Thermodynamics,+Phases+of+Matter

 Thermodynamics 3 Laws - Definitions - Equations Phases of Matter Colligative Properties - Examples Intermolecular Forces - Solids - Liquids - Gases - Examples London Dispersion Forces - [|Dipole-Dipole Forces] (provided by Kyla) - Hydrogen Bonding Phase Diagrams - Phase Changes - Vapor Pressure Heat of Vaporization and Fusion - Heating and Cooling Curves Sources


 * __The 3 Laws of Thermodynamics__ **

1st Law of Thermodynamics ΔE = q + w The basic premise of the first law of thermodynamics is that energy is always conserved. The measurable heat (q) and work (w) must equal total energy change. If the system cools, q will be positive and the process is endothermic. If the system heats up, the process is exothermic and q is negative. Work, or w, is positive if work is done on the system and negative if the system does work. So, in summary, the total change in the energy of a system is equal to the sum of the heat and the work transferred between a sytem and its surroundings.

2nd Law of Thermodynamics ΔS = ∑(S° x coefficients) of the products – ∑(S° x coefficients) of reactants Change Entropy is found by subtracting the entropy of formation, found in a table in units of entropy for 1 mol of a substance, of the products from that of the reactants. Entropy of formation for pure elements is always zero. In a spontaneous process, the total entropy increases when you consider both the system and the surroundings.

3rd Law of Thermodynamics States that at zero Kelvin, a completely ordered crystalline solid has an entropy of zero. Although this is merely theoretical, this law also tells us some general guidelines. Liquids have higher entropies than solids, while the entropies of gases are even higher than those of liquids. Also, as temperature increases, the entropy of the substance also increases.


 * __Definitions__ **

ΔG – Gibbs Free Energy. A measure of the energy available in a reaction, Gibbs free energy tells one if the reaction is spontaneous. If ΔG is negative then the reaction is spontaneous, if it is positive it is non-spontaneous. The important equation to find free energy is: ΔG ° = Δ H ° – T( Δ S ° )

ΔS – Entropy or disorder of a system. This is based on probability. Nature generally moves spontaneously from low to high probability. For more on entropy, see the third law of thermodynamics (above). For more on the spontaneity of reactions see ΔG (directly above).

The standard molar entropies (standard entropy per mole) for gases are usually higher because heat of melting and heat of vaporization must be included. The standard molar entropies for noble gases are: code He    Ne     Ar     Kr     Xe (all in gaseous state) 126.0 146.2  154.7  164.0  169.6 J (K mol)-1

code Note that the entropy increase as the atomic mass increase. The same trends is also found for the halogens, but the entropies for these diatomic gases are much greater than those of monoatomic noble gases. code H2    N2      O2  (all in gaseous state) 130.6 191.5  205.0  J K-1

F2    Cl2    Br2     I2 (all in gaseous state) 203.7 222.9  245.4  260.6  J K-1

code Standard entropy of some compounds have also been measured. For example, code H2O(l)  H2O(g)   NH3(g)   H2O2(l)  CH3OH(l)  CH3Cl(l)  CHCl3(l) 69.9   188.7    192.5    110.0    126.9    145.3     294.9 J (K mol)-1

CO      CO2      NO      NO2    N2O4    SO2  (all in gases state) 197.8   213.6   210.6   240.4  304.3   248.4 J (K mol)-1

CH4  C2H2  C2H4  C2H6O4 (all in gases state) 186  201  221  230 J (K mol)-1

code (Submitted By John Bowman)

ΔH – Change in Heat, or enthalpy. Enthalpy is the heat content of a chemical. ΔH is the difference in heat content of the products and the reactants.

ΔE – The internal energy of a system. It is found by calculating the sum of a substance’s potential and kinetic energies.

Calorimeter – Calorimeters measure the heat energy produced by a chemical reaction or a physical change. Calorimeters isolate the system that the reaction or change occurs in, ensuring that there is no heat transferred to or from the surroundings and enabling accurate measurements. Essential components include an accurate thermometer and some form of a stiring device for the water that is often used as an insulator. The entire calorimeter heat and cool during the rocess, so that any accurate calculations of specific heats, etc must take into account the influence of the various parts of the calorimeter. A specific sub-group of calorimeters are known as bomb calorimeters, which do not allow the volume to change during the reaction.

Thanks to [] for the photo!

Heat capacity of a calorimeter – or C, is also known as the calorimeter constant. It represents the sum of of the specific heat and the mass of all parts of the calorimeter.

Specific Heat – a measure of heat energy, which is the amount of heat need to raise the temperature of one gram of a substance by one degree Celsius. One calorie a measure of the specific heat of pure water, with a temperature change from 14.5 to 15.5 degrees Celsius. One calorie is also equal to 4.184 J/(g x Celsius). All other substances specific heat can be calculated once the specific heat of water is known.

