chapter 6 -chemical bonding a steroid alkaloid derived from skin secretions of the phyllobates and...
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Chapter 6 -Chemical BondingChapter 6 -Chemical Bonding
A steroid alkaloid derived from skin secretions of the Phyllobates and Dendrobates genera of South American poison-arrow frogs. It is one of the most potent venoms known.
Batrachotoxin
BondsBonds Forces that hold groups of atoms together and Forces that hold groups of atoms together and
make them function as a unit.make them function as a unit.
Ionic bonds– transfer of electronsIonic bonds– transfer of electrons
Covalent bonds– sharing of electronsCovalent bonds– sharing of electrons
Polar Covalent bonds – unequal sharing of electrons Polar Covalent bonds – unequal sharing of electrons that results in an that results in an unbalanced unbalanced distribution of chargedistribution of charge
ElectronegativityElectronegativity The ability of an atom in a molecule The ability of an atom in a molecule
to attract shared electrons to to attract shared electrons to itself.itself.
Linus Pauling1901 - 1994
e-e-
Table of Electronegativities
Shielding Effect
• Electrons in the inner energy levels block the attraction of the nucleus for the valence electrons.
• Shielding increases down a group.
• This causes EN values to decrease.
What is Nuclear Shielding?
P+
P+P+
P+
P+
P+P+
P+
P+
e-
e-
e-
The nucleus (+) pulls the electron (-) close to the core.
The further the electron is away from the nucleus, the weaker the nuclear pull.
e-
Inner electrons shield outer electrons from the nuclear pull.
Range of EN Values
3.3 2.0 0.5 0
Polar covalent- Electrons are shared, but unequally. There is some degree of ionic character
in these bonds.
Mostly Ionic Polar Covalent Mostly Covalent
∆∆EN ValuesEN Values
0.3 1.7
Nonpolar Covalent
Polar Covalent Ionic
EN ∆
Practice Problems
NaCl
Sodium’s EN = 0.9
Chlorine’s EN = 3.0
3.0 - 0.9 = 2.1
therefore it is mostly ionic
H2O
Hydrogen’s EN = 2.1
0xygen’s EN = 3.5
3.5 - 2.1 = 1.4
therefore it is polar covalent
Determine the bond type using EN values:
6.2 Covalent Bonds
Chapter 6 Chapter 6 Chemical Chemical BondingBonding
Covalent Bonds
Covalent Terms• Molecule: A neutral group of atoms that are held
together by covalent bonds• Diatomic Molecule: A molecule containing only two
atoms
• Molecular Compound: A chemical compound whose simplest units are molecules
• Chemical Formula: Indicates the relative numbers of atoms of each kind of a chemical compound by using atomic symbols and numerical subscripts
• Molecular Formula: Shows the types and numbers of atoms combined in a single molecule of a molecular compound
Example: BrINClHOF
Bonding Forces
Electron – Electron Electron – Electron repulsive forcesrepulsive forces
Proton – Proton repulsive forces
Electron – Proton attractive forces
Pure Covalent Bonding
Bond Length Diagram
Bond Energy
It is the energy required to break a bond.
It gives us information about the strength of a bonding interaction.
Electron Dot Electron Dot NotationNotation
The The OctetOctet Rule RuleChemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level.
Diatomic Fluorine
Hydrogen Chloride by the Octet Hydrogen Chloride by the Octet RuleRule
Formation of Water by the Octet Formation of Water by the Octet RuleRule
Comments About the Octet RuleComments About the Octet Rule
2nd row elements C, N, O, F observe the octet rule.
2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive.
3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals.
When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals.
Shows how valence electrons are arranged Shows how valence electrons are arranged among atoms in a molecule.among atoms in a molecule.
Reflects central idea that stability of a compound Reflects central idea that stability of a compound relates to noble gas electron configuration.relates to noble gas electron configuration.
Lewis Lewis StructuresStructures
CH
H
H
Cl..
.. .. ..
Completing a Lewis Structure CH3Cl
Add up available valence electrons:
C = 4, H = (3)(1), Cl = 7 Total = 14
Join peripheral atoms to the central atom with electron pairs.
