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Chapter 9:

Chapter 9:. Basic Concepts of Chemical Bonding NaCl versus C 12 H 22 O 11 . Types of Bonds. We can classify bonds based on the kinds of atoms that are bonded together. Tro: Chemistry: A Molecular Approach, 2/e. Types of Bonding. Tro: Chemistry: A Molecular Approach, 2/e.

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Chapter 9:

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  1. Chapter 9: • Basic Concepts of Chemical Bonding • NaCl versus C12H22O11.

  2. Types of Bonds • We can classify bonds based on the kinds of atoms that are bonded together. Tro: Chemistry: A Molecular Approach, 2/e

  3. Types of Bonding Tro: Chemistry: A Molecular Approach, 2/e

  4. Lewis Dot Symbols • Combines the element symbol plus the valence electrons as dots placed around symbol. • Dots are first placed on each of the four sides (N-S-E-W), then paired up after that. • Symbols for period 2.

  5. Octet Rule • Atoms tend to gain, lose, or share electrons until they are surrounded by eight electrons. • Ionic = gain or lose • Covalent = share • While there are exceptions to this rule, it is the important driving force for the formation of compounds.

  6. Ionic Bonding • When Na(s) and Cl2(g) are combined, a very violent, exothermic reaction results. • LEP #1

  7. Ionic Bonding

  8. Ionic Bonding • The energetics of ionic bond formation can be explained by a series of steps. • Step 1: Loss of electron by Na Na(s) Na(g) ; DH = +108 kJ Na(g)  Na+(g) + 1e- ; DH = +496 kJ • Step 2: Gain of electron by Cl ½ Cl2(g)  Cl(g) ; DH = +122 kJ Cl(g) + 1e-  Cl-(g) ; DH = -349 kJ

  9. Ionic Bonding • Lattice energy is the energy required to completely separate a mole of a solid ionic compound into gaseous ions. NaCl(s) Na+(g) + Cl-(g) ; DH = 788 kJ • Reverse this AND add it to the four previous reactions yields: Na(s) + ½ Cl2(g)  NaCl(s) ; DH = -411 kJ

  10. Lattice Energy • The electrostatic attraction of two charged particles is ruled by the equation: • Q1, Q2 are the magnitudes of the charges and d is the distance between the two nuclei.

  11. Lattice Energy • As the magnitude of the charge increases, the lattice energy will increase. • Na+1– Cl-1 U = +788 kJ • Sr+2 – Cl-1 U = +2127 kJ • Sr+2– O-2 U = +3217 kJ

  12. Lattice Energy • As the distance between the ions increases, the lattice energy decreases.

  13. Lattice Energy = −910 kJ/mol Lattice Energy = −3414 kJ/mol Summary of Lattice Energies • The force of attraction between oppositely charged particles is directly proportional to the product of the charges • Larger charge means the ions are more strongly attracted • larger charge = stronger attraction • stronger attraction = larger lattice energy • Of the two factors, ion charge is generally more important • LEP #2 Tro: Chemistry: A Molecular Approach, 2/e

  14. Covalent Bonding • When two electrons are shared by two atoms, this is known as a covalent bond. • Formation of H2

  15. Covalent Bonding • Formation of Cl2 – each Cl atom has seven valence electrons. Each has one unpaired electron that can pair up to make a bond.

  16. Lewis Structures • For molecules or ions containing three or more atoms, we can follow a set of rules to guide in the process. • Add up the total valence electrons from all the atoms in the compound. Ex) CF4 • Make a skeleton structure – the first element in the formula is usually the central atom – all others are then placed around this atom and connected with a bond.

  17. Lewis Structures • Fill the external atoms until they have an octet. • Compare total number of electrons used to step 1. If all are used, then go to step 5. If some are left, place on central atom as lone pair(s). • Check the central atom for an octet. If no octet, may need multiple bond(s) by moving non-bonding pair(s).

  18. Lewis Structures • General Guidelines • Group 7A as an external atom will NOT do multiple bonds. • Carbon will almost always have four bonds. • Group 2A, 3A as a central atom may be deficient of octet. • Group 5A, 6A, 7A, and Xe as a central atom may exceed the octet. • Oxygen will do up to two bonds.

  19. Electronegativity • Non-polar covalent bond – the electrons are shared equally. • Ex) F2 , Br2 , I2 , O2 , etc. • Polar covalent bond – electrons are not shared equally – one atom has a greater desire for the electron pair. • Electronegativity Scale

  20. Electronegativity

  21. Bond Polarity • If the electronegativity difference is zero, then the bond is non-polar covalent. • If there is a difference AND the two elements are non-metals, then the bond is polar covalent. • If there is a difference AND one element is a metal and the other is a non-metal, then the bond is ionic.

  22. Bond Polarity • Simple molecules like HCl have a polar covalent bond. • The more electronegative element will have a partial negative charge and the less electronegative element will have a partial positive charge.

  23. Dipole Moment • This is the quantitative measurement of the polar bond. • m = Q x d • m is measured in a unit called the Debye or Coulomb x meter. • LEP #4

  24. Resonance Structures • In some cases, a molecule or ion may be described by more than one Lewis Structure. • Ex) O3 • LEP #5

  25. Bond Strength • The energy required to break a covalent bond is its strength. • Some are fairly simple. • Cl2(g) 2 Cl(g) ; DH = 242 kJ • Others are more complicated. • CH4(g)  C(g) + 4 H(g) ; DH = 1660 kJ • Note: Energy is ALWAYS required to break a bond.

  26. Using Bond Enthalpies • One method for estimating the enthalpy of a reaction is: DH = S(Bonds Broken) – S(Bonds Made) • LEP #6

  27. Chemistry of Explosives • Many explosives are solids or liquids that contain the nitro (NO2) or nitrate (NO3) group. • The nitrogen atoms are typically weak (singly) bonded to carbon (293 kJ). • They then become N-N strong triple bonds (941 kJ).

  28. Chemistry of Explosives

  29. Bond Order • Bond order is the number of bonds connecting two atoms. • Can be 1, 2, or 3. C2H2

  30. Bond Length • As bond order increases, the bond lengths decrease. • Ex) N-N N=N NN 1.47Å 1.24Å 1.10Å • Ex) C-C C=C CC 1.54Å 1.34Å 1.20Å

  31. Resonance Structures • The bond orders and lengths for resonance structures must be averaged. • Ex) SO2 , NO3-1 , C6H6

  32. Formal Charges • Formal Charge is a fictitious charge assigned to each atom in a Lewis Structure. • It helps to evaluate the validity of competing structures. • Formal Charge = #Valence Electrons – #Nonbonding electrons – ½ #Bonding Electrons

  33. Formal Charges • General Rules: • The sum of all formal charges in a neutral molecule will equal zero. • The sum of all formal charges in an ion must equal the charge of the ion. • Small (+1 or -1 or 0) formal charges are preferred over larger ones. • When formal charges cannot be avoided, a negative formal charge will preferentially go on the more electronegative atom.

  34. Formal Charges • HCN molecule = 10 electrons. • Is skeleton structure H – C – N or H – N – C? • Which Lewis Structure for the cyanate (OCN-) ion is best?

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