On-Line Study Guide - OLSG
Chemistry 111

Chapter 3: Chemical Bonds

We have seen how to write electron configurations of atoms, in which atoms are placed in orbitals of increasing energy . Each orbital can hold two electrons. The particular geometric shape of the periodic table can be explained by how subsequent electrons in succeeding elements in the periodic table are placed into orbitals of higher energy. This principele leads to the arrangment of atoms in the periodic table into the s, p, d, and f block elements.

The electron configurations can also be used to understanding the chemical properties atoms. Noble gases in group VIIIA are unreactive, suggesting that a filled outer shell confers chemical inertness and nonreactivity. Other elements react to achieve an electron configuration similar to the nobel gases.

Ionic Bonds

Group 1A metals lose an electron to form a cation, a positively charged ion, which has the same electronic configuration of the previous noble gas element.; Group 7A atoms gain an electron to form an anion, a negatively charged particle with the electronic configuration of the next noble gas. These two oppositely charged ions attract each other tightly to form an ionic bond.

Examples of compounds formed between Gp 1A metals and Gp 7 nonmetals include sodium chloride, (NaCl), potassium iodide (KI), etc. Metals in Gp 2A typically lose 2 electrons to form cations with a postive 2 charge, which react with either nonmetals of Gp VI which gain two electron to form compounds, (Example MgO - magnesium oxide) or with 2 atoms of nonmetals of Gp VII which gain the two electron to form salts (Example MgCl₂ - magnesium chloride). Finally metals in group 1 can form salts with group VI nonmetals, such as Na₂O. H, a group 1 atoms, but a gas at room temperature, does not act like a metal, and does not usually give up an electron as do the other metals.

Animation: Formation of an ionic bond between Na and Cl.

Notice in this animation that the size of Na decreases as it loses an electron and forms Na⁺. In contrast, Cl^- is larger than an atom of Cl. Why?

Now take a quiz on the compounds formed from metals and nonmetals.

WebCT Quiz: Select Naming Ionic Compounds  as the Quiz

Covalent Bonds

In contrast, nonmetals normally share electrons and form covalent bonds. The shared electrons can be considered to part of the octet of electons around a nonmetal, or a duet of electrons around hydrogen. Consider carbon as an example. To achieve a full outer shell, it could either gain 4 electrons (to achieve the electronic configuration of Ne - 1s²2s²2p⁶) or lose 4 (to achieve the configuration of He). In the first case it would have a 4- charge and in the second a 4+ charge. Such high charges would be difficult to achieve as it would get increasingly hard, for example, to remove electrons successively from a more and more positively charged C ion. Instead C shares its 4 outer shell electrons with 4 outer shell electrons contributed by 4 H atoms to form 4 covalent bonds, each with 2 electrons. These bonds are represented by lines connecting the C and each H.

Carbon, in group 4, has 4 additional electrons to fill its outer shell. Hence it can form 4 bonds with H. N, in group 5, needs 3 electrons and forms 3 bonds with H. Likewise, O forms 2 and F forms 1. These characteristic number of bonds is true for compounds of these elements with atoms other than H. For instance nitrogen exists in the air as N₂, in which each N atoms is connected to the other N atoms through 3 covalent bonds, often called a triple bond. This is an example of an elemental form of N which exists as molecules (more than one atom bonded).

Double bonds are common in nature as well

The electrons in a covalent bond can be shared equally or unequally. If the two atoms aharing the 2 electrons are identical, the electrons are shared equally. If the atoms are different they may or may not be shared equally. Some elements, notably, F, O, and N "like" electrons more than the other elements. These elements (FON) have a high electronegativity - they attract shared electrons toward them more than other elements when covalenty bonded to the other element. When F, O, or N is covalently bond to any element like C, H, the bonds are polar covalent. The two shared electrons spend more of their time on F,O, or N, confering a partial negative charge on F, O, or N. The other element bonded to F, O, or N, then has a partial defiiciency in electrons and is partially positive. Which types of elements have high electronegativity? Hopefully you can deduce that metals don't. After all, they like to lose electrons and form ionic bonds. Nonmetals, in contrast, seem to "like" electrons since they gain electrons when forming bonds to metals. As a general rule, electronegativity increases from left to right within a period and from top to bottom in a group. The electronegativity of F>O>N.

We have visualized bonding in three ways. In nonpolar covalent bonds, the electrons are equally shared between atoms. In polar covalent bonds, electrons are shared unequally. In ionic bonds, the electrons are shared so unequally than in fact one atoms donates electrons completely to the other atom. - i.e. they aren't shared at all. In reality, bonds are neither all ionic or nonpolar covalent. They have characterisitics of both.

