sections 6.1 – 6.4. chemical bond = a link between atoms why does it occur? the nucleus of one...
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Sections 6.1 – 6.4
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Chemical Bond = A link between atoms
Why does it occur?The nucleus of one atom is attracted to the electrons of another.
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IONIC BONDIon = Atom which has gained or lost electron(s)Metal =
-LEFT side of Periodic TableWeak nucleus / Low Electronegativity
-LOSERS of electrons Become positively (+) charged ions
(Cations)
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Nonmetal =-RIGHT side of Periodic Table Strong nucleus / High Electronegativity-GRABBERS of electrons Become negatively (-) charged ions
(Anions)
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Atoms gain or lose valence electrons to become a NOBLE GAS CONFIGURATION
Right/Left Side?
Metal/ Nonmetal?
Lose/Gain? Noble Gas it resembles?
Mg
Li
Cl
O
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Ionic bond = A chemical bond between a cation (+) and an anion (-). Caused by a TRANSFER of electron(s).
Usually a metal + a nonmetal
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COVALENT BOND = A bond caused by a SHARING of electrons
Usually a nonmetal + a nonmetalNonpolar Covalent = Equal sharing of the
electrons. Atoms are close in strengthPolar Covalent = Unequal sharing of the
electrons. One atom is a little bit stronger than the other
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How do you tell which type of bond it is?-By ELECTRONEGATIVITY
A chart of electronegativity will be provided to you.
-The greater the difference in electronegativity – the more ionic the bond.
-Electrons spend more time closer to the element with higher electronegativity.
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If the ABSOLUTE VALUE of the electronegativity difference is:
GREATER THAN 1.7 = IONIC Bond LESS THAN 0.3 = NONPOLAR COVALENT
Bond 0.3 – 1.7 = POLAR COVALENT Bond
Examples:
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METALLIC BONDUsually metals only-The metal gives up valence electrons.-Electrons are free to move about.
Atom
Electron Sea
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Covalent Bond = A sharing of electronsMolecule = A group of atoms held by
covalent bonds (ex – water)Diatomic Molecule = Molecule with only 2
atoms (7 naturally occurring ones)Molecular Compound = Compound made of
moleculesMolecular Formula = The type and number of
atoms in a molecule (ex – H2O)
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Sharing electrons in a covalent bond makes the atoms more stable and decreases the energy of the atoms. Energy is released when a bond is FORMED.
Overlapping of Orbitals – Example H2:
H H
+
H2
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Atoms in a compound obtain the electron configuration of a NOBLE GAS to gain stability
Exceptions to the Octet Rule: Hydrogen Incomplete octets Expanded octets
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Surrounded by only 2 electrons◦ Can only form single bonds!!
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Too few valence electrons Rather rare:
◦ Boron, aluminum, beryllium Examples:
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Some central atoms form 5 or 6 bonds Elements with atomic #s greater than 10 Examples:
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-A picture of the covalent bonds in a molecule
-Basic Rules:◦ Determine a central atom
If C is present – always central H is never central Group 17 is never central
◦ Arrange to form skeleton (like a plus sign)◦ Place dots around each element and connect
dots
Examples:
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More Examples:
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Single Bond = 1 pair of electrons (2 e-s total) shared between two atoms (longest length; lowest bond energy)
Double Bond = 2 pairs of electrons (4 total e-s) shared between two atoms
Triple Bond = 3 pairs of electrons (6 total e-s) shared between two atoms (shortest length; highest bond energy)
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More Examples:
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Same as non-charged molecules, except:◦ For positive ions – subtract electrons to total # of
valence electrons◦ For negative ions – add electrons to total # of
valence electrons Examples:
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Ionic Bond = Bond formed by the attraction of a cation to an anion
Crystal Lattice = 3-Dimensional network of ions
Formula Unit = Simplest ratio of ions
NaCl
