1. Chapter 11 Chemical Bonds: The Formation of Compounds from Atoms Objectives: Describe the trends in the periodic table Describe the trends in the periodic table Know how to draw Lewis Structures of atoms Know how to draw Lewis Structures of atoms Understand and predict the formation of ionic bonds Understand and predict the formation of ionic bonds Understand and predict covalent bonds Understand and predict covalent bonds Describe electronegativity Describe electronegativity Know how to draw complex lewis structures of compounds Know how to draw complex lewis structures of compounds Understand the formation of compounds containing polyatomic ions Understand the formation of compounds containing polyatomic ions Describe molecular shape, including the VSEPR model Describe molecular shape, including the VSEPR model

2. Periodic Trends in Atomic Properties Periodic table designed to show trends Periodic table designed to show trends Use trends to predict properties and reactions between elements Use trends to predict properties and reactions between elements Trends include: Trends include: Metals, nonmetals, metalloids Metals, nonmetals, metalloids Atomic radius Atomic radius Ionization energy Ionization energy Electronegativity Electronegativity

3. Metals, Nonmetals and Metalloids Metals: Metals: Lustrous, malleable, good conductors of heat and electricity Lustrous, malleable, good conductors of heat and electricity Left-hand side of table Left-hand side of table Most elements are metals Most elements are metals Tend to lose electrons and form positive ions Tend to lose electrons and form positive ions

4. Metals, Nonmetals and Metalloids Nonmetals: Nonmetals: Nonlustrous, brittle, poor conductors Nonlustrous, brittle, poor conductors Right side of table Right side of table Tend to gain electrons and form negative ions Tend to gain electrons and form negative ions (Hydrogen displays nonmetallic properties under normal conditions but is UNIQUE element) (Hydrogen displays nonmetallic properties under normal conditions but is UNIQUE element)

5. Metals, Nonmetals and Metalloids Metalloids Metalloids Found along border between metals and nonmetals Found along border between metals and nonmetals Show characteristics of both Show characteristics of both Metal + Nonmetal Metal + Nonmetal Usually electrons are transferred from metal to nonmetal… Usually electrons are transferred from metal to nonmetal…

6. Atomic Radius Increases down each group Increases down each group Each step down = additional energy level Each step down = additional energy level More energy levels = greater distance from nucleus = large average size More energy levels = greater distance from nucleus = large average size Decreases from left to right across a period Decreases from left to right across a period Electrons added to the same energy level (also means protons added…) Electrons added to the same energy level (also means protons added…) Increase in positive charge = stronger pull on electrons = gradual decrease in atomic radius Increase in positive charge = stronger pull on electrons = gradual decrease in atomic radius

7. Atomic Radius

8. Ionization Energy The energy required to remove an electron from the atom The energy required to remove an electron from the atom More energy required to remove 2 nd, 3 rd, 4 th, 5 th, etc. electron More energy required to remove 2 nd, 3 rd, 4 th, 5 th, etc. electron Noble gas structure is stable so takes large amount of energy to remove an electron Noble gas structure is stable so takes large amount of energy to remove an electron

9. Ionization Energy Ionization energy in Group A elements decreases as you move down a group Ionization energy in Group A elements decreases as you move down a group Ionization energy increases from left to right across a period Ionization energy increases from left to right across a period Metals – some give up electrons more easily than others Metals – some give up electrons more easily than others Nonmetals – tend to gain electrons (rather than give them up) Nonmetals – tend to gain electrons (rather than give them up)

10. Ionization Energy

11. Lewis Structures Diagram that shows valence electrons Diagram that shows valence electrons American chemist Gilbert N. Lewis American chemist Gilbert N. Lewis Dots = number of s and p electrons in outermost energy level Dots = number of s and p electrons in outermost energy level Paired dots = paired electrons Paired dots = paired electrons Simple way of showing electrons Simple way of showing electrons Most reactions involve only outermost electrons Most reactions involve only outermost electrons

12. Lewis Structures When drawing: When drawing: Use electron configuration Use electron configuration Move in clockwise direction… Move in clockwise direction… “12” = s orbital “12” = s orbital “3, 6, 9” = p orbitals – fill each with ONE electron before filling with pairs… “3, 6, 9” = p orbitals – fill each with ONE electron before filling with pairs… Just like orbital filling diagram… Just like orbital filling diagram… Examples: draw Lewis Structures of B, N, F, Ne Examples: draw Lewis Structures of B, N, F, Ne

