Concepts of Chemical Bonding

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Concepts of Chemical Bonding

• Brown, LeMay Ch 8
• AP Chemistry
• Monta Vista High School

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8.1 Types of Inter-Atomic or Intra molecular
Bonding a.k.a bonding

• Ionic electrostatic attraction between
  oppositely charged ions. Ex. NaCl, K2SO4.
  Generally solids.
• Covalent sharing of e- between two atoms
  (typically between nonmetals). Ex. CO2, SO2.
  Generally gases and liquids
• Metallic sea of e- bonding e- are relatively
  free to move throughout the 3D structure
• Ex. Fe, Al, generally solids

IncreasingDiff. of EN
4. Covalent Network large number of
atoms/molecules bonded in a network through
covalent bonding. Ex. SiO2, Si, Ge, Diamond,
Graphite
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8.2 Ionic Bonding

• Results as atoms lose or gain e- to achieve a
  noble gas e- configuration is typically
  exothermic.
• The bonded state is lower in energy (and
  therefore more stable).
• Electrostatic attraction results from the
  opposite charges. Electrostatic attraction
  (force) determines the strength of ionic bond,
  just like electro negativity difference
  determines the strength of covalent bond.
• Occurs when diff. of EN of atoms is gt 1.7
  (maximum is 3.3 CsF)
• Can lead to interesting crystal structures (Ch.
  11).
• Use brackets when writing Lewis symbols of ions.
• Ex Draw the Lewis symbol of fluoride.
  Animation of Ionic/Covalent BondingAnimation
  of LiCl Crystal

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Lattice Energy

• Measurement of the energy of stabilization
  present in ionic solids
• DHlattice energy required to completely
  separate 1 mole of solid ionic compound into its
  gaseous ions

• Electrostatic attraction (and thus lattice
  energy) increases as ionic charges increase and
  as ionic radii decrease. Animation showing
  formation of a lattice

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• The lattice energy of NaCl is the energy given
  off when Na and Cl- ions in the gas phase come
  together to form the lattice of alternating Na
  and Cl- ions in the NaCl crystal (in this case
  lattice energy is negative) or it may be defined
  as the amount of energy required to break 1 mole
  of solid NaCl into its ions in gaseous state (in
  this case lattice energy will be positive).
• Na(g) Cl-(g) ? NaCl(s) ?Ho -787.3 kJ/mol

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• The lattice energies for the alkali metal halides
  is therefore largest for LiF (smallest size,
  smallest d) and smallest for CsI ( largest size,
  largest d), as shown below.
• Values of Lattice Energies of Alkali Metals
  Halides (kJ/mol)
• F- Cl- Br- I-
• Li 1036 853 807 757
• Na 923 787 747 704
• K 821 715 682 649
• Rb 785 689 660 630
• Cs 740 659 631 604 Data taken from Purdue
  website

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Understanding Check

• Ex Which has a greater lattice energy? Why?
• NaCl or KCl NaCl or MgS

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Covalent Bonding

• Atoms share e- to achieve noble gas configuration
  that is lower in energy (and therefore more
  stable).
• Occurs when diff. of EN of atoms is 1.7
• Polar covalent
• 0.3 lt diff. of EN 1.7 (e- pulled closer to
  more EN atom)
• Nonpolar covalent
• 0 diff. of EN 0.3 (e- shared equally)
• (ionic vs. covalent bonding in youtube video)
• Coordinate Covalent Shared pair contributed by
  only one of the two sharing species. Ex. Lewis
  acids and bases

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Cl-Cl Bond in Cl2 molecule
Animation of bonding in Chlorine Molecule
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23.5 Metallic bonding

• Metallic elements have low I.E. this means
  valence e- are held loosely.
• A metallic bond forms between metal atoms because
  of the movement of valence e- from atom to atom
  to atom in a sea of electrons. The metal thus
  consists of cations held together by
  negatively-charged e- "glue.

