Organic chemistry chapter 2

Draw a lewis dot structure for the azide ion, N3-

An alternative method for determining the Lewis dot structure of azide ion. Note that there are 3 contiguous charges on the 3 nitrogen atoms, though all have full octets. This is not ideal- see step 7A above. Is there an alternate structure with fewer charged atoms that may be preferable? See the problem set below. There may be more than one acceptable lewis structure (i.e. does not violate a-e of step 6: Final Structure Check, above.) In order to compare two or more structures to determine which is energetically the most favorable, prioritize as follows: First priority: any structures (i) showing 2nd row atoms C, N, O, F with full octets and (ii) without unpaired electrons will be more stable than those with atoms not containing full octets. Second priority: if all atoms have full octets, then the molecule possessing the fewer number of atoms with formal charges (or the same number but with more charge separation) is the more energetically stable and more significant. Additionally, if all atoms have full octets, placement of negative formal charges on relatively more electronegative atoms is preferable, as is having positive charges on relatively more electropositive atoms.

Summarize: how can one compare ionic, covalent, and polar covalent bonds?

Bond polarity and ionic character increase with increasing differences in EN between the bonding atoms. In ionic bonding, the electron attracting properties of each atom are so different that an electron may be fully transferred from one atom to the other. In NP covalent bonding, the electron attracting properties of each atom are so similar that the electrons in the bond(s) are shared essentially equally between the two atoms. In a polar covalent bond, the EN difference between atoms is great enough to result in more electron density residing in one of the two atoms, but is too small o allow a full electron transfer from one atom to the other. To qualitatively determine into which category a bond falser realize that bonds with high ionic character form mainly between the metals in columns 1 and 2 (Li, Na, K, Be, etc) and the nonmetals (eg C, N, O, F, Cl, Br, I). NP covalent bonds form between two of the same atoms such as in molecular hydrogen H2 and in carbon-carbon bonds. In addition, the carbon-hydrogen bond is considered essentially non-polar. All other bonds between different non-metal elements, such as H-Cl, C-O, B-H, C-F, N-O, etc., are considered polar covalent. Polar covalent bonds have modest EN differences between the bonded atoms, and therefore more electron density resides on the more EN atom, resulting in partial charges. These bonds have significant covalent as well as ionic character.

What is a covalent bond?

Carbon has four valence electrons, and would thus have to gain four more electrons to achieve an octet with neon's configuration. However, carbon gaining an octet via a transfer of 4 electrons from another atom would be energetically disfavored (eg a carbon atom with a 4- charge bonded to another atom with a 4+ charge) G.N. Lewis came up with an elegant solution to this problem by proposing the covalent bond. Lewis described the sharing of pairs of bonding electrons between two atoms. When two bonding atoms essentially do not differ in their EN values, such as C-C, H-H, etc., they form a n on-polar covalent bond. Note that a covalent bond is denoted by a line representing a pair of shared bonding electrons between the two atoms. When there is no difference in electronegativity values between the two bonded atoms, the bonding electrons are shared equally between them. Each atom is thus neutral, as it has neither gained nor lost any electron density. As we will see, covalent bonding can enable atoms within molecules to attain a full octet or closed shell noble gas configuration. What is considered a negligible difference in defining non-polar covalent bonding? Some sources state that the EN difference between the two atoms should be approx less than or equal to 0.4 on the Pauling scale. However, in organic chemistry it is important to understand trends than to memorize numbers. You will not be asked to quantitatively determine EN differences between atoms...

Why then do bonds form despite the gain in energy due to charge repulsion?

In order to understand the driving force for bond formation one needs to consider both potential and kinetic energy. Recall that potential energy is "stored" energy, resulting from an object's position or arrangement of its components. If an object moves, its potential energy transforms into kinetic energy, the energy resulting from its motion. The increase in electrostatic repulsion upon bonding results in a gain in potential energy. Therefore, the favorable energetics promoting bonding must be due to the lowering of the kinetic energy of the atoms upon bond formation. Quantum mechanical calculations have indeed verified that the degree of KE lowered upon bond formation will more than offset the PE gained, resulting in overall energy lowering and net stabilization.

Valence Shell Electron-Pair Repulsion (VSEPR) Theory

Lewis structures alone do not afford information about the three-dimensional structure of molecules. A model commonly used to compensate for this is VSEPR theory. VSEPR theory assumes that each atom in a molecule will achieve a geometry that minimizes the repulsion between valence electrons.

Clarifying a common misconception about why bonds form.

Most sources describe bonding between two atoms as energetically favorable due to a combination of (i) electrostatic attraction between the negatively charged electrons of one atom to the positively charged nucleus of its bonding partner, and (iii) shielding of the electrostatic repulsion between the two positively charged nuclei by their intervening shared electrons (the bonding electron pairs shown below in each molecule in red), specifically in the case of covalent bonding. However, repulsive forces actually increase upon bond formation. Despite bonding electron pairs shielding inter-nuclear positive charge repulsion, some repulsion will still occur and will be greater compared to zero repulsion between isolated non-bonded atoms. Moreover, the sharing of electrons in a covalent bond concentrates electron density relative to isolated atoms, resulting in increased electrostatic negative charge repulsion.

