Unit 11: Chemical Bonding

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Unit 11 Chemical Bonding
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Chemical Bonding Overview

• Chemical Bond
• attractive force is between the nucleus of one
  atom and the valence electrons of another atom.

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Review

• How do you find valence electrons?
• Hint there are two ways!
• Examples
• Mg ___
• O ___
• Ar ___
• Si ___
• Examples
• Mg
• O

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6
8
4
1s2 2s2 2p6 3s2
1s2 2s2 2p4
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Electron Dot Structures

• Depicts element symbol w/ valence e- shown as
  dots.

Na
Mg
Al
Si
Cl
Ar
O
N
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Purpose of Bonding

• Chemical bonding is driven by the octet rule
  atoms will lose, gain, or share electrons in
  order to achieve the electron configuration of
  the closest noble gas.
• Unfilled or partially filled valence orbitals are
  inherently unstable. (Unstable means possessing
  high potential energy.) Nature always strives for
  the lowest energy conditions

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Types of Chemical Bonds

• There are 3 types of chemical bonds
• Ionic Covalent Metallic
• Depends on the number of valence electrons of
  bonding atoms.
• Metallic bonds form between metal atoms only.
• Ionic bonds involve the transfer of electrons
  (lose/gain).
• Covalent bonds involve the sharing of electrons.
• Metallic bonds involve a mobile electron sea.

Flower Time
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Movie
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Ionic Bonds

• Occurs when ions of opposite charge (,-) attract
  each other.
• Metal ion Nonmetal ion
• Simplest attraction
• NaCl MgF2
• Polyatomic ions
• AlPO4 (NH4)2SO4

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Formation of Ionic Bond

• Cation- positive ion ()
• Forms when a metal atom loses e- to become
  stable.
• Anion- negative ion (-)
• Forms when a nonmetal atom gains e- to become
  stable
• An ionic bond is formed when e- are transferred
  between atoms and the resulting ions stick
  together.

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Examples

• Formation of NaCl
• Na Cl ? Na Cl - NaCl
• Formation of MgF2
• Mg F F ? Mg2 F - F - MgF2
• How would aluminum oxide form?

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Electron Configuration of Ions

• Cation example (metal)
• Ca atom 1s2 2s2 2p6 3s2 3p6 4s2
• Ca2 ion 1s2 2s2 2p6 3s2 3p6 Lost 2 electrons
  to obtain noble gas configuration (octet)

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Electron Configuration of Ions

• Anion example (nonmetal)
• N atom 1s2 2s2 2p3
• N3- ion 1s2 2s2 2p6
• Gained 3 electrons to obtain noble gas
  configuration (octet)

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Properties of Ionic Compounds
IONIC
e- transferred from metal to nonmetal
Bond Formation
Type of Structure
Crystal lattice
Physical State
Solid (hard and rigid)
Melting Point
high
Boiling Point
high
yes (solution or liquid)
Electrical Conductivity
Other Properties
brittle
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Electrolyte

• A substance that conducts electricity
• Because of ionic bonds ionic (charged) nature,
  ionic compounds conduct electricity in the molten
  or aqueous forms.

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Bonding Types

• Dog Analogy

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Bond Polarity

• Nonpolar
• Polar
• Ionic

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NON-POLAR COVALENT BONDS
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Polar Covalent Bonds Unevenly matched, but
willing to share.
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Ionic Bonds One Big Greedy Thief Dog!
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Bond Polarity

• Nonpolar Covalent Bond
• e- are shared equally
• symmetrical e- density
• usually identical atoms

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Bond Polarity

• Polar Covalent Bond
• e- are shared unequally
• asymmetrical e- density
• results in partial charges (dipole)

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Bond Polarity

• Most bonds are a blend of ionic and covalent
  characteristics.
• Difference in electronegativity determines bond
  type.

If DEN is Bond type is
lt 0.4 Nonpolar covalent
0.4 lt DEN lt 1.7 Polar covalent
gt 1.7 Ionic
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• A large electronegativity difference leads to an
  ionic bond.
• A small electronegativity difference leads to a
  polar covalent bond.
• No electronegativity difference between two atoms
  leads to a nonpolar covalent bond.

