2. Covalent Bonding Chapter 9

3. 1. Why do Atoms Bond? Lower energy states make an atom more stable. Gaining or losing electrons makes atoms more stable by forming ions with noble-gas electron configurations. Atoms also gain stability when they share electrons and form covalent bonds. Sharing valence electrons with other atoms also results in noble-gas electron configurations. Section - 9.1 Forming Compounds

4. 2. What is a Covalent Bond? Atoms in non-ionic compounds share electrons. The chemical bond that results from sharing electrons is a covalent bond.covalent bond A molecule is formed when two or more atoms bond.molecule

5. What is a Covalent Bond? (cont.) Diatomic molecules (H 2, F 2 for example) exist because two-atom molecules are more stable than single atoms. For example: As two fluorine atoms approach each other two forces become important. A repulsive force occurs between the like-charged protons of the two atoms. An attractive force also occurs between the protons of one fluorine atom and the electrons of the other atom. As the fluorine atoms move closer, the attraction of both nuclei for the other atom’s electrons until the maximum attraction is achieved. The most stable arrangement of atoms exists at the point of maximum net attraction, where the atoms bond covalently and form a molecule.

6. 3. Single Covalent Bonds When only one pair of electrons is shared, the result is a single covalent bond. The figure shows two hydrogen atoms forming a hydrogen molecule with a single covalent bond, resulting in an electron configuration like helium. The shared electron pair, often referred to as the bonding pair, is represented by either a pair of dots or a line. In a Lewis structure dots or a line are used to symbolize a single covalent bond.Lewis structure

7. Single Covalent Bonds (cont.) The halogens—the group 17/7A elements— have 7 valence electrons. To attain an octet, one more electron is necessary. Therefore, they will form single covalent bonds with atoms of other identical non-metals. Atoms in group 16 can share two electrons and form two covalent bonds. Water is formed from one oxygen with two hydrogen atoms covalently bonded to it. Atoms in group 15 form three single covalent bonds, such as in ammonia. Atoms of group 14 elements form four single covalent bonds, such as in methane.

8. Single Covalent Bonds (cont.) Single covalent bonds are also called Sigma bonds. Sigma bonds Sigma bonds occur when the pair of shared electrons is in an area centered between the two atoms. When two atoms share electrons, the valence atomic orbital of one atom overlaps or merges with the valence atomic orbital of the other atom. A sigma bond results if the atomic orbitals overlap end to end, concentrating the electrons in a bonding orbital between the atoms

9. 4. Multiple Covalent Bonds Double bonds form when two pairs of electrons are shared between two atoms. Triple bonds form when three pairs of electrons are shared between two atoms.

10. Multiple Covalent Bonds (cont.) The pi bond is formed when parallel orbitals overlap and share electrons.pi bond The shared electron pair of a pi bond occupies the space above and below the line that represents where the two atoms are joined together. A multiple covalent bond consists of one sigma bond and at least one pi bond.

11. 5. Strength of Covalent Bonds The strength depends on the distance between the two nuclei, or bond length. Determined by the size of the atoms and how many electrons pairs are shared. As the number of shared electron pairs increases, bond length decreases. As length increases, strength decreases, thus single bonds are weaker than triple bonds. The amount of energy required to break a bond is called the bond dissociation energy. The shorter the bond length, the greater the energy required to break it.

12. Strength of Covalent Bonds (cont.) Bond dissociation energy indicates the strength of a chemical bond. A direct relationship exist between bond energy and bond length. An endothermic reaction is one where a greater amount of energy is required to break a bond in reactants than is released when the new bonds form in the products.endothermic reaction An exothermic reaction is one where more energy is released than is required to break the bonds in the initial reactants.exothermic reaction

13. 6. Naming Binary Molecular Compounds 1.The first element is always named first using the entire element name. 2.The second element is named using its root and adding the suffix –ide. 3.Prefixes are used to indicate the number of atoms of each element in a compound. One exception to using these prefixes is that the first element in the formula never uses the prefix mono-. Many compounds were discovered and given common names long before the present naming system was developed (water, ammonia, hydrazine, nitrous oxide). Section 9.2 - Naming Molecules

14. 7. Naming Acids If the compound produces hydrogen ions in solution, it is an acid. Two common acids exist – binary acids and oxyacids. Naming Binary Acids Binary acids contain hydrogen and one other element. The first word has the prefix hydro- to name the hydrogen part of the compound. The rest of the name consists of a form of the root of the second element plus the suffix –ic followed by the word acid (hydrochloric acid is HCl in water).

