Chemical Bonding Chapter 6
Dec 28, 2015
Chemical Bonding
Chapter 6
Types of Chemical Bonds
Chemical Bond: mutual electrical attraction b/ the nuclei and valence e- of different atoms
Atoms make bonds b/c they become more stable
Types of Chemical Bonds
Ionic Bonding: Results from the electrical attraction
between cations and anions Atoms completely give up e-s to other
atoms
Types of Chemical Bonds
Covalent Bonding: Results from sharing of e- pairs
between two atoms Shared e- are “owned” equally by two
atoms
Determining the Type of Bond Bonds fall somewhere b/ purely Ionic
and purely Covalent Depends on electronegativity – how
much is the atom pulling the e-s Calculate the difference in
electronegativities
Determining the Type of Bond
Types: Ionic: electronegativity dif: 1.7 – 3.2 Polar-Covalent: dif: 0.4 – 1.6 Nonpolar-Covalent: dif: 0 – 0.3
Determining the Type of Bond
Ionic: One atom is so electronegative it strips
the other atom of electrons making cation and anion
NaCl
Determining the Type of Bond
Polar-Covalent: Bond where atoms have unequal
attraction for shared e-s More electronegative atom has
stronger attraction for e-s HCl
Determining the Type of Bond
Nonpolar-Covalent: Bond where atoms have equal sharing
of e-s, balanced distribution of charge Bonds b/ two atoms of the same
element, F2
C-H
Determining the Type of Bond
Determine the type of bond b/ these elements and sulfur, which is more electronegative element?
H, Cs, Cl
Covalent Bonding
Covalent Bonding
Molecule: neutral group of atoms that are held together by covalent bonds.
Molecular = “covalent”
Covalent Bonding
Covalent Bonding
Chemical formula: Ionic or Covalent NaCl, MgCl2, H2O
Molecular formula: only for molecules (covalent)
Covalent Bonding
Forming Covalent Bonds
Share e-s to get noble gas configuration
Octet Rule
Def: atoms gain, lose or share e-s to have octet (8) of electrons in outer energy level
H – exception, only needs 2 e-s Ex: F2
Exceptions to Octet
B – happy with 6 e-s in outer level Other elements can have more than
8 valence e-s – PF5, SF6 – d orbitals invovled in bonding
Electron-Dot Notation
Use element symbol and dots to indicate valence e-s.
Period 2:
1A 2A 3A 4A 5A 6A 7A8A
Li Be B C N O FNe
Lewis Structures
Def: Use electron-dot notation for molecules
H2
F2 – shared pair with dash (-) lone pair – unshared pair
Lewis Structures
Single bond: covalent bond where one pair of e- being shared b/ two atoms
H-H and HCl :
Lewis Structures
Draw Lewis Structure:1. Start with e-dot diagram of each2. Atom with most bonding sites in
middle3. Circle unpaired e-s to make bonds4. Replace circles with dashes
NH3
H2S
Multiple Covalent Bonds
C, N, and O can share more than 1 pair of e-s
Double bond: two pairs, 4 e-s, being shared
C2H4
Triple Bond: three pairs, 6 e-s, being shared
N2
Bond Lengths Bond length: average distance b/
two bonded atoms Forming bonds – atoms release energy Same amount of energy needed to
break bond bond energy (kJ/mol) Lengths of multiple bonds?
More bonds – shorter – more energy to break
p.187
Resonance Structures
Molecules can’t be shown with one Lewis structure
Ex: O3
Lewis Structure
CH2O
CH3Br
C2HCl
SO3
Recap Ch. 6
Bonds: Ionic and Covalent Ionic, Polar-Covalent, and Nonpolar-
Covalent Drawing Lewis Structures
C2H2
Type of bonds?
