Unit 10 Text Chemistry I CP 1

Transcription

1 Unit 10 Text Chemistry I CP 1 Your Key Chemistry Annotation Guide If you are NOT using the following annotation, put in your key to the left of each item. Mr. T s Key Items to be annotated Circle Box Write DEF Underline headings and subheadings. key content vocabulary next to definitions (sometimes #2 and #3 will be the same and in that case I expect to see a box AND DEF) important ideas (include captions for visuals). Mr. T s or Write Eq Underline Write Write in the margin next to an important equation. key steps to solving an example problem, or write those steps in the margin. any questions in the margin. a learning objective next to each subheading. Using this handout effectively The studying and annotations and other homework assignments are designed to require a maximum of 45 minutes to an hour of your time. This time limit, however, does not include things like test make-up points or extra credit activities. Students are expected to plan their time. If students wait until the last minute to complete assignments that are assigned multiple days in advance, then the time required to complete the assignment may require more than the 1 hour limit. This document is a work in progress. I have made an attempt to put in as much about the knowledge you need for this unit of study as I can. This document, however, DOES NOT contain everything that you need to know to make 100% on the Unit 4 test. You must also rely on the knowledge that you should have acquired in previous units of study, classroom explanations, your notes, the textbook sections that were assigned for study, and in some cases creative thinking and problem-solving skills. Thermochemistry Unit 10: Chemistry I Honors The Standards in This Unit According to the South Carolina Science Standard C-4, students will demonstrate an understanding of the types, the causes, and the effects of chemical reactions. According to the South Carolina Science Standard C-5, students will an understanding of the structure and behavior of the different phases of matter. Most of the material in this unit will come from these standards. Under South Carolina Science Standard C-4 and C-5 are lists of indicators and under each of those indicators are supporting documents. All the material in this section stems from these indicators and supporting documents. Indicator C-4.3 Indicator C-4.3 in the South Carolina Science Standards students should be able to analyze the energy changes (endothermic or exothermic) associated with chemical reactions. In physical science, students learned to Summarize characteristics of balanced chemical equations (including conservation of mass and changes in energy in the form of heat that is, exothermic or endothermic reactions). (PS-4.7)

2 Unit 10 Text Chemistry I CP 2 In this unit of chemistry students should: Show energy changes in chemical reactions (for example) Endothermic reaction Exothermic reaction 2H 2 O + energy 2H 2 + O 2 2H 2 + O 2 2H 2 O + energy Understand that the heat of reaction is the quantity of energy released or absorbed as heat during a chemical reaction. Heat of reaction should be expressed in terms of kj per mole, so quantitative application of this concept should not be introduced until after students have mastered the mole concept. Understand enthalpy as the energy absorbed or released as heat during a chemical reaction at constant pressure Explain the relationship between enthalpy change and the tendency of a reaction to occur. Understand that some endothermic reactions may be spontaneous due to an increase in entropy, a measure of the degree of randomness of the particles, such as molecules, in a system. Explain the relationship between entropy change and the tendency of a reaction to occur. Explain the relationship between various conditions and the entropy of the system Temperature Phase Formation of solutions OK, let s look at all these indicators one at a time: Show energy changes in chemical reactions (for example) Endothermic reaction Exothermic reaction Let s define a few terms: 2H 2 O + energy 2H 2 + O 2 2H 2 + O 2 2H 2 O + energy Write your learning objective for this section here: Endothermic This means that energy is being taken into the system as the reaction goes forward. Inside the system energy increases but that energy has to come from somewhere. The energy comes from the surroundings and the surroundings are where you and I are. So an endothermic reaction feel s cold to us. Note 1: We don t actually feel heat or cold. We feel heat energy leaving our bodies or heat energy entering our bodies and our brains interpret that as feeling heat or cold. Note 2: Bond breaking is always endothermic.

