Chemical Bonding

1. Chemical Bonding • H Chemistry I • Unit 4

2. Objectives #1-2: Introduction to Chemical Bonding • The Bonding Process • Chemical bonds form so as to lower the energy of each atom involved in the bond • Usually only involves the valence electrons in atoms • During the bonding process, electrons are shared or transferredin such a way so that each atom involved achieves the electron configuration of a noble gas!

3. Formation of H-H Bond Lowers Energy

4. Formation of the Noble Gas Configuration Through Bonding

5. Objectives #1-2: Introduction to Chemical Bonding • Types of Chemical Bonds A. Ionic Bond • Involves the transferof valence electrons from a metaltoa nonmetal; this transfer is permanent • Formula units are formed, NOT molecules! B. Pure Covalent aka NonpolarCovalent Bond • Involves the equal sharing of valence electrons between nonmetallic atoms • Molecules are formed • PolarCovalent Bond—Polar Bears!!  • Involves the unequal sharing of valence electrons between nonmetallic atoms • Polar molecules generally are formed, but there are exceptions! (Example diagrams for each of these follow on next slides)

6. Formation of Ionic Bond THINK BREAK!! Explain the differences in the ion size, based on trends. (think about U 3!!) Note also that each ion charge is a result of transferring electrons—an electrical attraction holds the bond together. The formula unit is the lowest ratio of positive & negative charges needed to form the bond.

7. Formation of Nonpolar Bond Note that only the SHARED electrons will form a new molecular orbital. Since atoms are of equal size, the shared elec. will spend equal time around each nucleus.

8. Formation of Polar Bond Note that electrons are notdistributed equally in this molecule, (known as a DIPOLE, as shown by the d-) so there will be areas that always have more of a negative charge than the central atom of sulfur. If the dipole is strong enough, it can influence the electron distribution of neighbouring molecules.

9. Objectives #1-2 Introduction to Chemical Bonding • Determining Expected Bond Type Through Electro –negativity • Definition: The ability of an atom within a compound to attract electrons • Differences in the electronegativity values of the atoms that bond together can help determine the type of bond that forms between them. • The following scale is used to determine expected bond type: 0.0 - 0.3 Nonpolar (low ionic character) 0.3 –1.7 Polar Covalent 1.7 – 4.0 Ionic (high ionic character)

10. Bonding is a Continuum!! NonPolar Polar Ionic Covalent Covalent Equal Unequal Transfer of SharingSharing Electrons 100% 0 % 50% Polar Covalent bonds are the “bridge” between nonpolar and ionic substances, and make excellent solvents!

11. Table of Electronegativities See Text P 161 for Full Table!

12. Determining Bond Type Thru Electronegativity • Examples: H and H N and H F and F Ca and O

13. Objectives #3-4: Lewis Structures and Covalent Bonding • Energy and Bonding • Once the repulsive forces between atoms are overcome, a stable bond between nonmetals can form • Bond formation decreasesthe overall energy of the atoms involved in the bond • The strength of a covalent bond can be expressed in its bond energy; the greater the bond energy the strongerthe bond will be; this same energy is also needed to breakthe bond • As the bond length between two atoms increases the bond energy decreases

14. Effects of Repulsive Forces on the Formation of a Chemical Bond

15. Bond Length vs. Bond Energy **Note that the bond length is a function of the number of electrons being shared between the atoms! Double and triple bonds have a shorter bond length & require more energy to break.

16. Objectives #3-4: Lewis Structures and Covalent Bonding • Drawing Lewis Structures • General formula : E • Examples: • Single Covalent Bonds • Single covalent bonds involve the sharing of 1 pair of electrons • Steps to drawing Lewis structures for molecules: 1. Add up the # of valence electrons

17. Objectives #3-4: Lewis Structures and Covalent Bonding 2. Place lone atom in the center. H & Halogens always around the outside as these only form single bonds. 3. Use single bonds to form connections; complete octets or duets as needed with lone pairs 4. If valence count is exceeded, reduce number of lone pairs and use multiple bonds • Lone pairs vs. Bonding pairs: Lone pairs occupy regions of molecule where bonds are not present in order to complete octets; These help determine shape!!

