Chapter 14 Liquids and Solids

Phase changes and temperature  Normally when heat is added the temperature goes up.  However when you hit a phase change point (melting/freezing, boiling/condensation)…  The temperature stays constant when heat is added (until the phase change is complete).

Why should you…  Turn the heat down once the water is boiling?  Recipes will always tell you to do this.  Heat the water to a boil. Add spaghetti, and turn the heat down.  Won’t your spaghetti cook faster if you turn the heat up?  No  The water can only get to 100 o C  Increasing the heat would increase how fast it boils off, but that water leaves.

So a graph would look like… Time vs. Temperature of water under constant heat temperature 100 o C 0 o C time melting point liquid boiling point gas solid

Changes in phase require energy  It takes more energy to completely turn water at 100 o C into steam than it does to take the same water from 0 o C to 100 o C.  It actually takes 10x more energy to convert 100 o C water to steam than it does to heat 0 o C water to 100 o C water.  Steam has a much higher heat energy content than 100 o water.  This is why steam burns are much worse than water burns (scalding).

Phase Diagram graphs  Phase changes normally occur with a temperature change.  However a change in pressure can also force a phase change.  Like the butane in a Bic lighter.  It is a liquid inside (higher pressure), but once released it is a gas (lower pressure).  No temperature change caused this

Terminology  Triple point is the point where the substance can exist in all three phases of matter. It is the meeting point of all three phases  Critical point is the temperature where no matter the pressure, the substance will always be a gas.

Phase Diagram Graph of H 2 O Temperature Pressure 0 o C 100 o C 101 kPa Gas Liquid Solid normal melting point Normal boiling point Triple point Critical point Water is odd since the liquid is more dense than the solid. This line normally veers the other way.

“Normal” Phase Diagram Temperature Pressure Gas Liquid Solid Triple point Critical point

Why is water more dense than ice?  Intermolecular forces- forces of attraction between molecules that forces them to come together to form solids or liquids.  Intermolecular Forces are collectively called Van der Waals Forces.  Don’t confuse these with bonds which are intramolecular forces or forces that hold a molecule together.

The forces between these two water molecules are intermolecular forces. The bonds holding hydrogen and oxygen together are intramolecular forces

Phase changes  When intermolecular forces are strong enough to hold particles in place you have a solid.  As you increase the amount of energy in the particles, they break free of Van der Waals forces and start to move around some. This is a liquid.  When the atoms break free of all significant intermolecular forces they become a gas.

Dipole-Dipole Attraction  There are several intermolecular forces that we are not discussing.  One specific intermolecular force is dipole-dipole attraction.  Remember we said some molecules have a dipole moment or positive and negative ends.  A dipole-dipole attraction is when the molecules arrange themselves so that the opposite ends face each other.

Before Dipole-Dipole Attraction Cl H H H

Dipole-Dipole Attraction Cl H H Now the negative side (chlorine) is next to the positive side (hydrogen) Cl H

A really strong dipole-dipole force  A strong dipole-dipole force occurs when you have a molecules that have hydrogen bonding with nitrogen, oxygen or fluorine.  This is called hydrogen bonding.  The name is a misnomer, it is not an intramolecular force (regular bond), it is an intermolecular force.  It is much weaker than a regular bond, but stronger than the average intermolecular force.

Hydrogen bonding in water

Why is liquid water more dense…  Hydrogen bonding.  In solid water, the molecules can’t rearrange themselves.  In liquid water, they are capable of moving around.  Normally random movement would increase the spaces between molecules, but with hydrogen bonding the molecules “purposefully” move to a position where they can be pulled in closer.

London Dispersion Forces  ~A short lived induced dipole-dipole attraction between atoms that don’t normally have a dipole moment.  An orbital is an area of probability of an electron.  The electron does not have a uniform motion, at least we don’t think it does. We don’t know what the motion of an electron is

Probability  Flip a coin 2 times, should you get 1 heads and 1 tails?  Not really. There is a chance you will but, but it is only the highest probability.  You have a 50% chance of getting 1 heads 1 tails, 25% chance of getting 2 heads and 25% chance of getting 2 tails.  Apply that to our atom…

Helium Positive nucleus with 2 electrons in a 1 s orbital. It is neutral because the negative electrons cancel out the positive charges. Imagine a line cutting the orbital in half. What is the probability the electron is on either side? 50/50, but just like the coin flip should we always expect to find 1 electron on either side? No. However, if we don’t…

Still Helium 2 electrons on this side. None on this side. This side is now negative. This side is now positive.  +  - Put this atom near another atom...  +  -  +  - The positive side will attract the electrons, increasing the chance of poles forming again. It not only forces another atom to have poles, but the “new” atom forces the original to keeps its poles.

