Chapter 11 Liquids and Intermolecular Forces
Multiply mc points by .56 to get an idea of how many bb points
Intermolecular forces
★The attractions between (inter) molecules are not nearly as strong as the intramolecular
(bonds, inside) attractions that hold compounds together
★Many physical properties reflect intermolecular forces
- Boiling points
- Melting points
- Viscosities
- Surface tension
- Capillary action
● The strongest force is the one that matters
● States of matter
Strength of intermolecular forces
★ A fundamental driver which gives the differences between states of matter
➔Solids have stronger intermolecular attraction
- The molecules are pulled closer together
- Stronger forces bring molecules closer together (on average)
- Kinetic energy keeps them apart and moving
- Average KE is related to temperature (in K)
Properties and states of matter
★ Solids and liquids are referenced to as condensed phases
- Relative strength of attractions (foreshadow)
➔ Intermolecular attractions are weaker than bonds (intra)
● FYI – hydrogen bonds are labeled as “bonds” but they are not
chemical bonds
- Types of intermolecular force between NEUTRAL (uncharged) molecules
- Weakest to strongest forces
- Dispersion forces (or London Dispersion Forces or induced dipole-induced dipole
interaction)
➔Everything pretty much has a London force
- Dipole-dipole forces
- Hydrogen bonding (a special dipole-dipole force)
The first 2 forces are also referred to as van der Waals forces
➔Dispersion forces
- A nonpolar particle (no dipoles) can temporarily be polarized to create
‘dispersion force’
- Can be spontaneously polarized
- The tendency of an electron cloud to distort is called it polarizability
➔ How dispersion forces arise
- If electrons move in one atom, then the electrons of another will respond to get its
electrons away from the newly negative charge
Molecules will always opt for the lowest energy state
Spontaneous changes on one molecule force a change on another
*Must be able to recognize which molecules have dipoles* *Lewis Structure does not go away*
4/9/2021
Chapter 11 Cont.
➔ Factors that affect the amount of dispersion forces in a molecule
- Size of atom or molecule/molecular weight
- Number of electrons in the molecule
- Shape of molecules with similar masses
More compact, less dispersion force
➔There’s less surface area, less contact area
- 2 linear molecules would have slightly more force between them than 2
compacted, spherical molecules
- Polarizability and boiling point (bp)
- If something is easier to polarize, it has a higher bp
There are stronger reactions when something is easy to polarize
- This means more intermolecular force between molecules
- Smaller molecule or atom: lower molecular weight, smaller size, fewer electrons
- List the substances CCl4, CBr4, CH4 in order of increasing boiling point CH4, CCl4,
CBr4
Dipole-Dipole Interactions
➔Polar molecules have a more positive and a more negative end – a dipole (two poles,
- Remember: partial charges; equal and opposite
You need to recognize structures to recognize dipoles
- The oppositely charged ends attract each other
- Stronger reaction than weak dispersion forces
Dipole forces are the next weakest
- For molecules of approximately equal mass and size, the more polar the
molecule, the higher its boiling point
How it works out in reality
- If 2 molecules have about the same size and shape, then dipole-dipole
interactions will likely be the dominating force
- It’s not a democracy, the strongest wins
- If one molecule is much larger than another, then dispersion forces
generally start more strongly influencing its physical properties
- BP trends
- In these specific groups, the period 3/4/5 elements have higher bp’s as the
group member gets larger
➔ Hydrogen bonding
- NOF
- These are dipole-dipole interactions experienced when H is bonded to N,
O, or F
They are unusually strong
➔ Stronger than dipole-dipole interactions
- Next step
- A hydrogen bond is an attraction between
- a H atom attached to a highly electronegative atom
- a nearby small sized, highly electronegative atom in another molecule or
chemical group
hydrogen bonds summarized
● hydrogen bonding arises in part from the high electronegativity of N, O, and F
● these atoms interact with a nearly bare hydrogen nucleus
ice compared to liquid water
● H bonding produces and anomaly in water
● H molecules actually makes the molecules farther apart in ice than in liquid water
Two consequences of H-bonding in water
- Weathering of stone and concrete/stucco
- Ice floats on liquid water, ice has the same mas, but greater occupied
volume than the equivalent mass of liquid water
- Survival of aquatic life during freezing cycles
Ion-dipole interactions
