basic concepts in chemical bonding n.
Download
Skip this Video
Loading SlideShow in 5 Seconds..
Basic Concepts in Chemical Bonding PowerPoint Presentation
Download Presentation
Basic Concepts in Chemical Bonding

Loading in 2 Seconds...

play fullscreen
1 / 63

Basic Concepts in Chemical Bonding - PowerPoint PPT Presentation


  • 122 Views
  • Uploaded on

Basic Concepts in Chemical Bonding. Chapter 8. What is a Bond?. A force that holds atoms together. Why? We will look at it in terms of energy. Bond energy the energy required to break a bond. Why are compounds formed? Because it gives the system the lowest energy.

loader
I am the owner, or an agent authorized to act on behalf of the owner, of the copyrighted work described.
capcha
Download Presentation

Basic Concepts in Chemical Bonding


An Image/Link below is provided (as is) to download presentation

Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author.While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server.


- - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - -
    Presentation Transcript
    1. Basic Concepts in Chemical Bonding Chapter 8

    2. What is a Bond? • A force that holds atoms together. • Why? • We will look at it in terms of energy. • Bond energy the energy required to break a bond. • Why are compounds formed? • Because it gives the system the lowest energy.

    3. Lewis Symbols and the Octet Rule • There are three types of chemical bonds (intramolecular force) • Ionic Bond - electrostatic attraction between ions of opposite charge (MgO) • Covalent Bond - sharing of electrons between two atoms (Br2) • Metallic Bond - sharing of electrons between several atoms (Cu).

    4. Keeping Track of Electrons • The electrons responsible for the chemical properties of atoms are those in the outer energy level. • Valence electrons - electrons in the outer energy level. • Core electrons -those in the energy levels below.

    5. Keeping Track of Electrons • Atoms in the same column • Have the same outer electron configuration. • Have the same valence electrons. • Easily found by looking up the group number on the periodic table. • Group 2A - Be, Mg, Ca, etc.- • 2 valence electrons

    6. Electron Dot diagrams (Lewis Symbols) • A way of keeping track of valence electrons. • How to write them • Write the symbol. • Put one dot for each valence electron • Don’t pair up until you have too (Hund’s rule). X

    7. Lewis Symbols and the Octet Rule Octet rule – Atoms tend to gain, lose or share electrons until they are surrounded by eight valence electrons.

    8. Ionic BondingEnergetics of Ionic Bonding As we saw in the last chapter, it takes 495 kJ/mol to remove electrons from sodium.

    9. Energetics of Ionic Bonding We get 349 kJ/mol back by giving electrons to chlorine.

    10. Energetics of Ionic Bonding • But these numbers don’t explain why the reaction of sodium metal and chlorine gas to form sodium chloride is so exothermic!

    11. Energetics of Ionic Bonding • There must be a third piece to the puzzle. • What is as yet unaccounted for is the electrostatic attraction between the newly formed sodium cation and chloride anion.

    12. Q1Q2 d Eel =  Lattice Energy • This third piece of the puzzle is the lattice energy: The energy required to completely separate a mole of a solid ionic compound into its gaseous ions. • The energy associated with electrostatic interactions is governed by Coulomb’s law:

    13. Energetics of Ionic Bond Formation • Coulomb’s equation: • Q1, Q2 = charge on ions • k = 8.99 x 109 J-m/c2 • d = distance between ions • Lattice energy is greater with more highly charged ions.

    14. Lattice Energy • Lattice energy, then, increases with the charge on the ions. • It also increases with decreasing size of ions.

    15. Energetics of Ionic Bonding By accounting for all three energies (ionization energy, electron affinity, and lattice energy), we can get a good idea of the energetics involved in such a process.

    16. Energetics of Ionic Bonding • These phenomena also helps explain the “octet rule.” • Metals, for instance, tend to stop losing electrons once they attain a noble gas configuration because energy would be expended that cannot be overcome by lattice energies.

    17. Electron Configurations of Ions of the Representative Elements • These are derived from the electron configuration of elements with the required number of electrons added or removed from the most accessible orbital. • Electron configurations can predict stable ion formation: • Mg: [Ne]3s2 • Mg+: [Ne]3s1not stable • Mg2+: [Ne]stable • Cl: [Ne]3s23p5 • Cl-: [Ne]3s23p6= [Ar]stable

    18. Transition Metal Ions and Polyatomic Ions • Transition Metal Ions • Lattice energies compensate for the loss of up to three electrons. • In general, electrons are removed from orbitals in order of decreasing n (i.e. electrons are removed from 4s before the 3d). • Polyatomic Ions • Polyatomic ions are formed when there is an overall charge on a compound containing covalent bonds. E.g. SO42-, NO3-.

    19. Covalent bonds • Nonmetals hold onto their valence electrons • They can’t give away electrons to bond • Still want noble gas configuration • Get it by sharing valence electrons with each other • By sharing both atoms get to count the electrons toward noble gas configuration.

