chapter 10: structure of solids and liquids chem 1110 figures: basic chemistry 3 rd ed., timberlake...
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Chapter 10:Structure of Solids and Liquids
Chem 1110
Figures: Basic Chemistry 3rd Ed., Timberlake and Timberlake
Electron-Dot Symbols
Electron-dot symbols:
• Show the valence electrons of an atom
• Electrons are arranged in s and p orbitals around the element symbol
• Can mix orbitals and create “hybrid” orbitals with equal energy (like in Chapter 6)
o sp, sp2, and sp3
Covalent Bond Revisited
Covalent Bond is a chemical bond that results from two nuclei attracting the same shared electron pair
• Shared electrons in the bond
For Example: F2
Bonding and Lone Pair Electrons
Bonding Pair are shared to create the covalent bond
Non-Bonding Pair or lone pair make up the octet
Bonding and Lone Pair Electrons
Bonding Pair are shared to create the covalent bond
Non-Bonding Pair or lone pair make up the octet
H2
HCl
H2O
Rules for Drawing Electron Dot Formulas
1. Find the total number of valence electrons contributed by all atoms
2. Write chemical symbols in the order that they are bonded
3. Place one covalent bond between each atom
• Determining the central atom is important
• First atom is usually the central atom
o EXCEPT when it is hydrogen (H)
Rules for Drawing Electron Dot Formulas
4. Add lone pairs to the bonded atoms first, then to central atom
5. Check the Octet Rule for ALL atoms
• May need to make multiple bonds
For Example: SF2
Rules for Drawing Electron Dot Formulas
1. Find the total number of valence electrons contributed by all atoms
2. Write chemical symbols in the order that they are bonded
3. Place one covalent bond between each atom
4. Add lone pairs to the bonded atoms first, then to central atom
5. Check the Octet Rule for ALL atoms
• May need to make multiple bonds
Electron Dot Formulas
EXAMPLES:
CO2
NH3
SiCl4
Exceptions to the Octet Rule
Boron is stable with a total of 6 electrons
• Boron will make compounds with only three bonds to B
BCl3
Period 3 (and above) non-metals can form compounds with expanded octets
(> 8 valence electrons): P, S, Se
Multiple Bonds
Multiple bonds involve the sharing of more than one pair of electrons:
Bond Order
Bond Order denotes the level of bonding in a molecule:
• Single bond: Share one electron pair
Bond Order of 1
• Double bond: Share two electron pairs
Bond Order of 2
• Triple bond: Share three electron pairs
Bond Order of 3
Bond Length and Bond Energy
As we add additional bonds between atoms the atoms are drawn closer together:
• Bond length decreases with more bonds:
Single > Double > Triple
As we add additional bonds between atoms the bonds get stronger:
• Bond energy increases with more bonds:
Triple > Double > Single
Bond Length and Bond Energy
Practicing organic compounds:
• Single bond (alkane): Ethane, C2H6
• Double bond (alkene): Ethene, C2H4
• Triple bond (alkyne): Ethyne (Acetylene), C2H2
Electron Dot Formulas
EXAMPLES:
O2
HCN
CS2
CO
Resonance Structures
Resonance structures are equivalent structures that show where multiple bonds may occur
• Electrons are delocalized over entire molecule• Resonance changes bond order, bond length,
and energy
Examples:
SO2
SO3
Draw the three resonance structures of carbonate, CO3
2-
Learning Check
Electron Dot Formulas
How can we draw electron dot structures for polyatomic ions?
NH4Br or Na2SO3
These compounds have both an ionic and covalent:
[NH4]+ [Br]-
2 [Na]+ [SO3]2-
Electron Dot Formulas
[NH4]+
[SO3]2-
Draw the electron dot formula for ClO3-
Learning Check
Write two resonance structures for nitrite.
