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When Atoms Meet

When Atoms Meet. Bonds. Forces that hold groups of atoms together and make them function as a unit. Bonding Forces. Electron – electron repulsive forces. Nucleus – nucleus repulsive forces. Electron – necleus attractive forces. Metals and Nonmetals.

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When Atoms Meet

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  1. When Atoms Meet

  2. Bonds • Forces that hold groups of atoms together and make them function as a unit. Bonding Forces • Electron – electron • repulsive forces • Nucleus – nucleus • repulsive forces • Electron – necleus • attractive forces

  3. Metals and Nonmetals

  4. Types of Chemical Bonding 1. Metal with nonmetal: electron transfer and ionic bonding

  5. Three models of chemical bonding Ionic Electron transfer

  6. Types of Chemical Bonding 1. Metal with nonmetal: electron transfer and ionic bonding 2. Nonmetal with nonmetal: electron sharing and covalent bonding

  7. Three models of chemical bonding Ionic Covalent Electron transfer Electron sharing

  8. Types of Chemical Bonding 1. Metal with nonmetal: electron transfer and ionic bonding 2. Nonmetal with nonmetal: electron sharing and covalent bonding 3. Metal with metal: electron pooling and metallic bonding

  9. Three models of chemical bonding Ionic Covalent Metallic Electron transfer Electron sharing Electron pooling

  10. Valence Electrons Group e- configuration # of valence e- ns1 1 1A 2A ns2 2 3A ns2np1 3 4A ns2np2 4 5A ns2np3 5 6A ns2np4 6 7A ns2np5 7 • The outer shell electrons of an atom • Participate in chemical bonding 9.1

  11. Lewis Structures Developed the idea in 1902. G. N. Lewis

  12. Lewis Dot Symbols Place one dot per valence electron on each of the four sides of the element symbol. Pair the dots (electrons) until all of the valence electrons are used. . . . : . . . : N . N . . N N : . . : . Nitrogen, N, is in Group 5A and therefore has 5 valence electrons.

  13. Lewis Dot Symbols

  14. The Octet Rule Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has eight electrons in its highest occupied energy level. The same number of electrons as in the nearest noble gas The first exception to this is hydrogen, which follows the duet rule. The second exception is helium which does not form bonds because it is already “full” with its two electrons

  15. - + Li+ Li Li+ Li 1s 2s 2p 1s 2s 2p F- + + F F F 1s 2s 2p 1s 2s 2p Ionic Bond [He] [Ne] 1s22s1 1s22s22p5 1s2 1s22s22p6

  16. E = k Q+Q- r lattice energy cmpd MgF2 2957 Q= +2,-1 MgO 3938 Q= +2,-2 LiF 1036 LiCl 853 Electrostatic (Lattice) Energy Lattice energy (E) is the energy required to completely separate one mole of a solid ionic compound into gaseous ions. Q+ is the charge on the cation Q- is the charge on the anion r is the distance between the ions Lattice energy (E) increases as Q increases and/or as r decreases. r F < r Cl

  17. How should two atoms share electrons? + 8e- 7e- 7e- 8e- F F F F F F F F lonepairs lonepairs single covalent bond single covalent bond lonepairs lonepairs Covalent Bond A chemical bond in which two or more electrons are shared by two atoms. Lewis structure of F2

  18. Distribution of electron density of H2 H H

  19. single covalent bonds H H H H or H H O 2e- 2e- O 8e- O C O C O O double bonds 8e- 8e- 8e- double bonds O N N triple bond N N triple bond 8e- 8e- Lewis structure of water + + Double bond – two atoms share two pairs of electrons or Triple bond – two atoms share three pairs of electrons or

  20. F H F H Polar Covalent Bond A covalent bond with greater electron density around one of the two atoms electron rich region electron poor region e- poor e- rich d+ d-

  21. Electron density distributions in H2, F2, and HF.

  22. Electronegativities (EN) Linus Pauling 1901 - 1994 The ability of an atom in a molecule to attract shared electrons to itself

  23. Increasing difference in electronegativity Covalent Polar Covalent Ionic partial transfer of e- share e- transfer e- Classification of Bonds Difference in EN Bond Type 0 Covalent  2 Ionic 0 < and <2 Polar Covalent

