Chemical Bonding

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Chemical Bonding

• Chapter 8 Sections 1-8

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• A chemical bond is a strong electrostatic force
  of attraction between atoms in a molecule or
  compound.
• Bonding between atoms occurs because it creates a
  more stable arrangement for the atoms.

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Lewis Symbols Dot Diagrams

• Convenient way to show the valence electrons

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Three types of bonding

• Metallic bonding results from the attraction
  between metal atoms and the surrounding sea of
  electrons
• Ionic bonding results from the electrical
  attraction between positive and negative ions.
• Covalent bonding results from the sharing of
  electron pairs between two atoms

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Ionic Bonding

• Many atoms transfer electrons and other atoms
  accept electrons, creating cations (positive
  metal ions) and anions (negative nonmetal ions).
• The resulting ions are attracted to each other by
  electrostatic force.

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Ionic Bonding

• The ions closely pack together in a crystal
  lattice.
• This arrangement maximizes the attractive forces
  among cations and anions while minimizing
  repulsive forces.

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• Because force is proportional to the charge on
  each ion, larger charges lead to stronger
  interactions.
• Because force is inversely proportional to the
  square of the distance between the centers of the
  ions, smaller ions lead to stronger interactions.

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Ionic bonding between Na and Cl
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Covalent bonding

• In many cases electrons do not completely
  transfer from one atom to another.
• The electrons between atoms are shared.

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Covalent bonding between H2

• Hydrogens electron configuration is 1s1
• Because both H atoms need 1 more electron to
  become isoelectronic with He, it is unlikely that
  either will give up an electron.

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Covalent bonding between H2

• ?

1s
?
1s
They share the two electrons. H H ?
H H
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Types of Covalent Bonds

• When electrons are shared equally the bond is
  called a NONPOLAR covalent bond. (i.e. H2)
• Sometimes the electrons between two atoms are NOT
  shared equally. The bond created is called a
  POLAR covalent bond.
• . . . .
• H Cl ? HCl
• . . . .

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Polar Covalent Bonding

• An example of this would be HCl.

HCl molecule Hydrogen atom
? 1s
Ne ?? ?? ?? ? 3s
3p
Chlorine atom
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How to classify bond types

• Electronegativity ability of an atom in a
  molecule to attract shared electrons to it
• Each element on the periodic table is assigned an
  electronegativity value (see page 353) that
  ranges from 0.7 to 4.0.
• The difference in the electronegativity
  determines the bonding type (ionic, polar
  covalent, or nonpolar covalent).

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Electronegativity Values
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If the electronegativity difference is

• 1.7 and higher ionic
• 0.3 to 1.7 polar covalent
• 0.0 to 0.3 nonpolar covalent

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What if I get an electronegativity difference
that is 0.3 or 1.7?

• These cut-off numbers are guidelines.
• It is a gradual change not stair-step.

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Ionic Character

• As the electronegativity difference increases,
  the ionic character increases as well!

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Practice Problems

• What type of bond will occur between iodine and
  the following elements cesium, iron, and sulfur?

Bonding between I and Electronegativity difference Bond Type
Cesium
Iron
Sulfur
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Determine the type of bond between the following
pairs.
Bonding between Electronegativity difference Bond type
Li Cl
S O
Ca Br
P H
Si Cl
S Br
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Other ways to determine bonding

• Electronegativity is not the only factor in
  determining bonding.
• Generally, bonds between a metal and nonmetal are
  ionic, and between two nonmetals the bonds are
  covalent.
• Examination of the properties of a compound is
  the best way to determine the type of bonding.

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Ionic Bonding

• Ionic compounds are formed to maximize stability.
  
• Nonmetal - will gain electrons to become
  isoelectronic with nearest noble gas called an
  anion
• Metal will lose electrons to become
  isoelectronic with noble gas called a cation

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Transition Metals

• Zinc
• Electron configuration is 1s22s22p63s23p64s23d10
• When it forms the 2 ion, it loses the 2 valence
  electrons in the 4s sublevel.
• Zn2 configuration is 1s22s22p63s23p63d10

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Practice Problems

• Write the electron configurations for the
  following ions.
• Fe2
• S2-
• Mg2
• Use electron configurations to explain why the
  most probable charge on the strontium ion is 2.

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Size of Ions
Does the size go up or down when losing an
electron to form a cation?

