Chapter 10 Outline The Shapes of Molecules
Introduction
Whether we consider the details of simple reactions, the properties of synthetic material, or the intricate life-sustaining processes of living cells, molecular shape is a crucial factor.
10.1 Depicting Molecules and Ions with Lewis Structures
Lewis structures, also called electron-dot structures or electron-dot diagrams, are diagrams that show the bonding between atoms of a molecule, and the lone pairs of electrons that may exist in the molecule. A Lewis structure can be drawn for any covalently-bonded molecule, as well as coordination compounds.
Using the Octet Rule to Write Lewis Structures
The octet rule tells us that all atoms want eight valence electrons (except for hydrogen, which wants only two), so they can be like the nearest noble gas. Use the octet rule to figure out how many electrons each atom in the molecule should have, and add them up. The only weird element is boron - it wants six electrons.
Lewis Structures for Molecules with Single Bonds
The atoms share a pair of electrons, and that pair is referred to as a bonding pair. The pairs of electrons which do not participate in the bond have traditionally been called "lone pairs". A single bond can be represented by the two dots of the bonding pair, or by a single line which represents that pair. The single line representation for a bond is commonly used in drawing Lewis structures for molecules.
· Hydrogen atoms form one bond.
· Carbon atoms four bonds.
· Nitrogen atoms form three bonds.
· Oxygen atoms form two bonds.
· Halogens form one bond when they are surrounding atoms; fluorine is always surrounding atom.
Lewis Structures for Molecules with Multiple Bonds
When two atoms share a single pair of electrons, the bond is referred to as a single bond. Atoms can also share two or three pairs of electrons in the aptly named double and triple bonds. The first bond between two atoms is called the σ (sigma) bond. All subsequent bonds are referred to as π (pi) bonds. In Lewis structures, two or three lines between the bonded atoms depict multiple bonds. The bond order of a covalent interaction between two atoms is the number of electron pairs that are shared between them. Single bonds have a bond order of 1, double bonds 2, and triple bonds 3. Bond order is directly related to bond strength and bond length. Higher order bonds are stronger and shorter, while lower order bonds are weaker and longer. Resonance: Delocalized Electron-Pair Bonding
Molecular geometry is a three-dimensional representation of atoms in a molecule. Bond lengths and angles must be determined experimentally. There is a simple procedure for predicting molecular geometry or structure called Valence Shell Electron Pair Repulsion (VSEPR). The main premise of this model is that the structure of the molecule can be determined by minimizing the electrostatic repulsion between electron pairs. In other words, bonding pairs and lone pairs try to get as far away from each other as possible.
In some cases after the atoms have arranged themselves according to electron pair repulsion, a double bond may be placed in more than one location. If this happens, then the electrons are considered to be localized. This is not an accurate representation since the electron will delocalize. Our method for depicting this action is to draw all of the allowed positions for the double bond. The term for this is resonance, or we say that it is a resonant structure and our Lewis Dot Structure consists of several drawings with the double bond being moved from position to position. Again, this is not an accurate representation since the electron density actually is spread evenly over all the possible positions. In this experiment we will be using carbonate (CO32-) which has one double-bonded oxygen and two single-bonded oxygen. Formal Change: Selecting the Best Resonance Structure
We use formal charges to select the best structure for a compound. An atom’s formal charge is its total number of atomic valence electrons minus the number of valence electrons it “owns” in the molecule: it owns all of its unshared valence electrons and half of its shared valence electrons (shared valence electrons are the electrons in a bond).

