combine so that each one is surrounded by eight electrons; this is called the octet rule. Unlike ionic bonds, covalent bonds are specific to the atoms involved. A group of atoms that are covalently bonded forms a true molecule.

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While the electrons are shared, they are only shared equally if the two atoms are exactly the same. Electrons that are shared equally are nonpolar bonds. Shared electrons that are more attracted to one atom than the other are called polar bonds. The more electronegative atom is the negative pole and the less electronegative atom is the positive pole. Extreme examples on polar bonds might actually qualify as ionic bonds. The line between an ionic bond and a covalent bond is not absolute.

The electrons involved in bonding are the outermost, or valence electrons. Lewis structures are a way of showing how s and p valence electrons are associated with an atom. The number of valence electrons can be determined by counting across the period, skipping transition metals if necessary. The number of valence electrons is also the group number. Since ions are made by adding or subtracting valence electrons, the charge must be considered when counting valence electrons. Lewis structures are a very convenient way to represent atoms, ions, and atoms involved in covalent bonding.

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For atoms and monatomic ions, the valence electrons are represented as dots, next to the atom symbol. The possible positions are left, right, over, or under the atomic symbol. The first two valence electrons are paired together (like s electrons). The remaining electrons are placed in the other positions with one electron in each position first, then doubling up until there are no more valence electrons. Which positions are used does not matter as long as the electrons are paired (or unpaired) appropriately.

For molecules and polyatomic ions, the number of valence electrons is the sum of the valence electrons of each atom and the charge (adding electrons for a negative charge and subtracting electrons for a positive charge). Two shared electrons are represented by a line connecting the appropriate atoms. If only two electrons are shared between the two atoms, it is called a single bond. If four electrons are shared, it is a double bond (two lines are used). If six electrons are shared, it is a triple bond (three lines are used). If the electrons are not shared, it is called a lone pair and is represented by a pair of dots.

Often it is possible to write a Lewis structure in more than one way. To determine which organization of atoms is correct, formal charges can be used. Formal charges compare the number of valence electrons contributed by the atom to the number of valence electrons assigned to the atom in the Lewis structure. Assigned electrons are all electrons in lone pairs associated with the atom and half of all electrons in the bonds associated with the atom. The sum of the formal charges will be zero for a molecule or the charge of the ion for a polyatomic ion. The appropriate Lewis structure is the one where all formal charges are closest to zero.

For some structures it is possible for a double bond to have more than one exactly equivalent position. These equivalent structures are called resonance structures. The true arrangement of atoms and electrons is not any one of these Lewis structures, but an average of all the resonance structures. (For example, if there are two resonance structures, at each location where the double bond could be located there is a 1 1/2 bond.)

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Another model of bonding is molecular orbital theory. In molecular orbital theory, atomic orbitals combine to make molecular orbitals. The number of orbitals and the energy of the orbitals are conserved. Therefore if two atomic orbitals combine, two molecular orbitals will be formed. One of these molecular orbitals will be lower in energy (more stable) than the original atomic orbitals. This is a bonding orbital and the electron density is located between the two atoms. The other molecular orbital is the antibonding orbital. It is as much higher in energy than the atomic orbital as the bonding orbital is lower in energy. The electron density of an antibonding orbital is located on the opposite side of the area between the atoms. Antibonding orbitals are designated with a star (*).

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When two s orbitals or two p_z orbitals (orbitals oriented along the z axis) combine, the electron density is centered along the axis between the atoms. These molecular orbitals are called sigma () bonds. When the other two p orbitals combine (p_x with p_x or p_y with p_y), the electron density is located on either side of the axis connecting the two atoms. These molecular orbitals are called pi () bonds. The energy change for bonds is smaller than it is for bonds. Figure 6.16 shows and bonding and antibonding orbitals.

Electrons fill molecular orbitals in the same way as they fill atomic orbitals. That is, two electrons spinning in opposite directions will fill the lowest-energy orbital first. If two or more orbitals are degenerate (have the same energy), parallel (spinning in the same direction) electrons will fill each orbital before doubling up.

A major strength of molecular orbital theory is that it explains magnetic properties of molecular compounds. An element or compound will be attracted to a magnetic field if it has an unpaired electron. Such elements and compounds are called paramagetic. An element or compound will be slightly repelled from a magnetic field if all electrons are paired. These elements and compounds are called diamagnetic. In Lewis structures, all electrons appear to be paired, although some (O₂ being the classic example) are magnetic.