Chemical Bonding

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Chapter 11

• Chemical Bonding

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Types of Chemical Bonds 11.1
Bond a force that holds groups of atoms of two
or more atoms together and makes them function as
a unit Bond Energy the amount of energy
required to break the bond
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Types of Bonds 4 TYPES

• Metallic

• Artists rendering of a metallic bond

• Cations packed in a sea of electronsMetals
• Metals consist of closely packed cations floating
  in a sea of electrons.
• All of the atoms are able to share the electrons.
• The electrons are not bound to individual atoms.

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Type 1 Metallic

• Properties of Metals
• Good conductors
• Ductile
• Malleable
• Electrons act as a lubricant, allowing cations to
  move past each other

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Metals have a Crystalline Structure

• Metals

• Example Body Centered Cubic (Chromium)

• Packed spheres of the same size and shape
• Body Centered Cubic
• Face Centered Cubic
• Hexagonal Close Packed

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More examples

• Face-Centered Cubic (gold)

• picture

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Last example

• Hexagonal Close-Packed (zinc)

• picture

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Type 2 IONIC

• IONIC

• picture

• Bond between closely packed, oppositely charged
  ions
• Bond between a metal and a nonmetal
• hard solid _at_ 22oC
• high mp temperatures
• nonconductors of electricity in solid phase
• good conductors in liquid phase or dissolved in
  water (aq)

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Covalent Bonding (2 types)

• Instead of gaining or losing electrons atoms can
  get stable by sharing electrons
• This is always between two non-metals.
• Two fluorine atoms, for example, can form a
  stable F2 molecule in which each atom has 8
  valence electrons by sharing a pair of electrons.
• In covalent bonds they can share
• more than two electrons

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Type 34 COVALENT

• COVALENT

• picture

• Electrons are shared
• Have low melting, boiling points
• Do not conduct electricity when melted or
  dissolved in water
• relatively soft solids as compared to ionic
  compounds at room temp

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Covalent bond subtype 1

• Non-polar Covalent

• picture

• When two of the same elements bond they are
  called diatomic molecules, some examples of this
  are Hydrogen H2, Oxygen O2 and Nitrogen N2.
• The atoms in these bonds would have the same
  electronegativities. This means that both atoms
  attract the shared electrons to that same extent.

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Covalent Bonds subtype 2

• POLAR COVALENT

• picture

• Unequal sharing of electrons

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Dipole Moment 11.3

• A molecule that has a center of positive charge
  and a center of negative charge
• Dipole often represented by an arrow
• Points towards negative charge center and its
  tail indicates the positive charge center

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Review 3 types of bonds thus far

• x

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Lewis Structures

• Section 11.6

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Lewis Dot Structures

• Show valence electrons
• Use group number to figure it out

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The Octet Rule

• The octet rule says that atoms tend to gain, lose
  or share electrons so they have eight electrons
  in their outer shell.
• There are some exceptions to the octet rule
  (imagine that)
• BF3
• BCl3
• PF5
• SF6

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Follow the interactive website!

• Ionic Bonding (this should be review)
• http//www.youtube.com/watch?vT40sM8-SXso
• Covalent Bonding
• http//www.wisc-online.com/objects/ViewObject.aspx
  ?IDGCH6404

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Drawing Lewis Structures

• Arrange the element symbols.
• Central atoms are generally those with the
  highest bonding capacity.
• Carbon atoms are always central atoms
• Hydrogen atoms are always peripheral atoms
• Add up the number of valence electrons from all
  atoms.
• For polyatomic ions, add one electron for each
  negative charge and subtract one for each
  positive charge.
• Draw a skeleton structure with atoms attached by
  single bonds.
• Complete the octets of peripheral atoms.
• Place extra electrons on the central atom.
• If the central atom doesnt have an octet, try
  forming multiple bonds by moving lone pairs.