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Size of Ions
Forming a cation.
Li,152 pm
3e and 3p

• CATIONS are SMALLER than the atoms from which
  they come.
• The electron/proton attraction has gone UP and so
  size DECREASES.

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Size of Ions

• Does the size go up or down when gaining an
  electron to form an anion?

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Size of Ions
Forming an anion.

• ANIONS are LARGER than the atoms from which they
  come.
• The electron/proton attraction has gone DOWN and
  so size INCREASES.
• Trends in ion sizes are the same as atom sizes.

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Trends in Ion Sizes
Figure 8.13
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Which is Bigger?

• Cl or Cl- ?
• K or K ?
• Ca or Ca2 ?
• I- or Br- ?

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Which is Bigger?

• Cl or Cl- ? Cl-
• K or K ? K
• Ca or Ca2 ? Ca
• I- or Br- ? I-

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Lattice Energy Effects

• The change in energy when separated gaseous ions
  are packed together to form an ionic solid.
• M(g) X-(g) ? MX(s)
• Lattice energy is negative (exothermic) from the
  point of view of the system.

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Lattice Energy

• To determine which compound will have the highest
  lattice energy, take into consideration the
  following
• The size of the ions in the compound
• The smaller the size, the greater the lattice
  energy
• The charge of the ions in the compound
• The greater the charge, the greater the lattice
  energy

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Calculating ?Hf

• We can take advantage of the fact the energy is a
  state function and break the reaction into steps,
  the sum of which is the overall reaction.
• Lets do 41 Na(s) ½ Cl2 (g) ? NaCl(s)
• Given the following
• Lattice energy -786 kJ/mol
• Ionization energy for Na 495 kJ/mol
• Electron affinity for Cl -349 kJ/mol
• Bond energy of Cl2 239 kJ/mol
• Enthalpy sublimation for Na 109 kJ/mol

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Process

• Step 1 Sublimation of Na
• Na(s) ? Na(g) 109 kJ/mol
• Step 2 Ionization of Na
• Na (g) ? Na (g) e- 495 kJ/mol
• Step 3 Dissociation of Cl2
• ½ Cl2 (g) ? Cl(g) 119.5 kJ/mol
• Step 4 Formation of Cl- (Electron Affinity)
• Cl (g) e- ? Cl-(g) -349 kJ/mol
• Step 5 Formation of NaCl
• Na(g) Cl-(g) ? NaCl(s) -786 kJ/mol

Na(s) ½ Cl2 (g) ? NaCl(s) -411.5 kJ/mol
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Drawing Lewis Structures Valence Electron Review

• Valence electrons are in outermost level
• You can use periodic table or electron
  configuration to determine valence electrons
• Example Phosphorus
• Located in Group 15 or 5A
• Electron configuration is 1s22s22p63s23p3
• Contains 5 valence electrons
• Complete Exercise 1 on worksheet

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Drawing Lewis StructuresOctet Rule

• Most useful rule for creating Lewis structures
• Every atom usually has 8 valence electrons
• Exception hydrogen is good with 2 (like He)
• Lines are used to link atoms together (same as
  using 2 dots)

Same as
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Steps to Drawing Lewis Structures

1. Count valence electrons.
2. Connect atoms together with bonds. In molecules
   with a single atom of one element and several
   atoms of another element, the single atom is
   generally in the center with the other atoms
   attached to it.
3. Add electrons around outside of atoms to give
   each atom 8 electrons (or 2 in the case of
   hydrogen).
4. Count electrons used. This number must be the
   same as valence electrons.

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PCl3
5(37)26 e-
Complete Exercise 2.
Bonding Pairs
Lone Pairs (a.k.a. nonbonding electrons)
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Helpful Hints

• Carbon atoms form 4 bonds.
• Nitrogen atoms form 3 bonds.
• Oxygen atoms form 2 bonds.
• Hydrogen atoms form 1 bond.
• Fluorine atoms form 1 bond.
• Other halogens (Cl, Br, and I) frequently form 1
  bond (but not always).

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Determining the Central Atom

•  In a molecule, the atom that typically forms the
  greatest number of bonds is in the center, with
  other atoms attached to it.
• Example CH3Cl
• Carbon forms 4 bonds
• Hydrogen forms 1 bond
• Chlorine forms 1 bond
• SO CARBON IS IN THE MIDDLE WITH HYDROGEN AND
  CHLORINE AROUND IT! Dont forget electrons on
  chlorine to make 8!
• Complete exercise 3.