- Ion-dipole interactions are found in solutions of ions
- The strength of these forces makes it possible for ionic substances to dissolve in polar
substances, like water
- Ion-dipole interactions are the strongest forces we’ve looked at so far
A few molecular properties affected by intermolecular forces
● Boiling point
● Viscosity
● Surface tension (and capillary action)
Viscosity
- The resistance of a liquid to flow through a capillary or hole
- It is related to the ease at which molecules can move past each other
- Viscosity increases with stronger intermolecular forces and decreases with higher
temperature
Surface tension
● Water acts like a drum head
● Water acts as if it has a “skin” on it due to extra inward forces on its
surface
● It causes water to “bead up” when in contact with nonpolar surfaces
Phase changes
- Energy is either added or released during a phase change
- Melting/freezing (solid to liquid/liquid to solid), vaporizing/condensing (liquid to gas/gas
to liquid), subliming/depositing (solid to gas/gas to solid)
- Endothermic process
Energy added to a substance
- Exothermic process
Energy released from a substance
Energy change and change of state
- The heat of fusion is the energy required to change a solid at its melting
point to a liquid
- The heat of vaporization is the energy required to change a liquid at its
boiling point to a gas
- The heat of sublimation is the energy required to change a solid directly to
a gas (CO2 or I2)
Heating curves
- A graph of T vs heat added
- Within a single phase
Q=(mass)*Cs*deltaT
Remember that deltaT in K is the same as deltaT in Celsius
- The temperature of a substance does not change during a phase change
- For the phase changes, the products of moles (or mass) and heat of fusion or vaporization
is heat needed/given up while changing phase
Calculate the Heat required/released in a phase change
- deltaH = (mass)*(1/MW)*(deltaHvap)
Vapor pressure
- At any T, some liquid molecules have enough energy to escape the surface and become a
gas
- As the T rises, the fraction of molecules that have enough energy to break free increases
- As more molecules escape the liquid, the P they exert in the gas phase (headspace)
increases
P = vapor pressure
- The liquid and vapor reach a state of dynamic equilibrium
Liquid molecules evaporate and vapor molecules condense at the same rate
- The boiling point of a liquid is the T at which its vapor pressure equals atmospheric
pressure
- The normal B.P. is the T at which the vapor pressure is 760 torr (1atm; 760 mm Hg)
Vapor Pressure Recap
- As more molecules escape the liquid, the P they exert in the gas phase (headspace)
increases
- The liquid and vapor reach a state of dynamic equilibrium: liquid molecules evaporate
and vapor molecules condense at the same rate
Vapor Pressure Trend (Just for general knowledge, not on test)
- The natural log of the vapor pressure of a liquid is inversely proportional to its T
- This relationship is quantified in the Clausius-Clapeyron equation
lnP = (-deltaHvap)/RT + C
Supercritical Fluids
o Gasses usually liquify when sufficient pressure is applied
o The T above which a gas cannot be compressed into a liquid is called its critical
temperature (Tc)
o The P needed to compress the liquid at Tc is called critical pressure (Pc)
o The state beyond this T is called a supercritical fluid
No clear distinction of gas vs liquid phases exists for a supercritical fluid – it is its own
unique phase of matter
o 1 atm of pressure is room pressure
Phase Diagrams – we’ll do lots of these
o A graph showing states of matter under conditions of T and P
o It also shows changes of state and the “triple point” and critical point
Triple point is the point where the matter is at equilibrium, there’s an equal amount of
gas, liquid, and solid; all the phases are in contact with each other
o A 2-D map of phase changes
Looks kind of like a map with counties or towns outlined
- Temperature on x axis
- Pressure on y axis
Roadmap for phase change
Gasses are on far right, solids on the left, liquid is the area in between
o Phase Diagram of Water
Note the high Tc and Pc (due to the strong H bonding forces)
o Phase Diagram of Carbon Dioxide
Unusual features
Can’t exist in the liquid state at P below 5.11 atm (triple point)
CO2 sublimes at normal pressure
o Dry ice
Final Recheck on the only new math in Ch. 11
o deltaH = (mass)*(1/MW)*(deltaHvap)
MW = molecular weight
DeltaHvap or melt or whatever it is, is given, not memorized
Chapter 12 Solids (Covering only parts of Chapter 12)
Focus on the solid state
o The particles are packed together, usually in a regular pattern (like tiling a 2-D floor or
stacking similar sized oranges in a grocery display)
o We categorize solids by answering two questions
What are these particles?