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Simple Rules

• 1. Figure out number of electrons by counting
  the TOTAL valence electrons in whole compound
• 2. Place the central element in the middle and
  surround it with the other elements
• 3. Place single bonds between elements
• 4. Place lone pairs around each element until
  there are a total of eight (Hydrogen only wants
  2)
• 5. Count total electrons surrounding the compound
  (dont forget the bonds count as 2 electrons)
• If electrons from 1 and 5 dont match. Erase
  electrons and put in double bond and recount

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Single, Double and Triple Bonds

• With Covalent bonds the elements can share two or
  more electrons
• A Single Bond is when 2 electrons are shared
  they are represented by a single line in bond
  diagrams
• A Double bond is when 4 electrons are shared they
  are represented by two lines in bond diagrams
• A Triple bond is when 6 electrons are shared
  they are represented by three lines in bond
  diagrams

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Lewis Dot Structures
H2CO
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Isomers multiple correct structures for a
single compound

• (requires breaking bond to make new compound)
• CH2Cl2

H Cl C
Cl H
Cl Cl C
H H
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Electronegativity and Polarity

• Section 11.2

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Electronegativity Values

• The electronegativity values can be found in the
  periodic table
• The higher the value the higher the
  electronegativity
• The Pauling scale is used to measure
  electronegativity. It is a relative scale running
  from 0.7 to 4.0 (hydrogen  2.2).
• The units for electronegativity are Pauling units.

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Electronegativity

• The ability of an atom to attract electrons when
  bonded
• Nonmetals have high electronegativity
• Metals have low electronegativity
• Electronegativity increases across a period and
  decreases down a group. WHY???

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Electronegativity Chart
Why would the metals have low electronegativity
numbers? Why dont the noble gases have
electronegativity numbers?
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Nonpolar Covalent Bond

• When electrons are shared between 2 atoms, a
  covalent bond is formed.
• If the atoms are identical, e.g. Cl2, the
  electrons are shared equally (nonpolar)
• Cl 3.0 therefore the ?EN 3.0-3.0 0
• ?EN electronegativity Difference
• 0 nonpolar

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Polar Covalent Bond

• If the electrons are shared between 2 different
  atoms, e.g. HBr, the sharing is unequal
• The bonding electrons spend more time near the
  more electronegative atom
• H 2.1 and Br 2.8 THEREFORE 2.8-2.1 0.7
• 0.7 a polar covalent bond

H
Br
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Bond Type by Electronegativity Value

• Remember the higher the atoms electronegativity
  value, the closer the shared electrons tend to be
  to that atom when it forms a bond
• Therefore, the polarity of a bond depends on the
  difference between the electronegativity values
  of the atoms forming the bond
• The greater the difference, the more polar the
  bond.

Electronegativity Difference Type of Bond Formed
0.0 to 0.2 nonpolar covalent
0.21 to 1.7 polar covalent
2.0 ionic
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Electronegativity Differences

• Why is there a gap between 1.7 and 2.0????
• If the two atoms are nonmetals polar covalent
  bond
• If nonmetal metal ionic bond

0 to 0.2 Nonpolar covalent
0.21 to 1.7 Polar covalent
2.0 Ionic
Electronegativity Difference
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Sample Problems

• Choose the bond that will be more polar
• H-P or H-C
• O F or O I
• N O or S O
• N H or Si - H

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Sample Problems

• Choose the bond that will be more polar
• H-P or H-C
• O F or O I
• N O or S O
• N H or Si - H

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Polar Molecules (overall polarity of the molecule)

• Note Not all molecules with polar bonds are
  polar molecules
• The dipoles in symmetrical molecules cancels out
• ? The bond is polar but the molecule is nonpolar

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How to determine polar molecules

• There are two important factors
• 1. The polarity of the individual bonds in the
  molecule
• 2. The shape or geometry of the molecule.
• Steps to take
• Determine if a given individual bond is polar,
  Look at the difference between electronegativity
  of the atoms in the perioidc table. If the
  difference is
• 0.2 lt non polar
• 0.2 - greater polar

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• b) Determine the shape of molecule. For now I
  will give them to you. Later you will figure out
  the shape yourself.
• i) if all bonds are non-polar, then the whole
  molecule is non-polar regardless of its shape.
• ii) If there is symmetry in the molecule so that
  the polarity of the bonds cancels out, then the
  molecule is non-polar. (symmetry arround the
  central atom)
• iii) If there are polar bonds but there is no
  symmetry the overall molecule is polar.

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Which molecules are polar?
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Which molecules are polar?
For these two molecules, even though there are
polar bonds the overall molecule is nonpolar
because the molecule is symmetrical