London Forces  This force is random and short lived, as the electrons do constantly move, and will eventually end the dipole moment.  It is also fairly weak.  You can tell it is really weak in helium because it stays a gas until o C.  Larger atoms or molecules (with more electrons) have stronger London forces.  With more electrons it is easier for the atom or molecule to have its electrons unbalanced and stay that way for an extended period of time.  Iodine (I 2 ) is a solid at room temperature.

Evaporation and Vaporization

Evaporation  Evaporation is a change in phase from liquid to gas, but is not the same as vaporization!  Vaporization requires you to heat the substance to its boiling point.  Evaporation can happen at much lower temperatures.  Volatility- A measure of how easily a liquid evaporates.

Is the vapor above the boiling point?  No it is not! (water vapor is not +100 o C)  It is possible to get matter in a phase that its temperature does not agree with.  It is like a solution (dissolved water in air)  It is also possible to get liquids above or below their freezing points. (supercooled or superheated liquids)

Evaporation works like this Liquids have molecules moving around in them Temperature is the average kinetic energy (which depends on the speed) of these molecules. Some are moving faster than others! If they are moving fast enough, at just the right angle, some will escape the surface of the liquid and turn into a gas. These evaporated! Molecules are held in by intermolecular forces.

Why are they a “gas”  Intermolecular forces determine whether something is a solid, liquid or gas.  In order to have intermolecular forces you need to have multiple particles.  The ones that escaped aren’t next to any other particles.  Since they have almost no intermoleluar forces they have to be a gas.

Where did they go?  They are in the air around the liquid.  They are called vapors, anything that naturally is a solid or liquid under standard conditions that is currently a gas at standard conditions.  If enough of them get together they will condense and reform a liquid.  As more of the molecules evaporate and fill the air around the liquid, the chance that some of them may condense increases.  Provided the vapors can’t escape, the liquid will reach a state where the rate of condensation and evaporation equal each other.

Evaporation works like this Liquids have molecules moving around in them Temperature is the average kinetic energy (which depends on the speed) of these molecules. Some are moving faster than others! If they are moving fast enough, at just the right angle, some will escape the surface of the liquid and turn into a gas. These evaporated! Molecules are held in by intermolecular forces.

Why are they a “gas”  Intermolecular forces determine whether something is a solid liquid or gas.  In order to have intermolecular forces you need to have multiple particles.  The ones that escaped aren’t next to any other particles.  Since they have almost no intermoleluar forces they have to be a gas.

Where did they go?  They are in the air around the liquid.  They are called vapors, anything that naturally is a solid or liquid under standard conditions that is currently a gas at standard conditions.  If enough of them get together they will condense and reform a liquid.  As more of the molecules evaporate and fill the air around the liquid, the chance that some of them may condense increases.  Provided the vapors can’t escape, the liquid will reach a state where the rate of condensation and evaporation equal each other.

Vapor Pressure  Vapor Pressure of a substance is the pressure of the vapor required for the rate of evaporation and condensation to be the same.  At this pressure the substance will reach a dynamic equilibrium.  Dynamic means changing, equilibrium means staying the same.  At a molecular level, constantly molecules are evaporating and condensing. However, since these cancel out, there is not net change.

After a liquid evaporates  The remaining liquid is cooler.  This is because the molecules with the most kinetic energy (heat) escaped.  Water has a “cooling” effect because it evaporates.  Sweat cools your body by evaporation.  Provided it is not humid out.  Humidity is a measure of the amount of water vapor present in the air.

Muggy (humid) weather  In humid weather, the water vapor in the air is closer to its vapor pressure.  Less net water can evaporate, and cool you off.  The rate of evaporation hasn’t changed, but more water vapor is condensing than normally.

Increasing Evaporation  Intermolecular forces play a big part.  Low molecular forces mean the substance will easily evaporate. These substances are volatile.  Evaporation occurs at the surface of a liquid so increasing the surface area will increase the rate of evaporation.  Allow evaporated vapors to escape so it can’t reach vapor pressure.  Heat the substance to increase kinetic energy.

Why do fans/wind feel cool?  The majority of the water vapor from your sweat is directly around you.  A fan or wind pushes air from somewhere else over to you, and the air that was around you somewhere else.  The water vapor that evaporated can’t condense back on you.  This only works if it isn’t extremely humid out.  If it is extremely humid the air from somewhere else contains a lot of water vapor that will condense on you.

Vaporization or Boiling  Evaporation occurs at the surface of a liquid.  As you continually heat a liquid, the particles inside move faster.  Eventually the particles move so quickly, that they break free of all intermolecular forces and form gas pockets inside of liquid.  These are always less dense than the liquid so the float to the surface and escape.  This is vaporization or boiling.

Evaporation and Vaporization Evaporation occurs at the surface If I get the substance hot enough

Evaporation and Vaporization I can force gas bubbles to form in the middle of the liquid This is vaporization or boiling

Types of Solids

Solids  Crystalline Solids- have a regular repeating arrangement of their particles.  Salts, Sugars, Metals  Amorphous Solids- have no regular repeating arrangement of their molecules  Common glass, several polymers.

