Bonding

1. http://www.youtube.com/watch?v=_M9khs87xQ8 Bonding

2. Chemical Bonding • Bonding:Force that holds atoms together • Bonding is determined by electronic configuraton: • How many valence electrons? Mg 12e- 1s2 2s2 2p6 3s2 Cl 17e- 1s2 2s2 2p6 3s2 3p5

3. Octet Rule • Atoms • Gain, Lose (transfer) • Or Share electrons To get a full shell of 8 valence electrons (except H and He are full at 2) Forms neutral compounds: both ionic & covalent

4. Don’t confuse valence electron # with the charge on the ion Ex: Oxygen Valence electrons: _6_ How many more to fill octet __2_ So charge: _-2_

5. Ionic Bonds: transfer electrons • Bond forms between Metal and Nonmetal + ion - ion Transfer electrons: creates ions • Metals lose electrons: form + ion (cation) • Nonmetals gain electrons: form – ion (anion) • Opposite charges attract to create bond • Like magnets

6. Properties of Ionic Bonding • Ionic bonds are one of the strongest types of bonding. • Ionic bonds are very strong, so ionic compounds are usually hard, brittle, with very high melting points and boiling points. • Solids at room temp

7. Properties of Ionic Bonding • Ions are packed into repeating patterns resulting in a crystal lattice structure.

8. Properties of Ionic Bonding • No single particle of an ionic compound → represented by the simplest ratio of ions, called a formula unit.

9. Properties of Ionic Bonding • Ionic compounds dissolve well in water and split up into their ions called dissociation.

10. Properties of Ionic Bonding • Electrolytes: ions conduct an electric current when dissolved in water.

11. Properties of Ionic Bonding • Goodconductors of electricity as a liquid or when dissolved in water

12. Forming Ionic Compounds • Metals (+) are always written first • Compounds are neutral, so… • Ions must come together in the ratio to balance (+) with (-)

13. Lewis Dot Diagrams for Ionic compounds • Superscript = charge H1+ O2- • Subscript = # of atoms(do not write subscripts of 1) H2O = 2 hydrogen atoms and 1 oxygen atom H2O2 = 2 hydrogen atoms & 2 oxygen atoms

14. If lithium and fluorine bond, Li+ and F- would make LiF, because the positive 1 charge balances a negative 1 charge. Li F Li1+ F1-

15. If lithium (Li+) and oxygen (O-2) bond, more positive lithiums are needed to balance out the larger negative of oxygen. It would take 2 lithiums for every 1 oxygen. To show two lithiums are needed, a subscript of “2” is written after the lithium, Li2O Li O Li1+ O2- Li1+ Li

16. If an ionic compound is made from Aluminum (Al+3) and Sulfur (S-2), the amounts of each element needed would be: Al+3 S-2 Al+3 S-2 S-2 Totals: +6 and -6 So the resulting compound would be Al2S3

17. Al S Al S S S-2 Al+3 S-2 Al+3 S-2

18. Al S Al S S S-2 Al+3 S-2 Al+3 S-2

19. Writing Ionic formulas • Write • Symbol • Charge as superscript • “Criss-cross” – to make subscripts • Reduce if you can

20. Writing Ionic formulas • Write • Symbol • Charge as superscript • “Criss-cross” – to make subscripts • Reduce if you can Mg+2 O-2 Mg2 O2 Mg2O2 Mg O

21. Try • Calcium oxide, • Ca2O2CaO • Potassium nitride, • K3N • Magnesium phosphide, • Mg3P2

22. Polyatomic Ions • Ions with 2 or more atoms • Acts as a single unit • Use parenthesis if subscript is more than 1 • Sulfate: SO4-2 Aluminum sulfate: Al+3 SO4-2 = Al2(SO4)3

23. Transition Metals • Can form more that 1 ion • d block electrons sometime act as valence • Roman numeral identifies the charge • Manganese (IV) = +4 charge • Iron (III) = +3 charge

24. Covalent Bonds • Forms between Nonmetaland Nonmetal • Share valence electrons to get full octets

25. Properties of Covalent Compounds • Smallest unit called molecules (remember, ionic compounds form formula units) • The bonds between the atoms in a molecule are strong, but the attraction between the molecules is relatively weak. These attractive forces are known as intermolecular forces, or van der Waals forces.

