1. Chapter 6 Chemical Bonds

2. Terms: Molecule- a neutral group of atoms held together by covalent bonds Molecular Compound- compound made of molecules Chemical Formula- indicates the type and # of atoms in a compound Ex: CH 4 = 1 Carbon & 4 Hydrogen atoms –Molecular Formula- chemical formula of a single molecule Diatomic Compound- contains only 2 atoms Ex: O 2 or NaCl

3. Chemical Bonds- the mutual attraction between the nuclei & valence electrons of different atoms that holds atoms together –Bonding creates a more stable arrangement of atoms

4. Ionic Bonds Ionic Bonds- result from the electrical attraction between cations (+) and anions (-) *Electronegativity of Ionic Bonds: ionic bonds form when the electronegativity difference between the atoms is 1.8 or more Ex: Na + Cl  NaCl Na = 0.9Difference = 2.1 = ionic bond Cl = 3.0

5. Covalent Bonds Covalent Bonds- result from sharing 1 or more electron pairs between 2 atoms * Electronegativity of Covalent Bonds: covalent bonds form when the electronegativity difference between two atoms is 1.7 or less Ex: H 2 + O  H 2 O H = 2.1 Difference = 1.4 = covalent bond O = 3.5

6. Covalent Bonds Polar Molecules- polar molecules have an uneven distribution of charge Non-Polar Covalent Bond- electrons are shared equally between 2 atoms Ex: H 2, O 2 – both atoms have equal electronegativities so electrons are shared equally Polar Covalent Bond- the bonded atoms have an unequal distribution of charge Ex: H 2 O - Oxygen(O) has a higher electronegativity than hydrogen(H) so the electrons are held closer to O and further from H

7. Polar Covalent Bonds

8. Bond Length- distance between the nuclei of two bonded atoms Bond Energy- The energy required to break a chemical bond Single bond = long, low energy Double bond = med length, med energy Triple bond = short, high energy * Increase bond length, Decrease bond strength Bond Properties

9. Octet Rule: Chemical compounds tend to form so that each atom through the loss, gain, or sharing of electrons has an octet (8) of valence electrons Ex: F 2 (fluorine molecule) Share these electrons to form a covalent bond

10. Exceptions to the Octet Rule: Some elements are stable with less than 8 valence electrons –H and He are full with 2 electrons because they only have a 1s orbital –Boron is usually surrounded with 6 electrons BF 3 –Expanded Valence: some elements such as sulfur (S) and phosphorus (P) may exist with more than 8 valence electrons (involve d orbital)

11. Electron Dot Notation: Electron dot notation represents valence electrons of an element using dots around the elements symbol Each dot represents a valence electron 12131415161718Group:

12. Electron Dot Notation

13. H Al Br Xe Fe

14. Lewis Structures: Lewis structures represent valence electrons and bonds between atoms Ex: HHorHH FForFF

15. Structural Formulas: Structural formulas do not show valence electrons, only the bonds between atoms Single Bond: A covalent bond in which one pair of electrons is shared between two atoms

16. How to Draw a Lewis Structure (ex. CH 3 I) 1.Determine the type & number of atoms 1 Carbon, 1 Iodine, 3 Hydrogen 2.Write the electron dot notation for each type of atom CIHCIH 3.Determine the total number of valence electrons 1 Carbon x 4 e - = 4 e - 1 Iodine x 7 e - = 7 e - 3 Hydrogen x 1 e - = 3 e - 14 e -

17. 4.Arrange the atoms to form a skeleton structure a.If Carbon is present, it is the central atom. Otherwise, the least EN atom is central. (*Exception – Hydrogen is never central!) H HCIHCI H How to Draw a Lewis Structure (ex. CH 3 I)

18. 5.Connect the atoms by electron pair bonds H HCIHCI H 6.Add unshared pairs of electrons to each nonmetal atom (except Hydrogen) so that each is surrounded by 8 electrons. H HCIHCI H How to Draw a Lewis Structure (ex. CH 3 I)

19. 7.Count the electrons in the structure to be sure the number of valence electrons used = the number of available electrons (from step 3) a.Be sure the central atom and other atoms (except H) have an octet! How to Draw a Lewis Structure (ex. CH 3 I)

20. NH 3 How to Draw a Lewis Structure - Example

21. Multiple Covalent Bonds Multiple covalent bonds form when more than one pair of electrons is shared between 2 atoms

22. Double Bonds Sharing 2 pairs of electrons between 2 atoms Ex: C 2 H 4 (ethene) Triple Bonds Sharing 3 pairs of electrons between 2 atoms Ex: N 2 (di-nitrogen) C 2 H 2 (ethyne)