3 Unit 10 Text Chemistry I CP 3 Exothermic This means that energy is given off by the system into the surroundings as the reaction goes forward. Inside the system the energy decreases. The energy leaves the system and enters the surroundings and the surroundings are where you and I are. So an exothermic reaction feel s warm or hot to us. Note: Bond forming is always exothermic. System In chemistry, the system is the chemical reaction. Surroundings Everything in the universe except for the system is the surroundings. If the chemical reaction is taking place in a solution, then even the water is part of the surroundings. The beaker in which the solution is held is part of the surroundings. You and I are part of the surroundings. Dorman High is part of the surroundings and on and on. The universe The universe is both the system and the surroundings. In the universe energy is neither gained nor lost so any energy that leaves the system enters the surroundings and energy entering the system comes from the surroundings. Energy get s moved around in the universe but it is never lost. Figure 1. An illustration of the interconnected concepts of the universe, the system, and the surroundings. Notice that in the examples of an endothermic reaction and an exothermic reaction above that energy appears in both reactions. In the endothermic reaction energy appears on the left side indicating that energy is being used as a reactant. As the reaction moves forward energy appears to disappear but in fact this is just a way to showing that energy is moved from the reactants and becomes part of the products. It is also acceptable (in fact, it s preferable) that energy be listed as a negative product in an endothermic reaction: Instead of this: 2H 2 O + energy 2H 2 + O 2 You can use this: 2H 2 O 2H 2 + O 2 energy This way the gain or loss of energy in the surroundings is always shown as a product. In fact, specific amounts of energy can be shown as entering or leaving the system this way: 2H 2 + O 2 2H 2 O kJ. The chemical equation above which includes the energy entering or leaving the surroundings is called a thermochemical equation. The kj value the amount of energy for each mole represented by the coefficients in the equation. In the thermochemical equation above kilojoules of energy is released to the surroundings for every 2 moles of hydrogen (H 2 ) and every 1 mole of oxygen (O 2 ) that react together. Also, kilojoules of energy is released to the surroundings for every 2 moles of water (H 2 O) that is formed.

4 Unit 10 Text Chemistry I CP 4 If you wanted to know how much energy is produced for each mole of hydrogen, you would cut everything in half: H 2 + ½O 2 H 2 O kJ If you wanted to know how much energy is produced for every 4 moles of hydrogen, you multiply everything by 2: 4H 2 + 2O 2 4H 2 O kJ The endothermic reaction would show the energy change in negative units: 2H 2 O 2H 2 + O kJ Understand that the heat of reaction is the quantity of energy released or absorbed as heat during a chemical reaction. 2H 2 O 2H 2 + O kJ The heat of reaction in the equation above is 483.6kJ. This is the energy absorbed during the chemical reaction described by the thermochemical equation above. If 10.00g of H 2 is reacted with oxygen from the air, you can calculate the amount of energy that would be released to the surroundings. You start by converting the mass to moles, then use the enthalpy from the reaction just as you would any other part of the equation to calculate energy released. It is probably easiest to see how to do this when the enthalpy is written as a product in the equation: 2H 2 + O 2 2H 2 O kJ Now we set up the stoichiometric equation (note that the molar mass of H 2 is g H 2 = 1 mole H 2 ): g H2 1 1 mol H g H kJ 2 mol H kJ kj. Heat of reaction can also be expressed in terms of kj per mole kj. mol Understand enthalpy as the energy absorbed or released as heat during a chemical reaction at constant pressure Another way of representing these same reactions is to pull the energy change out into a separate area to show the change of energy IN the system (as opposed to the surroundings). For this we use the symbol ΔH. You already know that Δ is the Greek letter delta which in chemistry means change. H used to mean heat, but not it s used for a broader term called enthalpy. ΔH means change in enthalpy which is the amount of energy absorbed by a system as heat during a chemical reaction (or some other process) at constant pressure. An equation that shows what s happening to energy in the surroundings looks like the one s you ve already seen: 2H 2 O 2H 2 + O kJ but an equation that shows what s happening to the energy in the system looks like this: 2H 2 O 2H 2 + O 2 ΔH = kJ Notice that the energy change in the last 2 equations are opposite of each other because the 1 st shows what s happening in the surroundings and the 2 nd shows what s happening in the system.

