Unit Five Chemical Bonding

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Unit FiveChemical Bonding
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I. There are many types of bondsand many ways to
classify them

• Intermolecular vs. Intramolecular bonds
• A. Intermolecular bonds are attractions between
  neighboring molecules
• Ex. The bond between 2 molecules of H2O
• Ex. Van der Waal or Hydrogen bonds

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Classifying continued

• B. Intramolecular bonds are within the molecules
  themselves
• They hold an individual molecule together
• Ex. Covalent, ionic, and metallic bonds

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Bonds are the forces that hold atoms together

• Intermolecular forces including
• dipole-dipole interactions
• hydrogen bonds
• are much weaker than
• Intramolecular forces like
• covalent bonds
• ionic bonds
• metallic bonds

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When Intermolecular forces change, only the state
of a substance is affected

• When a substance melts or boils intermolecular
  forces are broken
• Intramolecular bonds (i.e. covalent or ionic
  bonds) remain
• When a substance condenses (gas ? liquid)
  intermolecular forces are formed

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II. Energy and Bonds

• Forming bonds exothermic process
• Energy is released
• The compound formed has less potential energy
  (PE) than the starting substances
• The more stable the molecule formed, the lower
  the PE of the molecule
• Meaning a lot of energy was released

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Energy and Bonds cont

• Breaking bonds endothermic
• Requires energy
• The amount of energy needed to break the bond
  will tell you how strong the bond is

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III. Lewis Dot Structures

• Used to model the transferring/sharing of
  electrons during bond formation
• Also called the electron dot structure
• Composed of a positively charged kernel and small
  dots

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Pop QuizWhat is a kernel?

• A kernel is everything except the dots!
• It represents the nucleus plus all nonvalence
  electrons

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Pop QuizWhy are kernels positively charge?

• The protons outnumber the electrons because
  valence electrons are not included in the kernel.

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A. Lewis Dot diagrams

• Lets review
• Consist of kernel plus valence e-
• First electron _at_ 12 oclock
• 2nd _at_ 3, 3rd _at_ 6, 4th _at_ 9 oclock
• Maximum of valence e- ? 8
• Noble gases follow the octet rule

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B. Lewis diagrams of ions

• Differ from neutral atoms
• Remember, when atoms gain or lose e- they become
  ions
• To draw Lewis diagrams of ions
• Put the kernel in brackets
• Write the charge outside of the brackets
• Ex. Na Mg 2 Al 3

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Lewis diagrams of ions cont

• If the ion is positive, youre done!
• If the ion is negative, you must also include
  valence e-
• . . . . . .
• Ex. P -3 S -2 Cl -1
• . .
  . . . .
• Now, lets practice!!!
• Yippee!!!

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C. Lewis diagrams of compounds

• Lewis diagrams show how atoms combine
• When atoms form diatomic molecules we will see
  single, double, triple covalent bonds
• Single covalent bonds (share 2 e-)
• H2 ? HH single covalent bond
• A dash between two atoms is sometimes used to
  show single bonds ? H-H

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More Lewis diagrams of compounds

• Single covalent bonds continued
• Halogens (group 17 atoms) also share 2 e-and form
  single covalent bonds
• Ex.
  . . . .
• ? Cl Cl
• 
  . . . .
• Count the electrons does each atom have a
  complete octet?

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More Lewis diagrams of compounds

• Double covalent bonds (share 2 e- pairs)
• Ex. O O
• 
  . . . .
• Count the electrons does each atom have a
  complete octet?
• Triple covalent bonds (share 3 e- pairs)
• Ex. N N

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More Lewis diagrams of compounds

• The Octet Rule and Lewis diagrams
• When drawing diagrams of compounds remember the
  Octet Rule ? each atom must have 8 valence e- to
  be stable
• Exceptions ? H and He only need 2 valence e- to
  be stable

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More Lewis diagrams of compounds

• Now lets try drawing other compounds
• Consider CH3Cl and C2H2
• These compounds are more complex than diatomic
  molecules, however they follow the same Lewis
  diagram rules
• See Review Book page 81

