Physical Changes in State

 Problems:

 1. 72 grams of ice + 51 840 calories yield 72 grams of water vapor. How many calories must be removed from water vapor to condense it back to ice?

 2. If 1 gram of water at room temperature evaporates, about 600 calories are taken from the surroundings to convert the liquid to a gas. How many calories are `needed' to change 1001 grams of gas to a liquid?

 3. If 50 grams of water vapor loses 36 000 calories in turning to ice, how many calories would 1 gram of water vapor have to lose to be turned back to ice?

 Chemical Changes in State

 The molecular make-up (the specific arrangement of atoms) is changed, resulting in new substances being formed and energy changes occurring examples: oxidation - wooden splint burning --> heat, light, and gases like CO₂ and H₂O being given off with carbon and ashes left over electrolysis - splitting some compound (usually water) by running an electric current through it .

 Any chemical change that releases energy is EXOTHERMIC. examples: oxidation, burning H₂ in O₂, body reactions, dissolving Zn in HCI, mixing H₂O and H₂SO₄

 Any chemical change that absorbs energy is ENDOTHERMIC. examples: photosynthesis and electrolysis

 Exothermic reactions - the amount of heat released is greater than the amount of heat used to start the reaction

 Endothermic reactions - energy continues to be absorbed as long as the reaction continues

 The chemical change involving splitting or forming water takes about S times as many calories as the physical change of state. The reason is that atoms (or molecules) are bonded together in a compound; the stronger the bond the more energy holding the parts together, thus more energy required to break these bonds. A physical change needs far less energy to overcome intermolecular forces holding groups of molecules together. Much more energy is needed to break bonds within molecules than to overcome the forces between molecules.

 Physical changes -> strength of intermolecular forces increased or decreased

 Chemical changes -> bonds formed or broken

 Energy absorbed -> bonds broken or intermolecular forces overcome

 Energy released -> bonds formed or intermolecular forces strengthened

 Problems:

 Tell whether each of the following is a chemical or physical change and then add exothermic or endothermic for the chemical changes and energy absorbed or released for each physical change.

 1) dry ice vaporizes
2) CO₂ + H₂O  + sunlight  --> glucose
3) air in heated tire expands
4) burning coal
5) water frozen into ice
6) acid dissolves metal

Heating Curve Information:  This graph will aid in understanding the following information.

Phase Change Diagram

 This discussion deals with sublimation.  This is the direct change of a solid to a gas.  Examples to be discussed include the C0₂ fire extinguisher, moth balls (naphthalene), paradichlorobenzene, camphor, iodine.

 Liquid - Solid Phase Change

 This discussion deals with melting-freezing point. A complete discussion of this concept using ice and heat units will be completed in class.

See class discussion of ice cube. The addition of 1 calorie of heat to the ice cube at 0° C does not cause a change in the temperature of the ice cube though 1 calorie would change the temperature of 1 gram of water at 0° C.

It will take 80 calories just to melt the ice cube. That heat that is consumed in melting the solid is converted into potential energy. Freely moving molecules in liquids, with respect to intermolecular attraction, possess more energy than similar molecules bound rigidly in solids at the same temperature.

Remember that temperature is a measure of the average kinetic energy only while heat content is a measure of the total kinetic energy plus potential energy possessed by that body.

See class examples of the heat energy needed to change ice at any temperature to steam at any temperature.

 The melting-freezing point is defined as the temperature at which the solid and liquid phases are in equilibrium. This is the temperature at which a change of state between the solid and liquid phase can occur. Some of the solid will be melting and some of the liquid will be freezing.

When a solid is heated to its melting point, its atoms or molecules acquire enough energy to shift the bonds holding them together so they form separate clusters. This clustering in liquids is confirmed with X-ray studies but the clusters are constantly shifting their arrangement unlike the permanent arrangement in solids.  

When heat is added to a solid the temperature of the solid will increase till it reaches the melting-freezing point.   It will remain there until all the solid has melted and only then can the temperature of the liquid rise according to its specific heat.

Water molecules at 0° C contain more energy than the ice molecules at 0° C , not in the form of a faster more rapid motion but in the form of an ability to resist the attractive forces tending to pull them together.

Melting points also depend on pressure (though not as much as boiling points.) Ice is strange in that its melting point decreases with increasing pressure. Almost all other materials show increasing melting points with greater pressures. The pressure an ice skater exerts on the ice due to the small area of the skate blade is usually enough to melt the ice creating a thin film of water that acts as a lubricant. On unusually cold days the pressure may not be enough to melt the ice and thus skating would be impossible.

 Liquid - Gas Phase Change

 The change from a gas to la liquid is condensation. This is due to cooling and/or a pressure change.

 In liquids, the energy of the particles is raised by adding heat. When some molecules have enough K.E. they break away from the liquid surface and become vapor.

 If the temperature falls, there is a decrease in the energy of the moving molecules and the liquid may eventually freeze to the solid phase.

 Process of EVAPORATION: Molecules that have enough energy of motion (K.E.) break free from intermolecular forces and escape into the air as vapor. Some may return to the liquid is their energy is lost to other atoms.

 The liquid surface left behind is cooled. In evaporation the molecules that escape are the ones with the greatest velocity (heat) thus the average velocity and K.E. of the remaining particles is reduced. This results in cooling effects. Heat must be absorbed from the surroundings to continue the evaporation process.

 Adding heat increases evaporation because the VAPOR PRESSURE is increased. This is the pressure exerted by the vapor (gas) of a substance when it is in equilibrium with liquid or solid phase. The system is in equilibrium when the rate of evaporation equals the rate of condensation.

 The temperature at which the liquid's vapor pressure is equal to outside (atmospheric) pressure is that liquid's BOILING POINT. At this temperature the pressure of the vapor escaping from liquid equals the outside pressure.

 When the vapor pressure equals outside pressure bubbles of vapor form and push through to the surface. As they move into the gas phase we say this is boiling. Conduction of heat creates the gas, which rises because it is less dense than the liquid, as it strives for equilibrium.

 The boiling point varies with atmospheric pressure. In mountains, the boiling point is below 100°C because the pressure of the atmosphere is less.

 Cooking requires longer times at high altitudes because of low boiling point

 Pressure cookers make food cook more rapidly because the foods can be heated above the normal boiling point without actually boiling.

 Intermolecular Forces and Latent Heat

 if we heat a mixture of ice and water, we find that no matter how much heat is transferred to the mixture, the temperature remains at 0° C until the last of the ice is melted. Only after all the ice is melted is heat converted into kinetic energy, and only then can the temperature of the water begin to rise. Experiment shows that 80 calories of heat must be absorbed from the outside world in order that 1 gram of ice might be melted, and that no temperature rise takes place in the process. The ice at 0° C is changed to water at 0° C.

 But if the heat gained by the ice is not converted into molecular kinetic energy, what does happen to it? If the Law of Conservation of Energy is valid, we know it cannot simply disappear.

 The water molecules in ice are bound together by strong attractive forces that keep the substance a rigid solid. In order to convert the ice to liquid water (in which the molecules, as in all liquids, are free of mutual bonds to the extent of being able to slip and slide over, under, and beside each other) those forces must be countered. As the ice melts, the energy of heat is consumed in countering those intermolecular forces. The water molecules contain more energy than the ice molecules at the same temperature, not in the form of a more rapid motion or vibration but in the form of an ability to resist the attractive forces tending to pull them rigidly together.