Learning Check
Orbital Geometry and Molecular Shape
• Electrons in bonds and lone pairs want to get as far apart from each other as possible
• We have seen the tetrahedral shape of carbon
• Causes the molecule to conform to a set of shapes
Determining Orbital Geometry
• Electron pairs do NOT like to share space with other electron pairs
• Electrons will move away as far as possible to obtain “personal space”
o An electron pair can be in a bondo An electron pair can be a “lone pair”
(non-bonding)
VSEPR Theory: Valence-Shell Electron Pair Repulsion Theory
Areas of electron density around an atom want to maximize their “personal space”
• A lone pair of electrons is ONE area
• Two electrons in a single bond are ONE area
• Four electrons in a double bond are ONE area
• Six electrons in a triple bond are ONE area
VSEPR Theory
Determining Orbital Geometry:
• Count ALL areas of electron density to determine the base geometry:
• Two is linear (sp)
• Three is trigonal planar (sp2)
• Four is tetrahedral (sp3)
Orbital Geometry Linear (sp)
Orbital GeometryTrigonal Planar (sp2)
Orbital GeometryTetrahedral (sp3)
Orbital Geometry
Linear and Trigonal Planar Molecules
Tetrahedral Molecules
Molecular Shape is Altered by Lone Pairs
Two different “bent” structures which differ in their bond angle: 120° vs. 109.5°
• Lone pairs take up more room and squeeze bond angles together:
oTrigonal planar: 120°oTrigonal planar with lone pair: < 120°
oTetrahedral: 109.5°oTetrahedral with lone pair: < 109.5°
Orbital Geometry and Molecular Shape
REMEMBER:
1) Electrons repel each other
2) Electrons in bonds have a fixed location
3) Electrons in lone pair take up more space
4) Two atoms always define a line.
Molecular Shape
Let’s look at two seemingly similar molecules:
CO2 and H2O
We first have to draw the molecule to determine their shapes:
Molecular Shape
We can base molecular shape on orbital geometry:
• Linear: CO2
• Trigonal Planar: BF3, CH2O
• Tetrahedral: CH4 , NH3, H2O
Learning Check
Determine the shape of the N2O molecule:
Electron Dot Formulas
Practice drawing electron dot structures and identify the proper molecular shape for:
AlH3
NF3
SiCl4
HCN
CO
In the preceding structures we had sharing of electrons (covalent bonds) and charged species (transfer of electrons, ionic bonds):
How do we know what bond type we have?
• General Rule of Thumb:
o A metal with a non-metal forms an ionic bond
o A non-metal binding with a non-metal forms a covalent bond
WHY????
Electronegativity
Electronegativity
The type of bond interaction is determined by differences in electronegativity:
Electronegativity is the ability of an atom to attract shared electrons in a bond towards itself
• Time to revisit Periodic Trends…
Electronegativity
Increasing Electronegativity
Incr
ea
sing
Ele
ctro
ne
gativ
ity
Fluorine is the MOST electronegative
Cs/Fr are the LEAST electronegative
Electronegativity Values
• We can assign values for electronegativity:
Place the following in order of increasing electronegativity: O, K, and C
Learning Check
Determining Bond Type
Difference in Electronegativity (ΔEN) allows us to determine the type of bond formed:
1. Difference > 1.8 → ionic bond
2. Difference = 0.0 → pure covalent bond
3. Difference between 0.0 and 0.4
• Non-polar covalent
4. Difference between 0.4 and 1.8
• Polar covalent bond
Common Pure Covalent Bond Species
O2 Cl2
Br2 I2
H2 N2
S8 P4
Bond Polarity
Bonds which have an unequal sharing of electrons are said to be Polar:
• The electron cloud is pulled toward the more electronegative atom
• This sets up a separation of charges: dipole
• The more electronegative element attracts the cloud more strongly - partially negative
• The less electronegative element attracts the cloud less - partially positive
Bond Polarity
Blue: low electron density
Red: high electron density
Green: non-polar
Bond Polarity
Bond Polarity
A different way of looking at polarity:
EXAMPLES
Molecule
NH3
O2
NaCl
SO2
Δ EN and Bond Type
0.9 → polar covalent
0 → pure covalent
2.3 → ionic
1.0 → polar covalent
Electronegativity and Bond Types
Predicting Bond Types
Use electronegativity differences to classify each of the following bonds as nonpolar covalent (NP), polar covalent (P), or ionic (I):
A bond between:
1) K and N
2) N and O
3) Cl and Cl
4) H and Cl
Learning Check
Drawing Bond Dipoles
K – N
N – O
Cl – Cl
Molecular Shape and Polarity
Dipoles can add across a molecule
• Polar molecules – have an overall dipole
• Non-polar molecules – no overall dipole
To determine whether a molecule is polar or non-polar, we must determine the 3-D shape of the molecule
Molecular Polarity
• Additive dipoles across a molecule make it polar:
• If dipoles cancel, no net dipole and the molecule is non-polar:
Molecular Polarity
EXAMPLES:
CO2
NH3
CF4
Learning Check
Identify each of the following molecules as
(P) polar or (NP) nonpolar:
A. PBr3
B. HBr
C. Br2
D. SiBr4
Learning Check
Draw the Lewis Structure for SeCl2 and determine its shape and molecular polarity:
Forces in Matter
Intermolecular Forces: are attractive forces between molecules:
• Weaker than bonds (covalent or ionic)
• Help to determine the state of matter:
solid, liquid or gas
Intermolecular forces organize matter and are opposed by motion of molecules which disorganize matter
States of Matter
Solid
Attractive Forces >> Disruptive Forces
Liquid
Attractive Forces ≈ Disruptive Forces
Gas
Attractive Forces << Disruptive Forces
Intermolecular Forces
Three primary types of forces that help to hold covalent molecules together:
1. London Dispersion Forces
2. Dipole – Dipole Interactions
3. Hydrogen Bonding
Intermolecular Forces
London Dispersion Forces occur in non-polar molecules by forming a temporary dipole:
• Temporary dipoles induce dipoles in nearby molecules
• Lasts for milliseconds!