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Steps for drawing Lewis diagrams of compounds

• Determine the total number of valence e-
• Arrange the atoms to show bonds between them
  hint the central atom usually appears once in
  the formula AND has the least electronegativity
  (remember the trend?) Hydrogen cannot be central!
• Use a dash to represent covalent bonds between
  atoms each dash ? 2 e-

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Steps for drawing Lewis diagrams of compounds
cont

• Count the number of e- represented, are there any
  remaining e-?
• Distribute the remaining e- so each atom has a
  complete octet
• Left over e- are placed as double or triple
  bonds between atoms that are not complete

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Drawing Lewis diagrams

• Draw the diagrams for CH3Cl and C2H2 in your
  notes
• More practice
• Try drawing CO2, HF, H2O2, CH3OH, H2O
• http//www.stolaf.edu/depts/chemistry/courses/tool
  kits/121/js/lewis/

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IV. Ionic Bonds

• Who atoms that lose/gain electrons and therefore
  become positive/negative ions
• What a chemical bond formed by the attraction
  between () and (-) ions ?electrostatic
  attraction
• How a valence electron from one atom is
  transferred to the valence shell of another atom

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More on Ionic bondsMetals vs. Nonmetals

• METALS
• Lose electrons
• Become positive ions
• Ionic radii decreases

• NONMETALS
• Gain electrons
• Become negative ions
• Ionic radii increases

Remember the trends theyre back!!!
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Ionic bonds cont

• Ionic bonds form based on electronegativity
  values
• What is electronegativity?
• Atoms tendency to attract bonded e-
• The greater the difference between ion
  electronegativity, the more likely they will form
  an ionic bond that is polar
• Polar ? unequal (remember for later on)

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Ionic bonds cont

• If you subtract electronegativity values and the
  difference is greater than 1.7, then the bond is
  ionic (see ref. table for values)
• Atoms bond to one another in fixed ratios, an
  important characteristic of compounds
• When atoms form ions, the result is a noble gas
  configuration ? as we saw in our Lewis diagrams

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Lewis Diagrams of Ionic Compounds

• Draw each ion separate
• Use a symbol (x, ) to represent gained/lost
  electrons
• Looks like this
• 
  . .
  .
  
• Na Cl ? Na Cl-
• .
  . .
• See Review Book page 87

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Ionic bonds cont

• Ionic bonds usually form between metals and
  nonmetals
• If a polyatomic ion is involved there are both
  ionic and covalent bonds
• Covalent bonds are within the polyatomic ion
• Ionic bonds are between the polyatomic ion and
  the other ion
• Hydrogen Metal can also form an ionic bond

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Properties of Ionic Bonds

• Ionic compounds dissolve easily in water and
  other polar solvents
• In solution, ionic compounds conduct electricity
  due to movement of charges (electrolytes)
• Ionic compounds tend to have high melting points
  (due to intermolecular forces)
• Mostly solids (hard) at room temperature

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V. Metallic Bonds

• What bonds between metals
• Why metals and metallic atoms have loosely held
  electrons that can be taken away fairly easily
• These electrons are more or less free to move
  from one atom to another.
• Chemists often describe metals as
• metal ions floating in a sea of mobile
• electrons around a positive nucleus

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Properties of Metallic Bonds

• Good conductors of heat and electricity
• Ductile
• Malleable
• Have luster (shiny)
• High melting point
• Solid (hard) at room temperature
• Alloys homogeneous mixture of metals
• (ex. Steel)

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Metallic bonding cont

• Metallic bonding results in the formation of
  alloys rather than compounds because metal atoms
  do not combine in fixed ratios.
• Only ionic and covalent bonding results in the
  formation of compounds

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VI. Covalent Bonds

• What the strong attraction that holds
  nonmetallic elements together
• Where they are associated with a great variety
  of materials.
• found within elements and compounds
• found within molecules
• also found within polyatomic ions
• Essentially it is found in any material in which
  nonmetallic atoms are bonded together

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Covalent Bonds cont

• Why both atoms are trying to attract
  electrons--the same electrons.
• ? Therefore, the electrons are shared tightly
  between the atoms.
• The force of attraction that each atom exerts on
  the shared electrons is what holds the two atoms
  together.