23. Bond Length & Energy in Multiple Bonds The more bonds, the shorter the length, the higher the energy Single Double Triple

24. Steps 1-7 are the same. If you get to step 7 and there are too many valence electrons… 8.Subtract lone pairs and move them to become multiple bonds and fill the shells. How to Draw a Lewis Structure for Multiple Bonds –

25. How to Draw a Lewis Structure for Multiple Bonds – Example: CH 2 O

26. Resonance Structures(Hybrids) Resonance structures have 2 or more acceptable Lewis Structures Ex: O 3 (ozone)

27. Ionic Bonding & Ionic Compounds Ionic Compound- a compound composed of cations (+) and anions (-) with an overall neutral charge Ex: NaCl (sodium chloride) CaF 2 (calcium fluoride)

29. Ions & Group # Group #Charge 1+1 2+2 13+3 15-3 16-2 17

30. Formula Unit A formula unit is the simplest form of an ionic compound Written with the lowest possible whole # ratio between the ions Ex: Mg 3 Br 6  MgBr 2

31. Characteristics of Ionic Bonding Ionic bonds form orderly, 3-dimensional structures Many ionic compounds exist as crystalline solids (crystal lattice)

32. Lattice Energy Lattice energy is the energy released when 1 mole of an ionic crystalline compound is formed

33. Molecular Forces vs. Atomic Forces Ionic and Covalent Bonds are very strong (forces between atoms) Molecular Bonds are weak (forces between molecules) Solubility- ionic compounds are soluble in water (dissolve in water)

34. Polyatomic Ions Polyatomic Ions are charged groups of atoms held together by covalent bonds These molecules have an overall positive(+) or negative(-) charge Ex: NH 4 + (ammonium)

35. Metallic Bonding Metallic bonding results from the attraction metal atoms and the surrounding “sea of electrons” Sea of Electrons- mobile electrons roaming freely within a metal –Electrons can pass from metal atom to metal through overlapping orbitals

36. Properties of Metals Malleability- ability of a metal to bend without breaking Ductility- ability to be drawn into a fine wire Luster- shiny

37. Metallic Bond Strength Metallic Bond Strength = heat of vaporization Amount of heat energy required to vaporize a metal (solid  gas) Metallic bonds are weaker than ionic & covalent bonds

38. Molecular Geometry The 3D shapes of molecules are determined by the atoms and electrons within the molecule VSEPR Theory- Valence Shell Electron Pair Repulsion Theory - the VSEPR theory states that atoms will space themselves as far apart from one another when bonds are formed

39. Basic Molecular Shapes Linear(a): 2 atoms around a central atom Linear(b): diatomic molecules No Electron Pairs on Central Atom CO 2 HCN

40. More Molecular Shapes Trigonal Planer- 3 atoms around a central atom, no e - on central atom Trigonal Pyrmidal- 3 atoms around a central atom, e - on central atom Ex: SO 3 Ex: NH 3

41. Even More Molecular Shapes Tetrahedral- 4 atoms around a central atom Ex: CH 4

42. Unshared Electron Pairs Unshared electron pairs affect the shape of a molecule Ex: H 2 O Electron Pair Makes Shape BENT not Linear

43. VSEPR Geometry

45. Hybridization -Hybridization is the mixing of 2 or more atomic orbitals of equal energy to produce new orbitals -Ex: Carbon C: __ __ __ __ __ * Hybrid C: __ __ __ __ __ 1s2s2p

46. Hybrid Orbitals Hybrid Orbitals- can make additional bonds due to the splitting of paired electrons into empty orbitals

47. Types of Hybrid Orbitals sp- 2 orbitals = linear geometry sp 2 - 3 orbitals = trigonal planer geometry sp 3 - 4 orbitals = tetrahedral geometry 180* 120* 108*

48. Polarity of Molecular Shapes ShapePolarity LinearNon-polar Trigonal PlanerNon-polar TetrahedralNon-polar BentPolar Trigonal PyramidalPolar

49. Intermolecular Forces Intermolecular Forces are forces holding separate molecules together –Intermolecular forces are weaker than ionic and covalent bonds Dipole-Dipole: force of attraction between 2 polar molecules Ex: HCl (hydrochloric acid)

50. Intermolecular Forces Hydrogen Bonding: is a special type of dipole-dipole force between water molecules London Dispersion Forces: attraction between 2 non-polar molecules

51. END OF CHAPTER 6 NOTES!!!