Are they atoms, ions, or molecules?
What are the forces (or bonds) holding these particles together?
o By observation, two general categories
Hard, high melting point solids
Tmp > 200 degrees C
Completely interlocking strong connections
Soft, low melting point solids
Tmp < 200 degrees C
Independent units with weak connections
Broad classifications
o Metallic
Extended networks of atoms held together by metallic bonding (Cu, Fe)
o Ionic
Extended networks of ions held together by ion-ion interactions (NaCl, MgO)
o Covalent network
Extended networks of atoms held together by covalent bonds (C, Si)
o Molecular
Discrete molecules held together by intermolecular forces (HBr, H2O)
Characteristics of those two general categories of solids
o Hard, high Tmp solids
Particles are atoms or ions
Those particles are connected throughout by strong forces
Diamonds, iron, quartz, tungsten, salt
o Soft, low Tmp solids
Particles are molecules
Those particles are connected by weak forces
Water, carbon dioxide, wax, sugar, most plastics
Hard, high melting solids can be subdivided
o Covalent
Non-metal atoms
Covalent bonds
Diamond, the interlocking strong connections are covalent bonds
Covalent solids only contain non-metals, unfortunately so do molecular solids, so
keep an eye on the metalloid line, must be very close to the metalloid line (maybe
treat it like bio)
o Ionic
Ions
Electrostatic forces (ionic bonds)
Calcium Fluoride (CaF2), the interlocking strong connections are ionic bonds
Generally, a metal and a nonmetal, with some exceptions like salt and its derivatives
o Metallic
Metal atoms
“metallic bonds”
Tungsten, the interlocking strong connections are metallic bonds
Metallic bonds are like ionic bonds, only the electrons are delocalized (defined later
in the chapter)
Metallic solids are always easy to identify: they contain a metal
o Will need to predict the solids type from the chemical formula
Structure of Solids
o Amorphous
Without form (perhaps remember “amoral”)
o Crystalline
Repeating patterns
Terms
o Unit cell
A small part of what you’re looking at, the smallest set for molecules
o Crystal lattice
A lattice is a grid or reference frame
Having a lattice by itself does not define crystal structure
o Lattice points
o Vectors
Types of cubic solids
o Primitive
Lattice points only at the corners
o Body-centered
Lattice points at corners plus one lattice point in center of unit
Cell
o Face-centered
Lattice points at corners plus one lattice point at the center of each face
How the unit cell is “filled”
o Points must be identical
o Motif
Group of atoms; graphene
You can pack a motif, not just atoms
How is crystal structure known?