26. Weak bonds mean: Low melting points and boiling points b/c they’re easy to split apart from each other. • compounds aresoft. • Examples : H20 melts at 0.0˚C CF4 melts at –150˚ and boils at –129˚C • Tend to be begases and liquids at room temperature

27. Covalent compounds do notconduct electricity • Many are polar • The only purely covalent bond is between atoms of the same element

28. Polar bonds • Polar bonds- a covalent bond in which the electrons are not sharedequally. One atom has a greater attraction for the electrons (a greater electronegativity), so the electrons spend more time around that atom, creating a slightly negative charge. The other atom then has a slightly positive charge.

29. Polar bonds • Ex. H2O: big difference in electronegativity for oxygen and hydrogen. Oxygen pulls the electrons most of the time creating a slightly (-) charge, leaving the hydrogen with a slightly (+) charge

30. Example of a covalent bond Hydrogen and Bromine. Hydrogen has 1 valence electron, Bromine has 7. Both need 1 more electron to form a stable noble gas configuration…so they form a single covalent bond • Use Lewis dot diagrams to show electrons, and a line to show covalent bond H Br H – Br Or H Br

31. Oxygen and Hydrogen H O H H – O Or H O H H

32. Carbon and Chlorine: C ClCl ClCl – C – Cl ClCl Cl

33. Carbon and Chlorine: C ClCl Cl Or Cl C Cl ClCl Cl

34. http://www.youtube.com/watch?v=QXT4OVM4vXI

35. Rules: drawing Lewis dot diagrams • Find central atom (fewest in number or one to the left on periodic table) • Add up total number of valence electrons (all atoms plus any charge) • Attach other atoms to central atom with 1 bonding pair • Fill octet for each outside atom • Must use up all electrons • Check Octet Rule – if needed borrow a pair of electrons to make double or triple bond

36. Try • NH3 • N = 5 valence electron • H = 1 • H = 1 • H = 1 8 valence electrons total

37. Which element in the middle H – N – H H Start filling, using all 8 electrons: Remember: H valence shell is full at 2 So where will last pair be?

38. H – N – H or H

39. Multiple Covalent • Multiple Covalent bonds: sharing more than 1 pair of electrons between two atoms (double or triple bonds) Oxygen gas, O2 O O O=O Or O O

40. Nitrogen gas, N2 N N NN Or N N

41. Carbon dioxide, CO2 C O O O=C=O Or O C O

42. Try • SO2 • Both S & O have 6 valence electrons • S = 6 • O = 6 • O = 6 • 18 valence electrons total

43. Which element goes in the middle O – S – O Start filling with 18 electrons Oh, don’t have enough to satisfy Octet rule so borrow 1 pair to make a double bond

44. VSEPR: Covalent compound shapes • Valence Shell Electron Pair Repulsion • Geometric shape forms to keep electrons as far apart as possible • This shape will help determine whether molecule is polar or nonpolar • If molecule is symmetrical = nonpolar, even if individual bonds are polar.

45. VSEPR: Covalent compound shapes

48. Both Ionic and Covalent • Trying to fill outer shell (valance) • Compounds are neutral overall (no charge) • Can create many different compounds • Contains more than 1 electron/atom http://www.youtube.com/watch?v=_M9khs87xQ8

49. Metallic bonding • In metals the electrons are delocalized, which means they do not belong to any one atom but move freely from atom to atom. These electrons form a sea of electrons around the metal atoms. Metallic bonding is the attraction between metal atoms and the surrounding sea of electrons.

50. These freely moving electrons can act as conduct both heat and electricity which is why solid and liquid metals are good conductors.