Chapter 10 Chemical Bonding
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Transcript of Chapter 10 Chemical Bonding
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Chapter 10 Chemical Bonding
HIV-protease
Atoms interact with other atoms to form molecules, this is chemical bondingBonding theories – are models that predict how atoms bond together to form molecules
Bonding Theories are applied to design molecules that will interfere with the activesite of HIV-protease. This delays or inhibitsthe onset of AIDS.
protease inhibitor
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CHAPTER OUTLINE Chemical Bonds Ionic Bonds and Covalent Bonds Electronegativity Bond Polarity & Electronegativity Lewis Structures Resonance Molecular Shapes Molecular Polarity
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CHEMICAL BOND
Most matter in nature is found in form of compounds: 2 or more elements held together through a chemical bond.
Elements combine together (bond) to fill their outer energy levels and achieve a stable structure (low energy).
Noble gases are un-reactive since their energy levels are complete.
The nature and type of the chemical bond is directly responsible for many physical and chemical properties of a substance: (e.g. melting point, conductivity).
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CHEMICAL BOND
When the conductivity apparatus is placed in salt solution, the bulb will light.
But when it is placed in sugar solution, the bulb does not light.
This difference in conductivity between salt and sugar is due to the different types of bonds between their atoms.
Two common types of bonding are present: ionic & covalent.
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Gilbert Newton Lewis (1875 - 1946) was a famous American physical chemist known for the discovery of the covalent bond (see his Lewis dot structures and his 1916 paper "The Atom and the Molecule")
Other major contributions were his theory of Lewis acids and bases andLewis coined the term "photon" for the smallest unit of radiant energy.
Lewis is known for:Covalent bondLewis dot structuresValence bond theoryElectronic theory of acids and basesChemical thermodynamicsHeavy waterNamed photonExplained phosphorescence
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The Origin of Lewis Symbols of Atoms
Drawings of cubical atoms, the corners of the cube represented possible electron positions
Lewis later cited these notes in his classic 1916 paper on chemical bonding, as being the first expression of his ideas.
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LEWIS SYMBOLS OF ATOMS
Lewis structures use Lewis symbols to show valence electrons in molecules and ions of compounds.
Lewis symbols for the first 3 periods of representative elements are shown below:
In Lewis symbols, valence electrons for each element are shown as a dot.
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Lewis Bonding Theory
• atoms bond because it results in a more stable electron configuration
• atoms bond together by either transferring or sharing electrons so that all atoms obtain an outer shell with 8 electrons– Octet Rule– there are some exceptions to this rule – the
key to remember is to try to get an electron configuration like a noble gas
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Lewis Symbols of Ions
• Cations have Lewis symbols without valence electrons– Lost in the cation formation
• Anions have Lewis symbols with 8 valence electrons– Electrons gained in the formation of the
anion
Li• Li+1 :F: [:F:]-1
•
•• ••
••
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Ionic Bonds
• metal to nonmetal• metal loses electrons to form cation• nonmetal gains electrons to form anion• ionic bond results from + to - attraction
– larger charge = stronger attraction– smaller ion = stronger attraction
• Lewis Theory allow us to predict the correct formulas of ionic compounds
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IONIC BOND
Ionic bonds occur when electrons are transferred between two atoms.
After bonding, each atom achieves a complete shell (noble gas configuration).
Ionic bonds occur between metals and non-metals.
Metal Nonmetal
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IONIC BOND
Atoms that lose electrons (metals) form positive ions (cations).
Atoms that gain electrons (non-metals) form negative ions (anions).
The smallest particles of ionic compounds are ions (not atoms).
Cation
Anion
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Using Lewis Theory to Predict Chemical Formulas of Ionic Compounds
Predict the formula of the compound that forms betweencalcium and chlorine.
Draw the Lewis dot symbolsof the elements
Ca∙∙ Cl ∙∙∙
∙ ∙∙ ∙
Transfer all the valance electronsfrom the metal to the nonmetal,adding more of each atom as yougo, until all electrons are lost from the metal atoms and all nonmetal atoms have 8 electrons
Ca∙∙ Cl ∙∙∙
∙ ∙∙ ∙Cl ∙∙∙
∙ ∙∙ ∙
Ca2+
:: Cl
:: Cl
CaCl2
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Covalent Bonds• typical of molecular species• atoms bonded together to form molecules
– strong attraction
• sharing pairs of electrons to attain octets• molecules generally weakly attracted to
each other– observed physical properties of molecular
substance due to these attractions
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COVALENT BOND
Covalent bonds form when electrons are shared between two atoms.
