Covalent+Bonding

The octet rule states that atoms will bond in such a way to completely fill their valence shells. A full valence shell consists of eight electrons on the outermost electron shell of the atom. When a full valence shell is achieved, the atom loses energy and becomes more stable. There are some violations to the octet rule, however. In some cases, atoms can form covalent bonds and end up with too few electrons. When hydrogen forms a covalent bond, it only ends up with two electrons, a property that allows it to form special single bonds with other elements. In other cases, covalent bonds result in too many electrons around the central atom in a compound. These are //hypervalent// atoms. Cases such as these are the result of highly electronegative atoms (such as Cl, F, and O) connecting to atoms with lone pairs. Covalent bonding tends to occur between atoms that have similar electronegativities. Note that alkali metals and alkaline Earth metals do not form covalent bonds, only ionic bonds.
 * __I. Covalent Bonding__**
 * a.** Covalent bonding is a form of chemical bonding in which electrons are //shared// between atoms rather than transferred as they are in ionic bonding. This electron sharing produces an attraction between the atoms that holds them together as a compound. When sharing electrons, atoms tend to try to fill their outer electron shells, so most will bond according to the //octet rule//.


 * __Naming Binary Covalent Compounds__**

Binary covalent compounds contain two different types of nonmetals, so they can be recognized by their formulas. For example, N2O is a binary covalent compound because it contains only two different elements and they are both nonmetals. There are five basic steps to naming these compounds. Given N2O5:

1. Observe the subscript of the first element. If it is greater than one, add one of the prefixes to the beginning of the name. The subscript for the nitrogen in this compound is 2, so the name begins with //di-//

2. After the prefix, write the name of the first element: //dinitrogen//.

3. Observe the subscript of the second element. Even if it is one*, add one of the prefixes after a space to separate it from the name of the first element: //dinitrogen pent-//

4. Write the root name of the second element after this prefix: //dinitrogen pentox-//

5. Add //–ide// to the end of the name: //dinitrogen pentoxide//.

The name of the binary covalent compound N2O5 is dinitrogen pentoxide.


 * Prefixes:**

Subscript of element: Prefix:

1. mon(o) 2. di 3. tri 4. tetr(a) 5. pent(a) 6. hex(a) 7. hept(a) 8. oct(a) 9. non(a) 10. dec(a) 11. undec(a) 12. dodec(a)


 * Root Names for Nonmetals:**

N – nitr O – ox I – iod C – carb S – sulf H – hyd Br – brom Se – selen Cl – chlor F – fluor P – phosph As – arsen


 * When naming a binary covalent bond that contains a hydrogen atom and a halogen, it is understood that there is one bond because hydrogen atoms always form one covalent bond and halogens usually form one covalent bond. The prefix in front of the second element is not required. The compound HBr can either be called hydrogen bromide or hydrogen monobromide.


 * __Polarity of Molecules:__**

Polarity means that something has a positive and a negative charge. In a polar molecule, positive and negative charges build up on different ends of the molecule according to the flow of electrons. Vectors are drawn to represent the different charges, and their resultant, depending on its magnitude, determines whether or not the molecule is polar and in which direction the electrons are flowing. Polarity depends on the electronegativity of the elements in a compound. Electronegativity is the ability of an atom to attract electrons. If the electronegativity of one atom is sufficiently higher than the others in the compound, then a flow of electrons will be produced towards that atom. A dipole moment, or vector showing the magnitude and direction of the flow of electrons, will be drawn towards this atom.


 * Bond Types:**

A bond can either be polar covalent, non-polar covalent, or ionic. These bond types depend on the difference in electronegativity between the atoms.

1. For a polar covalent bond, the electronegativity difference is 0.0-0.5. There is a balanced sharing of electrons, and so no charge is building up. 2. For a non-polar covalent bond, the electronegativity difference is 0.5-1.7. The sharing of electrons is not balanced. There is a flow of electrons towards the atom with the higher electronegativity, and so there is a partial negative charge building up on this side of the molecule. A partial positive charge is building up on the atom with the lower electronegativity. 3. For an ionic bond, the electronegativity difference is 1.7-4.0. The electrons are captured by the atom with the greater electronegativity.