  24. Classify the following bonds as ionic, polar covalent, or covalent: The bond in CsCl; the bond in H2S; and the NN bond in H2NNH2. Classification of Bonds Cs – 0.7 Cl – 3.0 3.0 – 0.7 = 2.3 Ionic H – 2.1 S – 2.5 2.5 – 2.1 = 0.4 Polar Covalent N – 3.0 N – 3.0 3.0 – 3.0 = 0 Covalent

  25. Rules for Writing Lewis Structures • Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center. • Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge. • Use one pair of electrons to form a bond (a single line) between each pair of atoms. • Arrange the remaining electrons to satisfy an octet for all atoms (duet for H), starting from outer atoms. • If a central atom does not have an octet, move in lone pairs to form double or triple bonds on the central atom as needed.

  26. Write the Lewis structure of nitrogen trifluoride (NF3). F N F F Step 1 – N is less electronegative than F, put N in center Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5) 5 + (3 x 7) = 26 valence electrons Step 3 – Draw single bonds between N and F atoms. Step 4 – Arrange remaining 20 electrons to complete octets

  27. Write the Lewis structure of the carbonate ion (CO32-). O C O 2- O Step 1 – C is less electronegative than O, put C in center • Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4) • -2 charge – 2e- 4 + (3 x 6) + 2 = 24 valence electrons Step 3 – Draw single bonds between C and O atoms Step 4 - Arrange remaining 18 electrons to complete octets Step 5 – The central C has only 6 electrons. Form a double bond.

  28. O O O C C C O O O 2- 2- 2- - - - - O O O - - Resonance More than one valid Lewis structures can be written for a particular molecule The actual structure of the carbonate ion is an average of the three resonance structures

  29. Be – 2e- 2H – 2x1e- H Be H 4e- B – 3e- 3 single bonds (3x2) = 6 3F – 3x7e- F B F 24e- Total = 24 9 lone pairs (9x2) = 18 F Exceptions to the Octet Rule The Incomplete Octet BeH2 BF3

  30. N – 5e- S – 6e- N O 6F – 42e- O – 6e- 11e- 48e- F 6 single bonds (6x2) = 12 F F Total = 48 S 18 lone pairs (18x2) = 36 F F F Exceptions to the Octet Rule Odd-Electron Molecules NO The Expanded Octet (central atom with principal quantum number n > 2) SF6

  31. Covalent BondLengths Bond Lengths Triple bond < Double Bond < Single Bond

  32. Covalent BondEnergy H2 (g) H (g) + H (g) 436.4 kJ Cl2 (g) Cl (g) + Cl (g) 242.7 kJ HCl (g) H (g) + Cl (g) 431.9 kJ O2 (g) O (g) + O (g) 498.7 kJ O N2 (g) N (g) + N (g) 941.4 kJ O N N Bond Energies Single bond < Double bond < Triple bond The energy required to break a particular bond in one mole of gaseous molecules is the bond energy. Bond Energy

  33. 3 x 1012 3 x 108 3 x 104 Light-Matter Interactions 3 x 1020 3 x 1016 Frequency in Hz Dissociation Ionization Vibration Rotation

  34. Infrared light Vibrational Modes of Water

  35. Absorbance 3000 2000 1000 2000 1500 1000 Wavenumber (cm-1) Infrared Spectrum of Water Liquid Gas Reveal the interactions between molecules and their environments

  36. Infrared Spectrum of Caffeine Absorbance 2000 3000 1000 Wavenumber (cm-1) Identification of compounds

  37. Lab 1

  38. Acknowledgment Some images, animation, and material have been taken from the following sources: Chemistry, Zumdahl, Steven S.; Zumdahl, Susan A.; Houghton Mifflin Co., 6th Ed., 2003; supplements for the instructor General Chemistry: The Essential Concepts, Chang, Raymon; McGraw-Hill Co. Inc., 4th Ed., 2005; supplements for the instructor Principles of General Chemistry, Silberberg, Martin; McGraw-Hill Co. Inc., 1st Ed., 2006; supplements for the instructor NIST WebBook: http://webbook.nist.gov/ http://www.lsbu.ac.uk/water/vibrat.html http://en.wikipedia.org/wiki/Caffeine http://www.wilsonhs.com/SCIENCE/CHEMISTRY/MRWILSON/Unit%204%20Chemical%2 0Bonding%20Powerpoint1.ppt

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