15. Naming Acids (cont.) Naming Oxyacids An oxyacid is an acid that contains both a hydrogen atom and an oxyanion.oxyacid An oxyanion is a polyatomic ion that contains oxygen. Identify the oxyanion present. The name of an oxyacid consists of a form of the root of the anion, a suffix, and the word acid. If the anion suffix is –ate, it is replaced with the suffix –ic. When the anion ends in –ite, the suffix is replaced with the suffix –ous. These hydrogen-containing compounds are named as acids only when they are in water solution.

16. 8. Writing Formulas from Names The name of a molecular compound reveals its composition and is important in communicating the nature of the compound. Subscripts are determined from the prefixes used in the name because the name indicates the exact number of each atom present in the molecule The formula for an acid can be derived from the name as well

17. 9. Names and Formulas for Bases A base is an ionic compound that produces hydroxide ions when dissolved in water. Bases are named in the same way as other ionic compounds—the name of the cation is followed by the name of the anion. For example, aluminum hydroxide consists of the aluminum cation (Al 3+ ) and the hydroxide anion (OH – ). The formula for aluminum hydroxide is Al(OH) 3.

18. 10. Structural Formulas A structural formula uses letter symbols and bonds to show relative positions of atoms.structural formula The structural formula can be predicted for many molecules by drawing the Lewis structure. Section 9.3 - Molecular Structures

19. Structural Formulas (cont.) Drawing Lewis Structures 1.Predict the location of certain atoms. a.Hydrogen is always a terminal atom because it can share only one pair of electrons with one other atom b.The atom with the least attraction for shared electrons in the molecule is the central atom (lowest electronegativity) 2.Find the number of electrons available for bonding. 3.Determine the number of bonding pairs. 4.Place one bonding pair between the central atom and each terminal element. 5.Determine the number of valence electrons remaining. 6.Determine whether the central atom satisfies the octet rule. Atoms within a polyatomic ion are covalently bonded.

20. 11. Resonance Structures Resonance is a condition that occurs when more than one valid Lewis structure can be written for a molecule or ion.Resonance Resonance structures differ only in the position of the electron pairs, never the atom positions. The molecule or ion behaves as if it has only one structure. The bond lengths are identical to each other and intermediate between single and double covalent bonds. This figure shows three correct ways to draw the structure for (NO 3 ) -1.

21. 12. Exceptions to the Octet Rule Some molecules do not obey the octet rule. 1.A small group of molecules might have an odd number of valence electrons. Ex. NO 2 has five valence electrons from nitrogen and 12 from oxygen and cannot form an exact number of electron pairs.

22. Exceptions to the Octet Rule (cont.) 2.A few compounds form stable configurations with less than 8 electrons around the atom—a suboctet. Ex. Boron forms three covalent bonds with other nonmetallic atoms. Such compounds are reactive and can share an entire pair of electrons donated by another atom A coordinate covalent bond forms when one atom donates both of the electrons to be shared with an atom or ion that needs two electrons.coordinate covalent bond

23. Exceptions to the Octet Rule (cont.) 3.A third group of compounds has central atoms with more than eight valence electrons, called an expanded octet. Elements in period 3 or higher have an empty or partially filled d-orbital and can form more than four covalent bonds by expanding into the empty or partially filled d-orbital.