5 Unit 10 Text Chemistry I CP 5 Practice Problems: 1. ΔH for C 3 H 8 = 2220 kj/mole. How much energy is produced from burning grams of propane (C 3 H 8 ). 2. The heat of enthalpy for a reaction (ΔH) in which methane (CH 4 ) is burned (reacted with O 2 ) is kj/mole. How much energy would be produced by burning 20.0 grams of methane? Answers/Solution Guide to Practice Problems: kj Explain the relationship between enthalpy change and the tendency of a reaction to occur. Write your learning objective for this section here: The tendency of a reaction of occur is called spontaneity. There is an equation for determining spontaneity. ΔG = ΔH TΔS If ΔG is negative the reaction is considered to be spontaneous which means it will happen on its own. You already know that T stands for temperature. ΔH means the change in heat energy and ΔS means a change in entropy. Entropy means disorder. If the ΔH in the system is negative (loosing energy in the system and giving it off to the surroundings; i.e. exothermic) and ΔS is positive (the reaction results in increased disorder) then ΔG will be negative and the reaction will be spontaneous. As my mother would say, That s a lot to chew on. Let s look at this chart to help: ΔH ΔS ΔG Spontaneity + Always Always spontaneous. at lower temperatures Spontaneous at lower temperatures. + + at higher temperatures Spontaneous at higher temperatures. + Never These reactions are never spontaneous. So what was stated earlier matches the 1 st row: a negative ΔH and a positive ΔS will result in a negative ΔG and the reaction will be spontaneous. In the 2 nd row: a negative ΔH and a negative ΔS will result in a negative ΔG only at lower temperatures, so these reactions will be spontaneous only at lower temperatures. In the 3 rd row: a positive ΔH and a positive ΔS will result in a negative ΔG only at higher temperatures, so these reactions will be spontaneous only at higher temperatures. In the 4 th row: a positive ΔH and a negative ΔS will always result in a positive ΔG, so these reactions will never be spontaneous. Also, as you can see from the equation ΔG = ΔH TΔS, the more negative ΔH is, the more likely ΔG is to be negative and the more likely the reaction will be spontaneous.

6 Unit 10 Text Chemistry I CP 6 Understand that some endothermic reactions may be spontaneous due to an increase in entropy, a measure of the degree of randomness of the particles, such as molecules, in a system. Explain the relationship between entropy change and the tendency of a reaction to occur. The symbol for entropy change is ΔS. As you can see from the equation ΔG = ΔH TΔS, the more positive ΔS is the more likely ΔG is to be negative and the more likely the reaction is to be spontaneous. Explain the relationship between various conditions and the entropy of the system Temperature Also, as you can see from the equation ΔG = ΔH TΔS, if ΔS is positive the higher the temperature (T) is, the more likely ΔG is to be negative and the more likely the reaction will be spontaneous. So, generally increasing temperature increases the entropy of a system IF the change in entropy (ΔS) is positive. Phase Solids are more organized than liquids so a chemical change that results in changing from solid to liquid results in an increase in disorder which is an increase in ΔS and the larger the increase in ΔS the more a change is to be spontaneous. Liquids are more organized than gases so a chemical change that results in changing from liquid to gas results in an increase in disorder which is an increase in ΔS and the larger the increase in ΔS the more a change is to be spontaneous. Formation of solutions Pure substances are typically more organized than solutions so a change that results in changing from a pure substance to a solution results in an increase in disorder which is an increase in ΔS and the larger the increase in ΔS the more a change is to be spontaneous. Practice Problems: 3. If delta-h is positive, T is high, and delta-s is positive, is the chemical reaction spontaneous? 4. Use the Gibbs Free Energy equation to explain why ice melts at room temperature? 5. Use the Gibbs Free Energy equation to explain why steam condenses melts at room temperature? 6. Use the Gibbs Free Energy equation to explain why your desk remains a solid? Answers/Solution Guide to Practice Problems: 3. Yes. ΔH ΔS ΔG Spontaneity + + at higher temperatures Spontaneous at higher temperatures. 4. In order for ice to melt, energy must be transferred from the surroundings to the ice, so ΔH must be positive. The temperature is positive and liquid molecules are always