To review bonds, consider the follow analogy. Click on the links for images that correspond to this analogy. Two men each have a dollar bill (representing an electron in atoms) and wish to pool their resources and share the money. There are three ways for them to share the money. Each can pool their dollars and have equal access to their shared $2. This is analogous to a nonpolar covalent bond. In this case each person has contributed $1 but shares $2. Alternatively, one person can take most of the $2, giving that person a slight positive excess (δ+) over the money he initially contributed and the other a slight negative lack (δ-) over the money he contributed. This is analagous to a polar covalent bond. Finally one person can take the other persons $1. He will now have a ful 1+ excess of dollars, and the other person a full 1- lack of dollars. Now time for another self-study quiz.

WebCT Quiz: Select Ionic, Polar Covalent, and Nonpolar Covalent  as the Quiz

Lewis Structures

The periodic table gives us clues as to the reactivity of different elements. Nonmetals form covalent bonds with other nonmetals in which atoms obtains an octet of outer shell electrons (or duet for hydrogen) through sharing electrons. C, in group 4, needs 4 electrons to achieve an octet,so it forms 4 bonds with 4 different H atoms. In fact when C is bonded to any nonmetal to form an uncharged molecule, it has 4 bonds. Likewise, N usually has 3 bonds, O 2 bonds, and F one bond when they form neutral covalent bonds in neutral atoms.

Molecular structures give us a way to visualize the molecule and understand its' physical and chemical properties. What would be most useful is to draw a molecular structure which shows all the outer shell electrons of each atom. It is the outer shell electrons which are involved in the chemistry of the molecule. Such molecular structures are called Lewis Structures. The following rules explain how to draw these structues. which they reviewed the rules for drawing Lewis structures for molecules and molecular ions.

GENERAL RULES FOR WRITING "LEWIS" STRUCTURES FOR MOLECULES

1. Draw the initial "skeletal" structure with atoms connected by single bonds. It should be the most symmetric structure you can draw. Make sure that H has only 1 bond, and that the other atoms do not exceed their normal number (2 bonds to O, 3 bonds to N, 4 bonds to C), and that no atom in the molecule has more than 4 bonds.
2. Determine TOTAL NUMBER OF OUTER SHELL ELECTRONS in the structure by adding the number of OUTER SHELL electrons for all the atoms in an uncharged molecule. Add electrons if the species is negatively charged, or substract electrons if is positively charged (i.e. if they are ions).
3. Subtract 2 electrons used to form each single covalent bond in your structure and distribute the remaining electrons in pairs around the atoms. If there are not enough electrons to form an octet around each atom except H, complete the octet around F, O, or N first since these are the elements that most like electrons (i.e. they have the highest electronegativity).
4. Calculate the formal charge around each atom as follows: Count the number of electrons surrounding each atom. Unshared pairs are assigned to the atom on which the unshared pair is found. A shared pair (i.e. a covalent bond) contains two electrons. Arbitrarily assign one of these two electrons to each atom connected by the bond. Count the total number of electrons assigned to each atom. Compare that number to the usual number of electrons in the outer shell for the atom, which is equal to the group number in the periodic table. The atom will have a +1 charge if it has one less electron assigned to it than usually found in the outer shell of the atom, and a -1 charge if it has one additional electron. Higher charges are also possible
5. If any atom has an incomplete octet, change nonbonded electron pairs to bonded pairs to form double and triple bonds in order to complete the octet around that atom. This does not affect the number of electrons around the atom that donates the nonbonded pair. Recalculate formal charges.
6. If rule 5 creates additional charges, use structure 4.
7. If phosphorous (P) or sulfur (S) is in your structure, the octet rule may not apply to that atom. Those elements are in period 3 of the periodic table and can accommodate more electrons, since they are bigger. Form additional multiple bonds to remove as many formal charges as possible. P often has 5 bonds while S has 6.

