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covalent Bonding

covalent Bonding. Group. e - configuration. # of valence e -. ns 1. 1. 1A. 2A. ns 2. 2. 3A. ns 2 np 1. 3. 4A. ns 2 np 2. 4. 5A. ns 2 np 3. 5. 6A. ns 2 np 4. 6. 7A. ns 2 np 5. 7.

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covalent Bonding

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  1. covalent Bonding

  2. 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 Valence electrons are the outer shell electrons of an atom. The valence electrons are the electrons that participate in chemical bonding.

  3. - - - - Li+ Li Li Li+ + e- e- + Li+ Li+ + F F F F F F The Ionic Bond + 1s22s22p6 1s22s1 1s22s22p5 1s2 [He] [Ne]

  4. E = k Q+Q- r 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 compound Lattice energy (E) increases as Q increases and/or as r decreases. r F < r Cl

  5. o o o o o o DHoverall = DH1 + DH2 + DH3 + DH4 + DH5 Born-Haber Cycle for Determining Lattice Energy

  6. Chemistry In Action: Sodium Chloride Mining Salt Solar Evaporation for Salt

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

  8. 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- O N N N N triple bond Lewis structure of water + + Double bond – two atoms share two pairs of electrons or Triple bond – two atoms share three pairs of electrons or 8e- 8e-

  9. Lengths of Covalent Bonds Bond Lengths Triple bond < Double Bond < Single Bond

  10. F H Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms electron rich region electron poor region e- poor e- rich H F d+ d-

  11. X (g) + e- X-(g) Electronegativityis the ability of an atom to attract toward itself the electrons in a chemical bond. Electron Affinity - measurable, Cl is highest Electronegativity - relative, F is highest

  12. Increasing difference in electronegativity Covalent Polar Covalent Ionic partial transfer of e- share e- transfer e- Classification of bonds by difference in electronegativity Difference Bond Type 0 Covalent  2 Ionic 0 < and <2 Polar Covalent

  13. Bonding Triangle

  14. Classify the following bonds as ionic, polar covalent,orcovalent: The bond in CsCl; the bond in H2S; and the NN bond in H2NNH2. 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

  15. 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. • Complete an octet for all atoms except hydrogen • If structure contains too many electrons, form double and triple bonds on central atom as needed.

  16. F N F F Write the Lewis structure of nitrogen trifluoride (NF3). 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 and complete octets on N and F atoms. Step 4 - Check, are of e- in structure equal to number of valence e- ? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons

  17. O C O 2 single bonds (2x2) = 4 1 double bond = 4 8 lone pairs (8x2) = 16 O Total = 24 Write the Lewis structure of the carbonate ion (CO32-). 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 and complete octet on C and O atoms. Step 4 - Check, are # of e- in structure equal to number of valence e- ? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons Step 5 - Too many electrons, form double bond and re-check of e-

  18. H H C O H C O H formal charge on an atom in a Lewis structure total number of valence electrons in the free atom total number of nonbonding electrons ( total number of bonding electrons ) 1 - - = 2 Two possible skeletal structures of formaldehyde (CH2O) An atom’s formal charge is the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure. The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion.

  19. H C O H C – 4 e- O – 6 e- 2H – 2x1 e- 12 e- formal charge on an atom in a Lewis structure total number of valence electrons in the free atom total number of nonbonding electrons ( total number of bonding electrons ) 1 - - = 2 single bonds (2x2) = 4 2 1 double bond = 4 2 lone pairs (2x2) = 4 Total = 12 -1 +1 formal charge on C = 4 - 2 - ½ x 6 = -1 formal charge on O = 6 - 2 - ½ x 6 = +1

  20. C – 4 e- H O – 6 e- C O H 2H – 2x1 e- 12 e- formal charge on an atom in a Lewis structure total number of valence electrons in the free atom total number of nonbonding electrons ( total number of bonding electrons ) 1 - - = 2 single bonds (2x2) = 4 2 1 double bond = 4 2 lone pairs (2x2) = 4 Total = 12 0 0 formal charge on C = 4 - 0 - ½ x 8 = 0 formal charge on O = 6 - 4 - ½ x 4 = 0

  21. H C O H H C O H -1 +1 0 0 Formal Charge and Lewis Structures • For neutral molecules, a Lewis structure in which there are no formal charges is preferable to one in which formal charges are present. • Lewis structures with large formal charges are less plausible than those with small formal charges. • Among Lewis structures having similar distributions of formal charges, the most plausible structure is the one in which negative formal charges are placed on the more electronegative atoms. Which is the most likely Lewis structure for CH2O?

