What Are Chemical Bonds, and Why Do They Form?

　　　　A chemical bond is the result of an attraction between atoms or ions. The types of bonds that a molecule contains will determine its physical properties, such as melting point, hardness, electrical and thermal conductivity, and solubility. How do chemical bonds occur? As we mentioned before, only the outermost, or valence, electrons of an atom are involved in chemical bonds. Let’s begin our discussion by looking at the simplest element, hydrogen. When two hydrogen atoms approach each other, electron-electron repulsion and proton-proton repulsion both act to try to keep the atoms apart. However, proton-electron attraction can counterbalance this, pulling the two hydrogen atoms together so that a bond is formed. Look at the energy diagram below for the formation of an H–H bond.

　　As you’ll see throughout our discussion, atoms will often gain, lose, or share electrons in order to possess the same number of electrons as the noble gas that’s nearest them on the periodic table. All of the noble gases have eight valence electrons (s2p6) and are very chemically stable, so this phenomenon is known as the octet rule. There are, however, certain exceptions to the octet rule. One group of exceptions is atoms with fewer than eight electrons—hydrogen (H) has just one electron. In BeH2, there are only four valence electrons around Be: Beryllium contributes two electrons and each hydrogen contributes one. The second exception to the octet rule is seen in elements in periods 4 and higher. Atoms of these elements can be surrounded by more than four valence pairs in certain compounds.

