Course Outline

General Chemistry--Unit 1

04/25/09

�Before You Begin

To master this material you will need to know basic inorganic nomenclature and periodic law.

Compounds, Ions, and Molecules

An element is a type of matter which is in its simplest form composed of only one type of atom. A compound is a sample of matter that contains two or more elements chemically combined to form a new substance.  There are two fundamental types of compounds: ionic and molecular.

Ionic compounds are composed of positively and negatively charged ions held together by electrostatic attractions casually referred to as ‘ionic bonds.’ Molecular compounds are composed of molecules, which are groups of atoms joined together by pairs of shared electrons. Solid ionic compounds do not contain molecules. Instead, because opposite charges attract one another, the ions stack themselves in crystals such that positive ions are surrounded by negative ions, and vice versa.

The Chemical Bond

A bond is the interaction between two or more atoms that allows them to form a substance different from the independent atoms. This interaction involves the outer electrons of the atoms. These electrons are transferred from one atom to another or shared between them. Bonds may be between atoms of different elements to make a compound, like the two hydrogen atoms and one oxygen atom in a water molecule. But bonds can also be between atoms of a single element. Sulfur is an example of an element that has its most stable form as a small molecule, in this case a ring of eight sulfur atoms. Bonds can be interactions between a few atoms, as in a molecule of water or sulfur. However, bonds may hold hundreds or thousands of atoms together to form a large molecule like insulin or DNA.

There are three fundamental categories of bonds: ionic, covalent, and metallic. Ionic bonds are electrostatic attractions between oppositely charged particles more properly referred to as ionic structure. Covalent bonds are pairs of electrons that are shared more or less evenly between two atoms. Metallic bonds are loosely held outer electrons surrounding packed cations. Metallic bonds are most important in understanding the properties of pure metallic elements and in mixtures of metals (alloys) rather than in compounds, so we will save the discussion of metallic bonds for later.

Ionic Bonds(?)

Ionic compounds in the solid state are held together by electrostatic attractions between opposite charges. Sodium chloride (table salt), silver sulfide (silver tarnish), and hydrated iron (III) oxide (rust) are examples of ionic compounds.

In a compound like silver sulfide, the individual silver atoms have lost electrons and the sulfur atoms have gained electrons. These ions interact to form a solid compound because opposite ions attract one another. When a silver spoon tarnishes, the silver atoms lose electrons to the sulfur atoms in sulfur oxides in the air. This chemical process is called oxidation/reduction. The same compound can also be formed if we mix an aqueous solution of silver nitrate with an aqueous solution of sodium sulfide. This chemical process is called metathesis. In both cases, silver sulfide ionic compound is formed.

Is there such thing as a silver sulfide molecule? The answer to this question depends on who you ask. Most chemistry text books refer to the attractions between oppositely charged ions as ‘ionic bonds’ yet qualify this as referring to the forces that are involved in the formation of the overall three dimensional structure of an ionic crystal. Certainly, there isn’t anything similar to a molecule inside an ionic crystal.

The Covalent Bond

Electronegativity and Bonding

Covalent bonds are pairs of electrons shared between two atoms in a molecule. Pure covalent (also called non-polar) bonds are ones in which both atoms share the electrons evenly. By evenly, we mean that the electrons have an equal probability of being at a certain radius from the nuclei of either atom. Polar covalent bonds are ones in which the electrons have a higher probability of being in the proximity of one of the atoms.

Electronegativity is the periodic property that indicates the strength of the attraction an atom has for the electrons it shares in a bond. Atoms with high electronegativities tend to hold tightly to their electrons or to form negative ions. These elements are found to the upper right on the periodic table. Atoms with low electronegativities tend to have a lower attraction for their electrons and may form positive ions. These elements are found to the lower left on the periodic table.

A pure covalent or non-polar bond has difference of about 0.5 or less in the electronegativities of the two atoms. A pure covalent bond can form between two atoms of the same element (such as in diatomic oxygen molecule) or atoms of different elements that have similar electronegativies (such as in the carbon and hydrogen atom in methane). A polar covalent bond is a pair of electrons shared between two atoms with significantly different electronegativities (from about 0.5 to 2.0 difference). These bonds tend to form between highly electronegative non-metals and other non-metals, such as the bond between hydrogen and oxygen in water.

