States of Matter; Liquids and Solids Chapter 13 States of Matter; Liquids and Solids

Phase changes Section 13.1

Comparison of Gases, Liquids and Solids Gases are compressible fluids. Their molecules are widely separated. Liquids are relatively incompressible fluids. Their molecules are more tightly packed. Solids are nearly incompressible and rigid. Their molecules or ions are in close contact and do not move. Vapors term customarily used for the gasesous state of a substance that exists naturally as a solid or liquid at 25 C and 1 atmosphere States of Matter

Energy Requirements for Phase Changes Water (solid) water (liquid) = 6 kJ/mol Water (liquid) water (gas) = 41 kJ/mol What does this mean???? Takes more energy to convert water from a liquid to a gas then to convert water from a solid to a liquid. WHY?? Solid and liquid states of water are more similar then the liquid and gas states

Section 13.1 Water and its phase changes FYI!!! water is unique in that as it cools it EXPANDS WHY??? Explain why ice floats? Why pipes burst in the Winter? How potholes form?

Phase Transitions H2O(s)  H2O(l) Melting: change of a solid to a liquid. Freezing: change a liquid to a solid. Vaporization: change of a solid or liquid to a gas. Change of solid to vapor often called sublimation. Condensation: change of a gas to a liquid or solid. Change of a gas to a solid often called deposition. H2O(s)  H2O(l) H2O(l)  H2O(s) H2O(l)  H2O(g) H2O(s)  H2O(g) H2O(g)  H2O(l) H2O(g)  H2O(s)

Energy of Heat and Phase Change Temperature does not change during the change from one phase to another. notice: there is either a temperature change OR a phase change. You cannot have BOTH at the same time Increase in temperature = increase in kinetic energy (after all the definition of temperature is average kinetic energy). There is no change in the potential energy Increase in heat = increase in potential energy. There is no change in the kinetic energy because the temperature does not change!!!!

Vaporization vs Boiling vs Evaporation Boiling For water to boil, it must be 100 Celsuis. Boiling creates an actual gas. The substance changes phase. Evaporation also involves liquids become gaseous. However, the body of liquid does not need to be at the boiling temperature. It occurs because the molecules of a liquid are not tightly bound together, and so some escape with time. Vaporization is a blanket term referring to both boiling and evaporation. In the broadest sense, it is liquid becoming gas.

Boiling point vs vapor pressure Boiling point the temperature at which the vapor pressure of a liquid is equal to the pressure of the external atmosphere. Normal boiling point the temperature at which the vapor pressure of a liquid is equal to atmospheric pressure (1 atm). Vapor Pressure the pressure of the vapor over a liquid at equilibrium in a closed container

Vapor Pressure If a liquid is placed in a non- closed container, some of the molecules will escape the surface and evaporate. In a sealed container, some of a liquid still evaporates but cannot “leave” the container. The molecules move back and forth between liquid and gas phases until they establish an equilibrium and thus a pressure in the vapor phase. Vapor pressure: partial pressure of the vapor over the liquid measured at equilibrium and at some temperature.

Temperature Dependence of Vapor Pressures The vapor pressure above the liquid varies exponentially with changes in the temperature. What does this graph indicate about the relationship between vapor pressure and temperature?? There is a _______ relationship between vapor pressure and temperature. In other words as vapor goes UP the temperature goes _______ or as vapor pressure goes DOWN the temperature goes _______

Phase Diagrams Section 13.1

Phase Diagrams Graph of pressure-temperature relationship; describes when 1,2,3 or more phases are present and/or in equilibrium with each other. Lines indicate equilibrium state two phases. Triple point- Temp. and press. where all three phases co-exist in equilibrium. Critical temp.- Temp. where substance must always be gas, no matter what pressure. Critical pressure- vapor pressure at critical temp. Critical point- point where system is at its critical pressure and temp.

Phase Diagram

Phase Diagram of Water

Endothermic vs Exothermic Reactions Section 3.6

Temperature, energy and heat Temperature = average kinetic energy Energy – the ability to do work Heat – the transfer of energy Energy and Reactions Energy must be added to break bonds. Many forms of energy can be used to break bonds: heat - electricity sound - light Forming bonds releases energy. Example: When gasoline burns, energy in the form of heat and light is released as the products of the isooctane-oxygen reaction and other gasoline reactions form. Energy is conserved in chemical reactions. Chemical energy is the energy released when a chemical compound reacts to produce new compounds. The total energy that exists before the reaction is equal to the total energy of the products and their surroundings.

Exothermic vs Endothermic Reactions An exothermic reaction is a chemical reaction in which heat is released to the surroundings. An endothermic reaction is a chemical reaction that absorbs heat. The graphs to the right represent the changes in chemical energy for an exothermic reaction and an endothermic reaction.