o X-ray diffraction in x-ray crystallography
Unit cell
o The exact smallest representation that shows all the atoms in their place
Metallic Solids
o Metals held together with ‘metallic bonding’
o Properties: malleable and ductile
o Cubic structures of metals most common
Primitive
Body-centered
Face-centered
Types of ways atoms occupy cells
o Location, # of unit cells sharing atom, fraction of atom within unit cell
o Corner, 8, ⅛
o Edge, 4, ¼
o Face, 2, ½
o Anywhere else, 1, 1
Consider what is “inside” a unit cell
o Primitive
8 corners*1/8=1 atom
o Body-centered
8*1/8+1 in middle = 2 atoms
o Face-centered
8*1/8+6 sides*1/2 = 4 atoms
Metallic bonding
- Quite different from nonmetals
- New concept – the “electron sea” model
“shortage” of valence electrons (for trying to make a complete shell) and collective
sharing of those actuals electrons present makes for close packing
Metal ion (nucleus and core electrons) surrounded by a sea of mobile valence electrons
o This is a molecular orbital (MO) model – and explains metal properties better You only
need to know this one qualitative aspect
Ionic Solids
o Electrostatic attractions hold together
o High melting/boiling points
o May be brittle
o If examined in detail we would need to remember ion six relationships
Anion>cations (usually)
Structures of Ionic Solids
o Principles
Symmetric, close packed
Lower coordination numbers than metals
Ions of like charge must not touch
Note that we are packing motifs
General Qualitative Patterns and Reasons
o How you pack depends on size of cation compared to the size of the anion
o Relative number of cations vs anions will also influence coordination number
Molecular solids
o Remember the two basic questions
o The soft, low Tmp solids all have molecules as particles, and various weak forces
holding them together. Because the particles are molecules, they are called molecular
solids
o Held together by intermolecular (weak) forces
o Molecules or atoms
o Low melting points
o Dry ice is a molecular solid
BP elevation and FP depression – the math
o The change in T is directly proportional to molality (using the van’t Hoff factor) the
small “m”
Matters
o The van’t Hoff factor is the number of particles a substance becomes when it dissolves
deltaTb = Tb(solution) – Tb(solvent) = iKbm
deltaTf = Tf(solution) – Tf(solvent) = -iKfm
Observations on Molarity versus Molality
o When water is the solvent, dilute solutions have similar molality and molarity
o Molality does not vary with temperature (because mass does not change)
o Molarity varies with temperature (volume frequently changes with temperature)
Reminder Grams/Liter
o 13.2 grams of glucose is dissolved in 410 mL
3.22 g/L
Remember to change mL to L
Reminder Mass percentage
o Percent means out of 100
o For solutions where the solute concentration is asked, take the ratio of the mass of the
solute to the total solution mass
o Multiply by 100 to make it a percent
5 grams of NaCl in 495 grams of water 1.008%?
Osmosis
o Some substances form semipermeable membranes, allowing some smaller particles to
pass through but blocking larger particles
o The net movement of solvent molecules from solution of low to high concentration of
solute across a semipermeable membrane is osmosis. The applied pressure needed to
exactly stop this flow is called osmotic pressure.
Osmotic pressure
o Osmotic pressure is a colligative property
π=i(n/V)RT=iMRT
Types of solutions and osmosis
o Isotonic solutions: the same osmotic pressure
Example – osmosis and RBC
o Red blood cells (RBC) have semipermeable membranes
o If stored in a hypertonic solution, they will shrivel as water leaves the cell; this is called
Crenation
o If stored in a hypotonic solution, they will grow until they burst
Reaction Rates
o The speed at which reactions take place is called the reaction rate
o Reaction rate = (change in component “A” with change in time)
o The study of reaction rates is called chemical kinetics
o A step-by-step view of the change of reactants to products is called a mechanism
Factors that affect reaction rates
o Physical state of the reactants (liquid, solid, gas)
o Reactant concentrations (M, Px)
o Reaction temperature (in K)
o Presence of a catalyst
Physical State of the reactants
o The more readily the reactants collide, the more rapidly they may be able to react
o Homogeneous reactions (all gasses, or miscible liquids) are often faster
o Heterogeneous reactions that involve solids are slower
Examples/Consider
o AgNO3(aq) + KBr(aq) AgBr(s)
o 2AgNO3(aq) + CaCl2(s) 2AgCl(s)
o 2CO + Pt(s) + O2(g) 2CO2 + Pt(s)
Reactant Concentrations
o Increasing reactant concentration will generally increase reaction rate
o Why?
