Liquids, Solids, and Intermolecular Forces Chapter 11

Intermolecular Forces (IMF’s) The state of a sample of matter (solid, liquid, gas) depends on the magnitude of intermolecular forces between the particles relative to the amount of thermal energy in the sample Remember, thermal energy is kinetic (motion, vibrations, wiggling)

Review Solids Liquids Gases

Liquids Properties of liquids Liquids have high densities in comparison to gases Liquids have indefinite shape and assume the shape of their container Liquids have a definite volume and are not easily compressed

Solids Properties of solids Solids have high densities in comparison to gases Solids have a definite shape and do not assume the shape of their containers Solids have a definite volume and are not easily compressed Solids may be crystalline (ordered) or amorphous (disordered)

Properties of the Phases of Matter A comparison of IMF’s Properties of the Phases of Matter Phase Density Shape Volume Strength of IMF’s Gas Low Indefinite Weak Liquid High Definite Moderate Solid STRONG

Q?? Why does pressure only affect the (l) and (g)?? Phase Changes You can change the phase of a sample by changing the temp, pressure, or both Add heat to a solid  liquid … add more heat (or increase the pressure)  gas Take away heat from a gas (or decrease the pressure)  liquid … take away more  solid Q?? Why does pressure only affect the (l) and (g)??

IMF’s The strength of the intermolecular forces between molecules or atoms determine the phase (s, l, g) Strong IMFs result in solids and liquids (high melting and boiling points) Weak IMFs typically result in gases (low Tm & Tb)

IMFs vs Bonds IMFs originate from the interactions between charges, partial charges, and temporary charges on molecules / atoms / ions Similar concept as bonds but are MUCH WEAKER… they are not the bonds holding a compound / molecule together, they are the forces holding molecule A close to molecule B

Strengths Coulomb’s law: E = 1 * q1q2 4πɛo r Bonds are strong because of large charged attractions at a close range IMFs are much weaker due to weaker charges at longer distances

Water evaporates at 100⁰C But, to break the O-H bond would take thousands of degrees

Different Types of IMFs Dispersion Forces (London Force) Dipole-Dipole Forces Hydrogen Bonding Ion-Dipole Forces Three of these can potentially occur in all substances; H-Bonds can be found only in mixtures

#1: Dispersion Forces (London Forces) Found in all molecules and atoms Caused by the fluctuations in electrons around the atoms / within the molecules Causes an instantaneous dipole (+ side / - side) Think about your turn signal in your car…

Dispersion Forces CAN INDUCE and stick to / attract to other neighboring molecules (IMF’s)

Magnitude of Dispersion Forces The magnitude of the force depends on two things: The size of the molecule (simply put: the Molar Mass) –the number of electrons The shape / structure of the molecule

1. Size Typically, the larger the molecule, the more electrons and the more free “roaming” around they are able to do. This opens the possibility of more electrons ‘clumping’ up and causing the +/- As the electron cloud increases, the dispersion forces increase… causing a _________ B.P.

2. Structure The smaller the area of interaction, the lower the forces will be The larger the area, the higher the forces

#2: Dipole-Dipole Forces The dipole-dipole force exists in all molecules that are polar Polar molecules have permanent dipoles that interact with perm dipoles of neighboring molecules The (+) end of one molecule interacts with the (-) of a neighboring molecule

Polar Molecules Remember, all molecules have dispersion forces. Polar molecules have Dipole-Dipole forces IN ADDITION to these dispersion forces This raises their melting and boiling points The strength (magnitude) of a dipole moment is measured in “Dipole moment” (D) The greater the D, the larger the dipole-dipole forces

“LIKES DISSOLVE / MIX WITH LIKES” Miscibility The polarity of molecules also determines the miscibility of liquids The ability to mix without separating into two phases All polar liquids are miscible with other polar liquids All nonpolar liquids are miscible with other nonpolar liquids “LIKES DISSOLVE / MIX WITH LIKES”

