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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

Chemistry 100 Chapter 8

Chemical Bonding Basic Concepts

the valance electrons
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
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
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
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
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
Lattice Energies of Various Ionic Compounds

Determined using a thermochemical cycle -

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

covalent bonding
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
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
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

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

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
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
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
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
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

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
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
electronegativity and bond type
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
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.
slide24

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).
slide25

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
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
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
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
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
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
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.
slide32

Look at the dimerization reaction of NO2.

2 NO2 (g) ⇄ N2O4 (g) Keq = 210

valence shells having more than 8 electrons expanded octets
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

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.
slide35

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
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
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
slide38

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)
slide39

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.
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