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• Cesium and sulfur
• Carbon and oxygen
• Fluorine and fluorine

If ?EN is Bond type is
lt 0.4 Nonpolar covalent
0.4 lt ? EN lt 1.7 Polar covalent
gt 1.7 Ionic
Ionic Polar
Nonpolar
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Covalent Bonds

• Occurs when 2 nonmetals share pairs of electrons
  to become stable. Molecular compounds are
  formed.
• Examples
• H2O CO2 C6H12O6 PCl5

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Covalent Bonds

• Covalent bonds can be
• single (1 shared pair)
• double (2 shared pairs)
• or triple (3 shared pairs)
• Bond strength triple gt double gt single
• Bond length single gt double gt triple

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Lewis Structures

• Creating Lewis Structures
• Lewis structures are depictions of molecules that
  show valence electrons as __dots________.
• Shared pairs of electrons (i.e. _Valence
  electrons_) are drawn between the atoms sharing
  them.
• Unshared or ___lone_______ pairs of electrons are
  represented by dots located on one atom only

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8
2
NEED
4
8
8
8
8
6
8
30
8
1
Available
2
4
5
6
7
3
Figure out through e config
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Creating Lewis Structures

• Follow this system
• Example H2O
• 1) Draw a skeleton of the molecule. It
  generally works to place the different atom in
  the center.
• H O H

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Creating Lewis Structures

• Find the needed electrons (N) for each atom and
  add them up. N will be 8 for most elements, with
  these exceptions

• N 12
• H 2
• O 8
• H 2

• H gets 2 valence e-
• Be gets 4 valence e-
• B gets 6 valence e-

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• 3) Find the available (valence) electrons (A)
  for each atom and then add them up.
• A
• special note when completing a Lewis structure
  for a polyatomic ion, you will need to correct A
  by adding the absolute value of the charge if
  negative, and subtracting the charge if positive.
  For example, for the ion PO43-, you would add 3
  to A. For the ion NH4, you would subtract 1
  from A. (You do the opposite of the charge.)

H 1 O 6 H 1 Total A 8
N 12 A 8
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• 4) Find the shared (S) electrons for the entire
  molecule by this formula S N A
• S

S 12 8 4
N 12 A 8 S 4
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• 5) The shared electrons are the bonding
  electrons. Place all of the shared electrons
  between the atoms.
• H O H
• 6) You must place all of the available (A)
  electrons in the picture. The shared electrons
  are part of the available. See how many of the
  available electrons still need to be placed, and
  put them in the picture as lone pairs (unshared
  pairs) so that every atom gets an octet (remember
  H only needs 2).
• H O H

N 12 A 8 S 4
N 12 A 8 S 4
4
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• CH4

H H C H H
N A S
8(2x4) 16
4(1x4) 8
0
8
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• CF4

F F C F F
N A S
8(4x8) 40
4(4x7) 32
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8
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Cl Cl P Cl

• PCl3

N A S
8(3x8) 32
5(3x7) 26
20
6
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• CO2

O C O
N A S
8(2x8) 24
4(2x6) 16
8
8
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Polyatomic Ions

• To find total of valence e- (A)
• Add 1e- for each negative charge.
• Subtract 1e- for each positive charge.
• Place brackets around the ion and label the
  charge.

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Polyatomic Ions

• ClO4-

N A S
8(4x8) 40
O O Cl O O
7(4x6) 31
1 32
24
8
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Resonance

• Resonance structures are structures that occur
  when it is possible to write 2 or more valid
  Lewis structures for the same molecule or ion.

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• Experimental data indicate that the 2 bonds in
  ozone are the same __length_____, BUT double
  bonds are shorter than single bonds! So the
  explanation for this is that the actual bonds are
  __combination_______ of those in the 2 resonance
  structures. The extra electron pair in ozone is
  delocalized over the two bonding regions. So
  each bond spends about half the time being single
  and half being double.

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Expanded Octets

• Some molecules do not follow the octet rule. The
  central atom has more than 8 electrons, called an
  expanded octet.
• In these cases, the standard method for
  determining the Lewis structure will fail.

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How to figure out a Lewis structure for an
expanded octet

• 1. Calculate the available (A) number of
  electrons.
• 2. Give the surrounding atoms an octet, and
  assume only single bonds to the central atom
• 3. Place any remaining electrons as lone pairs
  on the central atom so that A electrons are
  included.
• not always the case, but you should still get
  the right answer

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Examples
F F P F F F

• PF5

N A S
5(5x7) 40
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VSEPR Theory

• Valence Shell Electron Pair Repulsion Theory
• Electron pairs orient themselves in order to
  minimize repulsive forces.

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A. VSEPR Theory

• Types of e- Pairs
• Bonding pairs - form bonds
• Lone pairs - nonbonding e-

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A. VSEPR Theory

• Lone pairs reduce the bond angle between atoms.

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B. Determining Molecular Shape

• Draw the Lewis Diagram.
• Tally up e- pairs on central atom.
• double/triple bonds ONE pair
• Shape is determined by the of bonding pairs and
  lone pairs.

