Section 6-2: Covalent Bonding and Molecular Compounds Coach Kelsoe Chemistry Pages

Section 6-2 Objectives Define molecule and molecular formula. Explain the relationships between potential energy, distance between approaching atoms, bond length, and bond energy. State the octet rule. List the six basic steps used in writing Lewis structures.

Section 6-2 Objectives Explain how to determine Lewis structures for molecules containing single bonds, multiple bonds, or both. Explain why scientists use resonance structures to represent some molecules.

Intro to Chemical Bonding Atoms rarely exist as independent particles in nature. Almost everything we encounter in nature consist of multiple atoms held together by chemical bonds. A chemical bond is a mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together.

Covalent Bonding Most chemical compounds are composed of molecules. A molecule is a neutral group of atoms that are held together by covalent bonds. A single molecule of a chemical compound is an individual unit capable of existing on its own. The reason it can is because of their neutrality. A molecule can be atoms of the same element or of many different elements.

Covalent Bonding A chemical compound whose simplest units are molecules is called a molecular compound. A chemical formula indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts. This is a broad category that includes molecular formula.

Covalent Bonding A molecular formula shows the types and numbers of atoms combined in a single molecule of a molecular compound. –Examples  H 2 O, C 6 H 12 O 6, NH 3 A diatomic molecule is a molecule containing only two atoms. –Examples  HCl, O 2, Cl 2

Formation of a Covalent Bond Nature favors chemical bonding because most atoms are at lower potential energy when bonded to other atoms. To help understand how this happens, let’s consider two hydrogen atoms. –Alone, the positively charged proton is responsible for keeping the negatively charged electron at a far enough distance. –Together, each atom has help because one atom’s nucleus attracts the other’s electron.

Research Notes, Page 180 Ultrasonic Toxic-Waste Destroyer Are we succeeding in eliminating hazardous wastes being produced by factories? Are we fighting for a lost cause by pushing for higher standards of toxic waste elimination? What are some benefits of Professor Hoffmann’s device? What are some drawbacks?

Characteristics of Covalent Bonds In covalent bonds, electrons of bonded atoms can be pictured as occupying overlapping orbitals, moving about freely, as in figure 6-7. The distance between two bonded atoms at their minimum potential energy, or the average distance between two bonded atoms is the bond length. When atoms bond, energy is released. Why? The amount of energy released equals the difference between the potential energy of the separated atoms and that of the bonded atoms.

Characteristics of Covalent Bonds The energy difference between the separated atoms and bonded atoms is the same amount of energy required to break the bonds. Bond energy is the energy required to break a chemical bond and form neutral isolated atoms. Bond energy is usually measured in KJ/mol. –For example, to break 1 mol of HF, it needs 569 KJ. In diatomic molecules, the longer the chemical bonds are, the weaker the bonds are. Look at table 6-1 in the columns on the left.

Characteristics of Covalent Bonds Remember Pauli’s exclusion principle: that no two electrons in the same atom can have the same set of four quantum numbers; primarily, electrons can not spin in the same direction. Therefore, two atoms that share electrons, the atoms must spin in opposite directions. For example, in hydrogen: H H H H _____ 1s _____ 1s _____ 1s _____ 1s

The Octet Rule Noble gases already possess a minimum of energy due to their electron config. Their stability is because their atoms’ outer s and p orbitals are completely filled with eight electrons. Other main-group (s and p block) atoms can fill their outermost s and p orbitals with electrons by sharing them through covalent bonding; such bond formation follows the octet rule.

The Octet Rule The octet rule states that chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet (8) of electrons in its highest occupied energy level. H _____ 1s F _____ 1s _____ 2s _____ 2p H _____ 1s F _____ 1s _____ 2s _____ 2p

Exceptions to the Octet Rule Most main-group elements tend to form covalent bonds according to the octet rule, but there are exceptions. Hydrogen, for example, forms bonds in which it is surrounded by 2 electrons. Boron tends to form bonds in which it is surrounded by only six electrons. Other elements can be surrounded by more than 8 electrons when they combine with fluorine, oxygen and chlorine. This is called expanded valence.

Electron-Dot Notation Covalent bond notation usually involves only the electrons in an atom’s outermost energy levels, or the valence electrons. Electron-dot notation is 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.

Electron-Dot Notation In general, an element’s number of valence electrons can be found by adding the superscripts of the element’s noble gas notation. Some examples of electron-dot notation: NaCFMg

Lewis Structures Electron-dot notation can also be used to represent molecules. The pair of dots represent the shared electron pair of the hydrogen-hydrogen covalent bond. HH

Lewis Structures For a molecule of fluorine, F 2, the electron-dot notations of 2 atoms of F: The shared pair of dots represent the shared electrons in a covalent bond. There are 3 pairs of unshared electrons around both atoms as well. FF

Lewis Structures An unshared pair, also called a lone pair, is a pair of electrons that is not involved in bonding and that belongs exclusively to one atom. FF The pair of dots representing a shared pair are usually replaced by a long dash.

Lewis Structures These drawings are called Lewis structures, formulas in which atomic symbols represent nuclei and inner-shell electrons, dot-pairs or dashes between two atomic symbols represent electron pairs in covalent bonds, and dots adjacent to only one atomic symbol represent unshared electrons. FF

Lewis Structures A structural formula indicates the kind, number, arrangement, and bonds, but not the unshared pairs of atoms in a molecule. For example: Lewis structures for many molecules can be drawn if one knows the composition of the molecule & which atoms are bonded. FFH ClHF

Lewis Structures A single covalent bond, or a single bond, is a covalent bond produced by the sharing of one pair of electrons between two atoms. FFH ClHF

Writing Lewis Structures Step 1: Determine the type and number of atoms in the molecule. –Example: iodomethane – CH 3 I Step 2: Write the electron-dot notation for each type of atom in the molecule. –See board Step 3: Determine the total number of valence electrons in the atoms to be combined. –14 e -

Writing Lewis Structures Step 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. –See board

Writing Lewis Structures Step 5: Add unshared pairs of electrons so that each hydrogen atom shares a pair of electrons and each other nonmetal is surrounded by 8 electrons. –See board Step 6: Count the electrons in the structure to be sure that the number of valence electrons used equals the number available. Dashes represent 2 electrons.

Multiple Covalent Bonds Atoms of some elements, especially carbon, nitrogen, and oxygen share more than one electron pair. A double bond, is a covalent bond produced by the sharing of two pairs of electrons between two atoms. A triple bond is a covalent bond produced by the sharing of three pairs of electrons between two atoms.

Multiple Covalent Bonds Example of a double bond: Example of a triple bond: How do you know if there is a triple bond? You don’t usually! CC H H H H NN

Multiple Covalent Bonds Double and triple bonds are referred to as multiple bonds. Double bonds in general have higher bond energies and are shorter than single bonds. Triple bonds are even stronger and shorter. Remember when doing Lewis structures for C, N, or O that multiple bonds can occur.

Sample Problems Draw the Lewis structure for methanal, CH 2 O, which is also known as formaldehyde. Draw the Lewis structure for C 2 H 4.

Resonance Structures Some molecules cannot be represented adequately by a single Lewis structure, i.e. ozone. (pg. 175) Scientists used to believe that these bonds would “resonate,” or alternate between both forms, but that is not the case. Resonance refers to bonding in molecules that cannot be correctly represented by a single Lewis structure.

Vocabulary Bond energy Bond length Chemical formula Diatomic molecule Double bond Electron-dot notation Lewis structures Lone pair Molecular compound Molecular formula Molecule Multiple bond Octet rule Resonance Single bond Structural formula Triple bond Unshared pair