chemical bonding. ***occurs when atoms of elements combine together to form compounds.*****
TRANSCRIPT
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CHEMICAL BONDING
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***Occurs when atoms of elements combine together
to form compounds.*****
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Formation of compounds Involves valence electrons. PE is lower in bonded atoms. Attractive force that develops is
called "chemical bond“ Occurs during chem. reactions
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Two (or three) methods• Ionic bonding - attraction of ions
• Covalent bonding - shared pairs of electrons
• Metallic bonding - alloys (metals) not compounds
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An ionic bond -(electrovalent)• a. Definition -
An attraction that forms between oppositely charged ions
• b. Pos. ions + Neg. ions → neutral compound
• ∆EN > 1.7
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Ionic Bonds• Cations: Metals lose electron(s) and
become positive ions(+). • Examples: Na+1, K+1, Mg+2, Ca+2, Al+3
• Anions: Nonmetals gain electron(s) & becomes negative ions(-).• Name ends in “–ide”• Examples: Cl1-, Br1-, O2-
• Choride, bromide, oxide . . .
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Ionic bonds - (crystals)
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Characteristics of Ionic Solids• Made from + and - ions• Metal-nonmetal (or polyatomic ions)• Compound is neutral.• Tend to be solid,• Brittle and crystalline• High MP and BP• Have strong attractions in all directions• Non-conductors as solids, but will
conduct when molten or dissolved• Some dissolve readily
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Ionic solids are brittledue to Crystal lattice
+ - + -+- +-
+ - + -+- +-
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Ionic solids are brittle
+ - + -
+- +-+ - + -
+- +-
• Strong Repulsion breaks crystal apart.
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Lattice Energy (ionic)
• More correct than bond energy for crystals.
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Metallic Bonding• Sea of Electrons• In metals, the valence electrons are
not bonded to any specific atom. (delocalized)
• Able to move freely over the positive centers.
• Causes unique properties.
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Characteristics of Metals
• A. Malleable: dent when hammered.• B. Ductile: draw into a wire• C. Conductivity: electricity and heat• D. Alloy: a blend of metals
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Sea of Electrons
+ + + ++ + + +
+ + + +
• Electrons are free to move through the solid.
• Metals conduct electricity and heat well.
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Malleable
+ + + ++ + + +
+ + + +
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Malleable
+ + + +
+ + + ++ + + +
• Electrons allow atoms to slide by.
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Alloys• Alloys: mixtures of two or more metals. • Important because their properties are often
better than the individual elements.• Examples:
– Bronze is made from copper & tin. It is harder than copper & more easily cast.
– Sterling silver: Ag (92.5%) & Cu (7.5%) – Stainless steel: Fe (80.6%), Cr (18%), C (0.4%), & Ni (1%)-- Brass, Pewter, and others.
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Covalent Bond • a. an attractive force that develops
between atoms that are sharing pairs of electrons
• ∆EN < 1.7• b. Hydrogen – H2
• H• + H• → H:H (dot diagram)• Structural formulas use a dash H - H
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Characteristics of Covalent Compounds
• Nonmetal-nonmetal combinations• Can be gases, liquids, or solids• Low to med. MP and BP• Insulators/Nonconductors (except for
acids)• Molecular (a few are crystalline)• Generally not soluble (some polar
exceptions)
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Bond energy• Energy required to
break a bond.
• Bonds form to lower PE, so breaking bonds will increase PE.
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Breaking Bonds is always endothermic
Energy is required.
ALWAYS.
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Why not ionic?• The difference in electronegativity is
less than 1.7• Electrons are not pulled away from
either atom.• They are shared.
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Fluorine – F2
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Ammonia -NH3
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Special Types of Covalent Bonds
• Multiple bonds – Occasionally atoms share more than one pair of electrons
– Double bond – two shared pairs Ex. O2 O=O or
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• - Triple bonds - two atoms share
three pairs of electrons. Ex. N2
N N or
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How to write Lewis Structures
• 1. Set out the atoms – think symmetry.• 2. Count all the valence e-
• 3. Insert single bonds first, then fill rest.• All the e- are paired.
• each nonmetal atom requires an octet.• H only requires 2 e-.
• Multiple bonds may be needed.• Readily formed by C, N, O, S, and P.
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Polarity
• An unequal sharing of electrons due to difference in electronegativity.
• Polar bond – Any bond with ∆EN 0.5 - 1.7• Polar molecule – Has a positive end and a
negative end. • Occurs in water and ammonia (**know
these two)• Causes intermolecular attraction to increase
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Molecular Substances• Covalently bonded substances – show more
variety in phases and properties. Tend to be insulators(nonconductors). • Nonmetal-nonmetal combinations
1. Nonpolar – tends to be gases at room temp. – have only dispersion (Vanderwaals) forces. Have low MP and BP and high VP(vapor pressure)
2. Polar –(dipoles) tend to be liquids or solids at room temp. Have ↑MP and BP and↓VP
3. H- bonding – very strong type of polar
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VSEPR
VALENCE SHELL ELECTRON PAIR REPULSION
A theory which describes the shapes of molecules based on the idea that pairs of electrons will repel each other as much as possible.
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Polar molecules are asymmetrical and have polar bonds.Nonpolar molecules are symmetrical and may or may not have polar bonds.
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Intermolecular Attractions (Vanderwaals forces)
• Attractive forces between molecules• Happens to covalent molecules.• Strongest - hydrogen bonds• Medium - dipoles/ polar • Weakest - dispersion or London
forces ( All molecules have London forces)
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Hydrogen bonding• In some highly polar compounds, a H
atom is attracted to, and forms a weak bond with, an adjacent molecule.
• Only occurs in compounds where there is:
• H-F (strongest)• H-O• H-N (weakest)
• (know these three)
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Hydrogen Bonding
HHO
d+
d-
d+
H HOd
+
d-
d+
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Hydrogen bonding
HHO H H
O
HH
O
H
H
OH
HO
H
HO HH
O
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H-bonding• Gives water its very unusual
properties.• High MP and BP• Holds the DNA molecule together.• Provides stability and shapes for
proteins, enzymes, etc.• Strongest type of intermolecular
attraction.
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London Dispersion Forces(Van der waals)
• Weakest type of intermolecular attraction. • Develops between nonpolar molecules
due to temporary shifts in the electron positions.
• Strength of attraction is directly proportional to the number of electrons
• (wax is a nonpolar molecule – but large so it is solid)
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Network Solids• Large arrays of covalently bonded
crystals. • Do not conduct, hardest solids • Very high MP and BP • Examples: diamond, graphite, SiO2
(very few – easiest to just memorize)• More on this in 2nd semester
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Bond length and energy• Bond length depends on the size of
the atoms/ions and the number of bonds between them • C-C is longer than C=C is longer than CC
Shorter bonds are stronger. • Measured in kJ/mol
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Breaking Bonds is always endothermic
Energy is required.
ALWAYS.
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Endothermic reactions
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Forming bonds is exothermic
Energy is released.
ALWAYS
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Exothermic reactions