To find the electronegativity of each of the atoms, you must use an electronegativity chart:


 * Electronegativity Differences:**

=
Electronegativity is defined as an atoms ability to attract electrons. Every element has a certain electronegativity. As you go across a period in the periodic table, electronegativity increases. As you go down a group, electronegativity decreases. Using the electronegativity values for the central atom and a branching atom, one can decide whether or not that particular bond is nonpolar, polar, or ionic. One can go on from there to determine whether or not the entire molecule falls in either of those three categories by analyzing the overall dipole that is formed.======

=
In order to decide whether or not a bond is polar, nonpolar, or ionic, subtract the electronegativity of the branching atom from that of the central atom. The classification of the molecule depends on the range that the difference falls in.======

Diatomic molecules are all nonpolar covalent because the difference in the electronegativities of the two atoms will be 0.
 * Electronegativity Difference (x) || Bond Type ||
 * 0.0≤x≤0.3 || Nonpolar Covalent ||
 * 0.3<x≤1.7 || Polar Covalent ||
 * 1.7<x≤4.0 || Ionic ||

Example: code Sodium Chloride (NaCl)

Electronegativity of Sodium: 0.9 Electronegativity of Chlorine: 3.0

3.0-.9 = 2.1

2.1 is greater than 1.7, NaCl is an ionic compound

code


 * Dipole Moments:**

If you can draw a molecular dipole (overall vector of electron flow), then you have a polar molecule. This molecular dipole is the resultant of the sum of each of the dipole moments in the molecule.

Example:

The molecule H2O (water) has a bent geometry with four regions of electron density. Two hydrogen atoms (electronegativity= 2.1) bond with a central oxygen atom (electronegativity = 3.5) at an angle of 109.5˚. The difference in the electronegativities of hydrogen and oxygen is 1.4 (3.5-2.1=1.4). The resulting molecular dipole from these two dipole moments goes towards the oxygen atom with a magnitude of 1.4, set between the two hydrogen atoms at an angle. This 1.4 is within polar range, so water is a polar molecule.

Sometimes a molecule contains polar covalent bonds because charges are building up between the atoms, but the dipole moments balance each other out enough to make the molecule non-polar overall.


 * Polarity and Chemistry:**

The forces of attraction between molecules in the liquid state depend on the polarity of the molecules. Because polar molecules are attracted by opposite charge, the polarity of molecules determines the amount that their kinetic energy will have to increase before they are able to break free from each other and enter the gas form. Polarity is therefore a key part of determining the boiling point of a particular compound. The polarity of the solvent and solute molecules in a solution affects solubility. Polar tends to dissolve polar. Non-polar tends to dissolve non-polar. Some substances have both polar and non-polar ends because they contain different compounds that are suspended within each other as an emulsion. Soap has this property. The non-polar ends of soap dissolve the non-polar dirt and oils in clothing, which in turn dissolves in water (a polar molecular) due to the polar ends. This allows the unwanted non-polar compounds to be carried away in the wash.

**Lewis Dot Diagrams**

When atoms bond covalently, they can be represented using Lewis Dot Diagrams. Consider these diagrams a very literal representation of the atom’s structure.

Atoms want exactly eight electrons to orbit its nucleus. This is called **The Octet Rule.** An atom with the right amount of electrons would look like this: When drawing a Lewis diagram, replace the center of the atom with its abbreviation from the periodic table. Now, as you will remember from the chapter on periodic table trends, chlorine is a negative ion. Cl-. The “-” signifies that chlorine needs one electron to make eight. In order to get the last electron, it can form a covalent bond with another atom, such as Na+. Each atom “lends” its unpaired electron to the other, resulting in a single bond.

The most notable exception to the octet rule is hydrogen. Hydrogen can only hold two electrons, and a single hydrogen atom has only one electron. While atoms such as carbon need to form multiple bonds to reach an octet, hydrogen can only form one single bond with one other atom. = H : H = However, single bonds aren’t the only way for atoms to bond. They can also form double and triple bonds, in which atoms share four or six electrons.
 * N::O::N:

Two double bonds.
 * C:::N:H

A single bond and a triple bond.

Ring Structures Here’s a ring structure.  The single lines represent a single bond, and the double lines double bonds. The nitrogen is represented by the “N,” but the intersections between bonds are each a CH. As every ring structure has at least one carbon atom, the carbons are often not written out. However, here is an alternate drawing of the same ring structure. The literal "ring" of carbon and nitrogen is the ring structure.