13. Lewis Structures N B Ne F

14. The Ionic Bond Ionic bond: the attraction between oppositely charged ions Ionic bond: the attraction between oppositely charged ions Transfer of electrons from one atom to another Transfer of electrons from one atom to another Ions form with + or – charges Ions form with + or – charges Attraction between electrostatic charges is a strong force which holds atomstogether Attraction between electrostatic charges is a strong force which holds atomstogether

15. The Ionic Bond

16. NOT A MOLECULE NOT A MOLECULE Bond not just between (for example) one sodium and one chloride Bond not just between (for example) one sodium and one chloride One sodium ion attracts 6 chlorine ions… One sodium ion attracts 6 chlorine ions…

17. The Ionic Bond Typically metal + nonmetal Typically metal + nonmetal Metals usually lose electrons Metals usually lose electrons Nonmetals usually gain electrons Nonmetals usually gain electrons

18. Predicting Formulas of Ionic Compounds In almost all stable chemical compounds of representative elements, each atom attains a noble gas electron configuration. This concept forms the basis for our understanding of chemical bonding. In almost all stable chemical compounds of representative elements, each atom attains a noble gas electron configuration. This concept forms the basis for our understanding of chemical bonding.

19. Predicting Formulas of Ionic Compounds How many electrons must be gained or lost to achieve noble gas configuration? How many electrons must be gained or lost to achieve noble gas configuration? Ba must lose 2 electrons [Xe]6s 2 Ba must lose 2 electrons [Xe]6s 2 Forms the ion Ba 2+ Forms the ion Ba 2+ S must gain 2 electrons [Ne]3s 2 3p 4 S must gain 2 electrons [Ne]3s 2 3p 4 Forms the ion S 2- Forms the ion S 2- So…must be ratio of 1 to 1 when barium and sulfur combine So…must be ratio of 1 to 1 when barium and sulfur combine BaS BaS

20. Predicting Formulas of Ionic Compounds Elements in a family usually form compounds with the same atomic ratios Elements in a family usually form compounds with the same atomic ratios Because they have the same number of valence electrons Because they have the same number of valence electrons Must gain or lose the same number of electrons Must gain or lose the same number of electrons See table 11.4 pg 233 See table 11.4 pg 233

21. Predicting Formulas of Ionic Compounds The formula for sodium oxide is Na 2 O. Predict the formula for The formula for sodium oxide is Na 2 O. Predict the formula for Sodium sulfide Sodium sulfide Sodium [Ne]3s 1 must lose one electron Sodium [Ne]3s 1 must lose one electron Sulfur [Ne]3s 2 3p 4 must gain two electrons Sulfur [Ne]3s 2 3p 4 must gain two electrons So…formula must be Na 2 S So…formula must be Na 2 S Sulfur is in same family as oxygen…so same ratio Sulfur is in same family as oxygen…so same ratio

22. Predicting Formulas of Ionic Compounds Rubidium Oxide Rubidium Oxide Rubidium [Kr]5s 1 must lose one electron Rubidium [Kr]5s 1 must lose one electron Oxygen [He]2s 2 2p 4 must gain two electrons Oxygen [He]2s 2 2p 4 must gain two electrons So…formula must be Rb 2 O So…formula must be Rb 2 O This makes sense b/c rubidium is in same family as sodium This makes sense b/c rubidium is in same family as sodium

23. The Covalent Bond A pair of electrons shared between two atoms A pair of electrons shared between two atoms Most common type of bond Most common type of bond Stronger than ionic bond Stronger than ionic bond Electron orbital expands to include both nuclei Electron orbital expands to include both nuclei most often found between two nuclei most often found between two nuclei Negative charges allow positive nuclei to be drawn close to each other Negative charges allow positive nuclei to be drawn close to each other

24. The Covalent Bond

25. Atoms may share more than one pair of electrons Atoms may share more than one pair of electrons Double bond – two pairs being shared Double bond – two pairs being shared Triple bond – three pairs being shared Triple bond – three pairs being shared Multiple bonds are stronger than single bonds Multiple bonds are stronger than single bonds Harder to break Harder to break Covalent bonding between identical atoms means electrons are shared equally Covalent bonding between identical atoms means electrons are shared equally Covalent bonding between different atoms leads to unequal sharing (polar covalent bond) Covalent bonding between different atoms leads to unequal sharing (polar covalent bond)

26. Electronegativity The attractive force that an atom of an element has for shared electrons The attractive force that an atom of an element has for shared electrons Atoms have different electronegativities Atoms have different electronegativities Electrons will spend more time near atom with stronger (larger) electronegativity Electrons will spend more time near atom with stronger (larger) electronegativity So…one atom assumes a partial positive charge So…one atom assumes a partial positive charge The other assumes a partial negative charge The other assumes a partial negative charge