Ionic Bonding
Ionic Compounds
Def: (+) and (-) ions that combine so charges balance out
Crystalline solids Formula Unit: simplest unit of
ionic compound where charges are balanced NaCl: Na+ Cl-
Video (68)
Forming Ionic Compounds
NaCl – use electron dot diagrams Compound with Ca and F:
Characteristics of Ionic Bonds
Ions in crystal lattice are more stable – lower potential energy
Lattice energy: energy released when 1 mole of gaseous ions form a lattice
More negative energy = more energy released = lower potential energy = more stable = stronger bonds
Ionic vs. Covalent
Ionic stronger bonds b/ formula units than b/
molecules in covalent compounds HIGHER melting and boiling points hard but brittle conduct electricity in molten or
dissolved state
Ionic vs. Covalent
Covalent Weak bonds b/ molecules Most compounds are gases at room
temp. LOW boiling and melting points
Polyatomic Ions
Def: charged group of covalently bonded atoms Result from excess or lack of electrons in
bonding
N HH
HH
+
N OO
O
Ammonium ion Nitrate ion
S OO
O
Sulfate ion
2
O
Metallic Bonding
Excellent conductors in solid state – due to highly mobile valence electrons
Filled outermost sublevel is s Vacant p and d orbitals overlap -
valence e-s are delocalized, do not belong to any one atom but move freely
Metallic Properties
High electrical and thermal conductivity
Malleability – hammered into sheets Ductility – drawn into wires
Molecular Geometry
3-D arrangement of molecules
VSEPR Theory
Valence-shell, electron-pair repulsion
Def: repulsion b/ valence e- pairs around atom causes them to be as far apart as possible
Shapes
NO lone pairs on CENTRAL atom
Symmetrical Linear Trigonal-Planar Tetrahedral Trigonal-
bipyramidal Octahedral
WITH lone pairs on CENTRAL atom
Non-symmetrical Trigonal-pyramidal Bent (angular)
Shapes – NO lone pairs on central atom
1. Linear (AB2):
• A – central atom B-bonded atoms- 3 atom molecules CO2
- 2 atom molecules, O2, HCl, etc.
- bond angles: 180o
Shapes – NO lone pairs on central atom
2. Trigonal Planar (AB3):
- BCl3- bond angles: 120o
Shapes – NO lone pairs on central atom
3. Tetrahedral (AB4):
- CCl4- bond angles: 109.5o
Shapes – NO lone pairs on central atom
4. Trigonal-bipyramidal (AB5):
- PCl5- bond angles: 120o and 90o
Shapes – NO lone pairs on central atom
5. Octahedral (AB6):
- SF6
- bond angles: 90o
Shapes – WITH lone pairs on central atom
6. Trigonal-Pyramidal (AB3E):
• A – central atom B – bonded atoms E – lone pair- NH3
- triangular sides - bond angles: 107o
Shapes – WITH lone pairs on central atom
7. Bent or Angular (AB2E2):
- H2O
- bond angles: 105o
Molecular Polarity
Polarity of each bond Molecular polarity
Molecular Polarity
1. Has ALL bonds NONPOLAR nonpolar molecule
2. Has bonds nonpolar AND polar polar molecule
3. Has ALL bonds POLAR depends on shape
Symmetrical shape (linear - octahedral) NONPOLAR
Non-symmetrical shape (bent & trigonal pyramidal) POLAR
Molecular Polarity Examples
CCl4 PH3
CBr3H
Intermolecular Forces
Intermolecular Forces
“between molecule” forces Generally weaker than bonds b/
atoms Boiling point – good to measure
intermolecular forces
Dipole-Dipole Forces
Dipole- equal but opposite charges separated by a short distance
Video 124H Cl
Dipole-Dipole Forces
Induced Dipole: polar molecule makes a dipole on a nonpolar molecule
Ex: O2 dissolved in H2O Weaker than regular dipole forces
Hydrogen Bonding
Type of dipole-dipole force Def: H-atom bonded to highly e-neg
atom is attracted to lone pair of the e-neg atom in nearby molecule
Ex: HF, H2O, NH3
Hydrogen Bonding
London Dispersion Forces
Def: constant motion of e-s and creation instantaneous dipoles
Video 133