Section 12-4 Section 12.4 Phase Changes (cont.) melting point vaporization evaporation vapor pressure boiling point Matter changes phase when energy is added or removed. freezing point condensation deposition phase diagram triple point

Section 12-4 Phase Changes That Require Energy: Melting, Vaporization, Sublimation Melting occurs when heat flows into a solid object. Heat is the transfer of energy from an object at a higher temperature to an object at a lower temperature.

Section 12-4 Phase Changes That Require Energy (cont.) When ice is heated, the temperature of ice will increase up until the ice reaches its melting point. At its melting point, the energy absorbed by the ice goes into breaking the hydrogen bonds that hold the water molecules together.

When the bonds break, the particles move apart and ice melts into water. This process will continue until all the ice is melted. The melting point of a crystalline solid is the temperature at which the forces holding the crystal lattice together are broken and it becomes a liquid.melting point

Section 12-4 Phase Changes That Require Energy (cont.) Particles with enough energy escape from the liquid and enter the gas phase.

Section 12-4 Phase Changes That Require Energy (cont.) Vaporization is the process by which a liquid changes to a gas or vapor.Vaporization Evaporation is vaporization only at the surface of a liquid.Evaporation Boiling: vaporization that occurs throughout the liquid.

In an open container of water, all the molecules of water will eventually evaporate. In a closed container, water vapor will collect over the liquid.

Section 12-4 Phase Changes That Require Energy (cont.) In a closed container, the pressure exerted by a vapor over a liquid is called vapor pressure.vapor pressure

Section 12-4 Phase Changes That Require Energy (cont.) The boiling point is the temperature at which the vapor pressure of a liquid equals the atmospheric pressure.boiling point At this point, molecules throughout the liquid have the energy to vaporize.

Section 12-4 Phase Changes That Require Energy (cont.) Sublimation is the process by which a solid changes into a gas without becoming a liquid. What do you know that goes through sublimation?

Section 12-4 Phase Changes That Release Energy: Freezing, Condensation, Deposition As heat flows from liquid water to the surroundings, the particles lose energy. As they lose energy, the molecules go back into their fixed positions. Freezing is the reverse of melting. The freezing point is the temperature at which a liquid is converted into a crystalline solid.freezing point

Section 12-4 Phase Changes That Release Energy (cont.) As water vapor loses energy, the velocity of the particles decreases. Water molecules bond, releasing energy as they do so. The process by which a gas or vapor becomes a liquid is called condensationcondensation Condensation is the reverse of evaporation.

Deposition is the process by which a gas or vapor changes directly to a solid, and is the reverse of sublimation.Deposition –Frost is an example of deposition. –Snowflakes are also formed in the upper atmosphere by deposition of ice crystals.

Section 12-4 Phase Diagrams A phase diagram is a graph of pressure versus temperature that shows in which phase a substance will exist under different conditions of temperature and pressure.phase diagram

Section 12-4 Phase Diagrams (cont.) The triple point is the point on a phase diagram that represents the temperature and pressure at which all three phases of a substance can coexist.triple point

Section 12-4 Phase Diagrams (cont.) The phase diagram for different substances are different from water. sXOM&feature=relatedhttp:// sXOM&feature=related

Heating Curve of a Liquid Kinetic Theory This type of graph is called a heating curve because it shows the temperature change of water as thermal energy, or heat, is added. Notice the two areas on the graph where the temperature does not change. At 0°C, ice is melting. At 100°C, water vaporizes.

Region (a) – part of curve where the solid is warming/cooling Region (b) – part of the curve where melting/freezing is occurring (temperature dose not change during a phase change) Region (c) – part of the curve where the liquid is warming/cooling Region (d) – part of the curve where vaporization/condensation is occurring (temperature dose not change during a phase change) Region (e) – part of the curve where the gas is warming/cooling

A.A B.B C.C D.D Section 12-4 Section 12.4 Assessment The addition of energy to water molecules will cause them to ____. A.freeze B.change to water vapor C.form a crystal lattice D.move closer together

A.A B.B C.C D.D Section 12-4 Section 12.4 Assessment The transfer of energy from one object to another at a lower temperature is ____. A.heat B.degrees C.conductivity D.electricity

Section 15-1 Section 15.1 Energy Define energy. temperature: a measure of the average kinetic energy of the particles in a sample of matter Distinguish between potential and kinetic energy. Relate chemical potential energy to the heat lost or gained in chemical reactions. Calculate the amount of heat absorbed or released by a substance as its temperature changes.

Section 15-1 The Nature of Energy Energy is the ability to do work or produce heat.Energy weightless, odorless, tasteless Two forms of energy exist, potential and kinetic. Potential energy is due to composition or position. Chemical potential energy is energy stored in a substance because of its composition. Kinetic energy is energy of motion.

Section 15-1 The Nature of Energy (cont.) The law of conservation of energy states that in any chemical reaction or physical process, energy can be converted from one form to another, but it is neither created nor destroyed.law of conservation of energy Heat is energy that is in the process of flowing (transferring) from a warmer object to a cooler object. q is used to symbolize heat.

Essentially all chemical reactions and changes in physical state involve either: release of heat, or absorption of heat Endothermic and Exothermic Processes

In studying heat changes, think of defining these two parts: the system - the part of the universe on which you focus your attention the surroundings - includes everything else in the universe Together, the system and it’s surroundings constitute the universe Endothermic and Exothermic Processes

Heat flowing into a system from it’s surroundings: defined as positive q has a positive value called endothermic system gains heat as the surroundings cool down Endothermic and Exothermic Processes

Heat flowing out of a system into it’s surroundings: defined as negative q has a negative value called exothermic system loses heat as the surroundings heat up

Section 15-1 Measuring Heat A calorie is defined as the amount of energy required to raise the temperature of one gram of water one degree Celsius.calorie Food is measured in Calories, or 1000 calories (kilocalorie). A joule is the SI unit of heat and energy, equivalent to calories.joule 1 calorie = J or 1 J = calories

Section 15-1 Measuring Heat (cont.)