Since there are more molecules, more collisions occur
o You no doubt notice this is a “kinetic” theory of reactions
Temperature
o Reaction rate generally increases with increased temperature (true no matter what T
units you use)
o Kinetic energy of molecules is related to temperature (KE up at T increases)
Ke = ½*mass*velocity2
o At higher temperatures, molecules move more quickly
Presence of a catalyst
o Catalysts affect reaction rate without being in the overall balanced equation
o Catalysts affect the underlying path by which the reactants reach the products
(mechanism; more later)
Catalysts in biological systems
o Bromelase in the stems of bromeliads
Reaction rate
o Rate is a change in concentration over a time period
Delta[x]/deltat
[] means molar concentration
o Types of rates measured
➔ Average rate (rate over total reactions)
➔ Instantaneous rate (at one given instant in time)
➔ Initial rate (rate at the very beginning of the reactions)
Following reaction rates
o Rate of a reaction is measured using the concentration change for a reactant or a
product over time
The rate equation for a generalized reaction
o 2A+BC
o Rate rxn = k*[A]x
[B]y
o Where
k is the “rate constant” for the reaction
The exponents are called the “order” or that reaction for that specific reagent
More about rate law
o The exponents tell the reaction order with respect to each reactant
o The order with respect to each reactant is q and the overall reaction is second order
(1+1=2) (from an example, add the order with respect to each reactant to get the
overall reaction order)
o k is a T dependent quantity (discussed in a few slides)
Reaction order does not equal reaction stoichiometry
o The order of the reaction must be determined experimentally, it is not necessarily
related to the balanced equation
Relative value of k
o The rate constant is often used to determine the relative rate of a reaction
o You should remember these rule of thumb
For reactions considered fast
- K is approximately 10^9 or higher
For reactions considered slow
- K is approximately 10^9 or lower
First order reactions
o Some rates depend only on one reactant to the first power
o These are first-order reactions
o The rate law becomes
Rate = k[A]
o The conversion of methyl isonitrile to acetonitrile is just a first order reaction
Relating k to [A] is a first-order reaction
o Rate = k[A]
o rate=-delta[A]/deltat
So: k[A]=-delta[A]/deltat
Rearrange to delta[A]/[A]=-kdeltat
Integrate to ln([A]/[A]0)=-kt
[A]0 is how much you started with
Rearranges to: ln[A] = -kt+ln[A]0
Memorize this equation and be able to use it
Note: this follows the equation of a line (y=mx+b)
o So, a plot of ln[A] vs t is linear
Finding the rate constant, k
o For this first order reaction we can find the rate constant from that plot of ln[A] vs time
o The correct plot will give a line, its slope will equal -k
Half-life
o The amount of time it takes for one half of a reactant to be used up in a chemical
reaction
(t1/2)
o For a first order reaction
Ln([A]/[A]0) = -kt
[A]/[A]0 = ½ at the halfway done point
Ln(1/2) = -.693, so .693 = k*t1/2
Second order reactions
o Some rates depend only on a single reactant to the second power
o These are second order reactions
o The rate law becomes
k[A]2
solving the second order reaction Rxn rate
o rate = - delta[A]/deltat
o 1/[A] = 1/[A]0+kt
o Notice the linear relationships that result for first-order and second-order reactions
differ
Half life and second order reactions
o Using the integrated rate law, we can see how t1/2 is derived for 2nd order rxns:
o MEMORIZE
1/(k*[A]0) = t1/2
.693 = k*t1/2
Zero-order reactions
o Occasionally, rate is independent of the concentration of the reactant:
Rate = k
o These are zero-order reactions
o These reactions are linear in concentration
[A]t=[A]0-kt
Factors that affect reaction rate
o T of the rxn
As T goes up so does rnx rate
o Frequency of collisions between reactants (T and conc)
o Orientation of reactant molecules relative to one another while colliding (more in O, P
and I Chem)
o “threshold” energy needed for the rxn to rake place (activation energy or Ea)
Frequency of collision
o The collision model is based on the kinetic molecular theory
o Molecules must collide to react
o So, if there are mole molecules, there are more collisions (reactions)
Temperature and rate
o Generally, as T increases, rate increases
o The rate constant is T dependent
o Rate constant doubles (very approximately) with every 10 degree Celsius rise (RULE
HEARD IN O. CHEM)
