Chapter Four: Forces Between Particles

1. Chapter Four: Forces Between Particles 2, 12, 14, 20, 22, 26-32, 36, 38, 48-58, 62, 66-74

2. Chemical Bonding Review • Compounds and Molecules are held together by chemical bonds • Three types of bonds • Ionic • Metals and non-metals • Covalent • Non-metal and Non-metal • Metallic • Between atoms of metals

3. Octet Rule • All atoms strive to have electronic configurations like the Noble Gases • Eight electrons in the outermost shell, highest principle quantum number (n) • Except H and He follow duet rule • Want two electrons in outermost shell • How do the atoms achieve an octet?

4. Taking or Giving and Sharing Electrons • Ionic Bonds • Atoms take or give electrons from other atoms • Covalent Bonds • Atoms share electrons between themselves • Metallic Bonds • Sea of electrons

5. n=1 1 1 2 3 4 5 6 7 8 n=2 1 2 3 4 5 6 7 8 n=3 1 2 3 4 5 6 7 8 n=4 n=5 1 2 3 4 5 6 7 8 n=6 1 2 3 4 5 6 7 8 1 2 3 4 5 6 7 8 n=7 Valence Electron Review 2

6. Lewis Dot Structures For Atoms/Ions • Symbol represents the nucleus and all electrons except for those in the valence shell • Give the Lewis Dot Structure for: Na F O2- • Species with the same number of electrons are isoelectric O2- F- Ne Na+ Mg2+ How many electrons does each species have?

7. Lewis Dot Structures • GN Lewis developed the theory of covalent bonding • Structures showing covalent bonds are called Lewis structures • Each line represents a shared pair of electrons (2 electrons) • Lone pairs of electrons are shown by a pair of dots

8. Drawing Lewis Structures • Decide on atom connectivity and placement • Hydrogen (never in the middle) is frequently bonded to oxygen • Oxygen is rarely the central atom • Oxygen will not bond to oxygen (except O2 or O3) • Carbon will be the central atom • Least electronegative atom is in the middle

9. Drawing Lewis Structures • Count the total number of valence electrons • An atom’s number of valence electrons is equal to its group number • Determine the total number of shared electrons electrons needed – valence electrons present • Connect the atoms with single bonds • A single bond is one shared pair of electrons • Use lone pairs and/or multiple bonds to give each atom an octet of electrons

10. Lewis Structure (Single Bonds) • Draw Lewis Structures for: • H2O • HCl • NH3

11. Lewis Structures (Multiple Bonds) • CO2 • N2

12. Ions • Definition: Ions are atoms or groups of atoms with an electrical charge • Cations: are positively charged, due to loss of electrons (Metals) • Anions: are negatively charged, due to gain of electrons (Non-Metals) • Number of electron’s gained or loss is due to atoms wanting Octet

13. Examples of Ions • Na • Ra • Al • Se • O • Cl • F

14. Ionic Compounds • Ionic compounds are held together by ionic bonds, or the attraction of oppositely charged ions • In the solid state, ionic compounds form crystalline lattices • Cations are attracted to all the neighboring anions, not just one • Thus, there are no discrete ionic “molecules” Ball and Stick Model

15. Transition Metal Cations • Most transition metals form more than 1 cation

16. Polyatomic Ions

17. Formulas of Ionic Compounds • The net charge on a formula unit must be zero S (+) charges = S (-) charges • Since there are no ionic “molecules” the formula of an ionic compound is the simplest ratio of cation to anion that gives an electrically neutral combination Al3+ and O2- Ca2+ and O2-

18. Writing Ionic Compound Formulas • Write the formula for each of the following pairs of ions • Na and Oxygen • Mg and Fluorine • Rb and Iodine

19. Nomenclature • Rules for naming compounds and molecules • Anions • Name the element, drop the ending leaving the root and add “ide” • Element – root + “ide” • Cl • O • N • S • I

20. Naming Ionic Compounds • Name the cation by naming the element • If the cation is a transition metal you need to distinguish the charge using Roman Numerals • Fe2+ is named Iron (II) • Pb4+ is named Lead (IV) • Name the anion • Can be an elemental anion or polyatomic • Combine them as two words

21. Naming Ionic Compounds K2O Li2CO3 K2SO4 NaHCO3 Cr2O3

22. Formulas from Names • What are the formulas of these compounds? calcium sulfide iron (III) acetate Chromium (III) sulfate

23. Naming Molecular Compounds • Name each element • Indicate how many of each element is present with a prefix multiplier • Mono =1; di =2; tr i=3; tetra =4; penta =5; hexa =6; hepta = 7; octa = 8; nona = 9; deca = 10 • Add the suffix “ide” to the last element • The prefix multiplier mono is left off of the first element in the compound