Complete octets on atoms other than hydrogen with remaining electrons
Make carbon the central atom
..
..
..
How to Make a Lewis Dot Structure for a Molecule
• Draw the Lewis Dot Diagram for each element• Place the atom with the lowest EN value in the
center• Attach the rest of the atoms to the central• If lone electrons exist on adjacent atoms, pair
them for multiple bonds.
Multiple Covalent Bonds:Multiple Covalent Bonds:Double BondsDouble Bonds
•Two pairs of shared electrons
•Higher bond energy and shorter bond length than single bonds
Multiple Covalent Bonds:Multiple Covalent Bonds:Triple bondsTriple bonds
•Three pairs of shared electrons•Higher bond energy and shorter bond length than single or double bonds
ResonanceResonanceOccurs when more than one valid Lewis structure can be written for a particular molecule.
These are resonance structures. The actual structure is an average of the resonance structures.
Resonance in OzoneResonance in Ozone
Neither structure is correct.
Resonance in a carbonate ion:
Resonance in an acetate ion:
Resonance in Polyatomic Ions
Covalent Network CompoundsCovalent Network CompoundsSome covalently bonded substances DO NOT form
discrete molecules.
Diamond, a network of covalently bonded carbon atoms
Graphite, a network of covalently bonded carbon atoms
ModelsModels Models are attempts to explain how nature operates Models are attempts to explain how nature operates
on the microscopic level based on experiences in on the microscopic level based on experiences in the macroscopic world.the macroscopic world.
Models can be physical as with this DNA model
Models can be mathematical
Models can be theoretical or philosophical
Fundamental Properties of ModelsFundamental Properties of Models
A model does not equal reality.
Models are oversimplifications, and are therefore often wrong.
Models become more complicated as they age.
We must understand the underlying assumptions in a model so that we don’t misuse them.
6.3 Ionic Bonding
Chapter 6 Chapter 6 Chemical Chemical BondingBonding
Ionic Bonds
Ionic BondsIonic Bonds Electrons are transferred
Electronegativity differences are generally greater than 1.7 The formation of ionic bonds is always exothermic!
Ionic Bonding
• Formula Unit: The simplest collection of atoms from which an ionic compound's formula can be established
• Lattice Energy: The energy released when one mole of an ionic crystalline compound is formed from gaseous ions
Na+ (g) + Cl- (g) → NaCl (s) + 787.5 kJ
• Formation of ionic compounds is ALWAYS exothermic
Sodium Chloride Crystal LatticeSodium Chloride Crystal LatticeIonic compounds form solids Ionic compounds form solids at ordinary temperatures.at ordinary temperatures.
Ionic compounds organize in Ionic compounds organize in a characteristic crystal lattice a characteristic crystal lattice of alternating positive and of alternating positive and negative ions.negative ions.
Formation of sodium chlorideNa = 3s1 Cl = 3s23p5
Na+ = 2s22p6 Cl- = 3s23p6
Polyatomic Ions
• A charged group of covalently bonded atoms• Creation of octets results in an excess or deficit
of electrons
A Comparison of Ionic and Molecular Compounds
6.4 Metallic Bonding
Chapter 6 Chapter 6 Chemical Chemical BondingBonding
The Metallic Bond Model
• The chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons
• Electron Delocalization in Metals• Vacant p and d orbitals in metal's outer energy
levels overlap, and allow outer electrons to move freely throughout the metal
• The valence electrons do not belong to any one atom!
Swim in the Sea of Valence Electrons…
• In metals, the valence electrons are held loosely. The vacant p & d orbitals overlap.
• Metal atoms DO NOT lose their valence electrons in metallic bonding, rather they release them into a “Sea of Electrons”
• Although the atoms are bonded together, they are not bonded to any one particular atom, it is more like a large network.
Sea of Valence Electrons…
Bonding Between Metals• Results in an interaction that hold metal atoms
together, however, it is not called a compound.