In compounds that have elements with very different electronegativities (greater than 2.0 difference), the electrons can be considered to have been transferred to form ions. We classify these compounds as being ionic. In the gas phase, pairs of these ions have electron densities similar to extremely polar bonds. This brings up the philosophical question of whether or not there is a fundamental difference between ionic and covalent compounds on the very smallest level (one more time we’re asking if there is such a thing as a silver sulfide molecule!). Many chemists prefer to call the attraction between oppositely charged ions in an ionic crystal an ‘ionic structure’ and reserve the word ‘bond’ for shared electrons. In that case, small clusters of gaseous ions have pairs of shared electrons and can be thought of as molecules, but solid ionic crystals do not.

Many of the properties of a compound, such as solubility and boiling point, depend, in part, on the degree of the polarity of its bonds. For a comparison of the properties of the different types of compounds in their solid state see The Solid State.

Recall from our study of quantum mechanics that we cannot know precisely both where an electron is and where it is going. But we can use quantum mechanics to find where the electrons are most likely to be. We can draw a diagram of a bond that looks like a contour map. The contour lines on the bond map correspond to bond density and can be thought of as the likelihood that a bonding electron will fall within that contour at a given instant in time.  For a pure (or non-polar) covalent bond, the contours fall symmetrically about the nuclei. For a polar covalent bond, the contour lines are more closely space near the atom with higher electronegativity.

Number and Kinds of Covalent Bonds: The Lewis Model

We do not need to use quantum mechanics for the most basic understanding of covalent bonds. Remember that a theory is an attempt to explain observations, and theories (also called models) are retained if they are useful, even if they do not adequately explain all observations. Molecules are so varied in their behaviors that several different models can be used to help us understand their different aspects. The earliest theory that tried to describe the electron arrangements in molecules was developed in the early 1900s by G. N. Lewis. This theory was contemporary with the Bohr atomic model. Like Bohr’s atomic theory, this model has been supplanted by quantum mechanics, but it can still be used to predict the number and kinds of covalent bonds in many molecules.

According to the Lewis model, molecules are held together by pairs of electrons shared between two atoms. The only electrons available for sharing are the valence, or outer shell, electrons. Atoms that have a full valence tend to be more stable than ones that do not have a full outer shell (think noble gases). When atoms bond together, their electrons are pooled to fulfill valence requirements and make the molecule stable.

Lewis Structures

Lewis dot structures where introduced by G. N. Lewis in 1912 as a means of determining the number of bonds needed to fulfill valence stability requirements of a molecule. These diagrams are still useful in making a first order approximation (or best first guess). The diagrams use the atomic symbol to represent the nuclei and core electrons. Valence electrons are represented by dots arranged in pairs at the top, bottom, left and right of the atomic symbol. The dot structure for the neon atom with eight valence electrons looks like this:

A hydrogen atom has one electron. A diatomic hydrogen molecule has two atoms and two electrons total. If molecules are held together by bonds, the two electrons must be shared between the two atoms. The Lewis structure of the hydrogen molecule looks like this:

with the pair of dots or the line representing a pair of shared electrons (a bond).  Note that a line represents a pair of electrons.

Atoms will share pairs of electrons so that each has a full outer shell, if possible. This is known as the octet rule, since the outer shell of most of the non-metals can hold at most eight electrons (there are many exceptions, including hydrogen molecule!). The fluorine atom has seven valence electrons. A diatomic fluorine molecule has two atoms and fourteen valence electrons total. If each atom has three electron pairs that it does not share and one electron pair that act as a bond, both atoms fulfill the octet rule (6 free electrons + 2 shared for each atom). The Lewis structure for the fluorine molecule looks like this:

Bond Multiplicity:

Atoms can share more than one pair of electrons. A single bond is one pair of electrons shared between two atoms. A double bond is two pairs of electrons shared between two atoms. The oxygen atom has six valence electrons, so the diatomic oxygen molecule has two atoms and twelve valence electrons. In order for the oxygen molecule to obey the octet rule, it must have a double bond:

A triple bond is three pairs of electrons shared between two atoms. The nitrogen atom has five valence electrons, so the diatomic nitrogen molecule has two atoms and ten valence electrons. In order for the nitrogen molecule to obey the octet rule, it must have a triple bond:

It is extremely rare for molecules to have higher multiplicity than triple bonds. For all practical purposes, single, double and triple bonds are the only possibilities, at least for beginning chemistry students.

