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1 When Atoms Meet. 2 Bonding Forces  Electron – electron repulsive forces  Nucleus – nucleus repulsive forces  Electron – necleus attractive forces.

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Presentation on theme: "1 When Atoms Meet. 2 Bonding Forces  Electron – electron repulsive forces  Nucleus – nucleus repulsive forces  Electron – necleus attractive forces."— Presentation transcript:

1 1 When Atoms Meet

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

3 3 Metals and Nonmetals

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

5 5 Three models of chemical bonding Electron transfer Ionic

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

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

8 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 9 Three models of chemical bonding Electron transferElectron sharingElectron pooling Ionic CovalentMetallic

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

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

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

13 13 Lewis Dot Symbols

14 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 15 Li + F Li + F - Ionic Bond 1s 2 2s 1 1s 2 2s 2 2p 5 1s 2 1s 2 2s 2 2p 6 [He][Ne] Li 1s2s2p F 1s2s2p + Li + 1s2s2p F-F- 1s2s2p +

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

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

18 18 Distribution of electron density of H 2 HH

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

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

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

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

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

24 24 Classify the following bonds as ionic, polar covalent, or covalent: The bond in CsCl; the bond in H 2 S; and the NN bond in H 2 NNH 2. Cs – 0.7Cl – – 0.7 = 2.3Ionic H – 2.1S – – 2.1 = 0.4Polar Covalent N – – 3.0 = 0Covalent Classification of Bonds

25 25 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. Rules for Writing Lewis Structures

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

27 27 Write the Lewis structure of the carbonate ion (CO 3 2- ). Step 1 – C is less electronegative than O, put C in center OCO O Step 2 – Count valence electrons C - 4 (2s 2 2p 2 ) and O - 6 (2s 2 2p 4 ) -2 charge – 2e (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2

28 28 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 OCO O -- OCO O - - OCO O - - Resonance 22 22 22

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

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

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

32 32 The energy required to break a particular bond in one mole of gaseous molecules is the bond energy. H 2 (g) H (g) kJ Cl 2 (g) Cl (g) kJ HCl (g) H (g) +Cl (g) kJ O 2 (g) O (g) kJ OO N 2 (g) N (g) kJ N N Bond Energy Bond Energies Single bond < Double bond < Triple bond Covalent Bond Energy

33 33 Frequency in Hz 3 x x x x x 10 4 Dissociation Ionization Vibration Rotation Light-Matter Interactions

34 34 Vibrational Modes of Water Infrared light

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

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

37 37 Lab 1

38 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., 4 th 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: 0Bonding%20Powerpoint1.ppt


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