The relationship between Ea and T
o Arrhenius notes the relationship between Ea and T
k = Ae^(-Ea/RT)
make sure to use Kelvin when solving problems like this
o activation energy can be determined graphically by math transforming and rearranging
that equation
ln(k) = ln(A) – Ea/RT
orientation of molecules
o molecules can often collide without forming products
o because bonds must be broken and made, so atoms need to b ein aligned in the proper
positions for rxn to occur (O Chem and P Chem)
energy needed for a reaction to take place – Ea
o the minimum energy needed for a reaction to take place called activation energy (Ea)
o an energy barrier must be overcome for a rxn to take place, much like the ball must be
hit to overcome the barrier in the golf figure shown on the powerpoint transition state
(activated complex)
- reactants gain energy as the reaction proceeds until the particles reach
maximum energy state
- the organization of the atoms at this highest energy state is called the
transition state (or activated complex)
- the energy needed to form this state is called the Ea
reaction progress
- plots are made to show the energy possessed by the molecules as the rxn proceeds
reaction progress figure
- products on one side, reactants on the other
- potential energy on y axis, reaction progress on x axis
● at the highest energy state, the transition state is formed
● rxns can be endothermic or exothermic after this point
● rate constant “k” depends on the magnitude of Ea
Math
- Done on separate sheet, attached to notes
Law versus theory in chemical kinetics
- Why do we experimentally observe these rate laws and orders of rxn that have been
discussed?
- Theories that attempt to explain that Q are called “mechanisms”
- A mechanism is a series of stepwise reactions that show how reactants become products
Reactions mechanisms – things chemists have learned in the past
- Reactions may occur all at once or through several discrete steps
- Each of these processes (whether one step or many discrete steps) is known as an
elementary reaction or elementary process
Definition in kinetics – “molecularity”
- The molecularity of an elementary reaction tells how many molecules are involved in that
step
Molecularity Elementary Reactions Rate Law
Unimolecular A➝products Rate=K[A]
Bimolecular A+A➝products Rate=K[A]²
Termolecular A+B➝products Rate=K[A][B]
Termolecular A+A+A➝products Rate=K[A]3
Termolecular A+A+B➝products Rate=K[A] [B]
2
Termolecular A+B+C➝products Rate=K[A][B][C]
Chapter 14 cont.
What limits the rate?
- The overall reaction cannot occur faster than the slowest reaction in the mechanism
- We call that the rate-determining step (“RDS”)
- So, what is required of a plausible mechanism (multiple elementary rxns to reach desired
overall rxn)
- The rate law must produce what occurs when determining the RDS
- The necessary balanced stoichiometry must be obtained when all steps are added up
- Each step must balance, like any equation
- All intermediates are made and used up
Intermediates
- An intermediate is not a reactant or a product
- It is also not the transition state
- Intermediates are stable
- In some instances, it is even possible to isolate or identify and intermediate
Reaction Example
Cl2 (g) + CHCl3 (g) HCl (g) + CCl4 (g)
Cl2 (g) 2Cl (g)
Cl (g) + CHCl3 (g) HCl (g) + CCl3 (g)
CCl3 (g) + Cl (g) CCl4 (g)
What are the intermediates?
● 2Cl and CCl3
● Catalysts
- Catalysts increase the rate of a reaction by decreasing the activation energy of the
reaction
- Catalysts change the mechanism by which the process occurs
Homogeneous catalysts
- Things just mix
- Heterogeneous catalysts
Enzymes
Homogeneous catalysts
- The reactant and catalyst are in the same phase
- Enzymes
- Enzymes are biological catalysts
- They have a region where the reactants attach that region is called the active site the
reactants are referred to as substrates
Chapter 15 Chemical Equilibrium
Reversible reactions
- Both the Ea and the delta energy are small enough for the reverse reaction to occur
- The more input per hour the more product per hourA reversible reaction is like having a
second factory that undoes what the first factory did
- Notice that the forward and reverse reaction will be interdependent
Writing the equation for an equilibrium reaction
- Since, in a system at equilibrium, both the forward and the reverse reactions are being
carried out, we write its equation with a double arrow
N2O4 (g) 21cc<alt><x>
The concept of equilibrium
- As a system approaches equilibrium