Example: A candy bar has 245 Calories. Convert this to calories and then to Joules of energy. 245,000 cal 1.03 x10 6 J

Section 15-1 Specific Heat The specific heat of any substance is the amount of heat required to raise one gram of that substance one degree Celsius.specific heat Heat Capacity - the amount of heat needed to increase the temperature of an object exactly 1 o C Depends on both the object’s mass and its chemical composition

Section 15-1 Specific Heat (cont.) Calculating heat absorbed and released –q = c × m × ΔT –q = heat absorbed or released (in Joules) –c = specific heat of substance –m = mass of substance in grams –ΔT = change in temperature in Celsius

Examples: How much heat does a 20.0 g ice cube absorb as its temperature increases from (-27.0 o C) to 0.0 o C? Give your answer in both joules and calories. q = c × m × ΔT Specific Heat of Ice = 2.03 J/g o C 1 calorie = J Specific Heat (cont.)

Example Cont.  q = ?  c = 2.03 J/g o C  m = 20.0 grams  ΔT = Final Temp (0.0 o C) – Initial Temp (-27.0 o C) = Change (27.0 o C)  q = c × m × ΔT q = (2.03 J/g o C)(20.0g)(27.0 o C)= 1.10 x 10 3 J 263 cal Specific Heat (cont.)

Example 2: A 5.00 gram sample of a metal is initially at 55.0 ºC. When the metal is allowed to cool for a certain time, 98.8 Joules of energy are lost and the temperature decreases to 11.0 ºC. What is the specific heat of the metal? What metal is it? q = c × m × ΔT J/g ºC Iron

For Water during a phase change: –The Heat of Fusion (melting) is 334 J/g –The Heat of Solidification (freezing) is -334 J/g They are the same value (energy in or out) –The Heat of Vaporization is 2260 J/g –The Heat of Condensation is J/g They are the same value (energy in or out) Measuring Heat

For Water during a phase change: –The Heat of Fusion (melting) is 6.02 KJ/mol –The Heat of Solidification(freezing)is KJ/mol They are the same value (energy in or out) –The Heat of Vaporization is 40.7 KJ/mol –The Heat of Condensation is KJ/mol They are the same value (energy in or out)

The solid temperature is rising from -20 to 0 o C (use q = mass x ΔT x C) The solid is melting at 0 o C; no temperature change (use q = mass x ΔH fus. ) The liquid temperature is rising from 0 to 100 o C (use q = mass x ΔT x C) The liquid is boiling at 100 o C; no temperature change (use q = mass x ΔH vap. ) The gas temperature is rising from 100 to 120 o C (use q = mass x ΔT x C) The Heat Curve for Water, going from -20 to 120 o C, 120

Example – Phase change Calculate the amount of energy needed to convert 55.0 grams of ice to all liquid water at its normal melting point. (55.0g)(334J/g) = J Using the same amount of water calculate the energy needed to completely vaporize the water at its normal boiling point. (55.0g)(2260J/g) = J Why is there such a large difference in energy needed? Many more bonds to break when vaporizing

Section 15-2 Calorimetry Calorimetry - the measurement of the heat into or out of a system for chemical and physical processes. A calorimeter is an insulated device used for measuring the amount of heat absorbed or released in a chemical reaction or physical process.calorimeter Based on the fact that the heat released = the heat absorbed

Section 15-2 Chemical Energy and the Universe (cont.) Chemists are interested in changes in energy during reactions. Enthalpy is the heat content of a system at constant pressure.Enthalpy Enthalpy (heat) of reaction is the change in enthalpy during a reaction symbolized as ΔH rxn.Enthalpy (heat) of reaction ΔH rxn = H final – H initial ΔH rxn = H products – H reactants

Changes in enthalpy =  H q =  H These terms will be used interchangeably in this textbook Thus, q =  H = m x C x  T  H is negative for an exothermic reaction  H is positive for an endothermic reaction

Section 15-2 Chemical Energy and the Universe (cont.)

Section 15-2 Chemical Energy and the Universe (cont.)

The products are lower in energy than the reactants Thus, energy is released. ΔH = -395 kJ –The negative sign does not mean negative energy, but instead that energy is lost. Exothermic

The products are higher in energy than the reactants Thus, energy is absorbed. ΔH = +176 kJ –The positive sign means energy is absorbed Endothermic

Chemistry Happens in MOLES  An equation that includes energy is called a thermochemical equation  CH 4 + 2O 2  CO 2 + 2H 2 O kJ 1 mole of CH 4 releases kJ of energy. When you make kJ you also make 2 moles of water and 1 mole of CO 2

Thermochemical Equations  The heat of reaction is the heat change for the equation, exactly as written The physical state of reactants and products must also be given. Standard conditions (SC) for the reaction is (1 atm.) and 25 o C (different from STP)

CH 4(g) + 2 O 2(g)  CO 2(g) + 2 H 2 O (l) kJ  If 10.3 grams of CH 4 are burned completely, how much heat will be produced? g CH g CH 4 1 mol CH kJ = 514 kJ ΔH = -514 kJ, which means the heat is released for the reaction of 10.3 grams CH 4 Ratio from balanced equation 1 Start with known value Convert to moles Convert moles to desired unit