• This results in excellent thermal electrical
  conductivity, ductility, and malleability.
• A combination of 2 metals is called an alloy.

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Free e- move rapidly in response to electric
fields, thus metals are excellent conductors of
electricity.
http//www.uwgb.edu/dutchs/EarthSC202Notes/mineral
s.htm
Free e- transmit kinetic energy rapidly, thus
metals are excellent conductors of heat.
Layers of metal atoms are difficult to pull apart
because of the movement of valence e-, so metals
are durable.
However, individual atoms are held loosely to
other atoms, so atoms slip easily past one
another, so metals are ductile.
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Covalent (a.k.a. covalent network) Atoms bonded in a covalent network Covalent bonds Very hard Very high MP Generally insoluble Variable conductivity C (diamond graphite) SiO2 (quartz) Ge, Si, SiC, BN
Diamond
Graphite
SiO2
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The Octet Rule

• Atoms tend to gain, lose, or share e- until they
  are surrounded by 8 e- in their outermost energy
  level (have filled s and p sub shells) and are
  thus energetically stable.
• Exceptions do occur (and will be discussed
  later.)
• Visualizing atomic orbitals

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Lewis symbols

• Valence e-
• e- in highest energy level and involved in
  bonding all elements within a group on P.T. have
  same of valence e-
• Lewis symbol (or electron-dot symbol)
• Shows a dot only for valence e- of an atom or
  ion.
• Place dots at top, bottom, right, and left sides
  and in pairs only when necessary (Hunds rule).
• Primarily used for representative elements only
  (Groups 1A 8A)
• Ex Draw the Lewis symbols of C and N.

Gilbert N. Lewis(1875 1946)
C
N
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• Transition metals typically form 1, 2, and 3
  ions.
• It is observed that transition metal atoms first
  lose both s e-, even though it is a higher
  energy subshell. Cr2, Cr3
• Most lose e- to end up with a filled or a
  half-filled subshell. Ex. Cu ion

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Lewis Structures

• Lewis structures are used to depict bonding pairs
  and lone pairs of electron in the molecule. 
• Step 1
• Total number of valence electrons in the system
  Sum the number of valence electrons on all the
  atoms . Add the total negative charge if you have
  an anion. Subtract the charge if you have a
  cation.
• Example CO32-

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• Step 2
• Number of electrons if each atom is to be happy
  Atoms in our example will need 8 e (octet rule)
  or 2 e ( hydrogen). So, for the ex.
•  Step 3
• Calculate number of bonds in the system Covalent
  bonds are made by sharing of e. You need 32 and
  you have 24. You are 8 e deficient. If you make 4
  bonds ( with 2 e per bond) , you will make up the
  deficiency. Therefore,
• of bonds ( e in step 2- e in step 1)/2
  (32-24)/2 4 bonds

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• Step 4
• Draw the structure The central atom is C (
  usually the atom with least electro negativity
  will be in the center). The oxygens surround it .
  Because there are four bonds and only three
  atoms, there will be one double bond.
• Step 5 
• Double check your answer by counting total number
  of electrons.
• Drawing Lewis Structures (youtube video)

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• Level 1 Practice for Lewis Structurs
• Draw the Lewis structures for the following using
  above steps. Show work! 
• A.Cl2 
• B. CH2Cl2
• C. NH3
• D. NaCl

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Level 2 practice on Lewis Structures

• SO42- HCN
• H2O2 CNS1-

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8.8 Exceptions to the Octet Rule

• Odd-electron moleculesEx NO or NO2 (involved
  in breaking down ozone in the upper atmosphere)
• Incomplete octet
• H2 He BeF2 BF3
• NH3 BF3 ? NH3BF3 (Lewis acid/base rxn)

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• Expanded octet occurs in molecules when the
  central atom is in or beyond the third period,
  because the empty 3d subshell is used in
  hybridization (Ch. 9)
• PCl5 SF6

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8.6 Formal Charge Movie on Formal Charge

• For each atom, the numerical difference between
  of valence e- in the isolated atom and of e-
  assigned to that atom in the Lewis structure.
• To calculate formal charge
• Assign unshared e- (usually in pairs) to the atom
  on which they are found.
• Assign one e- from each bonding pair to each atom
  in the bond. (Split the electrons in a bond.)
• Then, subtract the e- assigned from the original
  number of valence e-.