 Organic Naming

The naming of organic molecules is not difficult. It is like a code, where each part of the name means something in relation to the molecule.

Here is a structure and its name. 3-methylpentonic acid.

Let’s start with: 3-methyl**pent**anoic acid. This root word means “5.” It is used because the longest chain of carbons has 5 bonded together.

Here is a chart detailing each root word:
 * ** Number of Carbons ** ||   || ** Root Word **  ||


 * 1 ||  || meth ||


 * 2 ||  || eth ||


 * 3 ||  || prop ||


 * 4 ||  || but ||


 * 5 ||  || pent ||


 * 6 ||  || hex ||


 * <span style="text-align: center; display: block; font-family: 'Times New Roman'; font-size: 14pt; msobidifontsize: 9.0pt;">7 ||  || <span style="text-align: center; display: block; font-family: 'Times New Roman'; font-size: 14pt; msobidifontsize: 9.0pt;">hept ||


 * <span style="text-align: center; display: block; font-family: 'Times New Roman'; font-size: 14pt; msobidifontsize: 9.0pt;">8 ||  || <span style="text-align: center; display: block; font-family: 'Times New Roman'; font-size: 14pt; msobidifontsize: 9.0pt;">oct ||


 * <span style="text-align: center; display: block; font-family: 'Times New Roman'; font-size: 14pt; msobidifontsize: 9.0pt;">9 ||  || <span style="text-align: center; display: block; font-family: 'Times New Roman'; font-size: 14pt; msobidifontsize: 9.0pt;">non ||


 * <span style="text-align: center; display: block; font-family: 'Times New Roman'; font-size: 14pt; msobidifontsize: 9.0pt;">10 ||  || <span style="text-align: center; display: block; font-family: 'Times New Roman'; font-size: 14pt; msobidifontsize: 9.0pt;">dec ||


 * <span style="text-align: center; display: block; font-family: 'Times New Roman'; font-size: 14pt; msobidifontsize: 9.0pt;">11 ||  || <span style="text-align: center; display: block; font-family: 'Times New Roman'; font-size: 14pt; msobidifontsize: 9.0pt;">undec ||


 * <span style="text-align: center; display: block; font-family: 'Times New Roman'; font-size: 14pt; msobidifontsize: 9.0pt;">12 ||  || <span style="text-align: center; display: block; font-family: 'Times New Roman'; font-size: 14pt; msobidifontsize: 9.0pt;">dodec ||

<span style="font-family: Times New Roman; font-size: 14pt; mso-bidi-font-size: 12.0pt; msobidifontsize: 12.0pt;">Here is the next part of the name: 3-methylpent**an**oic acid. This means that all of the carbon molecules have single bonds, no double or triple. Were there double or triple bonds, it would be “pent**en**oic.”

3-methylpentan**oic acid**. There is a carboxylic acid at the end of the chain of five carbons. These acids will only ever appear at the end of the molecule and of its name.

3-**methyl**pentanoic acid. There is a second chain of carbons branching off of the first. This branch has one carbon (**meth)** and is followed by “**yl**” because it is a group and not the main chain.

** 3-methyl ** pentanoic acid. The one carbon group branches off of the third carbon in the chain of five. Therefore, the entire name begins with “**3.”** <span style="font-family: Times New Roman; font-size: 14pt; mso-bidi-font-size: 12.0pt; msobidifontsize: 12.0pt;"> This method for naming can be used on any organic molecule.

<span style="line-height: 115%; font-family: 'Palatino Linotype', 'Book Antiqua', Palatino, serif; font-size: 18pt;">**Valence-Shell Electron Pair Repulsion Theory (VSEPR)**

