1. Chapter 6 Chemical Bonds

2. 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

3. Ionic Bonds Ionic Bonds- result from the electrical attraction between large #’s of 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

4. 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

5. Polar Molecules 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

6. Polar Covalent Bonds

7. 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

8. Formation of Covalent Bonds Potential energy of each atom is used to create a covalent bond, resulting in an atom with less potential energy Forming bonds requires energy from each atom (an Endergonic reaction) As 2 atoms combine the distance between their nuclei fluctuates until equilibrium is reached

9. Diagram of Bond Formation:

10. 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

11. 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

12. Exceptions to the Octet Rule: Some elements are stable with less than 8 valence electrons –Elements 1-5 (H, He, Li, Be, B) –Elements in the s-block may also exist with less than 8 valence electrons

13. 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:

14. Electron Dot Notation

15. H Al Br Xe Fe

16. Lewis Structures: Lewis structures represent valence electrons and bonds between atoms Ex:

17. Drawing Lewis Structures

19. Structural Formulas: Structural formulas do not show valence electrons, only the bonds between atoms

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

21. Double Bonds Sharing 2 pairs of electrons between 2 atoms Ex: C 2 H 4 (ethene)

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

25. Resonance Structures(Hybrids) Ex: SO 2 (sulfur dioxide)

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

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

29. 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

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

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

32. 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)

33. 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)

34. Polyatomic Ions Ex: NO 3 - (nitrate) Ex: SO 4 -2 (sulfate)

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. 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)

49. Intermolecular Forces Hydrogen Bonding: is a special type of dipole-dipole force between water molecules

50. Intermolecular Forces Dipole Induced-Dipole: attraction between a polar and a non-polar molecule –Ex: London Dispersion Forces: attraction between 2 non-polar molecules –Ex:

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

52. END OF CHAPTER 6 NOTES!!!