What Are Chemical Bonds, and Why Do They Form?
A
chemical bond is the result of anattraction between atoms or ions. The types of bonds that a moleculecontains will determine its physical properties, such as melting point,hardness, electrical and thermal conductivity, and solubility. How dochemical bonds occur? As we mentioned before, only the outermost, or
valence,electrons of an atom are involved in chemical bonds. Let’s begin ourdiscussion by looking at the simplest element, hydrogen. When twohydrogen atoms approach each other, electron-electron repulsion andproton-proton repulsion both act to try to keep the atoms apart.However, proton-electron attraction can counterbalance this, pullingthe 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 thesame number of electrons as the noble gas that’s nearest them on theperiodic table. All of the noble gases have eight valence electrons (
s2
p6) 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 ofexceptions is atoms with fewer than eight electrons—hydrogen (H) hasjust one electron. In BeH2, there are only four valenceelectrons around Be: Beryllium contributes two electrons and eachhydrogen contributes one. The second exception to the octet rule isseen in elements in periods 4 and higher. Atoms of these elements canbe surrounded by more than four valence pairs in certain compounds.
Types of Chemical Bonds
You’ll need to be familiar with three typesof chemical bonds for the
SAT II Chemistry exam: ionic bonds, covalentbonds, and metallic bonds.
Ionic bonds are the resultof an electrostatic attraction between ions that have opposite charges;in other words, cations and anions. Ionic bonds usually form betweenmetals and nonmetals; elements that participate in ionic bonds areoften from opposite ends of the periodic table and have anelectronegativity difference
greater than 1.67. Ionic bonds are verystrong, so compounds that contain these types of bonds have highmelting points and exist in a solid state under standard conditions.Finally, remember that in an ionic bond, an electron is actually
transferredfrom 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
sharedbetween atoms rather than transferred from one atom to another.However, this sharing rarely occurs equally because of course no twoatoms have the same electronegativity value. (The obvious exception isin a bond between two atoms of the same element.) We say that covalentbonds are
nonpolar if the electronegativity difference between the two atoms involved falls between 0 and 0.4. We say they are
polarif the electronegativity difference falls between 0.4 and 1.67. In bothnonpolar and polar covalent bonds, the element with the higherelectronegativity attracts the electron pair more strongly. The twobonds 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 (s
), in which the electron density is concentrated along the line that represents the bond joining the two atoms.
However, double and triple bonds occurfrequently (especially among carbon, nitrogen, oxygen, phosphorus, andsulfur atoms) and come about when atoms can achieve a complete octet bysharing more than one pair of electrons between them. If two electronpairs 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 (p
).A pi bond is a bond in which the electron density is concentrated aboveand below the line that represents the bond joining the two atoms. Ifthree electron pairs are shared between the two nuclei, a
triple bond forms. 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 densitybetween two nuclei: they decrease nuclear repulsion while enhancing thenucleus-to-electron density attractions. The nuclei move closertogether, which means that double bonds are shorter than single bondsand triple bonds are shortest of all.
Metallic bonds exist only in metals,such as aluminum, gold, copper, and iron. In metals, each atom isbonded to several other metal atoms, and their electrons are free tomove throughout the metal structure. This special situation isresponsible for the unique properties of metals, such as their highconductivity.
Drawing Lewis Structures
Here are some rules to follow when drawingLewis structures—you should follow these simple steps for every Lewisstructure you draw, and soon enough you’ll find that you’ve memorizedthem. While you will not specifically be asked to draw Lewis structureson the test, you will be asked to predict molecular shapes, and inorder to do this you need to be able to draw the Lewis structure—somemorize these rules! To predict arrangement of atoms within themolecule
- Find the total number of valence electrons by adding up groupnumbers of the elements. For anions, add the appropriate number ofelectrons, and for cations, subtract the appropriate number ofelectrons. Divide by 2 to get the number of electron pairs.
- Determine which is the central atom—in situations where thecentral atom has a group of other atoms bonded to it, the central atomis usually written first. For example, in CCl4, the carbonatom is the central atom. You should also note that the central atom isusually less electronegative than the ones that surround it, so you canuse this fact to determine which is the central atom in cases that seemmore ambiguous.
- Place one pair of electrons between each pair of bondedatoms and subtract the number of electrons used for each bond (2) fromyour total.
- Place lone pairs about each terminal atom (except H, whichcan only have two electrons) to satisfy the octet rule. Leftover pairsshould be assigned to the central atom. If the central atom is from thethird or higher period, it can accommodate more than four electronpairs since it has d orbitals in which to place them.
- If the central atom is not yet surrounded by four electronpairs, convert one or more terminal atom lone pairs to double bonds.Remember that not all elements form double bonds: only C, N, O, P, andS!
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, rememberthe rules. Find the total number of valence electrons in the moleculeby adding the group numbers of its constituent atoms. So for C2H2, this would mean C = 4
2(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 10valence electrons. Next, the carbons are clearly acting as the centralatoms since hydrogen can only have two electrons and thus can’t formmore than one bond. So your molecule looks like this: H—C—C—H. So faryou’ve used up six electrons in three bonds. Hydrogen can’t support anymore electrons, though: both H’s have their maximum number! So yourfirst thought might be to add the remaining electrons to the centralcarbons—but there is no way of spreading out the remaining fourelectrons to satisfy the octets of both carbon atoms except to draw atriple 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 thiscompound, and you’ll see that it contains one double bond (between thetwo 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, sothere are five sigma bonds and one pi bond.
Exceptions to Regular Lewis Structures—Resonance Structures
Sometimes you’ll come across a structure thatcan’t be determined by following the Lewis dot structure rules. Forexample, ozone (O3) contains two bonds of equal bond length,which seems to indicate that there are an equal number of bonding pairson each side of the central O atom. But try drawing the Lewis structurefor ozone, and this is what you get:
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We have drawn the molecule with one doublebond and one single bond, but since we know that the bond lengths inthe molecule are equal, ozone can’t have one double and one singlebond—the double bond would be much shorter than the single one. Thinkabout it again, though—we could also draw the structure as below, withthe double bond on the other side:
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Together, our two drawings of ozone are resonance structures for the molecule.
Resonance structuresare two or more Lewis structures that describe a molecule: theircomposite represents a true structure for the molecule. We use thedouble-directional arrows to indicate resonance and also bracket thestructures or simply draw a single, composite picture.
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Let’s look at another example of resonance, in the carbonate ion CO32-:
`N-RVYw 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