Can a bond every be 100% ionic or 100% covalent?

No bonds are ever 100% ionic. This would mean that one of the two bonding atoms had a zero electronegativity value, which is not true even of the most electropositive elements. There is instead a continuum of partial ionic/partial covalent character, depending on the specific atoms involved. There can, however, be 100% covalent bonding when two of the same atoms bond, as in H2.

What is the octet rule?

Octet is a term used mainly for the row 2 elements B, C, N, O and F in their noble gas, or the Ne configuration with a filled valence shell of 8 electrons. Third and higher row elements, when using 8 electrons to create a filled outer valence shell, are also described as obeying the octet rule. Row 3 elements can exceed the octet rule due to their larger orbitals. Most commonly we will observe > an octet for P and S. Phosphorus can create 3 single bonds (neutral), 4 single bonds (+1 formal charge) and 5 single bonds (neutral). Sulfur can form 1-6 total bonds, including multiple bonds.

Drawing a lewis dot structure from a molecular formula.

One basic procedure for drawing the lewis dot structures of small molecules is shown in figure 1, using Lewis atomic symbols (described in chapter 1). Go over figure 1 to familiarize yourself with wiht the general concepts before proceeding to steps 1-7 and the other examples described in the text. The text goes into more detail than figure 1, covering various types of structures. The azide anion is used to illustrate an alternative simple method useful for molecules containing 1st-3rd row elements that will tend to obey the octet rule. Examples: C2H4 C2H2 ClO4- and N3- Step 1: Draw the lewis atomic symbols showing the valence electrons around each atom. The derivation of Lewis atomic symbols was described in chapter 1. The atoms needed for the example molecules are C, H, N, O and Cl, so we will need the following: Step 2: predict atom positions. Place carbon(s) at/near the center of your developing partial structure. Otherwise place the largest (or most electropositive) atom in the center. Also realize that H (always) and the halogens (X, almost always) form only 1 bond, so place them at th periphery of the structure (figure 1 part A). Step 3: Places ingle bonds between adjacent atoms. Step 4: check for octets: count all shared and lone pair electrons (or put a box around all shared and lone pair electrons, figure 2, left column) for each atom. If octets are complete and all electron are paired, skip step 5. Step 5: if there are unfilled octets, move nonbonding electrons to form multiple bonds. Keep electrons paired and keep octets full whenever possible. Pairing electrons can be accomplished via one or more of the following 3 moves: (a) pair unpaired electrons on adjacent atoms to create a multiple bond (b) if the molecule has an overall charge (see perchlorate and azide examples below), add or subtract the appropriate number of electrons and check formal charge. Step 6: check formal charge assignments Formal charge on an atom = periodic table column- (lone pair electrons + 1/2 the bonding electrons. A systematic method for formal charge determination is shown in figure 2 (right column), wherein one places a box around an atom's lone pair of electrons and one electron from each of its bonds to readily see the formal charge electron count. Compare the number of boxed electrons to the atom's periodic table column number to assign the formal charge, as described in figure 2. Step 7: final structure check (a) small molecule organic compound structures with charges on three or more atoms should be re-evaluated. (b) if the row 2 elements (B,C,N,O,F) exhibit a formal charge other than 0, +1, or -1, the structure should be re-evaluated. This is also generally true for other common non-metal atoms in columns 4-7. (c) due to their d-orbitals it is relatively common for S, Si and P to have > an octet of electrons. (d) among the common row 2 elements (B, C, N, O, F), only C and B may have < a full octet in a Lewis structure. This is because B and C are relatively electropositive compared to N, O and F, meaning that they can better "handle" an electron deficit than the more EN row 2 atoms. (e) None of the common row 2 elements (B,C,N,O,F) should ever have > an octet

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Return to practice.

Which has more covalent bond character, (a) O-Li or (b) O-Na. Note that each can be written correctly as either a covalent bond or as an ionic bond.

Since Na is less En than Li, the Na-O bond must be more ionic than the Li-O bond (there is more of an EN difference between the Na and O atoms compared to the Li and O atoms).

Why do the bond angles of NH4+, NH3, and H2O successively decrease?

The lone pairs occupy somewhat more space than the electrons in a bond, thereby repelling and compressing the angles between the H-N-H and H-O-H bonds to a slightly greater degree than bonding electrons. NH4+ has no lone pairs, ammonia has one and water has two, accounting for the observed trend. Note that each has four electron domains.

How can one predict the 3-dimensional array of the atoms in a molecule, as described in table 1?

The minimization of electron pair repulsion and resulting molecular shape can be formulated by first using the Lewis dot structure of a molecule to predict the total number of lone pairs and bonding partners of an atom. For the purpose of VSEPR theory lone Paris and bond partners are grouped together, referred to collectively as "electron domains" or "steric number" depending on the source. For ex, the central carbon atom in CO2 has two bonded atoms and no lone Paris, and thus two total electron domains. The farthest apart the electrons in the bonds to the oxygen atoms can be to best minimize electron repulsion is thus 180°, and therefore CO2 is linear. in the case of methane, CH4, the central carbon has four bonded atoms. The steric number of electron domains is thus four. A three-dimensional tetrahedral structure with H-C-H bond angles of 109.5° keeps the bonding electrons as far apart as possible. Boron in BF3 has three attached atoms.. The optimal structure keeping the bonding electrons as far apart as possible is trigonal planar, with F-B-F bond angles of 120°.