• Special Properties result from this interaction:– Malleable- pounded/ rolled into sheets (aluminum
foil)– Ductile- Drawn into wire. (copper wires)– Conductivity- the flow of electrons– Luster- Shiny-The narrow range of energy
differences between orbitals allows electrons to be easily excited, and emit light upon returning to a lower energy level
Metallic Properties
• Metals are good conductors of heat and light• Metals have luster (shiny)• The narrow range of energy differences between
orbitals allows electrons to be easily excited, and emit light upon returning to a lower energy level
• Metals are Malleable- can be hammered into thin sheets
• Metals are ductile- ability to be drawn into wire• Metallic bonding is the same in all directions, so
metals tend not to be brittle
Metallic Bond Strength• Heat of Vaporization• The ease with which atoms in a metallic solid can be
separated from one another into individual gaseous atoms is related to bond strength
6.5 Molecular Geometry
Chapter 6 Chapter 6 Chemical Chemical BondingBonding
VSEPR Model
The structure around a given atom is determined principally by minimizing electron pair repulsions.
The model for predicting molecular shapes
VSEPR Theory: Repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible
(Valence Shell Electron Pair Repulsion)
VSEPR and Unshared Electron Pairs
• Unshared pairs take up positions in the geometry of molecules just as atoms do
• Unshared pairs have a relatively greater effect on geometry than do atoms
• Lone (unshared) electron pairs require more room than bonding pairs (they have greater repulsive forces) and tend to compress the angles between bonding pairs
• Lone pairs do not cause distortion when bond angles are 120° or greater
Predicting a VSEPR StructurePredicting a VSEPR Structure
Draw Lewis structure of each atom.Draw Lewis structure of each atom.
Put pairs as far apart as possible.Put pairs as far apart as possible.
Determine positions of atoms from the way Determine positions of atoms from the way electron pairs are shared.electron pairs are shared.
Determine the name of molecular structure Determine the name of molecular structure from positions of the atoms.from positions of the atoms.
Molecular Shapes: Linear
• When only 2 atoms are connected, the only possible shape is a straight line, linear.
• Ex: Diatomics
Br Br
Molecular Shapes: Lone Pairs • When lone pairs of electrons are present
on the central atom, they may shift the shape of the molecule. These are pairs of electrons that are not involved in a bond.
O HH
• They occupy space and provide negative repulsion forces against other electrons. They can repel more strongly than a bonded pair.
Molecular Shapes: Repulsion
• Electrons want to spread out around the central atom.
• They want to be as far away from each other as possible, due to negative-negative repulsion.
• They want to maximize their distance.
e-e- e-e-
Molecular Shapes: Bent • In water, the shape formed
in called bent, which is a variation of a tetrahedron. A tetrahedral molecule has 1 central atom and 4 attached. The bent molecule has 1 central atom and 2 attached.
O
HH
• The other 2 places are left “vacant” for the 2 lone pairs of electrons on oxygen, which occupy space and cause repulsion on the 2 bonded hydrogen atoms & cause them to be pushed downward, or away.
Molecular Shapes: Linear
• Not all molecules with 3 atoms will have the bent shape.
• Carbon Dioxide has 2 double bonds; one to each oxygen. This gives CO2 a linear shape by placing the double bonds as far apart as possible.
O C O
Molecular Shapes: Pyramidal• When you have 1 central atom with 3 attached,
there are 2 possible arrangements.
• In ammonia, NH3, there are 3 hydrogen atoms bonded to a central nitrogen. Due to the lone pair on nitrogen, the hydrogen atoms are pressed downward to form the pyramidal shape.
N
HH
H
Molecular Shapes: Trigonal Planar• The other possible arrangement for 1 central atom
with 3 attached is called trigonal planar.
• In BCl3, the 3 chlorine atoms are bonded to a central boron atom. However, the boron atom does NOT have a lone pair. Therefore, there is not additional pressure placed on the 3 chlorine bonds, so they spread out equally around the central atom.
B
Cl Cl
Cl
Molecular Shapes: Tetrahedral
• When 4 atoms are bonded to 1 central atom, the shape is called tetrahedral.
• In methane, CH4, 4 hydrogen atoms are bonded to 1 central carbon.
