1. Covalent Bonding Chapter 9 ~ most compounds, including those in living organisms, are covalently bonded

2. What You’ll Learn  I will analyze the nature of a covalent bond  I will name covalently bonded groups of atoms  I will determine the shapes of molecules  I will describe characteristics of covalent molecules  I will compare and contrast polar and nonpolar molecules

3. The Covalent Bond (9.1)  I will apply the octet rule to atoms that bond covalently  I will describe the formation of single, double, and triple bonds  I will relate the strength of covalent bonds to bond length and bond dissociation energy

4. Why do Atoms Bond?  Noble gases  Stable  Unreactive  Full outer energy level  Metals + nonmetals  Form ionic compounds  Less common  Ionic bonds  Electrons transferred  1 atom gains  1 atom loses  Nonmetal + nonmetal  Form molecules  type of compound  More common  Covalent bonds  Electrons SHARED  BOTH atoms need to GAIN

5. What is a Covalent Bond?  Covalent Bond  the chemical bond that results in the SHARING of valence electrons  Shared electrons are part of the COMPLETE outer energy level of BOTH atoms involved  Occur when elements are relatively close to each other on the periodic table  Majority = nonmetal + nonmetal  Molecule  Formed when 2 or more atoms bond covalently

6. Formation of a Covalent Bond  Diatomic molecules  Hydrogen, H 2  Nitrogen, N 2  Oxygen, O 2  Fluorine, F 2  Chlorine, Cl 2  Bromine, Br 2  Iodine, I 2  More stable than individual atoms  Covalently bonded  Nucleus of one atom attracted to electrons of the other atom (and vice versa)  Electrons of both atoms repel each other  Nuclei of both atoms repel each other  Point of maximum attraction  attractive forces = repulsive forces  Most stable arrangement  Molecule forms

7. Formation of a Covalent Bond A: The atoms are too far away from each other to have noticeable attraction or repulsion B: Each nucleus attracts the other atom’s electron cloud, but the electron clouds repel each other C: The distance is right for the attraction of one atom’s protons for the other atom’s electrons to make the bond stable D: If the atoms are forced closer together, the nuclei and electrons repel

8. Single Covalent Bonds  Lewis Structures  Use electron-dot diagrams to show how electrons are arranged in molecules  Single covalent bonds represented by  A pair of dots or  A line  Two electrons shared by two nuclei belong to each atom simultaneously  BOTH atoms atoms have noble gas configuration (full outer shell) H:H H—H

9. How Many Covalent Bonds?  Group 5A  Form 3 covalent bonds  Share three electrons  Group 4A  Form 4 covalent bonds  Share 4 electrons  Group 7A  Form one single covalent bond  Share one electron  Group 6A  Form 2 covalent bonds  Share two electrons H—H

10. Multiple Covalent Bonds  Double Bond  Two pairs of electrons are shared between 2 atoms  Triple Bond  Three pairs of electrons are shared between 2 atoms  Molecules that share more than one pair of electrons between two atoms  Form multiple covalent bonds  Examples:  Double bonds  Triple bonds

11. Your Turn  Draw the Lewis structure for each of these molecules  PH 3  H 2 S  HCl  CCl 4  SiH 4

12. Strength of Covalent Bonds  # shared electrons increases = bond length decreases  Shorter bond length = stronger bond  Single bonds < double bonds < triple bonds  Single = weakest bonds  Triple = strongest bonds  Covalent bonds held together by balance of attraction & repulsion  Balance upset = bond broken  Bond Length  Distance between 2 bonded nuclei at the position of maximum attraction (repulsion = attraction)  Determined by  Size of atoms  # of electron pairs shared

13. Bond Dissociation Energy  Sum of the bond dissociation energy values for ALL bonds in a compound = amount of chemical potential energy available in a molecule of that compound  Related to bond length  Short bond length = greater bond energy needed to separate the atoms  The amount of energy required to break a specific bond  Always a positive value  Energy released = bond forms  Energy added = bond breaks  Hence…positive value!

14. Chemical Reactions  Exothermic  LESS energy added to break bonds (reactants) than to make the new bonds (products)  Reactants + energy added = products – ENERGY released  Endothermic  GREATER energy added to break bonds (reactants) than to make the new bonds (products)  Reactants + ENERGY added = products – energy released

15. Homework  Page 247  9.1 assessment  6-12  Skip 9

16. Naming Molecules (9.2)  I will identify the names of binary molecular compounds from their formulas  I will name acidic solutions

17. Naming Binary Molecular Compounds  Binary Molecular compounds  Composed of 2 different nonmetals  Do NOT contain metals or ions ① The first element in the formula is ALWAYS named first, using the entire element name. ② The second element in the formula is named using the root of the element and adding the suffix –ide. ③ Prefixes are used to indicate the number of atoms of each type that are present in the compound.

18. Naming Binary Molecular Compounds  The most common prefixes Prefixes in Covalent Compounds Number of Atoms PrefixNumber of Atoms Prefix 1Mono-6Hexa- 2Di-7Hepta- 3Tri-8Octa- 4Tetra-9Nona- 5Penta-10Deca-

19. Naming Binary Molecular Compounds  EXCEPTIONS ① The FIRST element in a formula NEVER uses the prefix mono- ② Drop the FINAL letter in the prefix when the element names begins with a vowel Example CO Carbon monoxide NOT carbon monooxide

20. Try it  Name the following binary covalent compounds  CCl 4  As 2 O 3  CO  SO 2  NF 3 carbon tetrachloride diarsenic trioxide carbon monoxide sulfur dioxide nitrogen trifluoride

21. Common Names Formulas and Names of Some Covalent Compounds FormulaCommon NameMolecular Compound Name H20H20Waterdihydrogen oxide NH 3 Ammonianitrogen tetrahydride N2H4N2H4 Hydrazinedinitrogen tetrahydride H20H20Nitrous oxide (laughing gas) dinitrogen monoxide NONitric oxidenitrogen monoxide

22. Naming Acids  Water solutions of some molecules are acidic and are named as acids  Two types  Binary acids  oxyacids

23. Naming Binary Acids  Binary acid  Contains hydrogen and 1 other element ① Use prefix hydro- to name the hydrogen part of the compound ② The root of the second element plus the suffix –ic ③ The word acid  Examples  HBrhydrobromic acid  HCNhydrocyanic acid

24. Naming Oxyacids  Oxyacid  An acid that contains an oxyanion  Hydrogen + oxyanion ① Identify oxyanion present ( from your polyatomic ion purple sheet) ② If the oxyanion ends in –ate, replace with the suffix –ic ③ If the oxyanion ends in –ite, replace with –ous ④ The word acid

25. Naming Oxyacids  The hydrogen of an oxyacid is NOT part of the name  Examples  HNO 3 nitric acid  HNO 2 nitrous acid

26. Try it  Name the following acids. Assume each compound is dissolved in water.  HI  HClO 3  HClO 2  H 2 SO 4  H 2 s Hydroiodic acid chloric acid chlorous acid sulfuric acid hydrosulfuric acid

27. Writing Formulas From Names  Subscripts are determined from the prefixes used in the name  Why?  The name indicates the exact number of each atom present in the molecule  The formula from an acid can be derived from the name as well

28. Homework  P 251  9.2 Assessment  23-29