AP Chemistry Unit 3

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AP ChemistryUnit 3

• Bonding
• Chapters 8 and 9

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Test
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Test

• Lets Talk
• Makeup Points

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Chemistography Presentations
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Bonding
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Ionic Bonding

• One atom steals the electrons of another
• Attraction between anions and cations
• Large difference in electronegativies
• ?EN gt 1.7
• Usually s block and nonmetal

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Ionic Bonding

• Na has EN of 0.9
• Cl has EN of 3.0
• ?EN of 2.1
• ?EN ENanion Encation
• Use table on page 353

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Covalent Bonding

• Sharing electrons
• Small difference in EN
• ?EN lt 1.2

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Covalent Bonding

• Cl2
• ?EN 0
• HCl
• H has EN of 2.1
• Cl has EN of 3.0
• ?EN of 0.9

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Polar Covalent Bonding

• Unequal sharing
• Moderate difference between atoms
• 1.2 lt ?EN lt 1.7

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Polar Covalent Bonding

• LiI
• Li has EN of 1.0
• I has EN of 2.5
• ?EN of 1.5

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Metallic Bonding

• Lots of cations (metals that have lost electrons)
  all have electrons flying around between them
• These arrangements make metals malleable,
  ductile, and good conductors
• Malleable ability to be hammered into sheets
• Ductile ability to be drawn into wires

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Bonding Visuals
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Bonding Visuals
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Bonding Visuals
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Practice

• What type of bonds are these?
• NaAt
• FrBr
• MgCl2
• AgBr2
• HgO
• H2O
• YN
• C2H2

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Bond Polarities
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Bond Polarities

• All covalent bonds have some small degree of
  polarity
• Except homonuclear diatomics
• A molecule with distinct positive and negative
  regions are dipolar, and have a dipole moment

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Bond Polarities

• Sum of dipoles gives the overall dipole of the
  compound
• Draw the dipole of each bond
• Equal dipoles in opposite directions cancel each
  other out

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Showing Dipole Moments

• Plus sign goes on the d atom with an arrow
  extending out and pointing to the d- atom

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Intermolecular Forces
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Intermolecular Forces

• Forces between compounds, not within
• Forces within a compound are intramolecular

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Dipole-Dipole

• Strongest force between polar molecules
• Uneven distribution of charges
• Polar covalent bonding

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Hydrogen Bonding

• A specific type of dipole-dipole force
• Only occurs in H-O, H-F, and H-N bonds
• Hydrogens are attracted to the lone pairs of
  another molecule

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London Dispersion Force

• Attraction due to momentary dipole created by the
  natural movements of electrons
• Very weak, occurs in all atoms and compounds

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Molecular Geometry
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VSEPR Theory

• Valence Shell Electron-Pair Repulsion Theory
• The repulsion of the electrons to each other
  makes the bonds between atoms desirable to be as
  far apart as possible.

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VSEPR Theory

• Lone pairs need to be accounted for as well
• Geometry can be determined based on the number of
  atoms and lone pairs bonded to the central atom
  (stearic ) and the number of lone pairs
• Larger molecules have multiple central atoms

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Example

• Draw the Lewis Structure for water (H2O)
• According to VSEPR Theory how do we get the
  electrons farthest apart?
• Hint Think 3-Dimensionally
• 4 bonds, 2 lone pairs
• Bent Structure

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Practice

• NH3
• CO2
• NaAt
• PCl3
• HCP
• SF6
• C2H5OH
• CH4

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Polyatomic Polarity

• Determined by 3D structure
• Write dipoles of each bond as an arrow pointing
  towards the negative region
• The sum of the charges equal the overall
  molecular polarity

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Examples

• Water
• CO2
• CH4
• NH3

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Bond Hybridization
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Hybridization

• Geometry determined by hybridization of the bond
• A hybrid bond is a combination of orbitals to
  create a new orbital

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Example

• Draw methanes (CH4) orbital notation
• 2s and 2p merge to form a 2sp3
• An sp3 sublevel has 4 orbitals, as many as there
  were before, just at a single energy

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Hybridization
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Valence Bond Theory
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Remember Lewis Structures?

• Showing covalent bonds using electron dot
  structures simplified to show bonds
• 1 flaw Assumes shared electrons are between the
  atoms
• Localized electron model
• Are they?

