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Chemistry 100 Chapter 8. Chemical Bonding Basic Concepts. The Valance Electrons. When atoms interact to form chemical bonds, only the outer (valance) electrons take part. Need a tool for keeping track of valence electrons, e.g., The Lewis dot symbol 1 v.E. 7 v.E’s

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Chemistry 100 chapter 8 l.jpg

Chemistry 100 Chapter 8

Chemical Bonding Basic Concepts


The valance electrons l.jpg
The Valance Electrons

  • When atoms interact to form chemical bonds, only the outer (valance) electrons take part.

  • Need a tool for keeping track of valence electrons, e.g., The Lewis dot symbol

    1 v.E. 7 v.E’s

  • When these two elements combine to form a compound

    2 Na (s) + Cl2 (g) ® 2 NaCl (s)


What s happening l.jpg
What’s Happening?

[Ne]3s1 [Ne]3s23p5

(g) ® Na+ (g) + e- (ionizes, loses e-)

  • an electron configuration of [Ne]

  • (g) + e- ® Cl- (g)

    • an electron configuration of [Ar]

  • In the crystal lattice,

    • Na+ and Cl- ions; strong electrostatic attractions



  • Ionic bonding l.jpg
    Ionic Bonding

    • Electrostatic attractions that hold ions together in an ionic compound.

    • The strength of interaction depends on charge magnitude and distance between them.

    • q1 magnitude of charge 1

    • q2 magnitude of charge 2

    • r  distance between the ionic

    • centres


    Stability of ionic compounds l.jpg
    Stability of Ionic Compounds

    • The stability of ionic compounds depends on two main factors

      • The electron affinity of one of the elements

      • The ionization energy of the other

    • Note

      • electron affinities and ionization potentials are gas-phase reactions?

    • How are they related to the stability of solid materials?


    The lattice energy l.jpg
    The Lattice Energy

    • A quantitative measure of just how strong the interaction is between the ionic centres (i.e., a measure of the strength of the ionic bond)

    • For the reaction

      KCl (s)  K+ (g) + Cl- (g) H = 718 kJ/mol

    • Lattice energy (DlatH).

      • The energy required to completely separate one mole of the solid ionic compound into its gas-phase ions.


    Lattice energies of various ionic compounds l.jpg
    Lattice Energies of Various Ionic Compounds

    Determined using a thermochemical cycle -

    the Born-Haber cycle (a Hess’s Law application)


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    Covalent Bonding

    • In a wide variety of molecules, the bonding atoms fulfill their valance shell requirements by sharing electrons between them.

    • Covalent bonds - a bond in which the electrons are shared by two atoms.

      H2 ® H-H, F2® F-F, Cl2 Cl-Cl

    • For many electron atoms (like F and Cl), we again to worry only about the outermost (valence) electrons.



    Examples of covalent bonding l.jpg
    Examples of Covalent Bonding

    • Let’s look at the Cl2 example.

    • Each Cl atom has 7 valence shell electrons 3 Lone pairs and one unpaired electron

    Lone pairs

    Unshared electron


    The cl 2 molecule l.jpg
    The Cl2 Molecule

    lone pairs (non bonding)

    • The structure we have just drawn are called Lewis structures.

    • The dash in between the atomic centres represents the bonding electrons

    • Redraw F2

    bonding electrons


    Slide13 l.jpg

    • Note both Cl2 and F2 satisfy their valence shell requirements by the formation of a single bond.

    • What about O2? How can we satisfy the octet rule for 2 O atoms?

    Valence shell requirements are satisfied by the formation of a double bond.


    Slide14 l.jpg

    • check out N2  :NºN: (triple bond)

    • Note that the octet rule works mainly for the second row elements.

    • Filled valence shells can have more than 8 electrons after Z=14 (Si). This is generally termed octet expansion.


    Covalent compounds l.jpg
    Covalent Compounds

    • Compounds that contain only covalent bonds are called covalent compounds.

