Covalent Bonding and Molecular Compounds Section 6.2 Covalent Bonding and Molecular Compounds

Introduction Most chemical compounds are molecules, a neutral group of atoms that are held together by covalent bonds. A molecule is a single unit capable of existing on its own.

Molecular Compound A chemical compound whose simplest units are molecules. They may be made up of one type of atom only (O2) or of two or more different atoms (H2O).

Diatomic Molecules A molecule containing only two atoms (O2 or CO).

Chemical vs. Molecular Formulas Chemical formula: Indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts. Molecular formula: Shows the types and numbers of atoms combined in a single molecule of a molecular compound.

So What’s The Difference? A chemical formula could either be a molecular formula or an empirical formula – describes the mole ratio of atoms in a particular compound.

So What’s The Difference? H20 is both the molecular formula and the empirical formula for water C6H12O6 is the molecular formula for glucose. The empirical formula is CH2O, which represents the simplest ratio of atoms in the compound.

Why Bond? Nature favors chemical bonding because most atoms are at a lower potential energy when bonded to other atoms than they are at as independent particles.

Formation of a Covalent Bond From before, it is a trade off between the attractive and repulsive forces between electrons and the nuclei of atoms. Attraction: Decrease in potential energy of the atoms. Repulsion: Increase in potential energy of the atoms.

The Potential Energy Well (see Fig. 6-5, p. 165) Relative strength of attraction and repulsion between the atoms depends on the distance separating them. When two atoms are far apart, they don’t influence each other.

The Potential Energy Well (see Fig. 6-5, p. 165) As the atoms approach each other, attractive forces dominate and potential energy is lowered. The distance at which the repulsive forces equal the attractive forces is the bond length, potential energy is at a minimum and a stable molecule forms.

Nature favors arrangements where potential energy is minimized. Bonding occurs at the energy minimum (the repulsion of like charges equals the attraction of unlike charges).

Characteristics of the Covalent Bond The molecule’s electrons can be pictured as occupying overlapping orbitals, moving about freely in either orbital.

Bond Length The distance between two bonded atoms at their minimum potential energy (the average distance between the two bonded atoms).

Bond Energy The energy required to break a chemical bond and form neutral isolated atoms. Really the same amount that was released when the bond formed, but with different sign.

The Lengths and Energies Vary With the types of atoms that have combined and may vary between the same two atoms, depending on what other bonds the atoms have formed.

This fills the s and p orbitals, Octet Rule Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet (8) electrons in its highest occupied energy level. This fills the s and p orbitals, just like a noble gas

This fills the s and p orbitals, Octet Rule Example: a molecule of Fluorine, F2 Bonding electron pair in overlapping orbitals F 1s 2s 2p F 1s 2s 2p This fills the s and p orbitals, just like a noble gas

Exceptions to the Octet Rule Hydrogen (why?) Lithium Beryllium Boron Molecules with an odd number of electrons (N-O).

Expanded Valence (More Than 8) Involves bonding in d orbitals as well as s and p orbitals. Usually occurs with bonding the highly electronegative elements F, O, and Cl.

Electron Dot Notation An electron configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the element’s symbol.

Write the electron dot notation for Phosphorus Xenon Aluminum Sodium You Try… Write the electron dot notation for Phosphorus Xenon Aluminum Sodium

Lewis Structures Using electron-dot notation to represent molecules. G. N. Lewis 1875 - 1946

Lewis Structures Formulas in which atomic symbols represent nuclei and inner-shell electrons dot-pairs or dashes represent electron pairs in covalent bonds dots adjacent to only one symbol represent unshared electrons.

Lone Pairs Also called unshared pairs. A pair of electrons that is not involved in bonding and that belongs exclusively to one atom.

Structural formula: Indicates the kind, number, arrangement, and bonds, but not the unshared pairs of the atoms in a molecule. Ex. F-F H-Cl

Single Bond A covalent bond produced by the sharing of one pair of electrons between two atoms.

Multiple Bonds Double and triple bonds are possible. This happens when atoms share 2 or 3 electron pairs to form covalent bonds to conform to the octet rule. They have progressively shorter bond lengths and higher bond energies than single bonds between the same atoms. Extremely common for C, N, and O. Example: N2

See Table 6-2, p. 173 As should be noted from the table, multiple bonds between pairs of C, N, and O are very possible. Multiple bonds are needed when there are not enough valence electrons to complete the octet.

Steps For Making Lewis Structures Not in the text: You will need lots of scratch paper (on assignments I don’t want to see all of the steps, just the final result).

#1 Determine the number of atoms in each element present in the molecule.

#2 Write the electron dot notation for each type of atom in the molecule.

#3 Determine the total number of valence electrons in the atoms to be combined.

#4 Arrange the atoms to form a skeleton structure for the molecule. If carbon is present, it is the central atom. Otherwise, the least electronegative atom is central (except for hydrogen, which is never central). Then connect the atoms by electron-pair bonds.

#5 Add unshared pairs of electrons so that each hydrogen atom shares a pair of electrons and each other nonmetal is surrounded by eight electrons.

This is the electron dot #6 Count the electrons in the Lewis structure to be sure that the number of valence electrons used equals the number available. This is the electron dot Lewis structure

#7 If too many electrons have been used, subtract one or more lone pairs until the total number of valence electrons is correct. Then move one or more lone pairs to existing bonds between atoms until they “see” 8 electrons. This forms multiple bonds.

Hag’s Last Step (Not in Text) Show all bonds as lines Show all lone pairs on the central atom(s) Show any lone pairs that may influence the polarity of the molecule Show all possible resonance structures

This Results In….. The structural Lewis diagram (also called Lewis structure or structural diagram). NOTE: All polyatomic ions and resonance structures are enclosed by brackets. The overall charge of an ion must be shown.

Resonance Bonding in molecules or ions that cannot be correctly represented by a single Lewis structure.

Experimental Evidence Shows that there is not a mixture of single and double bonds, but that all of the bonds are identical, so there is an average of the structures. We show all possible structures, using a double-headed arrow.

Try these examples... CH2O (formaldehyde) HCN CO2 HF NO3- (nitrate anion) NH4+ (ammonium cation)

Covalent-Network Bonding They do not contain individual molecules, but instead can be pictures as continuous, three-dimensional networks of bonded atoms. Diamond (C)

Assignment 2 Worksheets: Lewis Structures and 6.2 Review Text p. 175, #1-4 Due Thursday, BOP