*NOTE *
That the formal charges must sum to the actual charge on the species: zero for a neutral molecule or the ionic charge for an ion. There are three criteria to follow when trying to decide which is the “correct” Lewis structure:
Smaller absolute value of formal charges is preferable to larger ones.
Like charges on adjacent atoms are not desirable.
A more negative formal charge usually resides on a more electronegative atom. Lewis Structures for Exceptions to the Octet Rule
Atoms of the third period or higher may form extended octets by placing extra electrons into the d orbitals of the third energy level. In this way, phosphorus may form five bonds, sulfur and boron six, and the halogens as many as seven. Electron-Deficient Molecules The octet rule for drawing Lewis structures states that eight electrons must surround all atoms. However, boron and beryllium are sometimes stable in compounds with less than an octet. In compounds, Be is generally satisfied with 4 valence electrons and B is generally satisfied with 6. Odd-Electron Molecules There are also molecules in which the total number of valence electrons turns out to be an odd number. Some oxides of nitrogen fall into this category: nitric oxide, NO, and nitrogen dioxide, NO2, which are major components of air pollution and sources of acid rain. In cases like these, the nitrogen, the less electronegative element, is left with an incomplete octet with an unpaired electron.
Another consequence of having unpaired electrons is that certain chemical reactions can be readily understood and even predicted. Odd-electron species are called radicals and are generally very reactive. For example, nitrogen dioxide is observed to form dinitrogen tetroxide under appropriate conditions. This process occurs when the odd electrons on two NO2 molecules pair up to form the N2O4 molecule, a process called dimerization. Expanded Valence Shells Many molecules and ions have more than eight valence electrons around the central atom. An atom expands its valence shell to form more bonds, a process that releases energy. The only way to accommodate additional pairs of electrons is for the central atom to use its empty outer d orbitals in addition to its occupied s and p orbitals. Therefore, expanded octets occur around a central nonmetal from Period 3 or higher, those in which d orbitals are available. Let’s take a look at the Lewis structure for the sulfate ion, SO42-. However, because sulfur is a Period 3 element, it has empty d orbitals that it can use to expand its octet and decrease the formal charges throughout the ion. This is a better Lewis structure for the SO42‾ ion because it has fewer formal charges. Sulfur could expand its octet even further by bringing in two more pairs of electrons from the other two oxygen atoms, however, that would result in a formal charge of -2 on the sulfur, which is less electronegative than oxygen and so less likely to carry the negative charge. While this further expanded octet on sulfur would result in a possible Lewis structure for the sulfate ion, it is not the best Lewis structure.

10.2 Valence-Shell Electron-Pair Repulsion (VSEPR) Theory and Molecular Shape The VSEPR method is used to predict the shapes of molecules and polyatomic ions based on the mutual repulsions among valence-shell electron groups. An electron group is a single unpaired electron, an unshared pair of electrons, or the electrons in a single, double, or triple bond. The number of electron groups is found for each central atom in the molecule or polyatomic ion, and the geometric distribution of these electron groups is assessed. If all electron groups are bonding groups, the molecular geometry is the same as the electron-group geometry. If some of the electron groups are lone pairs, the molecular geometry is derived from (but is different from) the electron-group geometry.

Electron-Group Arrangements and Molecular Shapes The number of regions of high electron density, or VSEPR number, is the number of electron pairs surrounding the central atom; this includes lone pair electrons and bonding electrons. Each lone pair of electrons counts as one region of high electron density. Each bonding pair counts as one region of high electron density. However, multiple bonds count the same as a single bond and contribute only one region of high electron density. (Note: The number of regions of high electron density and VSEPR number refer to the same thing and are used interchangeably.) The Molecular Shape with Two Electron Groups (Linear Arrangement) Beryllium, on the other hand, forms only two pairs of valence electrons. These repel each other at cos-1(-1) = 180°, forming a linear molecule. An example is beryllium chloride, which has two chlorine atoms situated on opposite sides of a beryllium atom. This time, one 2s and one 2p orbital combine to form two sp hybrid orbitals. The two remaining p orbitals sit above and to the side of the beryllium atom (they are empty). Molecular Shapes with Three Electron Groups (Trigonal Planar Arrangement) Trigonal Planar geometry is a shape that is used by an atom that is attached to a combined total of three bonded atoms and nonbonded electron pairs. This shape is usually associated with a system that has two single bonds and one double bond. The structure contains three angles of 1200. All four positions in this geometric shape are located on the same plane. If the three external positions are moved into a different plane from the central atom, then the structure changes into a pyramid.

Molecular Shapes with Four Electron Groups (Tetrahedral Arrangement) Tetrahedral Geometry occurs around any atom that has a combined total of four bonded atoms or nonbonded electron pairs attached to it. The four groups arrange themselves in this geometric pattern to achieve maximum stability by maintaining minimum repulsion between the electron pairs. The angles in this shape consist of a total of six angles at 109.50.

Molecular Shapes with Five Electron Groups (Trigonal Bipyramidal Arrangement) The Trigonal Bipyramid is a geometric pattern that is used by any atom which has expanded its octet to maintain a combined total of five bonded atoms or nonbonded electron pairs. The shape consists of three equatorial positions and two axial positions. Within the structure there are three different size bond angles. The shape contains three angles of 1200, two angles of 1800 and six angles of 900.