7 Unit 10 Text Chemistry I CP 7 more disordered than solid ones, so as the reaction progresses from solid to liquid progresses, the ΔS grows larger, therefore ΔS is positive. Since both ΔH and ΔS are positive, ΔG must be negative and the reaction is spontaneous so long as the temperature is high enough to cause the product of T and ΔS to be larger than ΔH. ΔH ΔS ΔG Spontaneity + + at higher temperatures Spontaneous at higher temperatures. Indicator C-4.6 Indicator C-4.6 in the South Carolina Science Standards students should be able to explain the role of activation energy and the effects of temperature, particle size, stirring, concentration, and catalysts in reaction rates. In physical science, students learned to Explain the effects of temperature, concentration, surface area, and the presence of a catalyst on reaction rates. (PS- 4.11) This is addressed in a descriptive manner in physical science In this unit of chemistry students should: Analyze an energy level diagram (reaction energy diagram) Going from reactants to the top of the curve, you are going up the energy scale. Energy (heat) is being put in to break bonds in the reactants. At the top of the curve, the bonds in the reactants have been broken. The amount of energy put in to break these bonds is called the activation energy. The activation energy is the minimum amount of energy needed for the reaction to occur. Relate activation energy to heat of reaction Going from the top of the curve to the products, you are going down the energy scale. Energy (heat) is given out as bonds form in the products. The reactants are higher up the energy scale than are the products. The amount of energy (heat) you need to put in (the activation energy) is less than the amount of energy (heat) you get out. This is a typical exothermic reaction. The difference in energy levels between the reactants and the products is given the symbol H This is the amount of heat given out (or taken in) during the reaction. For an exothermic reaction, H is negative. Write your learning objective for this section here:

8 Energ y Activation Energy Unit 10 Text Chemistry I CP 8 For an endothermic reaction, H is positive. Analyze the effects of temperature, particle size, stirring, concentration, and catalysts on reaction rates For each factor students should be able to explain, in terms of the kinetic theory, how the factor influences the reaction rate in terms of Collision energy Collision frequency Activation energy OK, let s look at all these indicators one at a time: Analyze an energy level diagram (reaction energy diagram) Reaction Energy Diagram for an Endothermic Reaction Products ΔH Reactants Progress of reaction Figure 2. The change in energy can be plotted against the progress of a reaction, as the reactants turn into products in this illustration of the energy in an endothermic reaction.

9 Energy Activation Energy Unit 10 Text Chemistry I CP 9 Reaction Energy Diagram for an Exothermic Reaction Reactants ΔH Products Progress of reaction Figure 3. The change in energy can be plotted against the progress of a reaction, as the reactants turn into products in this illustration of the energy in an endothermic reaction. On both the endothermic and the exothermic Reaction Energy Diagrams going from reactants to the top of the curve (left to right), you are going up the energy scale. Energy (heat) is being put in to break bonds in the reactants. At the top of the curve, the bonds in the reactants have been broken. The most common models on the subject of bond breaking and forming of new bonds state that at the top of the energy curve not only are the bonds between the reactant atoms broken, but if bonds are formed in the reaction bonds that will form between products are already lining up. This old-bonds-breaking-and-new-bondslining-up-to-form is a kind of like what a cousin of mine would call a mish-mash. In chemistry theory, this mish-mash is called an activated complex. The amount of energy put in to break these bonds is called the activation energy. Sometimes energy needs to be added to a system of chemicals to get the reaction started. This is called activation energy. An example would be a fire where the fuel and oxygen are present in the correct ratios but the fire doesn t just start on its own; it s needs a spark of flame to initiate the reaction. This spark or flame can be called initiating energy. Activation energy is always required for a chemical reaction to occur but sometimes the activation energy is not obvious. All that you have to do to get zinc to react with hydrochloric acid is to put them together. That s because the atoms and ions are already moving around fast enough to provide the needed energy. Recall that all

10 Unit 10 Text Chemistry I CP 10 atoms, ions, and molecules are constantly in motion. If that motion is enough to provide the needed collision energy to start the reaction, then no energy need be added; the activation energy or initiating energy is present in the motion of atoms, ions, and molecules in the form we call ambient temperature: the temperature in the environment. The activation energy is the minimum amount of energy needed for the reaction to occur. Relate activation energy to heat of reaction Going from the top of the curve to the products, you are going down the energy scale. Energy (heat) is given out as bonds form in the products. Recall that bond formation is always an exothermic process. Bonds form because they are more stable than whatever existed before the bonds were formed. Higher stability means lower energy and the energy has to go somewhere. That energy goes to the surroundings and we feel that as heat. On the exothermic diagram, the reactant energy is higher up the energy scale than it is for the products. On the endothermic diagram, the reactant energy is lower on energy scale than is the product energy. This means that overall heat energy is lost by the reaction in the exothermic diagram and that heat energy is given to the surroundings which we feel as heat. Exothermic reactions usually feel warm or hot to us but the chemicals in the system are loosing heat energy. On the exothermic diagram, the amount of energy (heat) you need to put in (the activation energy) is less than the amount of energy (ΔH, change in enthalpy, or heat) you get out. This is a typical exothermic reaction. The difference in energy levels between the reactants and the products is given the symbol H This is the amount of heat given out (or taken in) during the reaction. For an exothermic reaction, H is negative. For an endothermic reaction, H is positive. Analyze the effects of temperature, particle size, stirring, concentration, and catalysts on reaction rates Temperature Recall that temperature is a relative measure of the average kinetic energy (or more correctly the average momentum) of the structural particles in a substance. An increase in temperature increases the speed of particles so particles hit each other with greater frequency and on average with greater force. A greater frequency and force of collision typically results in a greater reaction rate. Particle size If the reacting particles are clumped together (as folks from my part of the country used to say), that makes it harder for other particles to get at them to react. Before inner particles in these clumps to react, they have to wait for outer particles to react and get out of the way. Therefore, the smaller the clumps, the faster the reaction can move forward.

