When Atoms Meet 1 Bonds q Forces that

When Atoms Meet 1

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

Metals and Nonmetals 3

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

Three models of chemical bonding Ionic Electron transfer 5

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

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

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 8

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

Valence Electrons • The outer shell electrons of an atom • Participate in chemical bonding Group e- configuration # of valence e- 1 A ns 1 1 2 A ns 2 2 3 A ns 2 np 1 3 4 A ns 2 np 2 4 5 A ns 2 np 3 5 6 A ns 2 np 4 6 7 A ns 2 np 5 7 10

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

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. : Nitrogen, N, is in Group 5 A and therefore has 5 valence electrons. . . . N: : N. . : 12

Lewis Dot Symbols 13

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 14

Ionic Bond Li+ F - Li + F [He] 1 s 1 s 2[2 Ne] 2 s 22 p 6 1 22 s 22 p 5 1 s 22 s 1 s Li+ Li 1 s 2 s 1 s 2 p + + F 1 s 2 s 2 p 15

Electrostatic (Lattice) Energy Lattice energy (E) is the energy required to completely separate one mole of a solid ionic compound into gaseous ions. Q+QE=k r 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. cmpd Mg. F 2 Mg. O Li. F Li. Cl lattice energy 2957 Q= +2, -1 3938 Q= +2, -2 1036 853 r F < r Cl 16

Covalent Bond A chemical bond in which two or more electrons are shared by two atoms. How should two atoms share electrons? F + 7 e- F F F 7 e- 8 e- Lewis structure of F 2 single covalent bond lone pairs F F lone pairs single covalent bond 17

Distribution of electron density of H 2 H H 18

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

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

Electron density distributions in H 2, F 2, and HF. 21

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

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

Classification of Bonds Classify the following bonds as ionic, polar covalent, or covalent: The bond in Cs. Cl; the bond in H 2 S; and the NN bond in H 2 NNH 2. Cs – 0. 7 Cl – 3. 0 – 0. 7 = 2. 3 Ionic H – 2. 1 S – 2. 5 – 2. 1 = 0. 4 Polar Covalent N – 3. 0 – 3. 0 = 0 Covalent 24

Rules for Writing Lewis Structures 1. Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center. 2. Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge. 3. Use one pair of electrons to form a bond (a single line) between each pair of atoms. 4. Arrange the remaining electrons to satisfy an octet for all atoms (duet for H), starting from outer atoms. 5. 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. 25

Write the Lewis structure of nitrogen trifluoride (NF 3). Step 1 – N is less electronegative than F, put N in center Step 2 – Count valence electrons N - 5 (2 s 22 p 3) and F - 7 (2 s 22 p 5) 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 F N F F 26

Write the Lewis structure of the carbonate ion (CO 32 -). Step 1 – C is less electronegative than O, put C in center Step 2 – Count valence electrons C - 4 (2 s 22 p 2) and O - 6 (2 s 22 p 4) -2 charge – 2 e 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. 2 - O C O O 27

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 - 2 - O C O O - - 2 - O C O O 28 -

Exceptions to the Octet Rule The Incomplete Octet Be. H 2 BF 3 B – 3 e 3 F – 3 x 7 e 24 e- Be – 2 e 2 H – 2 x 1 e 4 e- F B H F Be H 3 single bonds (3 x 2) = 6 9 lone pairs (9 x 2) = 18 Total = 24 F 29

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

Covalent Bond Lengths Bond Type Bond Length (pm) C -C 154 C C 133 C C 120 C -N 143 C N 138 C N 116 Bond Lengths Triple bond < Double Bond < Single Bond 31

Covalent Bond Energy The energy required to break a particular bond in one mole of gaseous molecules is the bond energy. Bond Energy H 2 (g) H (g) + H (g) 436. 4 k. J Cl 2 (g) Cl (g) + Cl (g) 242. 7 k. J HCl (g) H (g) + Cl (g) 431. 9 k. J O 2 (g) O (g) + O (g) 498. 7 k. J O O N 2 (g) N (g) + N (g) 941. 4 k. J N N Bond Energies Single bond < Double bond < Triple bond 32

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

Vibrational Modes of Water Infrared light 34

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

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

Lab 1 37

Acknowledgment Some images, animation, and material have been taken from the following sources: Chemistry, Zumdahl, Steven S. ; Zumdahl, Susan A. ; Houghton Mifflin Co. , 6 th Ed. , 2003; supplements for the instructor General Chemistry: The Essential Concepts, Chang, Raymon; Mc. Graw-Hill Co. Inc. , 4 th Ed. , 2005; supplements for the instructor Principles of General Chemistry, Silberberg, Martin; Mc. Graw-Hill Co. Inc. , 1 st Ed. , 2006; supplements for the instructor NIST Web. Book: 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%20 Chemical%2 0 Bonding%20 Powerpoint 1. ppt 38