Covalent bonds form between twonon-metals.
The smallest particles of covalent compounds are molecules.
Electrons shared
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Single Covalent Bonds• two atoms share one pair of electrons
– 2 electrons
• one atom may have more than one single bond
F••
••
•• • F•••••••
HF••
••
•• ••
••F•••• H O
•• ••••
••
H•H• O••
• •
••
F F
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Double Covalent Bond
• two atoms sharing two pairs of electrons– 4 electrons
• shorter and stronger than single bond
O••••O••
••••••
O••
• •
••O••
• •
••
O O
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Triple Covalent Bond
• two atoms sharing 3 pairs of electrons– 6 electrons
• shorter and stronger than single or double bond
N••
• •
•N••
• •
•
N•••••••••• N
N N
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POLAR & NON-POLARBONDS
Two types of covalent bonds exist:
Non-polar covalent bonds occur between similar atoms.
In these bonds the electron pair is shared equally between the two protons.
Polar & Nonpolar
Electrons shared equally
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POLAR & NON-POLARBONDS
Polar covalent bonds occur between different atoms.
In these bonds the electron pair is shared unequally between the two atoms.
As a result there is a charge separation in the molecule, and partial charges on each atom.
+ H F
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Dipole Moments• A dipole is a material with positively and
negatively charged ends• Polar bonds or molecules have one end
slightly positive, +; and the other slightly negative, -
– not “full” charges, come from nonsymmetrical electron distribution
• Dipole Moment, , is a measure of the size of the polarity – measured in Debyes, D
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ELECTRONEGATIVITY
Electronegativity (E.N.) is the ability of an atom involved in a covalent bond to attract the bonding electrons to itself.
Linus Pauling derived a relative Electronegativity Scale based on Bond Energies.
Cs 0.7
F 4.0
Least electronegative
Most electronegative
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ELECTRONEGATIVITY
Electronegativity increases
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BOND POLARITY &ELECTRONEGATIVITY
The more polar the
bond formed
Polarity is a measure of the inequality in the sharing of bonding electrons
The more different the
electronegativity of the elements
forming the bond
The larger the electronegativity
difference(EN)
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POLARITY &ELECTRONEGATIVITY
As difference in electronegativity
increases
Bond polarity increases
Most polar
Least polar
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POLARITY &ELECTRONEGATIVITY
Electronegativity
differenceBond Type
EN = 0 Non-polar covalent
0 < EN <1.7 Polar covalent
1.7 < EN Ionic
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H H
Hydrogen Molecule
The molecule is nonpolar covalent
Electronegativity2.1
Electronegativity2.1
POLARITY &ELECTRONEGATIVITY
EN = 0
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H Cl
Hydrogen Chloride Molecule
Electronegativity2.1
Electronegativity3.0
The molecule is polar covalent
+ -
EN = 0.9
POLARITY &ELECTRONEGATIVITY
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29Sodium Chloride
Na+ Cl-
Electronegativity0.9
Electronegativity3.0
The bond is ionicNo molecule exists
EN = 2.1
POLARITY &ELECTRONEGATIVITY
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SUMMARYOF BONDING
Ionic Bond(large EN)
Covalent Bond(small to moderate EN)
Non-polar(similar electronegativities)
Polar(moderate EN)
EN > 1.7
EN = 0
0 < EN < 1.7
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Bonding & Lone Pair Electrons
• Electrons that are shared by atoms are called bonding pairs
• Electrons that are not shared by atoms but belong to a particular atom are called lone pairs– also known as nonbonding pairs
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LEWIS STRUCTURES
In a Lewis structure, a shared electron pair is indicated by two dots between the atoms, or by a dash connecting them.
Unshared pairs of valence electrons (called lone pairs) are shown as belonging to individual atoms or ions.
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LEWIS STRUCTURES
Covalent molecules are best represented with electron-dot or Lewis structures.
Structures must satisfy octet rule (8 electrons around each atom).