27. Electronegativity Electronegativity trends and periodic table Electronegativity trends and periodic table See table 11.5 page 237 See table 11.5 page 237 Generally increases from left to right Generally increases from left to right Decreases down a group Decreases down a group Highest is fluorine (4.0) Highest is fluorine (4.0) Lowest is francium (0.7) Lowest is francium (0.7)

28. Electronegativity

29. Electronegativity Polarity is determined by difference in electronegativity Polarity is determined by difference in electronegativity Nonpolar covalent Nonpolar covalent Electronegativities are equal Electronegativities are equal Electrons shared equally Electrons shared equally Polar covalent Polar covalent One electronegativity is larger…. One electronegativity is larger…. Ionic compound Ionic compound Electronegativity is so great that no sharing occurs Electronegativity is so great that no sharing occurs Electrons are lost or gained instead… Electrons are lost or gained instead…

30. Electronegativity

31. Electronegativity If the electronegativity difference is greater than 1.7-1.9 then the bond will be more ionic than covalent If the electronegativity difference is greater than 1.7-1.9 then the bond will be more ionic than covalent Above 2.0 = ionic bond Above 2.0 = ionic bond Below 1.5 = nonpolar covalent Below 1.5 = nonpolar covalent See Continuum on page 239 See Continuum on page 239

32. Electronegativity Polar bonds form between two atoms Polar bonds form between two atoms Molecules can also be polar or nonpolar Molecules can also be polar or nonpolar Dipole Dipole Electrically asymmetrical Electrically asymmetrical Oppositely charged at two points Oppositely charged at two points Polar Polar Usually molecules with only one polar bond Usually molecules with only one polar bond Part of molecule has + charge, part has – charge Part of molecule has + charge, part has – charge Nonpolar Nonpolar Molecules with multiple bonds Molecules with multiple bonds Dipoles cancel each other by acting in opposite directions Dipoles cancel each other by acting in opposite directions Overall has no charge Overall has no charge

33. Lewis Structures of Compounds Convenient way of showing ionic or covalent bonds Convenient way of showing ionic or covalent bonds Usually the single atom in a formula is the central atom Usually the single atom in a formula is the central atom

34. The Ionic Bond LEWIS STRUCTURES of ionic bonds LEWIS STRUCTURES of ionic bonds

35. The Covalent Bond LEWIS STRUCTURES of covalent bonds LEWIS STRUCTURES of covalent bonds Use dashes instead of dots… Use dashes instead of dots…

36. The Covalent Bond

37. Lewis Structures of Compounds 1) Obtain the total number of valence electrons 1) Add the valance electrons of all atoms 2) Ionic – add one electron for each negative charge and subtract one electron for each positive charge

38. Lewis Structures of Compounds 2) Write the skeletal arrangement of the atoms and connect with a single covalent vond 3) Subtract two electrons for each single bond 1) This gives you the net number of electrons available for completing the structure

39. Lewis Structures of Compounds 4) Distribute pairs of electrons around each atom to give each atom a noble gas structure 5) If there are not enough electrons then try to form double and triple bonds

40. Lewis Structures of Compounds Write the Lewis Structure for methane CH 4 Write the Lewis Structure for methane CH 4 1) Total number of valence electrons is eight 2) Draw skeletal structure 1) Dashes equal two electrons being shared 3) Subtract the eight electrons shown as dashes 4) Check that all atoms have a noble gas structure

41. Lewis Structures of Compounds Methane, CH 4 Methane, CH 4

42. Lewis Structures of Compounds Carbon Dioxide, CO 2 Carbon Dioxide, CO 2 Total valence electrons = 16 Total valence electrons = 16 Not Enough! Must try double bonds…

43. Complex Lewis Structures Some molecules and polyatomic ions have strange behaviors… Some molecules and polyatomic ions have strange behaviors… No single Lewis structure is consistent No single Lewis structure is consistent If multiple structures are possible the molecule shows resonance If multiple structures are possible the molecule shows resonance Resonance structures – show all possibilities Resonance structures – show all possibilities

44. Complex Lewis Structures Carbonate ion, CO 3 2- Carbonate ion, CO 3 2- Carbon only has 6 electrons – try double bonds – more than one location…..form resonant structures…