<span style="font-family: 'Palatino Linotype', 'Book Antiqua', Palatino, serif; font-size: 110%;"><span style="line-height: 115%; font-family: 'Palatino Linotype', 'Book Antiqua', Palatino, serif; font-size: 14pt;">- The VSEPR Theory is used to determine molecular geometry (the shapes of molecules) <span style="font-family: 'Palatino Linotype', 'Book Antiqua', Palatino, serif; font-size: 110%;"><span style="font-family: 'Palatino Linotype', 'Book Antiqua', Palatino, serif;">-<span style="font-family: 'Palatino Linotype', 'Book Antiqua', Palatino, serif; font-size: 110%;"> The VSEPR Theory states that the repulsion caused by valence shell electrons determines the molecule's shape. <span style="font-family: 'Palatino Linotype', 'Book Antiqua', Palatino, serif;">- <span style="line-height: 115%; font-family: 'Palatino Linotype', 'Book Antiqua', Palatino, serif; font-size: 14pt;">The number of  <span style="font-family: 'Palatino Linotype', 'Book Antiqua', Palatino, serif;"><span style="line-height: 115%; font-family: 'Palatino Linotype', 'Book Antiqua', Palatino, serif; font-size: 14pt;">__electron domains__ in a molecule and the number of __lone pair__ electrons determine the shape of a molecule. <span style="font-family: 'Palatino Linotype', 'Book Antiqua', Palatino, serif; font-size: 14pt;"> -<span style="line-height: 115%; font-family: 'Palatino Linotype', 'Book Antiqua', Palatino, serif; font-size: 16pt;"> <span style="line-height: 115%; font-family: 'Palatino Linotype', 'Book Antiqua', Palatino, serif; font-size: 14pt;">An __electron domain__ is the area around a central atom that is occupied by electrons. It is represented by either single bonds, double bonds, triple bonds or lone pairs. <span style="font-family: 'Palatino Linotype', 'Book Antiqua', Palatino, serif;">-  <span style="font-family: 'Palatino Linotype', 'Book Antiqua', Palatino, serif;">ATTENTION: A single bond, a double bond, a triple bond or a lone pair each count as one electron domain.

For example:

CO2 has two electron domains. Its molecular geometry is linear.

Here’s a table to help you…

<span style="line-height: 115%; font-family: 'Arial Black', 'sans-serif'; font-size: 14pt;">..................................................................................↓ Molecular Geometry <span style="line-height: 115%; font-family: 'Arial Black', 'sans-serif'; font-size: 14pt;">↓
 * Electron Domains || Electron Domain Geometry || Predicted Bond Angles || Hybridization of Central Atom || 0 Lone Pair || 1 Lone Pair || 2 Lone Pairs ||
 * 2 || Linear || 180° || sp || Linear || <span style="font-family: 'Kristen ITC'; background: yellow; color: yellow; font-size: 14pt; -moz-background-clip: border; -moz-background-origin: padding; -moz-background-inline-policy: continuous;">.................. || <span style="font-family: 'Kristen ITC'; background: yellow; color: yellow; font-size: 14pt; -moz-background-clip: border; -moz-background-origin: padding; -moz-background-inline-policy: continuous;">…… ……  ||
 * 3 || Trigonal Planar || 120° || sp2 || Trigonal Planar || Bent || <span style="font-family: 'Kristen ITC'; background: yellow; color: yellow; font-size: 14pt; -moz-background-clip: border; -moz-background-origin: padding; -moz-background-inline-policy: continuous;">…… …… ||
 * 4 || Tetrahedral || 109.5° || sp3 || Tetrahedral || Trigonal Pyramidal || Bent ||
 * 5 || Trigonal Bipyramidal || 90°, 120° || sp3d || Trigonal Bipyramidal || Seesaw || T-Shaped ||
 * 6 || Octahedral || 90° || sp3d2 || Octahedral || Square Pyramidal || Square Planar ||



<span style="font-family: 'Palatino Linotype', 'Book Antiqua', Palatino, serif; font-size: 130%;">Linear..................................Trigonal Planar.......................Bent...........................Tetrahedral <span style="font-family: 'Palatino Linotype', 'Book Antiqua', Palatino, serif; font-size: 130%;">Trigonal Pyramidal........Trigonal Bipyramidal.........Seesaw...............................T-Shaped



<span style="font-family: 'Palatino Linotype', 'Book Antiqua', Palatino, serif; font-size: 130%;">Octahedral...........................Square Pyramidal........................Square Planar <span style="font-family: 'Palatino Linotype', 'Book Antiqua', Palatino, serif; font-size: 150%;">
 * Hybridization**

__Hybridization__ is defined as the mixing of valence atomic orbitals to give a new set of orbitals that reflect the geometry of the molecule. In other words hybridization is the mixing of two or more orbitals of similar energies to produce new orbitals of equal energies. <span style="font-family: 'Palatino Linotype', 'Book Antiqua', Palatino, serif;"> Hybridization explains how the orbitals of an atom rearrange to form covalent bonds. It also helps determine the shape of the molecule and the bond angles that form within that particular molecular geometry.