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Valence Bond Theory

• Allows bond angles to be calculated
• Takes in to account the fact that electrons are
  not localized
• Thanks to Heisenberg

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Sigma and Pi Bonds
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Sigma and Pi Bonds

• When orbitals overlap they make 2 special types
  of bonds
• Sigma bonds (s bonds) are formed along a line
  between the nuclei
• Pi bonds (p bonds) are formed by overlapping
  side-by-side p orbitals

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s Bonding

• Can be formed by overlapping orbitals along the
  axis between the 2 nuclei overlapping
• s, p, or a hybrid

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p Bonding

• An overlap of either of the p orbitals
  perpendicular to the axis between the two atoms
• ONLY p orbitals

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Sigma and Pi Bonding
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Applying s and p Bonding

• A single bond is a s bond
• A double bond is a s bond and a p bond
• A triple bond is a s bond and 2 p bonds
• Only unhybridized bonds form s or p bonds
• This portion of Valence Bond Theory explains why
  we cant have more than a triple bond, or 6
  shared electrons

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Connecting it All
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Determining Bonding

• The formula can be translated in to a molecular
  structure, showing single, double, and triple
  bonds
• The type of bond is translated s and p bonds for
  each atom
• In addition, the orbital notation can tell the
  hybridization and s and p bonds

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Examples

• Draw the structure, then write the orbital
  notations, determine hybridization, and bonding
• CH4
• CO2
• C2H2

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Practice

• Draw the structure, then write the orbital
  notations, determine hybridization, and bonding
• N2
• H2O
• HCl
• HCN

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Double Bonds
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Double Bonds

• Double bonds have some special properties
• Double bonds usually only form between period 2
  nonmetals
• Double bonds crucial to all living organisms
• Silicon-based life cannot be as complex as
  carbon-based, even though silicon acts much the
  same as carbon

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Double Bonds

• They have a bond strength greater than a single
  bond, but less than 2 single bonds
• This is in part because p bonds are weaker than s
  bonds since they are only a side-by-side overlap
  on parallel axes rather than on the same axis

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Double Bonds

• Looking at molecules as 3D structures how can
  atoms move?
• In single bonds
• In double bonds

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Double Bonds

• Atoms cannot rotate around double bonds
• This is because p bonds dont allow rotations
• If groups larger than hydrogen are on the same
  side of a double bond it is referred to as cis-,
  if they are on opposite sides of a double bond
  they are called trans-

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Cis- and Trans-
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Practice

• Draw the cis- and trans- versions of each
• CH3CHCHOH
• C3H4Br2

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Molecular Orbital Theory
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Molecular Orbital Theory

• Based on quantum mechanics and Valence Bond
  Theory
• Lewis structures show lone pairs
• Not quite true

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Example

• Draw the Lewis structure of oxygen
• Are there lone pairs?

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Oxygen

• Oxygen is paramagnetic, which only happens with
  unpaired electrons
• Paramagnetism is the ability to be affected by a
  magnetic field, but do not hold their magnetic
  presence in the absence of an applied field
• Molecular Orbital Theory explains this

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Molecular Orbital Theory

• Developed in the 1920s by Robert Woodward and
  Roald Hoffman
• Also called MO Theory
• Describes the orbitals throughout a molecule and
  how they bond, draws on Valence Bond Theory

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MO Theory

• Electrons belong to the molecule, and shared not
  between any 2 atoms, but all of them attached in
  the molecule
• There arent really atom-specific orbitals, but
  more of molecular orbitals
• Electrons are delocalized, spread out over the
  whole molecule, not in bonds

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MO Theory

• MOs with a lower energy than the separate atomic
  orbitals are called bonding, MOs with higher
  energy are called antibonding
• The molecular orbital energy level diagram
  combines atomic energy levels and shows the
  created bonding and antibonding energy levels

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MO Diagram
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MO Diagram
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What does it mean?

• The lowest energy is s, which is located between
  the nuclei
• There is also a higher antibonding orbital
  formed, s
• The total number of orbitals is conserved

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Bonds

• For s orbital interactions only 1 bonding and 1
  antibonding orbital are created
• s and s

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Bonds

• Up to 3 bonding with p orbitals
• 1 s and 2 p
• Also 3 antibonding possible
• 1 s and 2 p
• If 2 atoms with p orbitals interact there are 6
  orbitals being combined

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Practice

• H2
• N2
• CH bond
• OH bond

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Practice

• CH3OH
• C6H6 (benzene)

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Special Cases of Lewis Structures
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Lewis Structures

• Remember what they are?
• Follow the Octet Rule - normally

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Example
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Practice

• Draw these Lewis Structures
• O2
• CH4
• C2H5OH
• HCl
• N2

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Exceptions to the Octet Rule

• Hydrogen and helium only need 2 valence electrons
• Boron and beryllium like 6
• This is called electron deficient
• Electron deficient elements are extremely
  reactive
• The more lone pairs another compound has, the
  more violent the reaction

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Practice

• Draw boron trichloride
• What is the geometry, steric ? What type of bond
  is it?
• Write the reaction showing the formation of boron
  trichloride. Was it energetic?

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Bond Energies
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Bond Energies

• Also called bond enthalpy (well do enthalpy
  second semester)
• The stronger the bond the more energy it has and
  the more is needed to break it

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Bond Order

• Bond order is just the amount of bonds that link
  a pair of atoms
• Bond order ½ (number of bonding electrons
  number of antibonding electrons)
• When determining bond order must show s and p MOs

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Example

• H2

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Practice

• N2
• O2
• F2

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Test Review

• Know 4 types of bonds
• Intermolecular forces
• VSEPR Theory and hybridization
• Sigma and pi bonds
• MO Theory (minimal, this is just getting you
  ready for college)

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End of Unit

• Questions?