Immiscibility Opposites generally do not mix together i.e. nonpolar + polar

Practice Which of the following molecules have dipole-dipole forces? CO2 (b) CH2Cl2 (c) CH4 Draw the structure and determine if there are polar bonds present… try to see if the molecule is overall polar to have dipole-dipole forces

Practice Which of the following molecules have dipole-dipole forces? C I4 (b) CH3Cl (c) HCl

#3: Hydrogen Bonding Polar molecules with hydrogen can bond to F,N,O CAN make Hydrogen bonds: HF NH3 H2O H-bonds are STRONG for IMFs!! LEADS TO HIGH MP’S AND BP’S! Sort of the “super dipole-dipole” force

Why?? The large Δ in EN between H and N,O,F makes a very (+) end on the H and a very (-) end on the N, O, or F Also, because these elements are still quite small, they are able to get pretty close to one another, increasing the strength of attraction (coulomb’s law)

What's going on here??

WARNING!! H-BONDS ARE NOT BONDS!!! STILL ONLY ABOUT 2-5% THE STRENGTH OF A TYPICAL COVALENT BOND!!

Practice #1 One of the following is a liquid at room temperature. Which one? Formaldehyde (CH2O) Fluoromethane (CH3F) Hydrogen peroxide (H2O2) **TRY drawing them out!!

Explanation to Practice #1 All three have similar molar masses Meaning similar dispersion forces All three are also polar BUT… the hydrogen peroxide has H-O bonds So, there are hydrogen bonding in the H2O2 Note: the other two DO have H and N,O,F… but they need to be bonded directly to the H Why??

Practice #2 What has the higher boiling point, HF or HCl?? Why??

THESE ARE THE STRONGEST INTERMOLECULAR FORCES #4 Ion-Dipole Force The ion-dipole force occurs when an ionic compound is mixed with a polar compound Occurs when something ionic is put into water and dissociates into its ions THESE ARE THE STRONGEST INTERMOLECULAR FORCES

Ion-Dipole Force Think about NaCl in water The Na+ and Cl- ions break apart and float around The water SURROUNDS and attaches to the respective opposite charges

IMF’s in Action Intermolecular forces is the reason and determination of phase (s, l, or g) They also are the cause of other very important (for life) properties Surface Tension Viscosity Capillary Action

Surface Tension Why does a fly fishermen’s fly float on the surface of the water… the hook is metal and more dense The Surface Tension of a liquid is the energy required to increase the surface are by unit amount Water has a surface tension of 72.8 mJ/m2 Meaning it takes 72.8 mJ of energy to increase the surface area of water by one square meter

Why?? Liquids always like to minimize their surface area This creates kind of a “skin” on the top of the liquid that resists penetration For the fishermen’s fly to sink it would have to slightly increase the surface area of the water… which is resisted by the water

Surface Tension and IMFs Surface Tension decreases with decreasing IMFs You can float a paper clip on water but cannot float a paper clip on benzene (a nonpolar organic solvent) because there are no H-bonds AND the dispersion forces are not strong enough to combat the increase in surface area

Water Droplets Surface Tension is also the reason why water “beads” up and forms almost perfect spheres of water – in space (with no gravity), large amounts of water will ball up perfectly. Why? The sphere is the shape with the smallest ratio of surface are to volume… minimizing the number of molecules at the surface… minimizing the potential energy of the system

Viscosity Viscosity is the resistance for liquid to flow Motor oil is more viscous than gasoline Viscosity is greater in substances with stronger intermolecular forces because the molecules are strongly attracted to each other, preventing them from flowing past one another

Viscosity and Temperature How does temperature affect viscosity? Viscosity decreases with an increase in temperature The added kinetic energy will temporary overcome the IMF’s and allow them to move more freely past one another.