    • There are two main of covalent compounds,

      • Molecular covalent compounds (CO2, C2H4)

      • Network covalent compounds (SiO2, BeCl2).

      • The network covalent compound are characterized by an extensive “3-D” network bonding


    Comparison between ionic and covalent compounds l.jpg
    Comparison between Ionic and Covalent Compounds

    • Ionic Compounds

      • usually solids with very high melting points

      • conduct electricity when molten (melted)

      • usually quite water soluble and they are electrolytes in aqueous solution

      • NaCl

    • Covalent Compounds

      • usually low melting solids, gases or liquids

      • don’t conduct electricity when molten

      • aren’t very soluble in water and are non electrolytes

      • CCl4


    The filled valence shell rule l.jpg
    The Filled Valence Shell rule

    • Filled Valence Shell rule

      • Atoms participate in the formation of bonds (either ionic or covalent) in order to satisfy their valence shell requirements.

    • Atoms other than H tend to form bonds until they end up being surrounded by 8 valence electrons (the noble gas configuration).

    • This is known as the octet rule.


    Electronegativity l.jpg
    Electronegativity

    • Electronegativity is defined as the ability of an atom to attract electrons towards itself in a molecule ( (pronounced ‘chi’))

    • Examine the H-F covalent bond

      +H-F

      •  denotes a partial “+” charge on the H atom

      • - denotes a partial “-“ charge on F atom


    Slide19 l.jpg


    Trends in the values l.jpg
    Trends in the  the ionization energy. Values

    • Across a row

      • The  values generally increase as we proceed from left to right in the periodic table.

    • Down a group

      • The  values generally decrease as we descend the group.

    • Transition metals

      • Essentially constant  values


    Plot of values l.jpg
    Plot of the ionization energy.  Values


    Electronegativity and bond type l.jpg
    Electronegativity and Bond Type the ionization energy.

    • Can we use the electronegativity values to help us deduce the type of bonding in compounds?


    An outline for drawing lewis structures l.jpg
    An Outline for Drawing Lewis Structures the ionization energy.

    • Predict arrangement of atoms (i.e., predict the skeletal arrangement of the molecule or ion).

      • The H is always a terminal atom, bonded to ONE OTHER ATOM ONLY. A halogen atom is usually a terminal atom.

      • Note that the central atom usually has the least negative electron affinity.


    Slide24 l.jpg

    • Count the ionization energy. total number of valence shell electrons (include ionic charges).

    • Place 1 pair electrons (sigma bond, ) between each pair of bonded atoms (i.e., the central atom and each one of the terminal atoms).

    • Place remaining electrons around the terminal atoms to satisfy the filled valence shell rule. (lone pairs).


    Slide25 l.jpg


    Formal charges l.jpg
    Formal Charges Atoms in the 3rd or higher row can have more than eight electrons around them.

    • Definition: formal charge on atom = number of valence electrons – number of non-bonding - ½ the number of bonding electrons.

    • Formal charge in a Lewis Structure is a bookkeeping “device”

      • keeps track of the electrons “associated” with certain atoms in the molecule vs. the valence e-‘s in the isolated atom!

    • How does it work?


    Rules for formal charges l.jpg
    Rules for Formal Charges Atoms in the 3rd or higher row can have more than eight electrons around them.

    • Neutral molecules ®S formal charges = 0

    • Ions ® S formal charges = charge of ion

    • For molecules where the possibility of multiple Lewis Structures with different formal charges exist

      • Neutral molecule - choose the structure with the fewest formal charges.

      • Structures with large formal charges are less likely than ones with small formal charges

      • Two Lewis Structures with similar formal charge distribution ® negative formal charges on more electronegative atom


    Resonance structures l.jpg
    Resonance Structures Atoms in the 3rd or higher row can have more than eight electrons around them.

    • Examine the NO3- anion.

    • The structures differ in the location of the N=O double bond.

    • They are said to be resonance structures.

    • The actual structure of the molecule is a combination of three resonance structures (the resonance hybrid).