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Covalent Bonding

• Multiple Bonds
• It is possible for more than one pair of
  electrons to be shared between two atoms
  (multiple bonds)
• One shared pair of electrons single bond (e.g.
  H2)
• Two shared pairs of electrons double bond (e.g.
  O2)
• Three shared pairs of electrons triple bond
  (e.g. N2).
• Generally, bond distances shorten with multiple
  bonding.

Octet in each case
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Resonance

• Occurs when more than one valid Lewis structure
  can be written for a particular molecule.

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Odd Number of Electrons
NO
Number of valence electrons 11
Resonance occurs when more than one valid Lewis
structure can be written for a particular
molecule (i.e. rearrange electrons)
NO2
Number of valence electrons 17
Molecules and atoms which are neutral (contain no
formal charge) and with an unpaired electron are
called Radicals
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Beyond the Octet

• Elements in the 3rd period or higher can have
  more than an octet if needed.
• Atoms of these elements have valence d orbitals,
  which allow them to accommodate more than eight
  electrons.

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More than an Octet
Elements from the 3rd period and beyond, have ns,
np and unfilled nd orbitals which can be used in
bonding
P (Ne) 3s2 3p3 3d0 Number of valence electrons
5 (5 x 7) 40
PCl5
S (Ne) 3s2 3p4 3d0 Number of valence electrons
6 (4 x 7) 34
SF4
The Larger the central atom, the more atoms you
can bond to it usually small atoms such as F,
Cl and O allow central atoms such as P and S to
expand their valency.
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Formal Charge

• Difference between the of valence electrons in
  the free atom and the of electrons assigned to
  that atom in the Lewis structure.
• FC formal charge G.N. Group Number
• BE bonding electrons LPE lone pair
  electrons
• If Step 4 leads to a positive formal charge on an
  inner atom beyond the second row, shift electrons
  to make double or triple bonds to minimize formal
  charge, even if this gives an inner atom with
  more than an octet of electrons.

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Formal Charge

• Not as good Better

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Molecular Shapes

• Lewis structures give atomic connectivity they
  tell us which atoms are physically connected
  together. They do not tell us the shape.
• The shape of a molecule is determined by its bond
  angles.
• Consider CCl4 experimentally we find all Cl-C-Cl
  bond angles are 109.5?.
• Therefore, the molecule cannot be planar.
• All Cl atoms are located at the vertices of a
  tetrahedron with the C at its center.

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Molecular Shape of CCl4
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VSEPR Theory

• In order to predict molecular shape, we assume
  the valence electrons repel each other.
  Therefore, the molecule adopts whichever 3D
  geometry minimized this repulsion.
• We call this process Valence Shell Electron Pair
  Repulsion (VSEPR) theory.

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Why is VSEPR Theory Important?

• Gives a specific shape due to the number of
  bonded and non-bonded electron pairs in a
  molecule
• Tells us the actual 3-D structure of a molecule
• In bonding, electron pairs want to be as far away
  from each other as possible.

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VSEPR and Resulting Geometries
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How does VSEPR THEORY work?

• We can use VSEPR theory using 4 steps.
• Draw the Lewis Structure for the molecule.
• Example SiF4

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How does VSEPR THEORY work?

• We can use VSEPR theory using 4 steps
• Draw the Lewis Structure for the molecule.
• Tally the number of bonding pairs and lone
  (non-bonding) pairs on the center atom.

Bonding pairs 4 Lone pairs on central atom 0
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How does VSEPR THEORY work?

• We can use VSEPR theory using 4 steps
• Draw the Lewis Structure for the molecule
• Tally the number of bonding pairs and lone pairs
  on the center atom.
• Arrange the rest of the atoms so that they are as
  far away from each other as possible.

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How does VSEPR THEORY work?