ANIMATION: Drawing Lewis Structure for Carbon Dioxide

Some key points to remember:

• The structures only show the number of outer shell electrons
• The structures (excluding the initial skeletal structure) should have the same number of electrons which equals to the sum of all the outer shell electrons plus or minus electrons for charged molecules.
• Make sure all atoms like C, N, O, F have an octet of electrons in the final structure and hydrogen has a duet. P, however, often has 5 bonds (10 electrons), while S has 6 bonds (12 electrons), since these period 3 atoms are larger, and can accomodate more electrons in unused d orbitals in the 3rd shell.
• There are two main ways to count electrons surrouding the atoms. Shared electrons (i.e. electrons in covalent bonds) count for each atoms when detemining if an atoms has an octet (or duet). However, shared elect rons are divided when determining formal charge. One electron from each shared pair is assigned to each of the two atoms atom connected by a covalent bond. The sum of the formal charges on each atoms must equal the overall charge on the molecule or molecular ion.
• Even though the formal charge ion atoms of neutral molecules is usually zero, the atoms might actual have a partial negative or positive charge if covalent bond through which they are connected to another atom is polar covalent. Any time F, O, or N is covalently attached to C or H, that bond is polar covalent, with the F, O, or N having a slight negative charge while the C or H have a slight positive charge. The sum of the partial charges in a neutral molecule must be zero.
• In chemical reactions, ionic or covalent bonds are broken. If a covalent bond is broken, the electrons either are split equally between the two separating atoms, or both leave with one of the atoms. If an O-H bond is broken, the electrons usually stay with oxygen, since O is electronegative and likes electrons more than H, and both atoms depart with a full outshell of electrons (O has an octet, and a - charge, while H leaves with no electrons (in a way a filled outer shell? and with a positive charge.)

ANIMATION: Counting Electrons- Octets, Formal Charge, Partial Charge

Having trouble with determining the formal charge on atoms?Take the following WCB quiz! 5 different molecular structures are shown below. Six different atoms are number and indicated in red in the structure. In the quiz you will be asked to determine the formal charge of those atoms.

WebCT Quiz: Select Formal Charges as the Quiz

In drawing Lewis structures we encounter molecules (not just atoms) that had a net charge. These are called molecular ions (or polyatomic ions). The ones you should know are carbonate (CO₃^2- -carbonate, HCO₃^- -bicarbonate, PO₄^3- -phosphate, SO₄^2- - sulfate, NO₃^- - nitrate, and NH₃⁺ - ammonium (see pg 99, Table 3.3 of the text) . In addition, please follow my rules for drawing Lewis structures and not the book's. For example, minimize formal charge even if you must violate the octet rule for S and P.

VSEPR theory

The utility of drawing Lewis structures is that you can use them to predict the chemical and physical properties of molecules. One such physical property is the geometry of the molecule. A simple theory can be used to this. It is called Valence (Outer) Shell Electron Pair Repulsion or VSEPR theory. In this model, the electron arrangement around the central atoms determine the geometry of the atoms. Electrons repel each other. Hence they will try to get as far away from each other as possible. Consider a simple model to understand this theory - a Styrofoam ball represents the nucleus of the atom and toothpicks represent PAIRS OF ELECTRONS (either a bonded or nonbonded pair). Now stick toothpicks in the ball so that they are as far away from each other as possible. You can figure out the angle made by any 2 toothpicks and the ball. The following geometries emerge:

ANIMATION: TOOTHPICKS IN BALLS - INTRO. TO VSEPR

If you can predict these geometries and angles, you can predict the geometry and angles around any atom in a molecule. Just treat nonbonded and bonded electrons as toothpicks. Also, consider a double or triple bond as a toothpick as well - i.e. as a single cloud of electron density between the nuclei of the bonded atoms.

First draw the correct Lewis structrues. To determine the geometry around a given atom, detemine the number of electron clouds surrounding that atom, counting multiple bonds as a single cloud of density. 2 clouds would be arranged in a linear fashion, with an angle of 180^o. If there are 3 clouds, they are arranged in a triangular or trigonal planar pattern, with angles of 120^o. 4 arranged in a tetrahedral fashion, with angles between clouds of 109^o. Now place the appropriate atoms at the end of clouds and determine the geometry of the actual atoms. See the animation below which shows the geometry of atom arrangement around a central atoms with 4 clouds of electrons.

ANIMATION: GEOMETRY OF ATOMS AT TETRAHEDRAL CENTER

Molecular Polarity

Now you can tell the geometry of the molecule and whether the bonds are polar or nonpolar covalent. In a polar covalent bonds, one atom (of greater electronegativity) has a δ - charge (not formal charge) while the other has a δ + charge. If the geometric center of all the δ - charges in an entire molecule is at the same point as the geometric center of all δ + charges - i.e. there is no separation of the centers of δ - and δ + charge, the whole molecule is nonpolar, even though it might have many polar covalent bonds. Conversely if the geometric center of all the δ - charges in an entire molecule is at different point than the geometric center of all δ + charges - i.e. there is a separation of the centers of δ - and δ + charge, the whole molecule is polar. In the later case, we say that the molecule has a fixed dipole.