　　Types of Chemical Bonds　　You’ll need to be familiar with three types of chemical bonds for the SAT II Chemistry exam: ionic bonds, covalent bonds, and metallic bonds. 　　Ionic bonds are the result of an electrostatic attraction between ions that have opposite charges; in other words, cations and anions. Ionic bonds usually form between metals and nonmetals; elements that participate in ionic bonds are often from opposite ends of the periodic table and have an electronegativity difference greater than 1.67. Ionic bonds are very strong, so compounds that contain these types of bonds have high melting points and exist in a solid state under standard conditions. Finally, remember that in an ionic bond, an electron is actually transferred from the less electronegative atom to the more electronegative element. One example of a molecule that contains an ionic bond is table salt, NaCl.　　Covalent bonds form when electrons are shared between atoms rather than transferred from one atom to another. However, this sharing rarely occurs equally because of course no two atoms have the same electronegativity value. (The obvious exception is in a bond between two atoms of the same element.) We say that covalent bonds are nonpolarif the electronegativity difference between the two atoms involved falls between 0 and 0.4. We say they are polar if the electronegativity difference falls between 0.4 and 1.67. In both nonpolar and polar covalent bonds, the element with the higher electronegativity attracts the electron pair more strongly. The two bonds in a molecule of carbon dioxide, CO2, are covalent bonds.　　Covalent bonds can be single, double, or triple. If only one pair of electrons is shared, a single bond is formed. This single bond is a sigma bond, in which the electron density is concentrated along the line that represents the bond joining the two atoms.　　However, double and triple bonds occur frequently (especially among carbon, nitrogen, oxygen, phosphorus, and sulfur atoms) and come about when atoms can achieve a complete octet by sharing more than one pair of electrons between them. If two electron pairs are shared between the two atoms, a double bond forms, where one of the bonds is a sigma bond, and the other is a pi bond. A pi bond is a bond in which the electron density is concentrated above and below the line that represents the bond joining the two atoms. If three electron pairs are shared between the two nuclei, a triple bondforms. In a triple bond, the first bond to form is a single, sigma bond and the next two to form are both pi. 　　Multiple bonds increase electron density between two nuclei: they decrease nuclear repulsion while enhancing the nucleus-to-electron density attractions. The nuclei move closer together, which means that double bonds are shorter than single bonds and triple bonds are shortest of all. 　　Metallic bondsexist only in metals, such as aluminum, gold, copper, and iron. In metals, each atom is bonded to several other metal atoms, and their electrons are free to move throughout the metal structure. This special situation is responsible for the unique properties of metals, such as their high conductivity.　　Drawing Lewis Structures　　Here are some rules to follow when drawing Lewis structures—you should follow these simple steps for every Lewis structure you draw, and soon enough you’ll find that you’ve memorized them. While you will not specifically be asked to draw Lewis structures on the test, you will be asked to predict molecular shapes, and in order to do this you need to be able to draw the Lewis structure—so memorize these rules! To predict arrangement of atoms within the moleculeFind the total number of valence electrons by adding up group numbers of the elements. For anions, add the appropriate number of electrons, and for cations, subtract the appropriate number of electrons. Divide by 2 to get the number of electron pairs. Determine which is the central atom—in situations where the central atom has a group of other atoms bonded to it, the central atom is usually written first. For example, in CCl4, the carbon atom is the central atom. You should also note that the central atom is usually less electronegative than the ones that surround it, so you can use this fact to determine which is the central atom in cases that seem more ambiguous. Place one pair of electrons between each pair of bonded atoms and subtract the number of electrons used for each bond (2) from your total. Place lone pairs about each terminal atom (except H, which can only have two electrons) to satisfy the octet rule. Leftover pairs should be assigned to the central atom. If the central atom is from the third or higher period, it can accommodate more than four electron pairs since it has d orbitals in which to place them. If the central atom is not yet surrounded by four electron pairs, convert one or more terminal atom lone pairs to double bonds. Remember that not all elements form double bonds: only C, N, O, P, and S!　　Example　　Which one of the following molecules contains a triple bond: PF3, NF3, C2H2, H2CO, or HOF?　　Explanation　　The answer is C2H2, which is also known as ethyne. When drawing this structure, remember the rules. Find the total number of valence electrons in the molecule by adding the group numbers of its constituent atoms. So for C2H2, this would mean C = 42 (since there are two carbons) = 8. Add to this the group number of H, which is 1, times 2 because there are two hydrogens = a total of 10 valence electrons. Next, the carbons are clearly acting as the central atoms since hydrogen can only have two electrons and thus can’t form more than one bond. So your molecule looks like this: H—C—C—H. So far you’ve used up six electrons in three bonds. Hydrogen can’t support any more electrons, though: both H’s have their maximum number! So your first thought might be to add the remaining electrons to the central carbons—but there is no way of spreading out the remaining four electrons to satisfy the octets of both carbon atoms except to draw a triple bond between the two carbons. 　　For practice, try drawing the structures of the other four compounds listed.　　Example　　How many sigma (s) bonds and how many pi (p) bonds does the molecule ethene, C2H4, contain? 　　Explanation　　First draw the Lewis structure for this compound, and you’ll see that it contains one double bond (between the two carbons) and four single bonds. Each single bond is a sigma bond, and the double bond is made up of one sigma bond and one pi bond, so there are five sigma bonds and one pi bond. 　　Exceptions to Regular Lewis Structures—Resonance Structures　　Sometimes you’ll come across a structure that can’t be determined by following the Lewis dot structure rules. For example, ozone (O3) contains two bonds of equal bond length, which seems to indicate that there are an equal number of bonding pairs on each side of the central O atom. But try drawing the Lewis structure for ozone, and this is what you get:　　We have drawn the molecule with one double bond and one single bond, but since we know that the bond lengths in the molecule are equal, ozone can’t have one double and one single bond—the double bond would be much shorter than the single one. Think about it again, though—we could also draw the structure as below, with the double bond on the other side: 　　Together, our two drawings of ozone are resonance structures for the molecule. Resonance structures are two or more Lewis structures that describe a molecule: their composite represents a true structure for the molecule. We use the double-directional arrows to indicate resonance and also bracket the structures or simply draw a single, composite picture.

　　Let’s look at another example of resonance, in the carbonate ion CO32-:

　　Notice that resonance structures differ only in electron pair positions, not atom positions! 　　Example　　Draw the Lewis structures for the following molecules: HF, N2, NH3, CH4, CF4, and NO+.　　Explanation

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