Larger Molecules

Molecules with only two atoms have limited possibilities for bonding. However, most molecules have several atoms, sometimes hundreds! The first dilemma when drawing the Lewis structure of a larger molecule is to determine the total number of electrons to use. Find the number of valence electrons for each atom and multiply by the number of atoms of that element in the molecule. For example, carbon dioxide has 4 + 2*6 = 16 valence electrons, and ammonia has 5 + 3*1 = 8 valence electrons. Keep track of the number of valence electrons and evaluate your Lewis structures to make sure that the number of electrons used matches the number available!

The second dilemma is to decide how to arrange the atoms. The blank paper can be very intimidating. However, even fairly small groups of atoms can be arranged many different ways. Use the following guidelines to reduce the number of likely possibilities then use formal charge to determine which of the possible Lewis structures is best:

1. Molecules with more than two atoms usually have one central atom to which the other atoms are joined. The atom with the lowest electronegativity is central. For example, the carbon tetrachloride molecule has the four chorine atoms bonded to the carbon atom rather than a line with the order C-Cl-Cl-Cl-Cl.
2. The exception to the above this guideline is hydrogen. Hydrogen has a rather low electronegativity. However, because it has only one electron, it can only form one bond. For this reason, hydrogen atoms will be on the edges or ends of molecules. For example, the ammonia molecule has a nitrogen atom in the center with three hydrogen atoms joined to it, even though nitrogen has a higher electronegativity.
3. Most atoms don’t form rings or chains except for carbon (sometimes silicon and sulfur). Unless you know differently, do not arrange atoms so that they form closed figures such as triangles and squares or long lines. For example, the ozone molecule, O₃, does not have its atoms arranged in a triangle, even though a valid Lewis structure can be drawn for that shape.

Concept Check: What is the Lewis structure for carbon dioxide?

Answer:

For the Lewis structure of the carbon dioxide molecule, we decide that the carbon atom has the lowest electronegativity and should be in the center of the diagram. The next problem is to decide where to place the oxygen atoms. Different students may come up with seemingly different possibilities.

The important thing to note is that all of the Lewis structures drawn from these beginnings are equivalent. The Lewis structure is not a real molecular shape; so, as long as both oxygen atoms are bonded to the carbon atom, the diagrams are conveying the same general information. On the other hand, a Lewis diagram in which the atoms are arranged carbon to oxygen to oxygen is not equivalent. In that arrangement, an oxygen atom rather than the less electronegative carbon atom is the central atom.

The next step is to fill in enough electrons to hold the molecule together. The Lewis diagram for carbon dioxide must have at least one bond between the carbon and each oxygen atom for the molecule to stay together.

When we compare the number of electrons used up to this point (four) with the number of valence electrons available (fourteen) we see that this Lewis structure is not finished. Besides, the atoms do not obey the octet rule. Fill in non-bonding electron pairs until the all the atoms have an octet OR until you have used all the electrons that are available.

If you run out of electrons before all of the atoms have an octet, the Lewis diagram should have multiple bonds. In this case, the structure lacks two pairs for the oxygen on the right to have an octet. This implies that the correct structure should have two multiple bonds.

Watch out, you can’t just draw in two more bonds. The other atoms already obey the octet rule in this sketch. Extra bonds will cause them to have too many. Start over. Arrange the bonds; then fill in the non-bonding electrons as before. This will give a molecule with two double bonds and two non-bonding electron pairs on each oxygen atom.

Polyatomic Ions

Polyatomic ions are groups of atoms joined together by covalent bonds, like molecules, except that they have a charge. The only difference between the process for drawing Lewis diagrams for molecules and for ions is in determining the number of available electrons. Count the valence as before then add an electron for each unit of negative charge or subtract one for each unit of positive charge. For example, the sulfite ion has the formula SO₃^2-. The number of valence electrons is 6 + 3*6 + 2 = 26.

The atom with the highest electronegativity in the sulfite ion is the sulfur, so it belongs in the center of the diagram with oxygen atoms bonded to it. Fill in enough bonds to hold the ion together then add non-bonding electrons until all atoms obey the octet rule or you have used all that are available.

Molecules That Do Not Obey the Octet Rule

Not all atoms in all molecules obey the octet rule. Some small atoms with few electrons may have fewer than eight electrons. Hydrogen, for example, always violates the octet rule because it can have at most one bond. Boron is another element that may violate the octet rule by having too few electrons. The correct Lewis structure for boron trichloride has three single bonds and zero non-bonding electron pairs on the central boron atom.