• ALL matter has London Dispersion Forces
• Let’s consider H2 (non-polar molecule, H–H)
Intermolecular Forces
London Dispersion Forces
• Increase with increasing size• Increase with increased number of electrons
Examples: CH4, CF4, CCl4, CBr4
• CBr4 will have the strongest London dispersion forces as it is the largest
• CH4 will have the weakest London dispersion forces
Intermolecular Forces
London Dispersion Forces
• Increase with increasing size• Increase with increased number of electrons
o Electrons on larger atoms are further from the nucleus therefore easier to induce temporary dipole
Examples: CH4, CF4, CCl4, CBr4
• CBr4 will have the strongest London dispersion forces as electrons on Br are further from nucleus
Intermolecular Forces
Dipole-Dipole Interactions occur between polar molecules where δ+ end attracts the δ- end of another molecule
• There is a permanent dipole in polar molecules: attraction of molecules
• Occur only between polar molecules
• Consider FCl:
Fluorine is more electroneagtive (δ-)
Intermolecular Forces
Dipole-Dipole Interactions
• Stronger than London dispersion forces
• Strength of interaction increases with increasing size of the permanent dipole
For Example:
FCl (ΔEN= 1.0), FBr (ΔEN= 1.2), FI (ΔEN= 1.5)
• FI will have the strongest dipole–dipole interaction
• FCl will have the weakest dipole–dipole forces
Intermolecular Forces
Hydrogen bonds are a special type of dipole-dipole force when we have a VERY large difference in electronegativity
• To have a Hydrogen Bond:
1. Need to have H
2. Need to have F, O, or N
3. H must be bonded to F, O, or N
4. At least one lone pair on the F, O, or N
Intermolecular Forces
Hydrogen bond: The hydrogen on one molecule is attracted to the lone pair on another molecule
• The Hydrogen bond is the STRONGEST intermolecular force!
• Let’s consider H2O!
Hydrogen Bonds in Water
Hydrogen Bonds
Hydrogen Bonds
Predicting strength of H-bonds:
1) CH3OCH3 vs. C2H5OH
2) CH3NH2 vs. CH3NHCH3 vs. (CH3)3N
Learning Check
Identify the major type of attractive force in each of the following substances:
A. NCl3
B. H2O
C. Br2
D. KCl
E. NH3
Equilibrium
• When the forces acting on matter are in balance, we have Equilibrium
• Or, Equilibrium occurs when two or more opposing forces balance each other
Equilibrium: Balance of Forces
Equilibrium occurs when two or more opposing forces balance each other, and is a dynamic process:
• Represented by the symbol: ⇄
• For Example (Changes of State):
Solid Liquid⇄ Liquid Gas⇄
Melting point Boiling point
Changes of State
Intermolecular Forces
Strength of intermolecular forces increase as previously described:
London < Dipole-Dipole < Hydrogen Bond
• As the strength of intermolecular forces increases, the properties of molecules change:
• Want to “stick together” more!
• Polar molecules tend to be liquids or solids at room temperature
• Requires increased kinetic energy (disruptive forces) to separate them!
Intermolecular Forces
Strength of intermolecular forces increase as previously described:
London < Dipole-Dipole < Hydrogen Bond
• As the strength of intermolecular forces increases, they change the properties of molecules:
oMelting Point IncreasesoBoiling Point Increases oVapor Pressure Decreases
Learning Check
Identify the compound in each pair that has the higher melting point. Explain your choice:
A. NCl3 or NH3
B. HBr or Br2
C. KCl or HCl
Balance of Forces
Let’s consider H2O:
• Hydrogen bonds are the strongest force
• Requires lots of energy to break apart intermolecular interactions
• Consequently, water has an increased melting point, decreased vapor pressure, and increased boiling point
Liquid Water