Chapter 1 Chemical Bonding and Chemical Structure.
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Transcript of Chapter 1 Chemical Bonding and Chemical Structure.
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Chapter 1Chemical Bonding and Chemical Structure
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Organic chemistry
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• The branch of chemistry that deals with carbon based compounds– Organic compounds may contain any number of
other elements, including hydrogen, nitrogen, oxygen, halogens, phosphorus, silicon, and sulfur
Methane Sucrose Morphine
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History• Vitalism: Only biological
systems (e.g., plants, animals) could produce organic compounds
• Wohler’s synthesis of urea (1828), began to undermine vitalism
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Why Study Organic Chemistry?
• Organic chemistry lies at the heart of the modern chemical industry
• Central to medicine and pharmacy• Interface of physical and biological sciences• Everyday applications: Plastics, textiles,
communications, transportation, food, clothing, cosmetics, etc.
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Review of Chemical Bonding
• Valence Electrons: Outermost electrons• s and p electrons for main group elements• Responsible for chemical properties of
atoms• Participate in chemical reactions
Core Electrons Valence Electron
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Octet Rule
• Octet Rule: the tendency for atoms to seek 8 electrons in their outer shells– Natural electron configuration of the Noble Gases– Done by gaining, losing, or sharing electrons– Increases stability– H and He seek a “Duet”
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Ionic Bonding• Ions: atoms that have a charge due to gain or loss of
electrons– Anion: (-) charged atom– Cation: (+) charged atom
• Ionic Bond: a bond formed through the transfer of one or more electrons from one atom or group of atoms to another atom or group of atoms
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Formula Unit
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• Ionic bonds are omni-directional• Can dissociate into free ions
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Covalent Compounds• Covalent Compounds: compounds composed of atoms
bonded to each other through the sharing of electrons• Electrons NOT transferred• No + or – charges on atoms• Non-metal + Non-metal• Also called “molecules”• Examples:– H2O– CO2
– Cl2
– CH4
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or H-H
or
Duet
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Covalent Bonds
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Electronegativity• The measure of the ability of an atom to
attract electrons to itself– Increases across period (left to right) and– Decreases down group (top to bottom)– fluorine is the most electronegative element– francium is the least electronegative element
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Electronegativity Scale
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Types of Bonding
1) Non-Polar Covalent Bond:• Difference in electronegativity
values of atoms is 0.0 – 0.4• Electrons in molecule are
equally shared• Examples: Cl2, H2, CH4
ENCl = 3.03.0 - 3.0 = 0
Pure Covalent
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2) Polar Covalent Bond:• Difference in
electronegativity values of atoms is 0.4 – 1.7/2.0
• Electrons in the molecule are not equally shared• The atom with the higher
EN value pulls the electron cloud towards itself
• Partial charges• Examples: HCl, ClF, NO
ENCl = 3.0ENH = 2.1
3.0 – 2.1 = 0.9Polar Covalent
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Electrostatic Potential Maps
• A graphical depiction of electron distribution
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3) Ionic Bond: • Difference in EN
above 1.7-2.0• Complete transfer
of electron(s)• Whole charges
ENCl = 3.0ENNa = 1.0
3.0 – 0.9 = 2.1Ionic
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Dipole Moment ()
• Depends on charge separation and distance• = qr (a vector quantity)
• q = magnitude of charge• r = vector from site of + charge to site of – charge
• Units = Debyes (D)
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Molecular Polarity
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Lewis Dot Structures1) Count the number of valence electrons present in
the molecule2) Determine the arrangement of atoms. Generally,
the atom that occurs least often is central. Join the terminal atoms to the central atom(s) using shared pairs of electrons (bonds)
3) Place any remaining electrons around the terminal atoms to satisfy the octet rule
• Exception: Hydrogen
4) Place any remaining electrons on the central atom(s) to satisfy the octet rule
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5) Check to make sure:• You’ve used the correct number of valence