24. 13. VSEPR Model 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. Once Lewis structure is drawn, you can determine the molecular geometry (shape). Molecular geometry can be determined with the Valence Shell Electron Pair Repulsion model, or VSEPR model.VSEPR model This model minimizes the repulsion of shared and unshared atoms around the central atom. Section 9.4 - Molecular Shape

25. VSEPR Model (cont.) Electron pairs repel each other and cause molecules to be in fixed positions relative to each other. The angle formed by any two terminal atoms and the central atom is called bond angle Lone pairs also determine the shape of a molecule. Because lone pairs are not shared between two nuclei, they occupy a slightly larger orbital than shared electrons Electron pairs are located in a molecule as far apart as they can be, but are pushed together slightly by lone pairs

26. 14. Hybridization Hybridization is a process in which atomic orbitals mix and form new, identical hybrid orbitals.Hybridization Carbon often undergoes hybridization, which forms an sp 3 orbital formed from one s orbital and three p orbitals. The number of atomic orbitals mixed to form the hybrid orbital equals the total number of pairs of electrons. The number of hybrid orbitals formed equals the number of atomic orbitals mixed Lone pairs also occupy hybrid orbitals.

27. 15. Electronegativity Difference and Bond Character Electron affinity measures the tendency of an atom to accept an electron. Affinity increases with increasing atomic number within a period, and decreases with increasing atomic number within a group. Noble gases are not listed because they generally do not form compounds. This table lists the character and type of chemical bond that forms with differences in electronegativity. Section 9.5 - Electronegativity and Polarity

28. Electronegativity Difference and Bond Character A covalent bond formed between atoms of different elements does not have equal sharing of the electron pair due to a difference in electronegativity Unequal sharing of electrons results in a polar covalent bond.polar covalent bond Bonding is often not clearly ionic or covalent. This graph summarizes the range of chemical bonds between two atoms.

29. 16. Polar Covalent Bonds Polar covalent bonds form when atoms pull on electrons in a molecule unequally. Electrons spend more time around one atom than another resulting in partial charges at the ends of the bond called a dipole (two poles). The more electronegative atom is located at the partially negative end, while the less electronegative atom is located at the partially positive end.

30. Polar Covalent Bonds (cont.) Covalently bonded molecules are either polar or non-polar, depending on the nature and location of the covalent bonds they contain. Non-polar molecules are not attracted by an electric field. Polar molecules align with an electric field. Compare water and CCl 4. Both bonds are polar, but only water is a polar molecule because of the shape of the molecule. The electric charge on a CCl 4 molecule measured at any distance from the center of the molecule is identical to the charge measured at the same distance on the opposite side.

31. Polar Covalent Bonds (cont.) Solubility is the physical property of a substance’s ability to dissolve in another substance. Bond type and the shape of the molecules present determine solubility Polar molecules and ionic substances are usually soluble in polar substances. Non-polar molecules dissolve only in non- polar substances. “Like dissolves like”

32. 17. Properties of Covalent Compounds Differences in properties are a result of differences in attractive forces. Covalent bonds between atoms (within a molecule) are strong, but attraction forces between molecules are weak. The weak attraction forces between molecules are known as intermolecular forces, or van der Waals forces. The forces vary in strength but are weaker than the bonds that join atoms in a molecule or ions in an ionic compound.

33. Properties of Covalent Compounds (cont.) There are different types of intermolecular forces. Nonpolar molecules exhibit a weak dispersion force, or induced dipole. The greater the number of electrons, the stronger the dispersion force. The force between two oppositely charged ends of two polar molecules stronger and is called a dipole-dipole force. The more polar the molecule the stronger the dipole- dipole force. A hydrogen bond is an especially strong dipole- dipole force between a hydrogen end of one dipole and a fluorine, oxygen, or nitrogen atom on another dipole.

34. Properties of Covalent Compounds (cont.) Many physical properties are due to intermolecular forces. Weak forces result in the relatively low melting and boiling points of molecular substances. Many covalent molecules are relatively soft solids. Molecules can align in a crystal lattice, similar to ionic solids but with less attraction between particles. Solids composed of only atoms interconnected by a network of covalent bonds are called covalent network solids. Quartz and diamonds are two common examples of network solids.

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