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Covalent Bonds cont

• When A covalent bond is formed when two nuclei
  share electrons to become a stable
  molecule/compound
• 2 Types of Covalent bonds
• small covalent and GIANT covalent
• A.Four types of small covalent bonds

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Small Covalent Bonds

• Multiple covalent bonds
• (weve already seen these!)
• The shared e- count for both atoms
• Double covalent share 2 pairs of e- 4
• Triple covalent share 3 pairs of e- 6
• Count all bonds to determine if the octet rule is
  satisfied or not

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Small Covalent Bonds cont

• 2. Nonpolar covalent bonds
• Form between atoms that have equal or very close
  electronegativity values
• Ex. Diatomic molecules or group 17 molecules (H,
  N, O, F, Cl, Br, I)
• Atoms have attraction to the shared e-
• Identical if flipped Ex. F-F ? F-F

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Small Covalent Bonds cont

• 3. Polar covalent bonds
• Atoms have different electronegativity values
• The e- are NOT shared equally
• The element/atom with the higher
  electronegativity value has a stronger attraction
  to the shared e-
• Not identical if flipped Ex. H-F ? F-H

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Small Covalent Bonds cont

• 4. Coordinate covalent
• A bond is formed when 2 e- come from a single
  atom
• Ex. H H2O ? H3O
• Ex. H NH3 ? NH4
• No different than any other covalent bond
• The 2 e- are shared

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Small Covalent Bonds cont

• B. Properties of small covalent bonds
• Soft
• Poor conductors of heat and electricity in all
  phases (solid, liquid, gas)
• Low melting and boiling points
• - Gases ? e- are least attracted
• - Liquids ? e- are more attracted

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Shapes of Small Covalent Bonds

• C. Molecular shapes 3 types
• Nonpolar nonpolar bonds, nonpolar molecule
• Diatomic molecules are nonpolar
• H2, N2, O2, Cl2, F2
• Shape is linear

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More shapes

• Polar bonds, nonpolar molecule
• Think tug of war
• The bonds are polar, but the overall shape is
  symmetrical
• There are equal attractions pulling in opposite
  directions
• Shapes include linear and tetrahedral

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More shapes

• Polar polar bonds, polar molecule
• If the central atom has a lone pair of e-, the
  shape is not symmetrical
• Shapes include bent (angular) and pyramidal
• Ex. H2O

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Steps for determining shape/polarity

• Start with a Lewis diagram for the molecule.
• If there are 2 atoms, they are either identical
  or not
• A. Identical ? linear nonpolar
• B. Not identical ? linear polar
• 2. If there are 3 atoms, are there any lone
  pairs of electrons?
• Yes ? bent polar
• No ? go to the next step

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Steps continued

• Are surrounding atoms identical?
• - Yes ? linear nonpolar
• - No ? linear polar
• If there are 4 or 5 atoms, it gets more
  complicated
• A. If there are 2 central atoms
• Are there lone pairs on the central atoms?
• Yes ? bent polar
• No ? Are the surrounding atoms identical?

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Steps continued

• - Yes ? linear nonpolar
• - No ? linear polar
• B. If there are not 2 central atoms, does the
  central atom have lone pairs?
• - Yes ? pyramidal polar
• - No ? Are the surrounding atoms identical?
• - Yes ? tetrahedral nonpolar
• - No ? tetrahedral polar

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Steps continued

• I think we need a flow chart
• for all this!

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D. Large Covalent Bonds

• Macromolecules and Network Solids
• What giant covalent structures or lattice
• formed when many atoms, usually non-metals, link
  together
• This produces a very strong 3-dimensional
  covalent bond ? network or lattice
• They are significantly different from the
  small/simple covalent molecules.

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Large Covalent Bonds cont

• Example Carbon has several macromolecules of
  allotropes huh?
• Allotropes different forms of the same element
  in the same physical state
• Allotropes of carbon are
• 1. Diamond
• 2. Graphite
• 3. Buckminster fullerene Bucky balls

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