Work – the force applied to an object as it moves a certain distance. Force is the pressure exerted over a certain area. Therefore, Work = Pressure x Area x Distance Moved. In a chemical reaction, work done is the product of pressure and the change in volume.

Hess's Law – The ΔH of a reaction is the same whether the reaction occurs all at once or in a series of steps. Thus, if a thermochemical equation can be broken down into the sum of multiple equations, the total ΔH is equal to the sum of the ΔH's of each of the steps.

System – a part of the universe that is under observation. Types of systems include open, closed and isolated systems. In an open system, energy and matter are able to travel between the system and the surroundings. In a closed system, only energy can be transferred between the system and its surroundings. In an isolated system, there is no exchange between the system and the suroundings of either energy or matter.

Surroundings – everything else in the universe besides the system.

UTF – Utilization of The Forces (that apply to all things). An important scientific term, this needs no definition. Should it be used as work on any assignment, it should require no explanation and should recieve full credit.

ST – A complex subject that no one fully understands, ST occurs at any point in a lesson when you have no idea WTF is going on. This term is also an acceptable explantation for an answer with no work shown.

 **__Other Random Key Equations__**

q = m x c x Δ t

ΔH = ΔE + ( Δ moles of gas)RT

The Dulong and Petit Law: The Specific heat x the molar mass = 25 J/(mol x °C)

 ** Colligative Properties ** Properties of a solution that are dependent on the numbers of solute molecules and not the solute's identity.

  ** Examples ** Vapor Pressure Freezing Point Deppression Boiling Point Elevation [|Osmostic Pressure]

[|Boiling Point Elevation/Freezing Point Depression Video] Contribution by Camille Sharpe >

When a pure solvent and a solution are seperated by a membrane that only allows to pass through, the solvent will try to pass through, the solvent will try to pass through the membrane to dilute the solution. The pressure that must be applied to stop this process is called the osmotic pressure. The greater the concentration of solute in the solution, the greater the osmotic pressure. The equation for osmotic pressure takes a form that is similar to the ideal gas equation.

λ = (//n//RT/V)//i// = MRT//i//

λ = osmotic pressure (atm) //n// = moles of solute R = the gas constant, 0.0821 (L•atm)/(mol•K) T = absolute temperature (K) V = volume of the solution (L) //i// = the Van’t Hoff factor, the number of particles into which the added solute dissociates M = molarity of the solution (M //n///V)

 __** Intermolecular Forces **__ The attraction between different molecules that hold a substance together.  Solids have strong IF so they remain close together. When energy is applied (or pressure is removed) to solids they the IF are weakened and once enough energy has been applied the solid becomes a gas.  Liquids have weaker IF so the molecules have more "wiggle room." As with solids, if enough energy is applied (or pressure is removed) to a liquid, the IF weaken and the liquid becomes a gas. On the other hand, if enough energy is removed from (or pressure is applied to) a liquid the IF become stronger and it will become a solid.  Gasses have the weakest IF so the molecules are the most free to move around. As with liquids, if enough energy is removed from (or pressure is applied to) a gas, the IF with strengthen and it will become a liquid.
 * Solids **
 * Liquids**
 * Gases**

More specific examples of the phase changes can be found here.

 London dispersion forces, the weakest type of intermolecular forces, happen between all molecules no matter if they are polar or nonpolar. These forces are only temporary, and they are sometimes referred to as induced-dipole forces. A dipole is formed when the electrons in an atom are unsymmetrically distributed (positive and negative side). Since electrons are always moving, atoms can become dipoles temporarily at any point in time. When two atoms form temporary dipoles, the London Dispersion Force is the resulting attractive force between them. Since electrons repel each other, if a dipole exists in one atom, that atom can cause a second atom to become a dipole. These two atoms become electrostatically attracted to each other, and dispersion forces occur when they are almost touching. When the temperature is lowered enough, the phase changes from nonpolar substance to liquid (condensation) and liquid to solid (freezing) take place because of London Dispersion Forces. Stronger dispersion forces are found in bigger and heavier atoms. Larger atoms are more likely to form temporary dipoles based on the location of their valence electrons (further from the nucleus). Polarizability is how easy it is to change the electron distribution around an atom. Molecules that are easily polarized tend to have stronger London Dispersion Forces, and vice versa.
 * Examples **
 * __London Dipersion (Van der Waals) Forces__** (by Cheryl Leighton)

Dipole-Dipole Forces Can only form between Hydrogen and Nitrogen, Oxygen, or Fluorine.
 * __Hydrogen Bonding__ ** Dipole-Dipole Forces that are stonger than other Dipole-Dipole Forces.



media type="file" key="Intermolecular Forces.wmv" width="461" height="316" Video contributed by Sammy Brennan.  
 * __Phase Diagrams__ **

Go to [|this site] for the Phase Diagram Tutorial (Contributions by Shannon Black)

A phase diagram is a type of chart that shows the realtionship between a substance in different states of matter and the effect of temperature and pressure. A phase diagram is exclusive to each substance or element.

