Covalent Bonding

1. Covalent Bonding Chapter 9

2. Section 9.1 The Covalent Bond

3. Worldwide, scientists are studying ways to increase food supplies, reduce pollution, and prevent disease. Understanding the chemistry of compounds that make up fertilizers, pollutants, and materials that carry genetic information is essential in developing new technologies in these areas. An understanding of the chemistry of compounds requires an understanding of their bonding.

4. Why do atoms bond? You have learned that all noble gases have particularly stable electron arrangements. This stable arrangement consists of a full outer energy level. A full outer energy level consists of two valence electrons for helium and eight valence electrons for all other noble gases. Because of this stability, noble gases, in general, don’t react with other elements to form compounds. You also learned that when metals and nonmetals react to form binary ionic compounds, electrons are transferred, and the resulting ions have noble-gas electron configurations. But sometimes two atoms that both need to gain valence electrons to become stable have a similar attraction for electrons.

5. Sharing of electrons is another way that atoms can acquire the electron configuration of noble gases. • The octet rule states that atoms lose, gain, or share electrons to achieve a stable configuration of eight electrons, or octet.

6. What is a covalent bond? • Atoms in these other compounds share electrons. • The chemical bond that results from the sharing of valence electrons is called a covalent bond. • In a covalent bond, the shared electrons are considered to be part of the complete outer energy level of both atoms involved. • Covalent bonding generally happens when elements are relatively close to each other on the periodic table. • The majority of covalent bonds form between nonmetallic elements. • A molecule is formed when two or more atoms bond covalently. You know that certain atoms, such as magnesium and chlorine, transfer electrons from one atom to another, forming an ionic bond. However, the number of ionic compounds is quite small compared with the total number of known compounds. What type of bonding is found in all these other compounds that are not ionically bonded?

7. Formation of a Covalent Bond • Hydrogen, oxygen, nitrogen, chlorine, bromine, iodine, and fluorine occur in nature as diatomic molecules because the molecules are more stable in pairs than singly. • HONClBrIF • Why do two atoms that do not give up electrons bond with each other?

8. Consider fluorine’s (F2) electron configuration. Each fluorine molecule has seven valence electrons and must have one additional electron to form an octet. As two fluorine molecules approach each other, two forces become important. A repulsive force occurs between the like-charged particles. An attractive force occurs between oppositely-charged particles. As they move closer, the nuclei become more attracted to the other’s valence electron until maximum attraction is achieved. At that point, the attractive forces balance the repulsive forces. A covalent bond now exists between the two atoms.

9. Single Covalent Bonds • When a single pair of electrons is shared, a single covalent bond forms. • Single covalent bonds are also called sigma bonds. • A sigma bond occurs when the electron pair is shared in an area centered between the two atoms.

10. Multiple Covalent Bonds • Many molecules achieve stability by sharing more than one pair of electrons between two atoms, forming a multiple covalent bond. • Carbon, nitrogen, oxygen, and sulfur most often form multiple bonds. • A double bond occurs when two pairs of electrons are shared. (O2) • A triple bond occurs when three pairs of electrons are shared. (N2)

11. The Pi Bond • A pi bond is formed when parallel orbitals overlap to share electrons. • A multiple bond consists of one sigma bond and at least one pi bond. • A pi bond always accompanies a sigma bond when forming double and triple bond.

12. Strength of Covalent Bonds • Some covalent bonds are broken more easily than others because they differ in strength. • Several factors control the strength. • How much distance separates the bonded nuclei (bond length) • The shorter the bond length, the stronger the bond. • The amount of energy required to break a bond is called bond dissociation energy. • Endothermic reactions occur when a greater amount of energy is required to break the bond(s) in the reactants than is released when the new bonds form in the products. • Exothermic reactions occur when more energy is released forming new bonds than is required to break bonds in the initial reactants.

13. Section 9.2 Naming Molecules

14. You know that many atoms covalently bond to form molecules that behave as a single unit. These units can be represented by chemical formulas and names that are used to identify them. When naming molecules, the system of rules is similar to the one you used to name ionic compounds.

15. Naming Binary Molecular Compounds The anesthetic dinitrogen oxide (N2O), commonly known as nitrous oxide, is a covalently bonded compound. Binary molecular compounds are composed of two different nonmetals and do not contain metals or ions. Although many of these compounds have common names, they also have scientific names that reveal their composition. Use these rules to name binary molecular compounds. The first element in the formula is always named first, using the entire element name. The second element in the formula is named using the root of the element name and using the suffix –ide. Prefixes are used to indicate the number of atoms of each type that are present in the compound.

16. Exceptions One exception to using these prefixes is that the first element in the formula never uses the prefix mono-. Also, to avoid awkward pronunciation, drop the final letter in the prefix when the element name begins with a vowel. For example, CO is carbon monoxide not monocarbonmonooxide.

17. Practice Problems Name the following binary covalent compounds: • CCl4 • As2O3 • CO • SO2 • NF3

18. Common Names of Some Molecular Compounds • Some compounds have “common names” • Can you think of one for dihydrogen monoxide?

19. Naming Acids Water solutions of some molecules are acidic and are named as acids. Acids are important compounds with specific properties that will be discussed later. If the compound produces hydrogen ions (H+) in solution, it is an acid. For example, HCl produces H+ in solution and is an acid. Two common types of acids exist– binary acids and oxyacids.