    20. Covalent Bonding • There are several electrostatic interactions in these bonds: • Attractions between electrons and nuclei • Repulsions between electrons • Repulsions between nuclei

    21. Lewis Structures Lewis structures are representations of molecules showing all electrons, bonding and nonbonding.

    22. Find the sum of valence electrons of all atoms in the polyatomic ion or molecule. If it is an anion, add one electron for each negative charge. If it is a cation, subtract one electron for each positive charge. PCl3 Writing Lewis Structures 5 + 3(7) = 26

    23. Writing Lewis Structures • The central atom is the least electronegative element that isn’t hydrogen. Connect the outer atoms to it by single bonds. Keep track of the electrons: 26  6 = 20

    24. Writing Lewis Structures • Fill the octets of the outer atoms. Keep track of the electrons: 26  6 = 20  18 = 2

    25. Writing Lewis Structures • Fill the octet of the central atom. Keep track of the electrons: 26  6 = 20  18 = 2  2 = 0

    26. Writing Lewis Structures • If you run out of electrons before the central atom has an octet… …form multiple bonds until it does.

    27. Multiple Bonds • Sometimes atoms share more than one pair of valence electrons. • A double bond is when atoms share two pair (4) of electrons. • A triple bond is when atoms share three pair (6) of electrons. • Generally, bond distances decrease as we move from single through double to triple bonds.

    28. Writing Lewis Structures • Then assign formal charges. • HUH???? Formal What???

    29. Formal Charge • It is possible to draw more than one Lewis structure with the octet rule obeyed for all the atoms. • To determine which structure is most reasonable, we use formal charge. • Formal chargeis the charge on an atom that it would have if all the atoms had the same electronegativity (i.e., if the electrons are shared equally between atoms).

    30. Using Formal Charge • A Lewis structure in which all formal charges in a molecule are equal to zero is preferable to one in which some formal charges are not zero. • If a Lewis structure has non-zero formal charges, the one with the fewest nonzero formal charges is preferred. • A Lewis structure with one large formal charge is preferable to one with several small formal charges. • A Lewis structure with adjacent formal charges should have opposite signs. • When choosing between several Lewis structures, the structure with negative formal charges on the more electronegative atom is preferable.

    31. Writing Lewis Structures • For each atom, count the electrons in lone pairs and half the electrons it shares with other atoms. • Subtract that from the number of valence electrons for that atom: The difference is its formal charge.

    32. Writing Lewis Structures • The best Lewis structure… • …is the one with the fewest charges. • …puts a negative charge on the most electronegative atom.

    33. Resonance Structures • Some molecules are not well described by Lewis Structures. • Typically, structures with multiple bonds can have similar structures with the multiple bonds between different pairs of atoms • Example: experimentally, ozone has two identical bonds whereas the Lewis Structure requires one single (longer) and one double bond (shorter).

    34. Resonance This is the Lewis structure we would draw for ozone, O3. + -

    35. Resonance Just as green is a synthesis of blue and yellow… …ozone is a synthesis of these two resonance structures.

    36. Resonance • But this is at odds with the true, observed structure of ozone, in which… • …both O—O bonds are the same length.

    37. Resonance • In truth, the electrons that form the second C—O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon. • They are not localized, but rather are delocalized.

    38. Resonance in Benzene • Benzene consists of 6 carbon atoms in a hexagon. Each C atom is attached to two other C atoms and one hydrogen atom. • There are alternating double and single bonds between the C atoms. • Experimentally, the C-C bonds in benzene are all the same length. • Experimentally, benzene is planar.

    39. Resonance in Benzene • We write resonance structures for benzene in which there are single bonds between each pair of C atoms and the 6 additional electrons are delocalized over the entire ring: • Benzene belongs to a category of organic molecules called aromatic compounds (due to their odor).

    40. Example • Draw the lewis dot structure for the following: • Nitrate ion • Sufite ion • Carbon dioxide • Thiocyanate ion • Carbonate ion • Formate ion

    41. Exceptions to the Octet Rule • There are three types of ions or molecules that do not follow the octet rule: • Ions or molecules with an odd number of electrons. • Ions or molecules with less than an octet. • Ions or molecules with more than eight valence electrons (an expanded octet).

    42. Odd Number of Electrons Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons.

    43. Fewer Than Eight Electrons • Consider BF3: • Giving boron a filled octet places a negative charge on the boron and a positive charge on fluorine. • This would not be an accurate picture of the distribution of electrons in BF3.

    44. Fewer Than Eight Electrons Therefore, structures that put a double bond between boron and fluorine are much less important than the one that leaves boron with only 6 valence electrons.

    45. Fewer Than Eight Electrons The lesson is: If filling the octet of the central atom results in a negative charge on the central atom and a positive charge on the more electronegative outer atom, don’t fill the octet of the central atom.

    46. More than an Octet • This is the largest class of exceptions. • Atoms from the 3rd period onwards can accommodate more than an octet. • Beyond the third period, the d-orbitals are low enough in energy to participate in bonding and accept the extra electron density.

    47. More Than An Octet • The only way PCl5 can exist is if phosphorus has 10 electrons around it.

    48. More Than Eight Electrons Even though we can draw a Lewis structure for the phosphate ion that has only 8 electrons around the central phosphorus, the better structure puts a double bond between the phosphorus and one of the oxygens.

    49. More Than Eight Electrons • This eliminates the charge on the phosphorus and the charge on one of the oxygens. • The lesson is: When the central atom is on the 3rd row or below it can expand its octet to eliminate some formal charges.

    50. Bond Polarity and Electronegativity • Although atoms often form compounds by sharing electrons, the electrons are not always shared equally. • Fluorine pulls harder on the electrons it shares with hydrogen than hydrogen does. • Therefore, the fluorine end of the molecule has more electron density than the hydrogen end.