The triple point is the point on phase diagrams where. The triple point marks conditions at which three different phases exist at once.

The critical point is the temperature and pressure at which a substance has too much energy to condense back into a liquid.

At high pressure and a low temperature, the substance is a solid. At low pressure and a high temperature, the substance is a gas. The liquid state is between these two. The lines that seperate the states of matters are where the substance can exist in two phases at the same time. The line between solid and liquid represents melting and freezing. The line between liquid and gas represents vaporization and condensation. The line between solid and gas represents sublimation and deposition.

__Phase Diagram for Water __

[|Explanation for Water's Phase Diagram having a negative slope.]  Liquid to Solid- Freezing Liquid to Gas- Vaporization Gas to Liquid- Condensation Solid to Gas- Sublimation Gas to Solid- Deposition These changes occur because of pressure and temperature changes. Phase Diagrams show the relationship between temperature and pressure and their effect on phases of matter. As a substance changes from a solid to a liquid, the molecules become further apart and the intermolecular forces weaken. Molecules move even further apart when a liquid vaporizes to a gas and the IF become even weaker. On the other hand, when a gas condenses into a liquid, the distance between molecules becomes smaller and the intermolecular forces become stronger. The same happens when a liquid freezes to a solid. The only exception is water. When water freezes it forms a lattice structure. The molecules remain further apart while in the lattice structure, ice, than in a liquid phase. This accounts for ice being less dense than liquid water; therefore, ice floats. [|Click Herrrrr for more info on phase changes]
 * __Phase Changes__ **Solid to Liquid- Melting

[|Phase Diagram Tutorial] (Courtney S.)



Because there is no change in temperature and only the bonds are being broken, a flat line is indicated on a heating curve.

[|Phase change video] - contribution by Molly Robinson

States of Matter by Cynthia Nuñez http://www.youtube.com/watch?v=QppS3JbHA7I

 __** Vapor Pressure **__ Some particles have enough energy to break away from the liquid or solid phase they are in and become a gas. This gas exerts a pressure as it moves from the solid/liquid called vapor pressure. AS temperature increases, the vapor pressure increases. When vapor pressure is equal to the atmospheric pressure, the substance is at its boiling point.

Vapor pressure is also a colligative property. When a solute is added to a solution, the vapor pressure decreases. Because the vapor pressure is lowered, the boiling point of the solution is raised. This vapor pressure of a solution after a solute is added can be found using Raoult's Law. P=XP º P= vapor pressure of solution P º= vapor pressure of the pure solvent X= the mole fraction of the solvent media type="file" key="Small Chem.wmv" width="300" height="300"

 Heating curves show how the temperature of a substance changes as it is heated and how the phases also change. As heat is added, the kinetic energy is increased. The line flattens out between solid and liquid and is known as the melting point of the freezing point. It is a flat line becase all of the energy is out towards breaking the intermolecular forces and that is why there is no increase in temperature of the substance. The line flattens out again between the liquid and gas phases and this represents the boiling point. The melting point line is proportional the the heat of fusion. The boiling point line is proportional to the heat of vaporization. It is easier to change the temperature of a substance with a low specific heat, which is the amount of heat required to raise the temperature of one gram of a substance one degree Celsius. The lower the specific heat, the greater the slope of the heating curve.  A cooling curve shows the temperature of a substance as it is cooled down. Cooling curves also have flat parts where it turns from liquid to solid.  
 * __Heating and Cooling Curves__ **

__** Heat of Vaporization and Fusion **__

__Heat of Fusion__ is the energy given off by a substance when it freezes or the energy needed to melt a substance from a solid. The energy breaks the forces holding the solid together. Energy is released when a substance freezes because the intermolecular forces are more stable in a solid and have lower energy than the intermolecuolar forces of a liquid.

__Heat of Vaporization__ is energy put into a liquid to turn it into a gas, boiling, or the heat given off by a substance when it condenses. Intermolecular forces become stronger when a substance condenses.

Sources Baron's AP Chemistry 2008 4th Edition. Neil D. Jesperson.

Chemistry: Principles and Reactions 5th Edition. Masterson and Hurley.

The Princeton Review: Cracking the AP Chemistry Exam 2008 Edition. Paul Foglino

Chemisty SparkCharts. SparkNotes. 2008 Spaks Publishing.

[] - bomb calorimeter photo