Molecules with an odd number of valence electrons always violate the octet rule. One of the electrons in the Lewis diagram will be un-paired. An example is nitrogen dioxide which has seventeen valence electrons.

If the central atom in a molecule is from period three or lower on the periodic table, it may violate the octet rule by having too many electrons (up to twelve), sometimes called an extended or expanded octet. How can you tell when a molecule has an extended octet? After determining the central atom, placing bonds, and filling in non-bonding pairs, you may have electrons left over. This signals that the molecule violates the octet rule (or that you miscounted the valence!). Place the remaining electrons, in pairs, on the central atom, if it is from period three or lower. How can they do this? We have not yet discussed the quantum mechanics involved in bond formation, but we can hint that atoms low on the periodic table have empty d orbitals in their outer shell.

!Warning! If the molecule seems to have too many electrons to obey the octet rule, it has an extended octet; if it has too few electrons, the molecule needs multiple bonds. Using the wrong remedy makes the situation worse rather than better.

Formal Charge

Formal charge is a method of determining which of several seemingly valid Lewis structures is best. The idea behind formal charge is that a molecule or polyatomic ion will be more stable if each of its atoms shares the fewest number of electrons possible without gaining or losing several electrons. The formal charge is an electron bookkeeping device. Each atom ‘owns’ its non-bonding electrons. To find the formal charge on an atom, count one electron for each bond. Subtract this number and the number of non-bonding electrons from the number of valence electrons. Repeat this process for all of the atoms in the Lewis structure. The best Lewis structure is one in which all the formal charges are as close as possible to zero.

For example, consider the diatomic chlorine molecule. Each atom has seven valence electrons. One possible Lewis structure is a molecule with one single bond between the two atoms. In that structure, each atom has a formal charge of zero (seven valence minus one bonded and six non-bonded electrons), which is good. Another possibility is for the atoms to form an ionic bond. One atom would gain an electron to form the stable chloride ion.  The other atom would lose an electron to form a very unstable positive one charged ion. The formal charges would be -1 and +1. The situation in which both atoms have formal charges of zero is better (reflects a more stable molecule) than the situation in which one has a -1 formal charge and the other has a +1 formal charge. 

Note that sum of the formal charges in a molecule will be equal to zero, since all of the electrons are assigned and counted. The sum of the formal charges of a polyatomic ion will equal the charge of the ion. An example is the sulfite ion, SO₃^-2, which will have a sum of formal charges equal to -2, no matter how high and low the individual formal charges are. One valid Lewis structure for the sulfite ion obeys the octet rule. It has three single bonds and a non-bonding pair on the central sulfur atom. Each of the oxygen atoms has a formal charge of 6 – 1 – 6 = -1 (valence minus one for each bond minus the number of non-bonding electrons). The sulfur has a formal charge of 6 – 3 – 2 = +1. The sum of these four formal charges equals -2, the same as the charge on the ion. A better Lewis structure has a double bond between the sulfur and one of the oxygen atoms (remember that it doesn’t matter which one in the diagram). The single bonded oxygen atoms still have a formal charge of -1. The double bonded oxygen has a formal charge of 6 – 2 – 4 = 0. The sulfur atom has a formal charge of 6 – 4 – 2 = 0. The sum of the four formal charges is still -2, but now two of the four atoms have formal charges of zero. This is much better, even if it violates the octet rule. Remember that atoms in period three or lower can have up to twelve electrons in an expanded octet.

What about giving the sulfite ion another double bond? An atom with an expanded octet can have up to twelve electrons. We can draw a dot structure with two double bonds and a single bond. It still has a non-bonding pair on the sulfur atom, which gives it a formal charge of -1. This is just as good as (but not better than) the structure with one double bond. We can also make valid Lewis structures in which the double bond is on the right, on the left, or below the sulfur atom. Which is best? To answer this question, we need to expand our model to include quantum mechanics, which we will do in the next unit.

!Warning! It is very easy to confuse formal charge with oxidation number. The oxidation number is used to determine if an element gains or loses electrons during an oxidation-reduction reaction. For bookkeeping purposes, the atom with higher electronegativity is assigned both electrons in a bond. Formal charge is used to determine which Lewis structure is better. For bookkeeping purposes, bond electrons are split evenly between atoms regardless of their relative electronegativity.

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