Hydrogen is one of the few exceptions and forms a doublet (2 electrons).
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LEWISSTRUCTURES
Bonding electrons can be displayed by a dashed line.
Non-bonding electrons must be displayed as dots.
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Polyatomic Ions
• The polyatomic ions are attracted to opposite ions by ionic bonds– Form crystal lattices
• Atoms in the polyatomic ion are held together by covalent bonds
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Lewis Formulas of Molecules
• shows pattern of valence electron distribution in the molecule
• useful for understanding the bonding in many compounds
• allows us to predict shapes of molecules
• allows us to predict properties of molecules and how they will interact together
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LEWISSTRUCTURES
More complex Lewis structures can be drawn by following a stepwise method:
1. Count the number of electrons in the structure.
2. Draw a skeleton structure.- most metallic element generally central- halogens and hydrogen are generally terminal- many molecules tend to be symmetrical- in oxyacids, the acid hydrogens are attached
to an oxygen
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LEWISSTRUCTURES
More complex Lewis structures can be drawn by following a stepwise method:
3. Connect atoms by bonds (dashes or dots).
4. Distribute electrons to achieve Octet rule.
5. Form multiple bonds if necessary.
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Example 1:
Write Lewis structure for H2O
H2O = 8 electrons 2 (1) + 6 = 8Step 1:
Step 2:
H O HSkeleton structure should be
symmetrical
Step 3:
4 electrons used4 electrons remainingStep 4:
Octet rule is satisfied
Hydrogen has doublet
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Example 2:
Write Lewis structure for CO2
CO2 = 16 electrons 4 + 2(6) = 16Step 1:
Step 2:
O C OSkeleton structure should be
symmetrical
Step 3:
4 electrons used12 electrons remaining
Step 4:
Octet rule is satisfied
10 electrons used6 electrons remaining
Octet rule is NOT satisfied
Step 5:
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Writing Lewis Structures forPolyatomic Ions
• the procedure is the same, the only difference is in counting the valence electrons
• for polyatomic cations, take away one electron from the total for each positive charge
• for polyatomic anions, add one electron to the total for each negative charge
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Example 3:
Write Lewis structure for CO32-
CO32- = 24 electrons 4+3(6)+2 = 24Step 1:
Step 2:
O C O
O
Step 3:
Step 4:
18 electrons remaining
12 electrons remaining6 electrons remaining0 electrons remaining
Octet rule is satisfiedOctet rule is NOT satisfied
Step 5:
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Example 4:
Determine if each of the following Lewis structures are correct or incorrect. If incorrect, rewrite the correct structure.
2(5) + 4(1) = 14
Structure is incorrect
Only 12 electrons shown
22
224
Structure has 14 electrons
Octets are complete
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Exceptions to the Octet Rule• H & Li, lose one electron to form cation
– Li now has electron configuration like He – H can also share or gain one electron to have
configuration like He
• Be shares 2 electrons to form two single bonds• B shares 3 electrons to form three single bonds• expanded octets for elements in Period 3 or
below
– using empty valence d orbitals
• some molecules have odd numbers of electrons– NO
:: ON
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Some molecules, such as SF6 and PCl5 have more than 8 electrons around a central atom in their Lewis structure.
SF6 and PCl5 can violate the octet rule through the use of empty d orbitals:
both S and P can utilize empty d orbitals to hold pairs of electrons that help
bond halogen atoms.
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Resonance
• we can often draw more than one valid Lewis structure for a molecule or ion
• in other words, no one Lewis structure can adequately describe the actual structure of the molecule
• the actual molecule will have some characteristics of all the valid Lewis structures we can draw
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Resonance• Lewis structures often do not accurately represent
the electron distribution in a molecule– Lewis structures imply that O3 has a single (147 pm) and
double (121 pm) bond, but actual bond length is between, (128 pm)
• Real molecule is a hybrid of all possible Lewis structures
• Resonance stabilizes the molecule– maximum stabilization comes when resonance forms
contribute equally to the hybrid
OO
+
O OO
+
O
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Resonance• we can often draw more than one valid
Lewis structure for a molecule or ion
• Real molecule is a hybrid
of all possible Lewis structuresO N
O
O·· ··
········
··
··
O N
O
O
·· ····
····
······
The three oxygens are chemically equivalent, so it makes nodifference to the ion which oxygen assumes the double bond.
represents resonance structures
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MOLECULARSHAPES
The three-dimensional shape of the molecules is an important feature in understanding their properties and interactions.