CHH
H
H
Table of Molecular Shapes# of
AtomsStructures
Bond Angles
Shape Name
2 180° Linear
3 105° Bent
4 120°Trigonal Planar
4 107° Pyramidal
5 109.5° Tetrahedral
Table – VSEPR Structures
VSEPR & The Water Molecule
VSEPR & The Ammonia Molecule
VSEPR & Xenon Tetrafluoride
Which one will it be???
VSEPR & Phosphorus Hexachloride
How to Determine Polarity of a Molecule
• Draw the molecule’s shape according to the Lewis Structure and VSEPR
• Fill in the EN values, and determine the difference for each bond
• Determine the bond type for each bond in the molecule
• Add in partial (δ) + and partial (δ) – symbols• Determine if there is a divisible axis to
separate the partial + from the partial -
Molecular Polarity • If your molecule has polar bonds, it may be a polar
molecule. • Polar molecules have specialspecial properties that result
from partial positive and negative ends of the entire molecule.
• Polar molecules can also be called dipoles. • These molecules are attracted to one another.
δ = partial
Molecular Polarity
• Water is a bent molecule, with 2 polar bonds.
• Since all of the negative charge is distributed on 1 side of the molecule, and all the positive charge is on the opposite side, water is a polar molecule.
O
HH
δ+
2 δ-
δ+
Molecular Polarity • Not all atoms with polar bonds are polar
molecules!
• Carbon Dioxide has all polar bonds. However, its geometry prohibits overall polarity since there are no distinct ends splitting the positive & negative charges.
O C O
δ- 2δ+ δ-
• CO2, therefore, is a non-polar molecule.
HybridizationHybridization
The Blending of OrbitalsThe Blending of Orbitals
We have studied electron configuration notation and the sharing of electrons in the formation of covalent bonds.
Methane is a simple natural gas. Its molecule has a carbon atom at the center with four hydrogen atoms covalently bonded around it.
Lets look at amolecule of methane, CH4.
What is the expected orbital notation of carbonin its ground state?
(Hint: How many unpaired electrons does this carbon atom have available for bonding?)
Can you see a problem with this?
Carbon ground state configuration
You should conclude that You should conclude that carbon only has carbon only has TWOTWO electrons available for electrons available for bonding. bonding. That is not enough for a full That is not enough for a full octet!octet!
How does carbon overcome this problem so thatHow does carbon overcome this problem so thatit may form four bonds?it may form four bonds?
Carbon’s Bonding ProblemCarbon’s Bonding Problem
The first thought that The first thought that chemists had was that chemists had was that carbon promotes one of its carbon promotes one of its 2s2s electrons… electrons…
…to the empty 2p orbital.
Carbon’s Empty OrbitalCarbon’s Empty Orbital
However, they quickly recognized a problem with such an arrangement…
Three of the carbon-hydrogen bonds would involve an electron pair in which the carbon electron was a 2p, matched with the lone 1s electron from a hydrogen atom.
This would mean that three of the bonds in a methane molecule would be identical, because they would involve electron pairs of equal energy.
But what about the fourth bond…?
Unequal bond energy
The fourth bond is between a 2s electron from thecarbon and the lone 1s hydrogen electron.
Such a bond would have slightly less energy than the other bonds in a methane molecule.
Unequal bond energy #2
This bond would be slightly different in character than the other three bonds in methane.
This difference would be measurable to a chemistby determining the bond length and bond energy.
But is this what they observe?
Unequal bond energy #3
The simple answer is, “No”.
Chemists have proposed an explanation calledHybridization.
Hybridization is the combining of two or more orbitals of nearly equal energy within the same atom into orbitals of equal energy.
Measurements show that all four bonds in methane are equal. Thus, we need a new explanation for the bonding in methane.
In the case of methane, they call the hybridization sp3, meaning that an s orbital is combined with threep orbitals to create four equal hybrid orbitals.
These new orbitals have slightly MORE energy thanthe 2s orbital…… and slightly LESS energy than the 2p orbitals.
sp3 Hybrid Orbitals
Here is another way to look at the sp3 hybridizationand energy profile…
sp3 Hybrid Orbitals
While sp3 is the hybridization observed in methane,there are other types of hybridization that atoms undergo.
