Chemistry 100 chapter 8 l.jpg
This presentation is the property of its rightful owner.
Sponsored Links
1 / 39

Chemistry 100 Chapter 8 PowerPoint PPT Presentation

  • Updated On :
  • Presentation posted in: General

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

Download Presentation

Chemistry 100 Chapter 8

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

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

  • When these two elements combine to form a compound

    2 Na (s) + Cl2 (g) ® 2 NaCl (s)

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

  • The NaCl Crystal

    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

    • 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

    • 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

    Determined using a thermochemical cycle -

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

    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.

    Covalent Bonding

    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 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

    • 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.

    • 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

    • 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

    • 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

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

    • Examine the H-F covalent bond


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

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

    • Electronegativity is related to the electron affinity and the ionization energy.

    • Compare the following elements.

      • Na  low I1, small negative E.A.  low 

      • F  high I1, large, negative E.A.,  high 

    Trends in the  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

    Electronegativity and Bond Type

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

    An Outline for Drawing Lewis Structures

    • 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.

    • Count 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).

    • All remaining electrons are assigned to the central atom. Atoms in the 3rd or higher row can have more than eight electrons around them.

      • If a central atom does not have a filled valence shell, use a lone pair of electrons from a terminal atom to make a pi () bond.

    Formal Charges

    • 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

    • 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

    • 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.

    • 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

    • 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

    • 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.

    • Look at the dimerization reaction of NO2.

      2 NO2 (g) ⇄ N2O4 (g) Keq = 210

    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

    • Look at HClO3

    • 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.

    • With the possibility of using the low lying d-orbitals on the Cl atom to accommodate extra electron pairs, we may write other Lewis structures

    • Note: the final three structures reduce the formal charges

    Bond Energies and Thermochemistry

    • 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.

    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

    H2O (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)

    • These are close but not quite exact. Why?

    • 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.

  • Login