electrons• Everyone has an octet (or duet)• Everyone is doing what they like to do
6) If the number of electrons around the central atom is less than 3, change the single bonds to multiple bonds
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What Things Like To Do1) Halogens
• Like to be terminal• Like to have one bonding pair
(two shared electrons) and 3 lone pairs (non-bonding electrons)
2) Carbon• Likes to have 4 bonding pairs
and no lone pairs• Likes to bond to other carbons• Likes to be central
3) Silicon• Likes to do what carbon does• Notice, it sits under C on the
periodic table
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4) Oxygen• Like to have 2 bonding pairs
and 2 lone pairs
5) Sulfur• Likes to do what O does
6) Nitrogen• Likes to have 3 bonding pairs
and 1 lone pair
7) Phosphorous• Likes to do what N does
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8) Hydrogen• Likes to be terminal with only 1
bond• Do not put lone pairs on H
9) Boron• Likes to have 3 bonds and no lone
pairs• Likes a sextet instead of an octet
(what everybody else besides Hydrogen likes)
10) *Note: • A double bond = 2 bonding pairs• A triple bond = 3 bonding pairs
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Problems
• Draw the Lewis Dot Structures for the following molecules
1) CO2
2) P2H4
3) O3
4) NO3-
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Drawing Resonance Structures
1. Draw first Lewis structure that maximizes octets
2. Assign formal charges3. Move electron pairs from atoms with (-)
formal charge toward atoms with (+) formal charge
-1
-1
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Formal Charge• Assigned charge for each atom in a molecule/ion– Electronic bookkeeping – may or may not correspond
to a real charge– Sum of formal charges on each atom must equal the
total charge on the molecule/ion
• FC = Valence e-’s – Lone Pair e-’s – ½ bonding e-’s
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Molecular Structures of Covalent Compounds
• Atomic connectivity: How atoms in a molecule are connected
• Molecular geometry: How far apart atoms are and how they are arranged in space– Bond lengths– Bond angles– Dihedral angles
OR
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Bond Length
• Distance between nuclei
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• Increases with atoms in higher rows• Decreases toward higher atomic number along a row• Decreases with increasing bond order
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Bond Angles
• Angle between each pair of bonds• Contribute to molecular shape• Determined by Valence-shell electron-pair
repulsion (VSEPR)• Use molecular models!• Line-and-wedge structures
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Drawing LDS With Correct Geometry
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Valence Shell Electron Pair Repulsion Theory
• VSEPR theory:– Electrons repel each other– Electrons arrange in a
molecule themselves so as to be as far apart as possible• Minimize repulsion• Determines molecular
geometry
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Defining Molecular Shape• Electron pair geometry: the geometrical
arrangement of electron groups around a central atom– Look at all bonding and non-bonding e-’s
• Molecular Geometry: the geometrical arrangement of atoms around a central atom– Ignore lone pair electrons
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Problems• Predict the approximate geometry in each of
the following molecules– BF3
– HCN– CO3
2-
• Estimate the bond angles and relative bond lengths in the following molecule
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Dihedral Angle
• Also known as the torsional angle • Rotation can occur along single bonds
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Valence Bond Theory
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Types of Bonds• A sigma () bond results when the bonding orbitals
point along the axis connecting the two bonding nuclei– either standard atomic orbitals or hybrids
• s-to-s, p-to-p, hybrid-to-hybrid, s-to-hybrid, etc.
• A pi () bond results when the bonding orbitals are parallel to each other and perpendicular to the axis connecting the two bonding nuclei– between unhybridized parallel p orbitals
• the interaction between parallel orbitals is not as strong as between orbitals that point at each other; therefore bonds are stronger than bonds
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Problems
• Write a hybridization and bonding scheme for acetaldehyde
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Molecular Orbital Theory
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Bond Order: ½ (# of electrons in bonding MO’s - # of electrons in antibonding MO’s)
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Problems
1) Draw an MO diagram to predict the bond order of N2
2) Draw an MO diagram to predict the bond order of CN-
3) Use MO theory to determine the bond order of Ne2