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Common Molecular Shapes
2 bond 0 lone
LINEAR 180
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Common Molecular Shapes
2 bond 1 lone
BENT lt120
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Common Molecular Shapes
2 bond 2 lone
BENT 109.5
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Common Molecular Shapes
2 bond 3 lone
Linear
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Common Molecular Shapes
3 bond 0 lone
TRIGONAL PLANAR 120
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Common Molecular Shapes
3 bond 1 lone
TRIGONAL PYRAMIDAL 107
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Common Molecular Shapes
3 bond 2 lone
T-shaped
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Common Molecular Shapes
4 bond 0 lone
TETRAHEDRAL 109.5
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Common Molecular Shapes
4 bond 1 lone
See-Saw
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Common Molecular Shapes
4 bond 2 lone
Square planar
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Common Molecular Shapes
5 bond 0 lone
Trigonal bipyramidal
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Common Molecular Shapes
5 bond 1 lone
Square pyramidal
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Common Molecular Shapes
6 bond 0 lone
Octahedral
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Examples

• PF3

3 bond 1 lone
PYRAMIDAL 107
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Examples

• CO2

2 total 2 bond 0 lone
LINEAR 180
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Examples

• Use VSEPR Theory to predict the shape of
• CO2 b) ClO3- c) NO3-
• d) SCl2 e) PCl3

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UNDERSTANDING MOLECULAR GEOMETRIES
PREDICTED BY VSEPR
H
H
C
H
H
FROM THE PERSPECTIVE OF
HYBRIDIZATION OF ATOMIC ORBITALS
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Electron Both wave and particle
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Electron Both wave and particle
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s, p, d Orbitals
p Orbitals
Px
Px
Pz
Pz
Py
Py
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Hybrid Orbitals

• __Hybrid_______ orbitals are formed when several
  atomic orbitals blend to form the same total
  number of equivalent orbitals. This process is
  known as hybridization.
• Hybridization explains the tetrahedral structure
  of methane. The s and px, py, and pz orbitals of
  the carbon atom blend to form four identical
  hybrid orbitals

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Number of regions of high electron density Orbitals that blend to make the hybrids Hybridization
2 s p sp
3 s p p sp2
4 s p p p sp3
5 s p p p d sp3d
6 s p p p d d sp3d2
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Now see if you can use what you have learned to
predict the hybridization of some other
compounds. Lets start with water.
Lewis Diagram
of electron clouds
four
sp3 hybridization
Hybridization
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Now try carbon dioxide
Lewis Diagram
of electron clouds
Two
sp hybridization
Hybridization
Remember double or triple bonds count as 1
electron cloud in VSEPR theory.
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Now try the sulfate ion (SO4-2)
Lewis Diagram
of electron clouds
Four
sp3 hybridization
Hybridization
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Now try the carbonate ion (CO3-2)
Lewis Diagram
of electron clouds
Three
sp2 hybridization
Hybridization
Remember double or triple bonds count as 1
electron cloud in VSEPR theory.
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Now try the sulfur hexaflouride (SF6)
Lewis Diagram
of electron clouds
Six
sp3d2 hybridization
Hybridization
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Sigma and Pi Bonds

• When two atomic orbitals combine along the axis
  connecting the 2 atomic nuclei (end-to-end
  overlap), a __sigma________ bond is formed. One
  sigma bond is formed in all bonds (single,
  double, triple).
• When two atomic orbitals combine with a
  side-to-side overlap, a ___pi________ bond is
  formed. A pi bond occurs in double and triple
  bonds only.

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Ethylene (Ethene) , C2H4
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Summary

• single bond
• composed of one _sigma___ bond
• double bond
• composed of one __sigma___ and one ____pi___ bond
• triple bond
• composed of one _sigma_____ and two ___pi____
  bonds

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Polarity of Bonds

• The bonding pairs of electrons in covalent bonds
  are located between the ____elements_______ of
  the atoms sharing the electrons.
• When the pull of each nucleus for the electrons
  is equally strong, the electrons are shared
  equally and are located on average halfway
  between the two nuclei. Recall that this type of
  bond is a ___non polar__________ covalent bond.
• When the pull of one nucleus for the electrons is
  stronger than the other, the electrons spend more
  time closer to the more electronegative nucleus.
  Recall that this type of bond is a
  __polar________ covalent bond.

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• The more electronegative atom acquires a partial
  ___negative___________ charge. The less
  electronegative atom therefore acquires a partial
  __positive___________ charge.
• These partial charges are indicated by the
  following symbols
• d - for the more electronegative element.
• d for the less electronegative element.

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Ex H2O
d
d
d -
The water molecule contains two dipole moments,
sometimes just called dipoles. The dipole
moments can be drawn in with arrow notation
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Polarity of Molecules

• Note that just because a molecule contains polar
  bonds, it may or may not be classifed as a polar
  molecule.
• The polarity of a molecule will depend on
• (1) existence of polar bonds
• (if none of the bonds are polar, the molecule is
  nonpolar)
• (2) shape of the molecule
• (3) orientation of the polar bonds

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• Note to help determine if the molecule will be
  polar, look for lines of symmetry.
• If the molecule is symmetrical
• then the dipoles will cancel and the molecule is
  nonpolar.
• If the dipoles do not cancel
• the molecule is polar.

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Examples

• HCl CO2
• CF4 NH3