    Experimental evidence for resonance l.jpg
    Experimental Evidence for Resonance. Atoms in the 3rd or higher row can have more than eight electrons around them.

    • The resonance structures for benzene C6H6

    • We would expect to find two different bond lengths in benzene (C=C and C-C bonds).

      • C= C ® bond length = 133 pm = 0.133 nm

      • C- C ® bond length = 0.154 nm

    • Experimentally, all benzene carbon-carbon bond lengths are equivalent at 0.140 nm


    Exceptions to the filled valence shell rule l.jpg
    Exceptions to the Filled Valence Shell Rule Atoms in the 3rd or higher row can have more than eight electrons around them.

    • Be compounds  BeH2, BeCl2,

    • Boron and Al compounds  BF3, AlCl3, BCl3

    • BF3 is stable Þ The B central atom has a tendency to pick up an unshared e- pair from another compound

      BF3 + NH3® BF3NH3

    • the B-N bond is an example of a coordinate covalent bond, or a “dative” bond ® i.e. a bond in which one of the atoms donates both bonding electrons.


    Odd e molecules l.jpg
    Odd e- molecules Atoms in the 3rd or higher row can have more than eight electrons around them.

    • These molecules have uneven numbers of electrons \ no way that they can form octets.

    • Examples

      • NO and NO2. These species have an odd number of electrons.


    Slide32 l.jpg


    Valence shells having more than 8 electrons expanded octets l.jpg
    Valence Shells having more than 8 Electrons (Expanded Octets)

    • A central atom having more than 8 valance shell electrons is possible with atomic number 14 and above.

    Reason - elements in this category can use the energetically low-lying d orbitals to accommodate extra electrons


    Slide34 l.jpg

    • High formal charge on the electronegative Cl atom (f.c.(Cl) = 7-2-1/2 (6) = +2)

    • This resonance structure would make a very small contribution to the overall resonance hybrid.


    Slide35 l.jpg

    • Note: the final three structures reduce the formal charges


    Bond energies and thermochemistry l.jpg
    Bond Energies and Thermochemistry the Cl atom to accommodate extra electron pairs, we may write other Lewis structures

    • Look at the energy required to break 1 mole of gaseous diatomic molecules into their constituent gaseous atoms.

      H2 (g) ® H (g) + H (g) DH° = 436.4 kJ

      Cl2 (g) ® Cl (g) + Cl (g) DH° = 242 kJ

    • These enthalpy changes are called bond dissociation energies. In the above examples, the enthalpy changes are designated D (H-H) and D (Cl-Cl).


    For polyatomic molecules l.jpg
    For Polyatomic Molecules. the Cl atom to accommodate extra electron pairs, we may write other Lewis structures

    CO2 (g) ® C (g) + 2 O (g) DH = 745 kJ

    • Denote the DH of this reaction D(C=O)

    • What about dissociating methane into C + 4 H’s?

      CH4 (g) ® C(g) + 4 H (g) DH° = 1650 kJ

    • Note 4 C-H bonds in CH4\ D (C-H) = 412 kJ/mol


    Slide38 l.jpg

    H the Cl atom to accommodate extra electron pairs, we may write other Lewis structures2O (g) ® 2 H (g) + O (g) DH° = 929 kJ/mol H2O

    • It takes more energy to break the first O-H bond.

      H2O (g) ® H (g) + OH (g) DH° = 502 kJ/mol H2O

      HO (g) ® H (g) + O (g) DH= 427 kJ/mol H2O

    • Note: we realize that all chemical reactions involve the breaking and reforming of chemical bonds.

      • Break bonds  add energy.

      • Make bonds  energy is released.

    • rxnH° S D(bonds broken) - S D(bonds formed)


    Slide39 l.jpg

    • These are close but not quite exact. Why? the Cl atom to accommodate extra electron pairs, we may write other Lewis structures

    • The bond energies we use are averaged bond energies, i.e.,

    • This is a good approximate for equations involving diatomic species.

    • We can only use the above procedure for GAS PHASE REACTIONS ONLY.