Catalysts Catalysts provide alternative pathways for a reaction, usually with a lower activation energy. With this lower energy threshold, more collisions will have enough energy to result in a reaction. An enzyme is a large organic molecule that folds into a unique shape by forming intermolecular bonds with itself. The enzyme’s shape allows it to hold a substrate molecule in the proper orientation to result in an effective collision. The rate of a chemical reaction is the change in the amount of reactants or products in a specific period of time. Increasing the probability or effectiveness of the collisions between the particles increases the rate of the reaction.

Calculating Energy Changes Section 13.2 and 3.7

Calculating energy changes Specific Heat Capacity: the amount of energy required to change the temperature of 1 gram of a substance 1 celsius degree (in j/g ◦C) Heat of vaporization: heat (energy) needed for the vaporization of a 1 mol of a liquid. H2O(l) H2O(g) DH = 40.6 kJ/mol Heat of fusion: heat (energy) needed for the melting of a I mol of a solid substance. H2O(s) H2O(l) DH = 6.02 kJ/mol EQUATION Q = sm∆t Q= energy required (joules) s = specific heat capacity (given in table page 70 in book) M = mass of water (grams) ∆t = change in temperature (Celcius)

Molar heat of fusion and vaporization

Example: calculate the energy (in kJ) required to heat 25g of water from 25C to 100C and change it to steam at 100C. The specific heat capacity of water is 4.18J/g and the molar heat of vaporization of water is 40.6kJ/mol Step 1: Q = s m ∆t = 4.18J x 25g x (100C – 25C) g = 7.8x103J Convert to kJ 7.8kJ Step 2: now use the molar heat of vaporazation to calculate heat energy required to vaporize 25g of water at 100C. Heat of vaporation is given in mols to we must convert 25g water to mols of water 25g H2O x 1 mol H2O = 1.4 mol H2O 18 g H2O Now calcualte energy need to vaporize the water 40.6 kJ x 1.4 mol H2O = 57kJ mol Step 3: TOTAL THE ENEGY NEEDED: 7.8kJ + 57 kJ = 65kJ

FYI One calorie is the amount of heat (or energy) needed to raise the temperature of 1 gram of water by 1°C. 4.184 joules = one calorie. 1,000 calories = 1 Calorie

Intermolecular Forces Section 13.3

Intra vs. inter molecular forces Intramolecular Forces: The attractive forces between atoms and ions within a molecule Ex: ionic bonds, covalent bonds STRONG Intermolecular Forces The attractive forces between molecules Ex: London dispersion forces , dipole-dipole forces and hydrogen bonding WEAK I.e. much less energy to melt H2O (inter) than for it to decompose into H2 and O2 (intra)

Comparison of Energies for Intermolecular Forces Interaction Forces Approximate Energy within Bond or Attraction between molecules Intermolecular London 1 – 10 kJ Dipole-dipole 3 – 4 kJ Hydrogen bonding 10– 40 kJ Chemical bonding Ionic 100 – 1000 kJ Covalent

London dispersion forces Aka: van der waals forces Attractive forces between ALL molecules temporary attractive force that results when the electrons in two adjacent atoms occupy positions that make the atoms form temporary dipoles Exist with nobles gases and nonpolar molecules, H2, N2, I2

Intermolecular Forces London Dispersion Forces Induced Dipole – Induced Dipole Weakest of all intermolecular forces. It is possible for two adjacent nonpolar molecules to affect each other. The nucleus of one molecule (or atom) attracts the electrons of the adjacent molecule (or atom). This attraction causes the electron clouds become distorted. In that instant a polar molecule (dipole) is formed (called an instantaneous dipole).

Forces of attraction between oppositely charge ends of polar molecules Dipole-dipole Forces of attraction between oppositely charge ends of polar molecules Line up + and - ends

Strong/Special type of dipole-dipole force IT IS NOT A BOND. Hydrogen bonding Strong/Special type of dipole-dipole force IT IS NOT A BOND. Between the positive H atom attached to an N, O or F and the negative N, O, or F of another molecule

London Dispersion < Dipole-Dipole < H-bonds Relative strength London Dispersion < Dipole-Dipole < H-bonds

Determining Intermolecular forces STEP ONE STEP TWO Does the compound contain N-H, O-H, or F-H Bonds??? NO = dipole- dipole and london dispersion forces YES = hydrogen bonding, dipole- dipole and london dispersion forces Is the compound polar or nonpolar??? Nonpolar = ONLY london dispersion forces Polar = go to step 2

CH4 is nonpolar: london dispersion forces. SO2??? Examples: What type(s) of intermolecular forces exist between each of the following ? HBr??? HBr is a polar molecule: dipole-dipole forces. There are also london dispersion forces between HBr molecules. CH4??? CH4 is nonpolar: london dispersion forces. SO2??? SO2 is a polar molecule: dipole-dipole forces. There are also london dispersion forces between SO2 molecules.

Determining Relative Boiling points Step ONE STEP TWO Determine the types of intermolecular forces Different forces = H > D > L Stronger force = higher boiling point b. Same forces = go to step 2 Look at the size of the compounds Larger molecules = higher boiling point

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