18. Single Bond Examples:

19. Objectives #3-4: Lewis Structures and Covalent Bonding • Representations for Covalent Substances: Molecular Formula: Symbols & Subscripts; Group like atoms together. Complete Structural Formula: Arrangement with all bonds shown Condensed Structural Formula: Shows sequence of atoms that are bonded, but not lone pairs

20. Formation of Double Bonds • Double covalent bonds involve the sharing of 2pairsof electrons. • Examples:

21. Objectives #3-4: Lewis Structures and Covalent Bonding • Formation of Triple Bonds • Triple covalent bonds involve the sharing of 3 pairs of electrons • Examples:

22. Objectives #3-4 : Lewis Structures and Covalent Bonding • Other Types of Organic Functional Groups • The functional group is the chemically active site on a carbon containing organic molecule. • Area of chemical interaction with another chemical.

24. Objectives #5-7: Formation of Ions & Ionic Bonds andProperties of Ionic vs. Covalent Substances • Applications of the Octet Rule in Ionic Bonding • Formation of Cations • Metals achieve octets by losing electrons • Formation of Anions • Nonmetals achieve octets by gaining electrons

25. Examples: Cation & Anion Formation

26. Objectives #5-7: Formation of Ions & Ionic Bonds andProperties of Ionic vs. Covalent Substances • Formation of Ionic Bond • The product of ionic bonding is the formation of a formula unit which shows the chemical formula of the compound in its lowest ratio terms. • The arrangement of ions in an ionic crystal is called the crystal lattice. The strength of the ionic bond is determined by the lattice energy.

27. Lewis Structures for Ionic Compounds • BaCl2 • NaI • Al2S3

28. Objectives #5-7: Formation of Ions & Ionic Bonds /Properties of Ionic vs. Covalent Substances • Energy changes needed to form crystal lattice: (NaCl) • Formation of cation: Na  Na+1 + e- requires addition of I.E., endothermic • Formation of anion: Cl + e- Cl-1 requires removal of E.A., exothermic • Formation of crystal lattice: NaCl requires removal of lattice energy, exothermic IE: Ionization Energy EA: Electron Affinity

29. III. Characteristics of Ionic and Covalent Substances

30. IV. Crystalline Solids • Solids usually exist as two types; amorphous solids which lack a definitecrystalline structure such as wax or glass and crystalline solids which contain a definite crystalline structure called a crystal lattice such as in sodium chloride • The four major types of crystal structures and their properties are as follows:

31. Example of Ionic Crystal Properties: Hard, brittle, high melting points, good insulators (NaCl) Structure: + and – ions arranged in a regular pattern

32. Example of Covalent Network Crystal Properties: Hard, brittle, high melting points, nonconductors or semiconductors (sand, diamond) Structure: A network of covalently bonded atoms forming giant molecules

33. Example of Metallic Crystal Structure: Metal cations surrounded by electron “sea” Properties: High malleability and high conductivity (Co, Bi)

34. Covalent Molecular Crystal Structure: Covalently bonded molecules held together with intermolecular forces Properties: Low melting points, soft, good insulators (C6H12O6)

35. Objectives #8-9 VSEPR Theory and Molecular Polarity • VSEPR = Valence Shell Electron Pair Repulsion • Steric Number is the sumof the numberof atoms attached to the central atom plus the number of lone pairs attached to the central atom • A molecule is POLAR if the valence electrons are NOT evenly distributed or if different atoms are attached to a central atom.

36. VSPER Examples!

37. Objectives #8-9 VSEPR Theory and Molecular Polarity • Relationship of number of lone pairs and resulting bond angles: • As the number of lone pairs on the central atom increases, the bond angle decreases.

38. Objective #10 Intermolecular Forces • Intermolecular vs. Intramolecular forces: Inter:attractive forces that operate betweenmolecules Intra:attractive forces that operate within a molecule; shape the electron cloud H2O vs. 2 H2O O OO H HHHHH

39. Objective #10 Intermolecular Forces • London Dispersion Forces • Involves interactions between adjacent nonpolarmolecules; occurs in all molecules as all have electrons; may result in temporary dipoles; Magnitude of the force depends on number of electrons and the mass of molecule • Example:

40. Objective #10 Intermolecular Forces • Dipole-Dipole Interaction Forces • Involves interactions between polar molecules; permanent pole areas in the molecule; Magnitude of force depends on EN difference between atoms in molecule—the greater the difference, the greater the dipole. • Example:

41. Objective #10 Intermolecular Forces Hydrogen bond • Hydrogen Bonding Forces • Involve interactions between polarmolecules and the element hydrogen; usually only involves polar molecules containing the elements F, O, and N; strongest of the forces. • Diagrams and Examples of IM forces:

42. Objective #10 Intermolecular Forces • Influence of Intermolecular Forces on Boiling and Melting Points • Consider the following molecular BPs: -39oC, -84oC, -164oC Which temperature belongs to which molecule? (Hint: Draw the molecule!!)

43. IMFs & States of Matter