Perchlorate anion example...

The next two solved examples, the perchlorate and azide anions, involve two of the less straight forward analyses. After working through these latter examples and applying the procedure and checks above, one will have encountered most of the common challenges involving the creation of lewis structures.

Why use the term "formal" charge of an atom? Why not just use the word "charge"?

The reason is that in specific cases the charge assigned to an atom does not physically "reside" on that atom, but rather is dispersed onto the atoms it is bonded to. This occurs when an EN atom such as O and N have positive formal charge or when electropositive elements such as B are assigned a negative formal charge. Because this situation is a common occurrence and potential point-of-confusion, it is explained in more detail below: In the example shown in figure 5, in the more stable Lewis dot structure, the oxygen has a positive formal charge. Since oxygen is too electronegative to have less than an octet, wouldn't it also be too electronegative to have a formal positive charge? Electronegativity refers to how much an atom in a bond "wants" both bonding electrons. This leads to bond dipoles, or, for example, partial charges on each atom (delta minus on the more EN atom). Apart from the electron count used in the formal charge calculations, how can one physically justify an EN atom bearing a positive charge? The common examples in this course are: oxygen with 3 bonds and nitrogen with 4 bonds. Conversely, how can. a relatively electropositive atom such as boron have a negative charge (when it has 4 bonds)? Formal charges are used to determine the overall charge associated with an entire molecule. When a formal charge is assigned to an individual atom, the charge can be viewed as localized on that atom except when there is seemingly an inconsistency between the formal charge and atom's electronegativity. Examples include a (+) charge on electronegative O or N or a (-) charge on electropositive B, as noted above. In these latter cases, the charge is not viewed as localized on the specific atom that it is assigned to, but instead as dispersed onto the atoms surrounding the charged atom. For ex, in BH4- the hydrogens are actually polarized as hydrides. or H-. Likewise, in the left structure in figure 5, the H and C atoms attached to the positively charged oxygen actually have enhanced partial positive charge character. However, when an electronegative atom is drawn with a negative charge the charge is indeed viewed as being located on the electronegative atom. Note: Carbon is in the "middle" of row 2, with 4 or 8 (half) the valence electrons of the row 2 noble gas neon. It can hold either a negative or a positive formal charge, and in each case the charge is mainly localized on the carbon atom. This ability to possess either a positive or a negative charge, in addition to its being the only row 2 element capable fo having as many as 4 covalent bonds without creating charge, is further evidence of carbon's uniqueness and versatility.

Does one need to know exact EN values for distinguishing ionic, covalent, or polar bonds?

There are EN values assigned to each atom, enabling on etc assign quantitative values; however, numbers vary according to sources. As mentioned above, but worth repeating, in this course, unlike gen chem for ex, it is much more practical (and necessary) to be able to make qualitative predictions.

Example:

Two different Lewis dot structures for CH3O are shown in figure 5. Neither violates a-e of step 6. In the left structure, carbon and oxygen have full octets. In the right structure, carbon has less than a full octet. The structure on the left is thus the more stable (significant), despite the more EN atom (O) having a positive charge, since the octet rule is 1st priority.

What is an ionic bond?

When two atoms differ considerably in their EN values, such as K and F, the more electropositive atom can transfer an electron to the more electronegative atom. This results in a cation and an anion. The ensuing electrostatic attractive force between the oppositely charged ions is an ionic bond. In other words, an ionic bond results upon the transfer of an electron from a metal (far left side of the periodic table) to a non-metal (right side). Example: K loses an electron (K is oxidized) to F (F gets reduced) resulting in the salt K+F- Importantly, note that when this compound forms, K loses its valence electron to have the 18 electron configuration of Ar. In the case of F, it gained an electron to complement its original seven valence electrons, giving it a full valence shell of eight, giving it the same configuration as the noble gas Ne. These atoms bond in a manner that enables them to fill their outer (valence) shells to attain stable noble gas configurations, exactly as described in chapter 1. The more electropositive atom (K) can readily achieve the Ar configuration by losing one electron. Ionic bonds and ionic compounds generally have electronegativity differences of approx greater than or equal to 2.0. Electronegativity values can be found in chat 1, figure 6.

What is a polar covalent bond?

When two atoms differ in their electronegativity values, but are relatively close to each other in the periodic table, they form a polar covalent bond. Because there is some difference in EN between the atoms, the bonding electrons are shared unequally. Each atom is partially positive (less EN) and partially negative (more EN). An example is the C-O bond. We note this partial charge by using the symbol "delta" d as follows, placing the partial negative charge on the more EN oxygen to show that it has more electron density than the carbon to which it is bonded: The EN difference between C and O is 0.89. Polar covalent bonds will have EN differences that are in between those of ionic and NP bonds, or approx >0.4 and <2.0.