Crystalline Structure

Amorphous

Amorphous solids  Amorphous solids, due to a lack of arrangement of molecules, will actually flow, slowly.  If you look at very old windows, you will find there is more glass at the bottom than at the top. That is because the glass flowed down.  You can also see the same effect with Silly putty.

Making solids…  Technically, anything can be made amorphous.  A rapid cooling from liquid to solid makes it amorphous. The particles just don’t have time to arrange themselves in a pattern.  A slower cooling or heat treatment can make some amorphous solids crystalline.

Safety Glass  Cars don’t use common glass for their windshield because it breaks into dangerous shard when it breaks.  Instead they use a heat strengthened glass, one that is slowly cooled to a solid to allow for a better arrangement of molecules, so that when it breaks it breaks into less dangerous “dice”.

Glass Safety Glass

Back to crystalline solids  Crystalline solids can be made up of 3 different things  Ionic Solids –made of ions  Molecular Solids- made of molecules held together by covalent bonds  Atomic Solids- Made of atoms

Ionic Compounds  Ionic Compounds have very high melting points.  Sodium Chloride melts at 801 o C  That is because every single negative particle is attracted to every single positive particle and vice versa.  This is in essence a very strong intermolecular force.

Ionic Solids  Ionic solids are brittle. When they break their crystal structure shows, as they break into similar shapes.  NaCl breaks into CaCl 2 into cubes spheres.

Conduction of electricity  Electricity is a flow of electrons  Anything that allows electrons to easily pass through will be a good conductor of electricity.  While solids, electrons can only jump from ion to ion.  This is a very slow process so solid ionic compounds are not good conductors.

Melts and solutions  If you melt an ionic compound, then the ions can move. Electrons can now easily move through the substance.  If you dissolve an ionic compound, the ions are also free to move.  Therefore, liquid ionic compounds and ionic solutions are good conductors.

Molecular Compounds  Molecular Compounds have much lower melting points.  Several are liquids (water) or gases (carbon dioxide) at room temperature.  Molecular compounds are not good conductors of electricity.

Atomic Solids/Elements  Solid nonmetals and metalloids commonly form very large molecules.  A diamond (any size) could actually be viewed as one molecule of all carbon.  These solids are called network solids.  They have high melting points and don’t conduct electricity.

Allotopes of Carbon

Nonmetal Gases  Noble gases and diatomic elements (except bromine, and iodine)  These all have only London dispersion forces.  These are very weak intermolecular forces.  They all have very low melting points, obviously since they are gases.  None are good conductors

Bromine and Iodine  These act the same as the other diatomic elements but since the atoms are larger the London dispersion forces are greater.  That is why they are a liquid (bromine) or a solid (iodine) at room temperature.

Metals  Metals have high melting points and are good conductors of electricity.  Metals are held together by metallic bonds.  Similar to ionic bonds these are somewhere in between intramolecular forces and intermolecular forces.

Metallic Bonding  Bonds between metals  Metallic bonds only occur with the same metal not with other metals.  Ca can bond with other Ca atoms, but not Ba.

Metallic Bond  In metallic bonds the valence electrons become community property, traveling anywhere they want to throughout the element.  This “Sea of Electrons” is why metals are such good conductors of electricity and heat.

Model of Metallic Bonds Ca The “sea of electrons” is kind of like bees (valence electrons) swarming around a few flowers (rest of the atoms). All of the electrons move like this. Calcium has 2 valence electrons

Properties  The nuclei inside the “sea of electrons” are movable without breaking the structure.  This is why metals are malleable and ductile.  Electrons can easily move through so they are great conductors of electricity.  Heat is the speed of the particles. If I heat up electrons at one end they quickly hit the slower moving ones and speed them up. So the whole material gets hot. That is why they conduct heat.

Alloys  ~a substance that is mixture of elements and has metallic properties.  Alloys are mixtures so they can be separated without chemical reactions  Steel is an alloy. It is made of iron and % carbon.  The carbon makes it harder, stronger, and less malleable than normal iron.  More carbon makes it stronger.

Interstitial Alloy  Steel is an interstitial alloy because the carbon atoms fit into the “holes” between the iron atoms in the crystal structure.

Substitutional Alloy  A substitutional alloy is when a metal atom of similar size replaces the host metal.  Brass (copper and zinc), sterling silver (silver and copper), white gold (gold, palladium, silver, and copper) are all substitutional alloys.  This changes the properties of the metal.

Both substitutional and interstitial alloys  Stainless Steel is iron and carbon (interstitial) mixed with chromium and nickel (substitutional).  It resists corrosion.  Slightly changing the presence of any of these drastically changes the properties of the final metal.