Capillary Action Capillary Action – the ability of a liquid to flow against gravity up a narrow tube Caused by two forces: Adhesion – Cohesion –

Capillary Action Capillary Action – the ability of a liquid to flow against gravity up a narrow tube Caused by two forces: Adhesion – the attraction between liquid molecules and the sides of the glass, root, etc. (to other things) Cohesion – the attraction between molecules within a liquid (to itself)

QUESTION: What is a meniscus?? Why does it happen??

Properties Affected by IMFs Phase Changes… Vaporization – the process by which thermal energy can overcome IMFs and produce a phase change from liquid to gas Removing energy will decrease their average kinetic energy and bring them back down as a liquid (condensation)

Open System Both vaporization and condensation occur at EQUAL rates (as some go into the gaseous phase, others fall back in the liquid phase) Average kinetic energy (some are lower than the average, others are higher) In an open beaker / container The evaporation occurs more… the gases escape into the atm and do not come back

Closed System In the same scenario, the beaker or bottle is not capped / sealed The two processes still occur at the same rate Reach a happy “dynamic equilibrium”

Other Liquids What happens if you increase the temperature? What happens to the vaporization process / temperature if the liquid is not water… but something with much weaker IMFs?

Volatility If a liquid evaporates easily it is volatile Those that do not evaporate easily are called nonvolatile Example: Acetone (nail polish remover) is more volatile than water. Motor oil is essentially nonvolatile at room temp Does not evaporate much or at all at room temp.

**Summary** The rate of vaporization increases with increasing temperature The rate of vaporization increases with increasing surface area!! The rate of vaporization increases with decreasing strength of IMFs

Endo- or Exo- ?? Vaporization vs. condensation Q: endo- or exo- ? If no additional energy is put into a system, what happens to the E as the liquid is let to evaporate?

Heat of Vaporization The amount of heat required to vaporize one mole of a liquid to gas is called the heat of vaporization (ΔHvap) The ΔHvap is always positive because the process is always endothermic It is temperature / pressure dependent

Heat of Vaporization The ΔHvap for water at 100 ⁰ C is +40.7 kJ/mol When a liquid condensates, that same amount of energy is released The ΔH for water condensing at 100⁰C is: ΔH = - ΔHvap = -40.7 kJ/mol (@100⁰C) (releases 40.7 kJ per mole of H2O)  (-) means leaves the system

Practice #1 Calculate the mass of water (in grams) that can be vaporized at its boiling point with 155 kJ of heat. 68.6 grams H2O

Practice #2 (pg 480) Calculate the amount of heat (in kJ) required to vaporize 2.58 kg of water at its boiling point.

More Practice (good TQ) Suppose that 0.48 g of water at 25⁰C condenses on the surface of a 55-g block of aluminum that is initially at 25⁰C. If the heat released during condensation goes only toward heating the metal, what is the final temperature (in ⁰C) of the metal block? The specific heat capacity of aluminum is 0.903 J/g ⁰C)

Vapor Pressure – Back to a SEALED System!! Once the liquid/gas is allowed to reach this dynamic equilibrium: The pressure of the gas with its liquid is called vapor pressure The vapor pressure depends on the IMFs!! The weaker the IMFs, the more volatile the liquid, the more molecules in the gaseous state, the higher the vapor pressure!!

Volume and Vapor Pressure What happens to the vapor pressure in these scenarios?

Systems WANT Equilibrium!! When a system in dynamic equilibrium is disturbed, the system will respond so to minimize the disturbance and reestablish its state of equilibrium

Practice What happens to the vapor pressure of a substance when its surface area is increased at constant temperature? The vapor pressure increases The vapor pressure remains the same The vapor pressure decrease B!! The VP is independent of surface area… with the increase in surface area DOES cause an increase in vaporization… but also an increase in the condensation!!

Vapor Pressure and Temp When the temperature of a substance is increased, the vapor pressure rises The thermal energy enables more molecules to break free and enter the gaseous state A small change in T will cause a big increase… but this curve relationship does level out and flatten What do we call this?? The boiling point!!!