Molecular Shapes with Six Electron Groups (Octahedral Arrangement) The Octahedron is a geometric pattern that is used for any atom that has expanded its octet to contain a total of six electron pairs. The structure is characterized as being highly symmetrical. It consists of 3 - 1800 bond angles and 12 -900 bond angles. Using VSEPR Theory to Determine Molecular Shape Use Lewis structures and the VSEPR theory: Valence Shell Electron Pair Repulsion, bonding electrons and lone pairs (the valence electrons) are placed on a sphere as far apart as possible.
Use Lewis structures and theValence bond(VB) theory : Bonding electrons and lone pairs are accommodated in hybridized orbitals, as far apart as possible in three dimensional space. Molecular Shapes with More than One Central Atom
Compounds which need more than one electron to complete the octet will share as many electrons as necessary in order to complete the octet. This 10.3 Molecular Shape and Molecular Polarity In chemical bonds, polarity refers to an uneven distribution of electron pairs between the two bonded atoms—in this case, one of the atoms is slightly more negative than the other. But molecules can be polar too, and when they are polar, they are called dipoles. Dipoles are molecules that have a slightly positive charge on one end and a slightly negative charge on the other. Look at the water molecule. The two lone electron pairs on the oxygen atom establish a negative pole on this bent molecule, while the bound hydrogen atoms constitute a positive pole. In fact, this polarity of water accounts for most of water’s unique physical properties. However, molecules can also contain polar bonds and not be polar. Carbon dioxide is a perfect example. Both of the C—O bonds in carbon dioxide are polar, but they’re oriented such that they cancel each other out, and the molecule itself is not polar. Bond Polarity, Bond Angle, and Dipole Moment The distinction is between Bond Polarity and Molecular polarity. The total polarity of a molecule is measured as Dipole Moment. The actual calculation of dipole moment isn't really necessary so much as an understanding of what it means. Frequently, a guesstimate of dipole moment is pretty easy once you understand the concept and until you get into the more advanced organic chemistry, exact values are of little value.
Basically, the molecular polarity is, essentially, the summation of the vectors of all of the bond polarities in a molecule. The Effect of Molecular Polarity To get a sense of the influence of molecular polarity on physical behavior, consider what effec a molecular dipole might hace when many polar milecules lie near each other, as they do in a liquid. In the liquid state, molecules, so more energy is needed to overcome these strong forces. Therefore the cis isomer has a higher boiling point.

Glossary Lewis Dot Formula (Electron Dot Formula) Representation of a molecule, ion or formula unit by showing atomic symbols and only outer shell electrons.
Resonance The concept in which two or more equivalent dot formulas for the same arrangement of atoms (resonance structures) are necessary to describe the bonding in a molecule or ion.
Electronic Geometry The geometric arrangement of orbitals containing the shared and unshared electron pairs surrounding the central atom of a molecule or polyatomic ion.
Free Radical A highly reactive chemical species carrying no charge and having a single unpaired electron in an orbital.
Lone Pair Pair of electrons residing on one atom and not shared by other atoms; unshared pair.
Bond Order Half the numbers of electrons in bonding orbitals minus half the number of electrons in antibonding orbitals.
Bonding Orbital A molecular orbit lower in energy than any of the atomic orbitals from which it is derived; lends stability to a molecule or ion when populated with electron
Bonding Pair Pair of electrons involved in a covalent bond.
Molecular Geometry The arrangement of atoms (not lone pairs of electrons) around a central atom of a molecule or polyatomic ion.
Tetrahedral A term used to describe molecules and polyatomic ions that have one atom in center and four atoms at the corners of a tetrahedron.
OctahedralA term used to describe molecules and polyatomic ions that have one atom in the center and six atoms at the corners of a octahedron.
Octet Rule Many representative elements attain at least a share of eight electrons in their valence shells when they form molecular or ionic compounds; there are some limitations.
Square Planar A term used to describe molecules and polyatomic ions that have one atom in the center and four atoms at the corners of a square.
Square Planar Complex Complex in which the metal is in the center of a square plane, with ligand donor atoms at each of the four corners.

Questions 1.How Lewis structures depict the aroms and their bonding and lone electron pairs in a molecule of ion?
2.How resonance and electron delocalizarion explain bond properties in many sompounds with double bonds adjacent to single bonds?
3.The meaning of formal chare and how it wis used to select the most iportant resonance structure; the difference between formal charge and oxidation number?
4. The octet rule and its three major exceptions- molecules with a central atom that has an electron defiiciency, an odd number of elctrons, or an expanded valence shell?
5. How electron group repulsions lead to molecular shapes?
6.The five electron group arrangements and their associated molecular shape?
7.Why double bonds and lone pairs cause deviation from ideal bond angles? Get your own web address.
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