Understanding Ionic and Covalent Bonding in Chemistry

1. Bonding Between Atoms

2. Why do Atoms Form Bonds? • To get a stable octet of valence electrons. • Called a “noble gas configuration”

3. Changes in Energy • Energy is released when bonds form. (exothermic) • Results in lower energy, more stability • Know example…

4. Bonding Type #1 Ionic Bonding

5. Atoms Become Ions • Lose electrons, become a (+) ion (cation) • Ex: Mg is 2-8-2 • Loses 2 electrons to become Mg+2 2-8

6. Gain electrons, become a (–) ion (anion) • Ex: Cl is 2-8-7 • Gains 1 electron to become Cl-1 2-8-8

7. Metals lose electrons (form + ions) • Nonmetals gain them (form – ions)

8. Electron Transfer • As ions form, an exchange or transfer of electrons happens. Lithium (metal) 2-1 Fluorine (nonmetal) 2-7

9. Electronegativity • Atoms in ionic compounds have a large difference in their EN values > 1.7

10. Ex: EN Values • Metal Na = 0.9 • Nonmetal Cl = 3.2 • Nonmetal with higher EN “takes” electron(s) from the metal

11. A positively charged ion and a negatively charged ion attract each other. • This attraction forms an IONIC BOND - ion + ion

12. Ionic compounds have a uniform crystalline lattice structure.

14. Classic Example of Ionic Bond Formation • http://youtu.be/xTx_DWboEVs

15. Properties of Ionic Compounds

16. High Melting Point Attraction between the ions is very strong. Requires a lot of heat energy to separate ions and make solid crystal melt.

17. Solubility in Water • Most ionics will dissolve in water, or be “soluble”. • When dissolved in water they are “aqueous” Ex: NaCl (aq)

18. Water is a “polar molecule”. • Acts like a magnet to pull ions apart and into solution. • Ions are now “dissociated” or “hydrated” ions.

20. http://youtu.be/gN9euz9jzwc • http://youtu.be/EBfGcTAJF4o

21. Conductivity • Ionic compounds conduct when ions are “mobile” or free to move about.

22. Conductivity • Ionic compounds conduct when: • Molten (melted/liquid) (l) • Aqueous (aq) Note: They DO NOT conduct when solid as the ions are locked in place.

23. Types of Ionic Compounds • Binary: Contain 2 elements Ex: MgCl2, Al2O3, NaCl

24. Ternary: Contain 3 elements • Polyatomic ion present. • Contain both ionic and covalent bonds!! Ex: NaNO3, Ca3(PO4)2, NH4Cl Elements inside the polyatomic ion are covalently bonded (all nonmetals).

25. Ionic Compounds are Neutral • Criss-Cross charges if necessary to balance the formula

26. Roman Numeral (Stock System) • Roman numeral is used in the name of ionic compounds in which the metal can have more than one possible charge. Ex: NiBr2 Nickel II Bromide NiBr3 Nickel III Bromide

27. Bonding Type #2

28. Covalent Bonding • Share valence electrons between atoms • Electron clouds overlap

29. Happens between nonmetals Ex: H2O CH4 C6H12O6 NH3 CO2

30. Electronegativity • Difference in EN is smaller than in ionics and is usually < 1.7 • Ex: HCl H = 2.1 Cl = 3.2 Difference = 1.1

31. Molecular Formulas • All covalent compounds are called molecules. • Molecular formulas: show actual number of atoms of each element present in compound Ex: H2O 2 hydrogen atoms and 1 oxygen

32. Empirical Formulas • Empirical formulas: • Show simplest whole number ratio of atoms or ions in the compound. • All ionic compounds have empirical formulas • Ex: MgCl2 1 : 2 ion ratio

33. You can simplify some molecular formulas to make them empirical ratios Ex: C6H12O6 Simplest ratio of atoms CH2O

34. Structural Formulas • Show how the atoms are bonded together in a molecule. • Use “lines” to show covalent bonds

35. Molecule vs. Ionic Crystal CH4 = 5 atoms in molecule NaCl = 1:1 ion ratio

36. Ionic Character • Note: The greater the EN difference is between atoms the more “Ionic Character” the bond has.

37. Single, Double, Triple Bonds • Atoms can share single double or triple bonds between them. • Each bond represents a shared pair of electrons. • http://youtu.be/1wpDicW_MQQ

38. Bond Polarity • Polar Bond: when there is a difference in EN values. (unequal sharing) • Ex: H Cl EN=2.1 EN=3.2

39. NonPolar Bond: no difference in EN values. (equal sharing) Ex: O2, N2, Cl2, H2 (all the diatomics!)

42. Shapes of Molecules

43. VSPER • Valence Shell Electron Pair Repulsion • Valence electrons will orient themselves around the “central” atom to be as far apart from each other as possible. • This influences the “shape” of the molecule.

44. Polarity of a Molecule • Polar Molecules: • Have polar bonds and are not symmetrical • Nonpolar Molecules • Have nonpolar bonds OR • Have polar bonds and are symmetrical

45. Tetrahedral • Has 4 atoms bonded (no free pairs)

46. Symmetry? Depends on what atoms are attached. Can be polar (asymmetrical) or nonpolar (symmetrical)

47. Pyramidal • Three atoms bonded (one free pair)

48. Symmetry? • All pyramids are asymmetrical. • These molecules are always POLAR!

49. Bent • Two atoms attached (2 free pair) The 2 free pair make it bent and not linear. These are always asymmetrical so are always polar. H2O

50. Hey, Water is Polar!!!!! Never forget this!!!