All binary molecules have a linear shape since they only contain two atoms.
More complex molecules can have various shapes (linear, bent, etc.) and need to be predicted based on their Lewis structures.
A very simple model , VSEPR (Valence Shell Electron Pair Repulsion) Theory, has been developed by chemists to predict the shape of large molecules based on their Lewis structures.
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MOLECULARSHAPES
Based on VSEPR, the electron pair groups in a molecule will repel one another and seek to minimize their repulsion by arranging themselves around the central atom as far apart as possible.
Electron pair groups can be defined as any one of the following:
bonding pairs non-bonding pairs
multiple bonds
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SUMMARY OFVSEPR SHAPES
Number of electron pair groups around central atom
MolecularShape
BondAngle
ExamplesBonding Non-bonding
2 0 Linear 180 CO2
3 0 Trigonal planar 120 BF3
2 1 Bent 120 SO2
4 0 Tetrahedral 109.5 CH4
3 1 Pyramidal 109.5 NH3
2 2 Bent 109.5 H2O
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MOLECULARSHAPES
Molecules with 2 electron pair groups around the central atom form a linear shape.
2 electron pairs around the
central atom
Shape is linear
Linear molecules have polar bonds, but are usually non-polar.
Bond angle is 180
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MOLECULARSHAPES
Molecules with 3 electron pair groups around the central atom form a trigonal planar shape.
Trigonal planar molecules have polar bonds, but are usually non-polar.
Bond angle is 1203 electron pairs
around the central atom
Shape is trigonal planar
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MOLECULARSHAPES
Molecules with 2 bonding pairs and 1 non-bonding pair groups around the central atom form a bent shape.
Bent molecules have polar bonds, and are polar.
Shape is bent
2 bonding pairs around the
central atom
1 Non-bonding pair
Bond angle is 120
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MOLECULARSHAPES
Molecules with 4 electron pairs groups around the central atom form a tetrahedral shape.
Tetrahedral molecules have polar bonds, and are usually non-polar.
4 bonding pairs around the
central atomShape is
tetrahedral Bond angle is 109.5
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MOLECULARSHAPES
Molecules with 3 bonding pairs and 1 non-bonding pair groups around the central atom form a pyramidal shape.
Pyramidal molecules have polar bonds, and are polar.
Shape is pyramidal
3 bonding pairs around the
central atom
1 Non-bonding pair
Bond angle is 109.5
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MOLECULARSHAPES
Molecules with 2 bonding pairs and 2 non-bonding pair groups around the central atom form a bent shape.
Bent molecules have polar bonds, and are polar.
Shape is bent
2 bonding pairs around the
central atom
2 Non-bonding pair
Bond angle is 109.5
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SUMMARY OFMOLECULAR SHAPES
Linear
Trigonal planar
Tetrahedral
Bent
Pyramidal
Symmetrical shapesPolar bonds
Non-polar molecules
Unsymmetrical shapes
Polar bondsPolar molecules
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Polarity of Molecules
• For a molecule to be polar it must
1) have polar bonds, symmetrical shape, and
different terminal atoms
2) have polar bonds• electronegativity difference - theory• bond dipole moments – measured
3) have an unsymmetrical shape• using vector addition
• polarity effects the intermolecular forces of attraction
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:: OCO
polar bonds,but nonpolar moleculebecause pulls cancel
OH H
polar bonds,and unsymmetrical
shape causes moleculeto be polar
Dipole moment is the measured polarity of a polar covalent bond. It is defined as the magnitude of charge (electrons) on the atoms and the distance between the two bonded atoms.
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CH2Cl2
= 2.0 D CCl4
= 0.0 D
C
Cl
ClCl
Cl
C
Cl
ClH
H
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Adding Dipole Moments
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COMPARING PROPERTIESOF IONIC & COVALENT
COMPOUNDS
Ionic Covalent
Structural Unit Ions Molecules
Melting Point High Low
Boiling Point High Low
Solubility in H2O High Low or None
Electrical Cond. High None
Examples NaCl, AgBr H2, H2O
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THE END