• We can use VSEPR theory using 4 steps
• Draw the Lewis Structure for the molecule
• Tally the number of bonding pairs and lone pairs
  on the center atom.
• Arrange the rest of the atoms so that they are as
  far away from each other as possible
• Give the type of geometry the molecule has

Tetrahedral
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Another Example
To determine the electron pair geometry 1) draw
the Lewis structure 2) count the total number
of electron pairs around the central
atom. 3) arrange the electron pairs in one of
the geometries to minimize e--e-
repulsion. 4) multiple bonds count as one
bonding pair for VSEPR
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The VSEPR Model
Predicting Molecular Geometries
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The VSEPR Model
Predicting Molecular Geometries
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The VSEPR Model
Difference between geometry and
shape Geometry We determine the geometry only
looking at electrons. All the atoms that obey the
octet rule have the same tetrahedral-like
geometry. Shape We name the shape by the
positions of atoms. We ignore lone pairs in the
shape.
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The VSEPR Model
Predicting Shape
Shape
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The VSEPR Model
Predicting Shape
Shape
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The VSEPR Model
The Effect of Nonbonding Electrons and Multiple
Bonds on Bond Angles By experiment, the H-X-H
bond angle decreases on moving from C to N to
O Since electrons in a bond are attracted by
two nuclei, they do not repel as much as lone
pairs. Therefore, the bond angle decreases as the
number of lone pairs increase.
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The VSEPR Model
The Effect of Nonbonding Electrons and Multiple
Bonds on Bond Angles Similarly, electrons in
multiple bonds repel more than electrons in
single bonds.
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The VSEPR Model
Molecules with Expanded Valence Shells Atoms that
have expanded octets have AB5 (trigonal
bipyramidal) or AB6 (octahedral) electron pair
geometries. Examples PF5 trigonal
bipyramidal SCl6 octahedral
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The VSEPR Model
Molecules with Expanded Valence Shells
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The VSEPR Model
Molecules with Expanded Valence Shells
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The VSEPR Model
Molecules with More than One Central Atom In
acetic acid, CH3COOH, there are three central
atoms. We assign the geometry about each central
atom separately.
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Hybrid Orbitals

• In bonding, s and p orbitals are used in bonding.
  It is easy to tell which ones are used by
  looking at our molecule.
• For example, CH4. Looking again at the Lewis
  structure, we see that there are 4 bonds. We
  call this sp3 hybridized.

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Hybrid Orbitals

• Regions of electron density-EACH BOND AND LONE
  PAIR OF ELECTRONS ON THE CENTRAL ATOM IS KNOWN AS
  A REGION OF ELECTRON DENSITY.
• 2 regions of electron density-sp hybridized
• 3 regions of electron density-sp2 hybridized
• 4 regions of electron density-sp3 hybridized

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Hybridization
sp Hybrid Orbitals The two lobes of an sp hybrid
orbital are 180? apart.
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Hybrid Orbitals
sp2 Hybrid Orbitals Important when we mix n
atomic orbitals we must get n hybrid
orbitals. sp2 hybrid orbitals are formed with one
s and two p orbitals. (Therefore, there is one
unhybridized p orbital remaining.) The large
lobes of sp2 hybrids lie in a trigonal plane. All
molecules with trigonal planar electron pair
geometries have sp2 orbitals on the central atom.
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Hybridization
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Hybridization
sp3 Hybrid Orbitals sp3 Hybrid orbitals are
formed from one s and three p orbitals.
Therefore, there are four large lobes. Each lobe
points towards the vertex of a tetrahedron. The
angle between the large lobes is 109.5? All
molecules with tetrahedral electron pair
geometries are sp3 hybridized.
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Hybridization
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Hybrid Orbitals
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Hybrid Orbitals

• Summary
• To assign hybridization
• Draw a Lewis structure.
• Assign the geometry using VSEPR theory.
• Use the geometry to determine the hybridization.
• Name the shape by the positions of the atoms.

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Hybridization and Multiple Bonds

• Multiple bonds overlap differently and are
  called s-bonds and p-bonds
• All single bonds are s
• Double bonds contain 1 s and 1 p bond
• Triple bonds contain 1 s and 2 p bonds

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Bond Energy
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Covalent Bonding Orbital Overlap

• As two nuclei approach each other their atomic
  orbitals overlap.
• As the amount of overlap increases, the energy of
  the interaction decreases.
• At some distance the minimum energy is reached.
• The minimum energy corresponds to the bonding
  distance (or bond length).

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Covalent Bonding Orbital Overlap

• As the two atoms get closer, their nuclei begin
  to repel and the energy increases.
• At the bonding distance, the attractive forces
  between nuclei and electrons just balance the
  repulsive forces (nucleus-nucleus,
  electron-electron).

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Bond Energies

• Bond breaking requires energy (endothermic).
• Bond formation releases energy (exothermic).
• ?H ?D(bonds broken) ? ?D(bonds formed)

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Bond Energies