20. Naming Acids Naming Binary Acids Naming Oxyacids Contains an oxyanion and hydrogen Consists of a form of the root of the anion, a suffix, and the word acid -ate is replaced with –ic -ite is replaced with –ous Examples: HNO3 is nitric acid HNO2 is nitrous acid • Contains a hydrogen and one other element • Use the prefix hydro- • Root name of second element plus suffix -ic • Follow with word acid • Example: • HClis hydrochloric acid

21. Writing Formulas from Names

22. Section 9.3 Molecular Structures

23. You can now identify atoms that bond covalently and name the molecular compounds formed through covalent bonding. In order to predict the arrangement of atoms in each molecule, a model, or representation is used. Several different models can be used:

24. Structural Formulas • Predict the location of certain atoms. • Hydrogen is always terminal and connects to only one atom. • The atom with the least attraction for shared electrons is usually central atom • Find the total number of electrons available for bonding. This total is the number of valence electrons in the atoms in the molecule. • Determine the number of bonding pairs by dividing the number electrons available by 2. • Place one bonding pair between the central atom and each of the terminal atoms. • Subtract the number of pairs you used in step 4 from the number of bonding pairs you determined in step 3. The remaining electron pairs include lone pairs as well as double and triple bonds • If the central atom is not surrounded by four electron pairs, it does not have an octet. You must convert one or two of the lone pairs on the terminal atoms to a double bond or triple bond between the terminal and central atom. One of the most useful molecular models is the structural formula, which uses letter symbols and bonds to show relative positions of atoms. The structural formula can be predicted for many molecules by drawing Lewis structures, but more involved structures are needed to help you determine the shapes of molecules.

25. Example Problem 9-3 (pg 253) Ammonia is a raw material for the manufacture of many materials, including fertilizers, cleaning products, and explosives. Draw the Lewis structure for ammonia (NH3). Carbon dioxide is a product of all cellular respiration. Draw the Lewis structure for carbon dioxide. Example Problem 9-4 (pg 254)

26. Resonance Structures • Using the same sequence of atoms, it is possible to have more than one correct Lewis structure when a molecule has both a double and single bond. • Resonance is a condition that occurs when more than one valid Lewis structure can be written for a molecule or ion.

27. Draw the Lewis resonance structures for the following: • SO3 • SO2 • O3 • NO2-

28. Exceptions to the Octet Rule Three Reasons: • Odd number of valence electrons cannot form octet around each atom (NO2) • Form with fewer than 8 electrons present around an atom (BH3) • Some have central atoms that contain more than 8 electrons (PCl5)

29. Section 9.4 Molecular Shape

30. The shape of a molecule determines many of its physical and chemical properties. Molecular shape, in turn, is determined by the overlap of orbitals that share electrons. Theories have been develop to explain the overlap of bonding orbitals and are used to predict the shape of the molecule.

31. VSEPR Model • Many chemical reactions, especially those in living things, depend on the ability of two compounds to contact each other. • The shape of the molecule determines whether or not molecules can get close enough to react. • Once a Lewis structure is drawn, you can determine the molecular geometry, or shape, of the molecule. • The model used to determine the molecular shape is refereed to as the VSEPR model. • This model is based on an arrangement that minimizes the repulsion of shared and unshared pairs of electrons around the central atom. Valence Shell Electron Pair Repulsion

33. Hybridization A hybrid results from combining two of the same type of object, and it has characteristics of both. Atomic orbitals undergo hybridization during bonding. Let’s consider the bonding involved in the methane molecule (CH4). The carbon atom has 4 valence electrons and electron configuration [He]2s22p2. You might expect the two unpaired p electrons to bond with other atoms and the 2s electrons to remain in a loan pair. However, carbon atoms undergo hybridization, a process in which atomic orbitals are mixed to form new, identical hybrid orbitals. Each hybrid orbital contains one electron that it can share with another atom.

34. Determine the molecular geometry, bond angle, and type of hybridization for the following: • BF3 • NH4+ • OCl2 • BeF2 • CF4

35. Section 9.5 Electronegativity and Polarity

36. You now know that the type of bond that forms when two electrons react depends on which elements are involved. What makes one type of bond form when carbon burns and another type form when iron corrodes? The answer lies in how much attraction each type of atom has for electrons.

37. Electronegativity Difference and Bond Character • Electron affinity is a measure of the tendency of an atom to accept an electron. • Excluding noble gases, electron affinity increases as the atomic number increases within a given period and decreases with an increase in atomic number within a group. • The scale of electronegativity allows a chemist to evaluate the electron affinity of specific atoms when they are incorporated into a compound. • The character and type of chemical bond can be predicted using the electronegativity difference of the elements that are bonded.

39. For identical atoms which have an electronegativity difference of 0, the electrons are equally shared between the two atoms and the bond is considered non-polar covalent. • Unequal sharing results in a polar covalent bond. • Electronegativity of about 1.7 or greater typically means the bond is ionic and not covalent.

40. A polar molecule has a partial negative charge on one side and a partial positive charge on the other. • The molecule is a dipole because to the two partial charges. • Symmetry also plays a role in polarity. • Asymmetrical molecules tend to be polar as long as their bonds are polar. • Symmetrical molecules tend to be nonpolar even if their bonds are polar.

41. Covalent Network Solids • A number of solids are composed only of atoms interconnected by a network of covalent bonds. • These solids are often called covalent network solids. • Quartz is a network solid, as is diamond. • They are typically brittle, nonconductors of heat or electricity, and extremely hard.