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Dot structures for Ionic Compounds:-Want to reach noble gas configuration“Math equation”-Draw an ARROW to show the transfer of e--Draw as many of each ion as neededExamples:
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Molecular Ionic
Bond Type Covalent Ionic
Structure Individual Molecules Crystal Lattice
Strength of Bond Strong VERY strong
Mp/bp Low High
Drawing Lewis Structures “Math Eqn”
Other --- Conducts electricity when in water
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Do electronegativity difference first!! Examples:
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Metals have LOW electronegativity – Will LOSE electrons
The steps:-Donate valence electrons to electron sea-Electrons free to move about-All electrons in sea are shared by all atoms
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Properties of Metals:1. Good conductors of heat – e- sea shakes2. Good conductors of electricity – e- in sea
can move3. Malleable – atoms can be pushed closer4. Ductile – atoms can be pushed closer5. Luster – light bounces off e- sea
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Section 6.5
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VSEPR = Valence Shell Electron Pair Repulsion Theory
Valence electrons move as far away from each other as possible
1. Draw Lewis Structure2. Look at Central Atom3. Count electron areas (shared areas & lone
pairs)4. Use chart info
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Name Shape Shared Areas
Lone Pairs Bond Angles
Linear 2 0 180°
Bent 2 1 120°
Trigonal Planar
3 0 120°
Tetrahedral 4 0 109.5°
Trigonal pyramidal
3 1 107°
Bent 2 2 104.5°
Trigonal bipyramidal
5 0 90°, 120°
Octahedral 6 0 90°
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Examples:
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Examples:
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Examples:
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Additional Handout!! Not in note packet!! Partial Charges
◦ In a polar bond ONLY!!◦ A tug of war occurs – one atom is “stronger”
than the other! δ+ and δ-
◦ Greek letter delta
◦ Compare EN: Higher EN value of the two = δ-
Lower EN value = δ+
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Examples:◦ H2O
◦ CH4
◦ NH3
◦ CO2
◦ And the hardest – CH3Cl
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Mixing a set of atomic orbitals◦ forms a new set of atomic orbitals with the same total electron capacity◦ properties and energies intermediate between those of the original
unhybridized orbitals.
Three types: sp (triple bonds), sp2(double bonds), sp3(single bonds)
Carbon:
C
BECOMES
C
1s22s22p2 four sp3 hybrid
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1. DipoleDipole = Molecule with overall charge2. NonPolar With Polar Sites (NPWPS)NonPolar With Polar Sites (NPWPS) =
Molecules with area of charge which cancel out
3. Nonpolar Nonpolar = Molecule with no areas of charge
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How do you tell the difference?-Ask yourself these questions…
Is the molecule polar or nonpolar (Difference in EN)
Polar (∆EN = 0.3-1.7) Can charge be sliced?
YES = Dipole NO = NPWPS
Nonpolar (∆EN < 0.3)
One straight line so all positive charge is separated from all
negative charge; through BONDS only!!
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AKA – EXTERNAL BONDSThe attraction BETWEENBETWEEN MoleculesTypes of External Bonds:1.1. Dipole-Dipole InteractionsDipole-Dipole Interactions
-Occur due to attraction between partial charges-Occur between:
• Two dipoles (strongest)• Dipole to NPWPS• Two NPWPS (weakest)
Hydrogen Bond = Special case of a dipole-dipole external bond that involves a hydrogen atom
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2.2. London ForceLondon Force-Occurs between nonpolar molecules-Very weak connectionThe Steps:
A. Electrons in one molecule shift instantaneously to one side
B. Instantaneous charge resultsC. Electrons in another molecule are repelledD. Very weak attraction results
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1. State of Matters > l > g
**This means that solids have strongest external bonds; gases have weakest bonds
2. Evaporation (Volatility – tendency of a substance to vaporize)slow > fast
**Those compounds that evaporate very slow have stronger bonds than those that evaporate quickly
3. Thickness (Viscosity – a measurement of resistance to flow)thick > thin
**Substances that are “thicker” have stronger bonds
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4. Wetness (Adhesion – the force of attraction)To feel wet the substance must bond to your skin (to the Na+Cl-
**If you feel wetness, the substance is bonding to your skin
5. DissolvingLIKE DISSOLVES LIKE
**Polar dissolves in polar; nonpolar dissolves in nonpolar
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