24. Naming Molecular Compounds: Examples • IBr • NI3 • N2O4

25. Formulas from Names • Sulfur dioxide • Diphosphorous pentoxide • Carbon tetrachloride

26. Molecular Compounds: Common Names • These compounds have common (non-systematic names) • Water (H2O) • Dihydrogen monoxide • Ammonia (NH3) • Nitrogen trihydride • Methane (CH4) • Carbon tetrahydride • Nitrous oxide (N2O) • Dinitrogen monoxide • Hydrazine (N2H4) • Dinitrogen tetrahydride

27. Acids • Acids are compounds that can donate an hydrogen ion (H+ ion) • Acids fall into two categories • Binary Acids HX • Oxoacids HXOn • Polyatomic anions

28. Binary Acids • Most binary acids result from dissolving the corresponding molecular compound in water • Binary acids are named as hydro (stem name of X) ic acid HCl(g) HCl(aq) HCN(g) HCN(aq)

29. Oxoacids • Oxoacids are named based on the oxoanion • “Ate” anion => ic acid • “Ite” anion => ous acid

30. Oxoacids Polyatomic Anion

31. Review of What We Know • We can write formulas • We can name compounds and molecules • We can draw Lewis Structures • But what do these molecules look like?

32. VSEPR Theory • VSEPR: Valence Shell Electron Pair Repulsion • Like charges repel and want to be as far apart as possible • Therefore a given combination of electrons will form into a specific shape

33. VSEPR • Draw the Lewis Structure • Assign the central atom (A) • Determine the number (n) of atoms bonded to (A) designate them (Xn) • Determine the number of lone pairs on (A) designate them (Em) • Put together the AXnEm notation

34. X + E = 2 X + E = 3 X + E = 4 X + E = 5 X + E = 6

35. VSEPR Examples • What is the geometry of • CO2 • BF3 • H2O • NH3

36. Electronegativity • Linus Pauling developed the electronegativity scale • Electronegativity is a measure of an atom’s affinity for electrons • Fluorine is the most electronegative element (EN=4.0) • The closer an atom is to fluorine, the more electronegative it is

38. Polar Covalent Bonds • If two atoms of identical electronegativity are bonded together, the bond is non-polar • If two atoms of different electronegativity are bonded together, the bond is polar, and the electrons spend more time around the more electronegative atom • This creates partial charges • The greater the difference in EN between two atoms, the more polar the bond • The limiting example of this is the ionic bond

40. Example • The bond in hydrogen is • The bond in hydrogen chloride is

41. Molecular Polarity • Bond dipoles are vectors • The vectoral sum of the bond dipoles gives the molecular dipole • Based on the shape of the molecules you can predict if the dipoles will cancel each other or if they will create a dipole moment • If a dipole moment exists then the molecule is said to be polar • If no dipole moment exists then the molecule is said to be non-polar

42. Molecular Polarity Examples • Is carbon dioxide polar or non-polar? • Is water polar or non-polar? • Is boron trifluoride polar or non-polar?

43. Intermolecular Forces • These are attractive forces between molecules or atoms or ions • Immensely important • These forces hold DNA molecules in a helix and and are the mechanism for DNA transcription

44. d- d- d+ d+ H H Cl Cl Dipole Dipole Attraction • This is the attraction between the opposite (partial) charges of polar molecules

45. Hydrogen Bonding • This is generally stronger than dipolar attractions • Hydrogen bonding occurs between a hydrogen atom and O, N or F. • For H-bonding to happen the H must be directly bonded to a O, N or F. O—H O—H H H This is an attraction not really a bond

46. London Forces • Also called Van der Waal’s forces, these are created by instantaneous dipoles • London forces are much weaker than either dipole-dipole or H-bonding • London forces get stronger with larger atoms/molecules

47. He He He He d- d+ d- d+ He He He He d- d+ He d- d+ He d- d+ London Forces Between Helium Atoms At a given instant, the electrons on an atommay be non-symmetrically distributed. This leads to creation of a temporary dipole. For the merest fraction of time, there is adipole-dipole attraction between the atoms. This dipole induces temporary dipoles on neighboring atoms. As the electrons re-distribute, the dipoles andand the attraction vanishes.

48. Ion Dipole Attraction • This is the attraction between an ionic charge and a polar molecule • This attraction allows ionic solids to dissolve in water • The strength of this force varies widely and depends on the magnitude of the dipole moment of the polar species and the size of the ion

49. A Sodium Ion and a Chloride Ion Hydrated by Water Molecules

50. Effects of Intermolecular Forces • More intermolecular forces mean: • Higher boiling and melting points • More viscous liquids • IM Forces also affect solubility • ‘like dissolves like’