These include sp hybridization, in which one s orbital combines with a single p orbital.
Notice that this produces two hybrid orbitals, whileleaving two normal p orbitals
sp Hybrid Orbitals
Another hybrid is the sp2, which combines two orbitals from a p sublevel with one orbital from an s sublevel.
Notice that one p orbital remains unchanged.
sp2 Hybrid Orbitals
Predicting the Geometry of Hybridized Orbitals
Intermolecular Forces
• Forces of attraction between molecules• Generally weaker than bonds that join atoms in
molecules• Boiling point gives a rough estimate of
intermolecular forces• high bp = large attractive forces• low bp = small attractive forces
Molecular Polarity and Dipole-Dipole Forces
• Dipole- Created by equal but opposite charges that are separated by a short distance
• A dipole is represented by an arrow with a head pointing toward the negative pole and a crossed tail situated at the positive pole
PolarityPolarity A molecule, such as HF, that has a center of A molecule, such as HF, that has a center of
positive charge and a center of negative positive charge and a center of negative charge is said to be polar, or to have a dipole charge is said to be polar, or to have a dipole moment.moment.
+
FH
This symbol means “partial”
Polar Covalent Bonding
• δ = partial
Dipole-Dipole forces
• The negative region of one molecule is attracted to the positive region of another molecule
• A polar molecule can induce a dipole in a nonpolar molecule by temporarily attracting its electrons
Dipole-Dipole Dipole-Dipole AttractionsAttractions
Attraction between oppositely charged regions of neighboring molecules.
The Water Dipole
The Ammonia Dipole
Polarity• Water is a polar molecule. • (A molecule with partially
positive & negative ends due to differences in electronegativity)
• Its electrons are shared unequally so it is a dipole.
• Causes strong intermolecular attractions.
Hydrogen Bonding
• The intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom (F, O, N) is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule
• It is usually represented by dotted lines
Hydrogen Bonds
• When the partial positive, or hydrogen end of a molecule is attracted to the partial negative, or oxygen end of another molecule we call it a Hydrogen bond.
• It is an attraction! • Hydrogen, and Fluorine, Oxygen
& Nitrogen. (F-O-N)
(high electronegativity values)
H-Bonds
• Diagram of H-bonds
• Solid lines are normal bonds
• Dotted lines are H-bonds
H-Bonds• Molecules containing
hydrogen bonds, like water, have very high points.
• Why might this occur?• Hydrogen bonds are
strong -holds molecules together.
• It takes a lot of energy to break apart these attractions to liberate each molecule into the gaseous state.
Hydrogen BondingHydrogen Bonding
Hydrogen bonding is used in Kevlar, a strong polymer used in bullet-proof vests.
Bonding between hydrogen and more electronegative neighboring atoms such as oxygen and nitrogen
Hydrogen Hydrogen Bonding Bonding in Waterin Water
Hydrogen Bonding Between Ammonia & Water
London Dispersion ForcesLondon Dispersion Forcesaka Van Der Waals Forcesaka Van Der Waals Forces
The temporary separations of charge The temporary separations of charge that lead to the London force that lead to the London force attractions are what attract one attractions are what attract one nonpolar molecule to its neighbors.nonpolar molecule to its neighbors.
Fritz LondonFritz London1900-19541900-1954
•All molecules experience London All molecules experience London forces forces •London forces increase with the size London forces increase with the size of the molecules.of the molecules.
• London forces are the only forces of attraction London forces are the only forces of attraction among noble-gas atoms, nonpolar, and slightly among noble-gas atoms, nonpolar, and slightly polar moleculespolar molecules
London Forces in HydrocarbonsLondon Forces in Hydrocarbons
Comparison of Boiling Points & Bond Types
Relative Magnitudes of ForcesRelative Magnitudes of ForcesThe types of bonding forces vary in their The types of bonding forces vary in their strength as measured by average bond strength as measured by average bond energy.energy.
Covalent bonds (400 kcal)
Hydrogen bonding (12-16 kcal )
Dipole-dipole interactions (2-0.5 kcal)
London forces (less than 1 kcal)
Strongest
Weakest