VALENCE e- in free atom NON-BONDING e-
½(BONDING e-) FC
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• Used to select most stable (and therefore most
  likely structure) when more than one structure
  are reasonable according to the rules.
• The most stable
• Has FC on all atoms closest to zero
• Has all negative FC on most EN atoms.
• FC does not represent real charges it is simply
  a useful tool for selecting the most stable Lewis
  structure.

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Examples Draw at least 2 Lewis structures for
each, then calculate the FC of each atom in each
structure.

• SCN1-
• N2O BF3

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8.7 Resonance Structures

• Equivalent Lewis structures that describe a
  molecule with more than one likely arrangement of
  e-
• Notation use double-headed arrow between all
  resonance structures.
• Ex O3
• Note one structure is not better than the
  others. In fact, all resonance structures are
  wrong, because none truly represent the e-
  structure of the molecule. The real e-
  structure is an average of all resonance
  structures.

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Bond Order

• An indication of bond strength and bond length
• Single bond 1 pair of e- shared
• Ex F2

Longest, weakest
F-F

• Double bond 2 pairs of e- shared
• Ex O2

• Triple bond 3 pairs of e- shared
• Ex N2

Shortest, strongest
N N
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Bond Order Resonance Structures

• To determine bond order with resonance
  structures
• Determine the bond order at one position in one
  resonance structure, and add it to the bond
  orders at the same bond position in all other
  resonance structures.
• Divide the sum by the number of resonance
  structures to find bond order.

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Examples draw the Lewis structure and determine
the bond S-O, C-C, and C-H bond orders

• SO3 C6H6

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8.9 Bond enthalpy

• Amount of energy required to break a particular
  bond between two elements in gaseous state. Given
  in kJ/mol. Remember, breaking a bond always
  requires energy!
• Bond enthalpy indicates the strength of a bond.
  
• Bond enthalpies can be used to figure out ?Hrxn .
• Ex CH4 (g) Cl2 (g) ? CH3Cl (g) HCl (g)
  DHrxn ?
• 1 C-H 1 Cl-Cl bond are broken (per mole)
• 1 C-Cl 1 H-Cl bond are formed (per mole)
• ?Hrxn ? (Hbonds broken) - ? (Hbonds formed)
• Note this is the opposite of Hess Law where
• ?Hrxn DHproducts Dhreactants
• Bond Enthalpy link

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• Ex CH4 (g) Cl2 (g) ? CH3Cl (g) HCl (g)
  DHrxn ?
• Bond Ave DH/mol Bond Ave DH/mol
• C-H 413 Cl-Cl 242
• H-Cl 431 C-Cl 328
• C-C 348 CC 614
• ?Hrxn ? (Hbonds broken) - ? (Hbonds formed)
• ?Hrxn (1(413) 1(242) 1(328) 1(431)
• ?Hrxn -104 kJ/mol
• ?Hrxn -99.8 kJ/mol (actual)
• Note 2 C-C ? 1 CC
• 2(348) 696 kJ ? 614 kJ

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Ex CH4(g) Cl2(g) ? CH3Cl(g) HCl(g) DHrxn?
CH3(g) H(g) 2 Cl(g)
Absorb E, break 1 C-H and 1 Cl-Cl bond
Release E, form 1 C-Cl and 1 H-Cl bond
H
CH4(g) Cl2(g)
CH3Cl (g) HCl (g)
DHrxn
?Hrxn ? (Hbonds broken) ? (- Hbonds
formed) ?Hrxn ? (Hbonds broken) - ? (Hbonds
formed)