45. Compounds Containing Polyatomic Ions Polyatomic ion: stable group of atoms that has a positive or negative charge Polyatomic ion: stable group of atoms that has a positive or negative charge Behaves as a single unit in many chemical reactions Behaves as a single unit in many chemical reactions Sodium carbonate (Na 2 CO 3 ) Sodium carbonate (Na 2 CO 3 ) Carbonate ion (co 3 ) has covalent bonds Carbonate ion (co 3 ) has covalent bonds Sodium atoms are ionically bonded to carbonate ion Sodium atoms are ionically bonded to carbonate ion

46. Compounds Containing Polyatomic Ions Easier to dissociate ionic bond than break covalent bond Easier to dissociate ionic bond than break covalent bond More in chapters 6 and 7 More in chapters 6 and 7

47. Molecular Shape Three-dimensional shape of molecule important Three-dimensional shape of molecule important Explains molecular interactions Explains molecular interactions Helpful to know how to predict the geometric shape of molecules… Helpful to know how to predict the geometric shape of molecules… Linear? Linear? V-shaped? V-shaped? Trigonal planar? Trigonal planar? Tetrahedral? Tetrahedral?

48. The VSEPR Model Valence Shell Electron Pair Repulsion Model Valence Shell Electron Pair Repulsion Model Make predictions about shape from Lewis structures Make predictions about shape from Lewis structures Electron pairs will repel each other electrically Electron pairs will repel each other electrically Try to minimize this repulsion Try to minimize this repulsion Arranged as far apart as possible around a central atom Arranged as far apart as possible around a central atom

49. The VSEPR Model Linear Structure Linear Structure Two pairs of electrons surrounding the central atom Two pairs of electrons surrounding the central atom 180 o apart 180 o apart

50. The VSEPR Model Trigonal Planar Trigonal Planar Three pairs of electrons around the central atom Three pairs of electrons around the central atom 120 o apart 120 o apart

51. The VSEPR Model Tetrahedral structure Tetrahedral structure Four pairs of electrons on central atom Four pairs of electrons on central atom 109.5 0 apart 109.5 0 apart When drawing: When drawing: Wedged line to show atom protruding from page; dashed line to show atom receding from page Wedged line to show atom protruding from page; dashed line to show atom receding from page

52. The VSEPR Model Pyramidal shape Pyramidal shape Four pairs of electrons on central atom BUT only three shared… Four pairs of electrons on central atom BUT only three shared… Electrons are tetrahedral but actual shape is more of a pyramid Electrons are tetrahedral but actual shape is more of a pyramid

53. The VSEPR Model Electron pairs determine shape BUT name for shape is determined by position of atoms Electron pairs determine shape BUT name for shape is determined by position of atoms

54. The VSEPR Model V-shaped or bent V-shaped or bent Four electron pairs but only two shared Four electron pairs but only two shared Electron arrangement is tetrahedral Electron arrangement is tetrahedral But, moledule is “bent” But, moledule is “bent” Water Water Helps explain some properties Helps explain some properties

55. The VSEPR Model Predict the shape for CF 4, NF 3, and BeI 2. Predict the shape for CF 4, NF 3, and BeI 2. Draw the Lewis Structure Draw the Lewis Structure Count the electron pairs and determine the arrangement that will minimize repulsions Count the electron pairs and determine the arrangement that will minimize repulsions Determine the positions of the atoms and name the structure Determine the positions of the atoms and name the structure

56. The VSEPR Model CF 4 CF 4 Lewis Structure (on board) Lewis Structure (on board) Number of electron pairs: 4 Number of electron pairs: 4 Electron pair arrangement: tetrahedral Electron pair arrangement: tetrahedral Molecular shape: tetrahedral Molecular shape: tetrahedral NF 3 NF 3 Lewis Structure (on board) Lewis Structure (on board) Number of electron pairs: 4 Number of electron pairs: 4 Electron pair arrangement: tetrahedral Electron pair arrangement: tetrahedral Molecular shape: pyramidal Molecular shape: pyramidal

57. The VSEPR Model BeI 2 BeI 2 Lewis Structure (on board) Lewis Structure (on board) Number of electron pairs: 2 Number of electron pairs: 2 Electron pair arrangement: linear Electron pair arrangement: linear Molecular shape: linear Molecular shape: linear

58. Homework Questions: #1-6; 14-16 Questions: #1-6; 14-16 Paired Exercises: #19-29 odd; 33-57 odd Paired Exercises: #19-29 odd; 33-57 odd Additional Exercises: #59; 65-66; 69-72 Additional Exercises: #59; 65-66; 69-72