Chapter 8

1. Chapter 8 Covalent Bonding VSEPR Theory Molecular Shape Polar or NonPolar Properties of Molecular Substances

2. Why Do Atoms Bond? • The stability of an atom, ion or compound is related to its energy • lower energy states are more stable. • Metals and nonmetals gain stability by transferring electrons (gaining or losing) to form ions that have stable noble-gas electron configurations. • Ionic Bonding • Another way atoms can gain stability is by sharing valence electrons with other atoms, which also results in noble-gas electron configurations. • Covalent Bonding

3. The most stable arrangement of atoms exists at the point of maximum net attraction, where the atoms bond covalently and form a molecule.

4. The Covalent Bond • Atoms will share electrons in order to form a stable octet. • Covalent bond : the chemical bond that results from the sharing of valence electrons • also called a molecular bond

5. The Molecule • formed when two or more atoms bond covalently • The smallest piece in a covalent compound • Formed when the proton of one atom is attracted to the electron cloud of another atom.

6. Models Molecules

7. Single Covalent Bonds • In a single covalent bond a single pair of electrons is shared • This can be represented with a Lewis structure • A single line represents a single covalent bond • A single pair of electrons

8. Bonding pair: a pair of electrons shared by two atoms • Lone pair: an unshared pair of electrons on an atom

9. Formation of Water

10. Group 17 elements will form one covalent bond.

11. Group 16 elements will form two covalent bonds.

12. Group 15 elements will form three covalent bonds.

13. Group 14 elements will form four covalent bonds.

15. Sigma Bonds • Single covalent bonds are also called sigma bonds: • the electron pair is centered between two atoms.

17.  Multiple Covalent Bonds     When more than one pair of electrons is shared, a multiple covalent bond is formed Multiple bonds are made up of sigma bonds and pi bonds: formed when parallel orbitals share electrons.

18. Double Covalent Bond Two pairs of electrons are shared Contains one sigma and one pi bond.

19. Triple Covalent Bond Three pairs of electrons are shared Has one sigma and two pi bonds.

20. Strength of Covalent Bonds • The strength of covalent bonds is determined by the bond length: • distance between the bond nuclei • Bond length is determined by: • The size of the atoms involved—larger atoms have longer bond lengths • How many pairs of electrons are shared—the more pairs of electrons shared, the shorter the bond length is.

21. Bond Dissociation Energy • the amount of energy required to break a bond • Indicates the strength of a covalent bond • When a bond forms, energy is released; • When a bond breaks, energy must be added • Each covalent bond has a specific value for its bond dissociation energy.

22. Bond Energy and Bond Length • A direct relationship exists between bond energy and bond length • Shorter Bond • Stronger Bond • Higher Bond Dissociation Energy • Longer Bond • Weaker Bond • Lower Bond Dissociation Energy

23. Energy Changes 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. An exothermic reaction is one where more energy is released than is required to break the bonds in the initial reactants.

24. Naming Molecules • Molecular Formula • Shows what atoms and how many are in a molecule • Examples: • Nonmetal-Nonmetal Combinations

25. Naming (Binary Compounds) The first element is always named first using the entire element name The second element is named using its root and adding the suffix -ide Prefixes are used to indicate the number of atoms of each element that are present in the compound

26. Common Names Many compounds were discovered and given common names long before the present naming system was developed (water, ammonia, hydrazine, nitric oxide).

27. Binary Acids An acid that contains hydrogen and one other element Ex. HCl ion ends –ide. Name the acid with hydro-root of the anion-ic HCl (hydrogen and chloride ) becomes hydrochloric. HCl in a water solution is called hydrochloric acid.

28. Naming Acids • Acids contain hydrogen as the first element • Binary Acids: H bonded to one other element • An ion that ends –ide • Name the acid with hydro-root-ic • Example: HCl • Hydrogen ion and chloride ion • Hydrochloric acid

29. Oxyacids • An acid that contains both a hydrogen atom and an oxyanion. • Example: HNO3 • Identify the oxyanion present. • name ends with the suffix –ate, replace it with the suffix –ic. • If the name ends with suffix –ite, replace it with suffix –ous, • NO3 is the nitrate ion so the acid is nitric acid.

30. Acid Naming Summary

32. Structural Formulas A structural formula uses letter symbols and bonds to show relative positions of atoms.

33. Lewis Structures Used to predict the structural formula Show arrangement of the atoms and un-bonded electrons

34. Five steps to draw Lewis structures: • Count the total number of valence electrons in all atoms involved. • Decide how the elements are arranged in the structure and draw it out. • Hydrogen is always an end atom. • Central atom is usually written first in compound • Central atom has least attraction for the electrons • Usually closer to left on periodic table • Subtract the # of electrons used in the bonds.

35. Satisfy the octets of the terminal atoms. Place any remaining electrons around the central atom to satisfy its octet. If the central atom cannot be satisfied, make a multiple bond using a lone pair from the terminal atoms. Check your work 

36. Examples

37. Drawing Lewis structures for polyatomic ions is very similar to drawing Lewis structures for covalent compounds EXCEPT in finding the number of electrons available for bonding • Count the total number of valence electrons in all atoms involved. • If the polyatomic ion is negatively charged, ADD the charge to the number of valence electrons. • If the ion is positively charged, SUBRACT the charge from the number of valence electrons. • Follow the rest of the steps to drawing Lewis structures.

38. Examples

39. Resonance Structures When a molecule or polyatomic ion has both a double bond and a single bond, it is possible to have more than one correct Lewis structure:

40. Resonance a condition that occurs when more than one valid Lewis structure can be written for a molecule or ion. The structures are called resonance structures. A molecule that undergoes resonance behaves as if it has only one structure.

41. Exceptions to the Octet Rule Three Ways Molecules Might Violate the Octet Rule

42. Odd Number of Valence Electrons • Some molecules have an odd number of valence electrons and cannot form an octet around each atom • Example: NO2

43. Sub Octet • Some compounds form with fewer than 8 electrons present around an atom. • Boron • BF3

44. Coordinate covalent bond • when one atom donates an entire pair of electrons to be shared with atoms or ions that need two more electrons. • Boron compounds often do this

45. Expanded Octet • Some elements can have more than eight electrons in their valence shell • Because of d-level electrons • PCl5

47. How? • The d orbital starts to hold electrons. • This occurs in atoms in Period 3 or higher. • When you draw Lewis structures for these compounds, extra lone pairs are added to the central atom OR the central atom will form more than four bonds.

48. Things to Remember • Any exceptions to the Octet Rule are on the central atom

49. Molecular Shape How a molecule “looks” Determines properties The shape of a molecule determines whether or not two molecules can get close enough to react We describe shape using the VSEPR model

50. VSEPR This model is based on the fact that electrons pairs will stay as far away from each other as possible Valence Shell Electron Pair Repulsion