Boiling Point The boiling point is the temperature at which its vapor pressure equals the external pressure. This is the temperature where molecules, EVEN THOSE IN THE MIDDLE, have the energy to break free (overcome the IMFs between them and their neighbors) …. Not just those on the surface bouncing and springing out

The Normal Boiling Point The normal boiling point is the temperature that the vapor pressure equals 1 atm (standard) This changes!! How? Increasing or decreasing the pressure (deviating from the P = 1 atm) will change this boiling point (no longer “normal” boiling point)

Supercritical Points Sealed system… but NOW increasing the Temperature The critical temperature (Tc) and critical pressure (Pc) are the conditions needed to cause this change – SUPERCRITICAL FLUID

Supercritical Fluids Supercritical Fluids have properties of both liquids and gases. Kind of an in between phase They are particularly good solvents… dissolving a large range of molecules. Supercritical CO2 dissolves caffeine, making it easy to be removed for de-caf

Phase Changes Involving Solids Think about a block of ice The thermal energy is significantly less than that of liquids or gases However, at the surface, molecules may still have enough energy to break free from their neighboring IMFs holding them together and float off into the gaseous phase What is this called?

Sublimation Sublimation - The molecules on the surface of a solid that have enough energy to go directly to the gaseous phase Some of the water vapor molecules may also collide with the surface of the ice and be ‘pulled’ in / captured by the IMFs of the solids… this process is known as deposition

Open System Sublimation usually occurs at a faster rate than deposition because the gaseous particles float away and never come back Think about ice cubes left in the freezer for too long… they shrink… why? Notice the gradual growth of ice crystals on the walls around the freezer??

Interesting Freezer burn is when foods are left in the freezer too long and dry out The frozen water molecules are left to sublime and never return

Heating Curves Thinking back to our block of ice… If we heat the cube up (from, say -10⁰C) the molecules move faster and faster and the temperature will increase However, you soon realize the temperature stops increase even though you are adding in more heat energy. What is happening?

Heating Curve for Water, H2O

Terms for Heating Curves Fusion – the phase change between solid and liquid (aka: melting and freezing) Vaporization – (we already know) is the change between liquid and gas (aka: evaporation and condensation)

Heat of Fusion (ΔHfus) The fastest way to cool a drink is by adding ice cubes to it They melt and pull absorb energy (endothermic) H2O(s)  H2O(l) ΔHfus = 6.02 kJ/mol Freezing is the opposite (exothermic) – releases energy to cool and freeze itself H2O(l)  H2O(s) ΔH = -ΔHfus = - 6.02 kJ/mol

Compare ΔHfus and ΔHvap Typically, the ΔHfus is much less than the ΔHvap This is because the solid and liquid phases are much closer together and require less energy to exhibit this phase change. Gases are much more energized and will usually require a much higher ΔH (in this case, ΔHvap)

Mathematical Relevance

Five Parts When there is NOT a phase change taking place, you can calculate the heat (q) using q = mCΔT Like Parts 1, 3, and 5

q = nΔHfus Like Part (2) q = nΔHvap Like Part (4) Five Parts When there IS A CHANGE taking place, you need to use the ΔH fus/vap Remembering, these values are PER MOLE q = nΔHfus Like Part (2) q = nΔHvap Like Part (4)

Phase Diagrams Phase Diagrams are maps of the phases of a substance as a function of pressure (y-axis) and temperature (on the x-axis)

Important Parts Regions – solid, liquid, or gas Lines – each line represents a T and P that the two phases are at dynamic equilibrium. The two phases are equally stable and coexist. The Triple Point – a unique set of conditions where all three phases are actually equally stable and in equilibrium The Critical Point – the set of P and T above which a supercritical fluid exists

Phase Diagram for CO2

Phase Diagram for Iodine

WATER IS SPECIAL!!