  22. What are the resonance structures of the carbonate (CO32-) ion? - - + + O O O O O O O O O C C C O O O - - - - O O O - - A resonance structureis one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure (after formal charge has been determined!). The true structure is an AVERAGE of all the possible structures. More possible structures gives the overall structure more validity.

  23. BF3 SF4 Violations of the Octet Rule Usually occurs with B and elements of higher periods and most nonmetals. Common exceptions are: Be, B, P, S, Xe, Cl, Br, and As. How do you know if it’s an EXPANDED octet? • More than 4 bonds • Formal Charge doesn’t work out with just 8 Be: 4 B: 6 P: 8 OR 10 S: 8, 10, OR 12 Xe: 8, 10, OR 12

  24. 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

  25. 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

  26. DH0 = 436.4 kJ H2 (g) H (g) + H (g) DH0 = 242.7 kJ Cl2 (g) Cl (g) + Cl (g) DH0 = 431.9 kJ HCl (g) H (g) + Cl (g) DH0 = 498.7 kJ O2 (g) O (g) + O (g) O DH0 = 941.4 kJ N2 (g) N (g) + N (g) O Bond Energies Single bond < Double bond < Triple bond N N The enthalpy change required to break a particular bond in one mole of gaseous molecules is the bond energy. Bond Energy

  27. Bond Energies (BE) and Enthalpy changes in reactions Imagine reaction proceeding by breaking all bonds in the reactants and then using the gaseous atoms to form all the bonds in the products. DH0 = total energy input – total energy released = SBE(reactants) – SBE(products)

  28. 2H2 (g) + O2 (g) 2H2O (g) H2 (g) + Cl2 (g) 2HCl (g)

  29. Use bond energies to calculate the enthalpy change for: H2 (g) + F2 (g) 2HF (g) Type of bonds broken Number of bonds broken Bond energy (kJ/mol) Energy change (kJ) 1 436.4 436.4 1 156.9 156.9 Type of bonds formed Number of bonds formed Bond energy (kJ/mol) Energy change (kJ) H H F H F F 2 568.2 1136.4 DH0 = SBE(reactants) – SBE(products) DH0 = 436.4 + 156.9 – 2 x 568.2 = -543.1 kJ

  30. Molecular Geometry

  31. VSEPR: Valence Shell Electron Pair Repulsion theory. Most important factor in determining geometry is relative repulsion between electron pairs. Molecule adopts the shape that minimizes the electron pair repulsions.

  32. Some Common Geometries Linear Tetrahedral Trigonal Planar

  33. VSEPR charts • Use the Lewis structure to determine the geometry of the molecule • Electron arrangement establishes the bond angles • Molecule takes the shape of that portion of the electron arrangement • Charts look at the CENTRAL atom for all data! • Think REGIONS OF ELECTRON DENSITY rather than bonds (for instance, a double bond would only be 1 region)

  34. Structure Determination by VSEPR Water, H2O The electron pair geometry is TETRAHEDRAL 2 bond pairs 2 lone pairs The molecular geometry is BENT.

  35. Ammonia, NH3 The electron pair geometry is tetrahedral. The MOLECULAR GEOMETRY— the positions of the atoms — is TRIGONAL PYRAMID.

  36. Bond Polarity HClisPOLAR because it has a positive end and a negative end. (difference in electronegativity) Cl has a greater share in bonding electrons than does H. Cl has slight negative charge (-d) and H has slight positive charge (+ d)

  37. Bond Polarity • This is why oil and water will not mix! Oil is nonpolar, and water is polar. • The two will repel each other, and so you can not dissolve one in the other

  38. Bond Polarity “Like Dissolves Like” • Polar dissolves Polar • Nonpolar dissolves Nonpolar

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