<span style="font-family: 'Palatino Linotype', 'Book Antiqua', Palatino, serif; font-size: 120%;"><span style="font-family: 'Palatino Linotype', 'Book Antiqua', Palatino, serif; font-size: 110%;">The new orbitals that are formed during hybridization are called __hybrid orbitals__. The number of hybrid orbitals produced is the same as the number of orbitals that have combined.

<span style="font-family: 'Palatino Linotype', 'Book Antiqua', Palatino, serif; font-size: 120%;">Let's see an example:

Methane, CH4 -When CH4 forms, carbon has to hybridize its orbitals. -Carbon has four valence electrons. It has two in the 2s orbital and two in the 2p orbital. -Carbon's one 2s orbital and three 2p orbitals (including the empty 2p orbital) combine to form four sp3 orbitals. -Four sp3 orbitals are formed because the number of hybrid orbitals produced is the same as the number of orbitals that have combined. -The hybrid orbitals are called sp3 because one s orbital was used and 3 p orbitals were used.

<span style="font-family: 'Palatino Linotype', 'Book Antiqua', Palatino, serif; font-size: 110%;">-<span style="font-family: 'Palatino Linotype', 'Book Antiqua', Palatino, serif; font-size: 120%;">The sp3 orbitals are 25% s-shaped and 75% p-shaped. There new shape looks like this... - <span style="font-family: 'Palatino Linotype', 'Book Antiqua', Palatino, serif; font-size: 120%;">The four sp3 orbitals seen here make a tetrahedral shape with 109.5 degree bond angles. -<span style="font-family: 'Palatino Linotype', 'Book Antiqua', Palatino, serif; font-size: 120%;">These sp3 orbitals have the same energy. However their energy is more than the 2s orbital but less than the 2p orbital. -Whenever hybridization happens the hybrid orbitals always have an energy that is between the level of orbitals that have combined.

These videos explain it even better....

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 * Sigma and Pi Bonds:**

In covalent bonding, there are three major types of bonds: Sigma bonds, Pi bonds, and coordinate covalent bonds.

Sigma Bond (<span style="font-family: Times New Roman, serif;">σ):

A sigma bond is the name for the bond that exists when any two atoms share electrons. A sigma bond can be created out of shared electrons in overlapping p, s, or hybrid orbitals between the atoms. The bond physically exists in the space directly between the two nuclei of the atoms. A molecule with sigma bonds is CH <span style="font-family: Times New Roman, serif;">4. CH<span style="font-family: Times New Roman, serif;">4 has four single bonds; overlapping regions of electron density.

Most covalent bonds contain a sigma bond. An exception to this are covalent bonds made through coordinate covalent bonding. In this case, one atom shares two or more electrons, while the other atom shares none. An example of a molecule with a coordinate covalent bond is NH<span style="font-family: Times New Roman, serif;">4+. NH<span style="font-family: Times New Roman, serif;">3 has one lone pair. When an H<span style="font-family: Times New Roman, serif;">+ ion with no electrons comes near it, the H+ ion will be attracted to NH<span style="font-family: Times New Roman, serif;">3. This will form NH<span style="font-family: Times New Roman, serif;">4+ <span style="font-family: Times New Roman, serif;"> Sigma bonds

Pi Bond (π)

Every double and triple bond is composed of not only a sigma bond, but pi bonds as well. Pi bonds are formed when two p orbitals on each atom stretch and overlap, surrounding the sigma bond that is also formed. In a double bond, two pi bonds are formed and sandwich the sigma bond, while in a triple bond, four pi bonds are formed that surround the sigma bond. Pi bonds act as reinforcements to the sigma bond, making the entire bond stronger. However, pi bonds are not as strong as sigma bonds. <span style="font-family: Times New Roman, serif;">



<span style="font-family: Times New Roman, serif;">Pi and Sigma Bonds