Reading Notes for Chapter 4

These are Dr. Bodwin's reading notes for Chapter XX of "Introduction to Chemistry". I am using a local .pdf copy that was downloaded in August 2020.

Chapter Summary:

In Chapter 3, we looked at ionic compounds. In this chapter, we look at the other major classification of substances: covalent compounds. Whereas ionic compounds lose or gain electrons to give charged particles that interact with one another, covalent compounds share electrons in chemical bonds. The electrons they share are the ourtermost shell of valence electrons, so because the atoms in these compounds share electrons with their neighbors, they are called covalent.

Covalent Bonds:

When elements are far apart on the Periodic Table, one of the atoms probably has a very strong attraction to electrons and the other probably does not, so it makes sense to lose and gain electrons to form an ionic compound. But what if the two atoms have a very similar (or even identical) attraction for electrons? For either of them to lose or gain an electrons does not result in an energetically favorable result. In cases like this, it would be more energetically favorable to allow both atoms to "see" the valence electrons - to share rather than transfer.
A covalent bond is the pair of electrons that is being shared between two atoms.

NOTE: There are cases where chemists might talk about bonds that consist of only 1 electron or 3 electrons. Although interesting, those special cases are beyond the scope of this course so we won't be talking about them beyond this note. For the purpose of this class, all covalent bonds are 2-electron bonds.

Bonding pairs vs. Non-bonding pairs - in a covalent molecule, the bonding electrons are those valence electrons that are being shared with another atom. If there are other lone pairs of electrons on an atom, they are called non-bonding electrons. Non-bonding electrons help define the molecular shape of a molecule.
What it it's BOTH?!?!?! We've actually already seen this in the previous chapter. Polyatomic ions are held together by covalent bonds, but interact with external ions to form ionic compounds. As an example, let's consider potassium cyanide, KCN. The C & N share SIX electrons to form a triple (covalent) bond and there is a non-bonding lone pair on each atom for a total of 6+2+2=10 valence electrons. Neutral carbond has 4 vlaence electrons and neutral nitrogen has 5 valence electrons (for a ttoal of 9), so there is an extra electron that gives cyanide its negative charge. The negative cyanide ion reacts ionically with potassium ions to form the ionic compound KCN.

NOTE: We typically classify compounds based upon their strongest interaction. Since ionic interactions are generally stronger than covalence interactions, compounds containing polyatomic ions are classified as "ionic" even though there are covalent interactions within the polyatomic ion.

Covalent Compound Formulas & Names:

Naming covalent compounds is a blessing and a curse - many compounds tell us very clearly what they are (diphosphorus pentoxide), BUT there are a LOT of "common" names that you just have to memorize, like "water". If you didn't know what "water" was, it's pretty unlikely that you would come up with H₂O based upon the name, but you could probably make a pretty safe guess that diphosphorous pentoxide is P₂O₅.
For formulas and names, compare and contrast the formula-writing and naming conventions between ionic compounds and covalent compounds. There are many similarities, and some key differences that can help you identify different compounds and the interactions present in them.

Multiple Bonding:

Multiple bonds form to decrease the overall energy of a structure. ("lower energy" = "more stable") Don't try to cram double bonds into everything, just use them where they give your structure some benefit - like when a double bond allows atoms to complete their octet.
Hydrogen never forms a double bond.
Carbon tends to form 4 bonds. If tthere are not enough other atoms, try using a double bond between carbon atoms OR between carbon and oxyegn OR between carbon and nitrogen.
Fluorine never forms double bonds.

Bond Lengths:

When we talk about "bond length", we're looking at the nucleus-to-nucleus distance.
Bonds are overlapping orbitals that allow electrons to be shared. Larger atoms can overlap farther apart from each other, so their bond legths are longer.
Look at Table 4.2 in your textbook and compare some typical bond lengths. H-H is about 74pm, H-N is about 100pm, and H-I is 161pm. As we go down the P.T., atomic size increases because there are more and more shells of electrons. Now look at C-C, C-N, C-O, N-N, and O-O... those atoms are all very lcose to the same size, and their bond lengths are very similar.
Multiple bonds are always shorter than single bonds between the same elements.

Electronegativity & Polar Bonds:

This is the big gray area between the extremes of "ionic" and "covalent".
A perfect, 100% ionic interaction involves the complete transfer of one or more electrons between atoms/chemical species.
A perfect, 100% covalent interaction involves the equal sharing of one or more valence electrons between atoms/chemical species.
So Cs-F is a great example of an ionic interaction, and a C-C bond is a great example of a covalent interaction. But what about something like C-F? Fluorine only needs 1 additional electron to complete its valence octet, so it would make sens that it is really attracted to the electron that carbon is sharing with it. Carbon doesn't really want to lose any electrons, but it does not hold the electrons that it is sharing with fluorine quite as tightly as fluorine does. That means that the electrons are more strongly drawn to fluorine, so (on average) there is a little more electrons density around fluorine (as compared to a neutral fluroine atom) than there is on carbon (as compared to a neutral carbon atom). This makes the fluorine end of that bond look a little bit negatively charged and the carbon end of that bond look a little bit positively charged. This is a polar covalent bond and because it has a slightly positive end and a slightly negative end, it can interact with other C-F bonds very strongly compared to covalent bonds that are not polar.
The polarity of a bond is measured by a quantity called electronegativity with values shown in Figure 4.4 of your textbook. Fluorine is the most electronegative element, and the farther we get from fluorine, the less electronegative an element is.
A couple words of caution about electronegativity...

Electronegativity is a calculated quantity. There are a number of different systems for calculating electronegativity... your book uses Pauling Electronegativities. This is by far the most common electronegativity scale.
Electronegativity is only really useful as a comparative value when looking at covalent bonds. the electronegativity of sulfur is 2.5. Is that a large EN value or a small EN value? It depends upon what we're comparing it to! If we're looking at a S-F bond, then 2.5 seems kind of small, and in that bond, sulfur would be the positive end of the bond; if we're looking at a Ge-S bond, then 2.5 seems pretty big, and in that bond, sulfur would be the negative end of the bond!
As a general rule of thumb, bonds are only appreciably polar if their Pauling EN values differ by 0.5 or more. The place this becomes most important is when we look at carbon-containing compounds... we can treat C-H bonds as if they are non-polar (in most cases) even though there is a difference of 0.4 in their Pauling Electronegativities.

Polar Molecules:

Just because a bond is polar, doesn't mean that the resulting molecule will be polar. If the polarity is of equal size and opposite direction, it will cancel itself out and yield a non-polar molecule.
We can predict molecular shapes using Valence Shell Electron Pair Repulsion (VSEPR) Theory. Since electrons are negatively charged, and negative charges repel one another, the vlaence shell electrons (both bonding pairs and non-bonding pairs) with spread out as much as possible.

Simple Organic Molecules:

In this case, "simple" means molecules that contain only 4 or fewer carbon atoms.
Hydrocarbons are named using a root that tells you how many carbon atoms there are in the longest chain and a suffix that tells you how many multiple bonds there are.
Meth- means 1 carbon
Eth- means 2 carbons
Prop- means 3 carbons
But- means 4 carbons
-ane means the structure contains only single bonds (it's "saturated")
-ene means there is a double bond
-yne means there is a triple bond

Methane is CH₄ and is the principle component of natural gas.
Ethyne (commonly known as acetylene) is C₂H₂ and is a common gas used for welding.
Propane is C₃H₈ and is the gas usually used in gas grills and some portable heaters.
Butane is C₄H₁₀ and is the liquid you see inside disposable cigarette lighters.

Alcohols - contain a C-OH unit (hydroxyl group) in place of a C-H unit. Their names end in -ol, as in methanol or ethanol.
Ethanol is the "alcohol" present in alcoholic beverages like beer and wine.
iso-Propanol is rubbing alcohol.

Carboxylic Acids - contain a -COOH unit (carboxyl group) in place of a CH₃ unit. Their names and in -anoic acid as in propanoic acid.
Ethanoic acid (common name: acetic acid) is the acid found in household vinegar.

Hydrocarbons are non-polar (only C-C nad C-H bonds) and are often gases, liquids, or greasy low-melting solids
Both hydroxyl and carboxyle groups are polar, so they increase the strength of interactions between molecules.
Methane is a gas that turns liquid at -160°C. Methanol is a liquid that boils at +65°C. Polar interactions are MUCH stronger than non-polar interactions.

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