11 Energy Unit 10 Text Chemistry I CP 11 Smaller reactant clumps or crystals, or groups the faster the reaction can proceed because reacting particles can hit each other with greater frequency. Stirring stirring works a lot like increased temperature in that stirring increases the speed of particles so particles hit each other with greater frequency and on average with greater force. A greater frequency and force of collision typically results in a greater reaction rate. Concentration Recall that concentration is a measure of the number of particles of solute in a given measurement of either the solute of the solution. An increase in concentration results in a greater frequency of collisions and typically more a greater collision rate results in a higher rate of reaction. Catalysts Catalysts reduce the amount of activation energy required for the reaction to move forward. See the illustration below. Recall that temperature is relative a measure of AVERAGE particle speed, energy, or momentum (and momentum is the best term to use). This means that some particles are moving close and some fast but most are moving around some average speed. If we reduce the energy needed for a reaction to occur then the number of particles moving fast enough to react is increased just by lowering the need for high activation energy and the reaction proceeds more quickly. Reaction Energy Diagram for an Exothermic Reaction with Catalyst Activation energy with catalyst Effect of a catalyst on activation energy Reactants ΔH Products Progress of reaction Figure 4. The change in energy can be plotted against the progress of a reaction, as the reactants turn into products in this illustration of the energy in an endothermic reaction and the effect of a catalyst on the activation energy. For each factor students should be able to explain, in terms of the kinetic theory, how the factor influences the reaction rate in terms of

12 Unit 10 Text Chemistry I CP 12 Collision energy Increase collision energy (the force with which one particle strikes another) you increase the activation energy and make a reaction more likely. This is why increasing temperature and stirring tend to increase reaction rates. Collision frequency Increase collision frequency (how often reacting particle strike each other) and you increase how often a reaction is likely between particles making it likely that the reaction will move forward more rapidly. This is also why increasing temperature and stirring tend to increase reaction rates. Activation energy Indicator C-4.9 This is where catalysts come into play and why diamonds are forever. Catalysts reduce the amount of activation energy needed to make a reaction proceed forward and therefore will tend to speed up the reaction. Diamonds are less thermodynamically stable than graphite, but diamonds do not readily become graphite. This is because the activation energy required to get a diamond to change into graphite is huge so it take a huge amount of energy to get diamonds over that energy hill in the energy reaction diagram. Indicator C-4.9 in the South Carolina Science Standards students should be able to summarize the concept of chemical equilibrium and Le Châtelier s principle. In this unit of chemistry students should: Understand that equilibrium is a dynamic condition in which two opposing changes occur at equal rates in a closed system. Illustrate that equilibrium as it applies to Reversible chemical reactions Solubility Phase change Understand and apply La Châtelier s Principle in reference to the following stresses A change in concentration A change in temperature A change in pressure OK, let s look at all these indicators one at a time: Write your learning objective for this section here: Understand that equilibrium is a dynamic condition in which two opposing changes occur at equal rates in a closed system.

13 Unit 10 Text Chemistry I CP 13 We start the process of teaching chemistry as if chemical reactions move in only one direction. That s because it s easy to see and understand, but some of the most important chemical reactions around us don t really behave that way. Respiration (your breathing and using oxygen and exhaling carbon dioxide) would not work in the chemistry wasn t reversible. The production of ammonia from the gases in the air though the Haber process is also a reversible reaction. Without the Haber process it is likely that millions of people would die of starvation due to a lack of man-made fertilizers to increase food crop yields. In these reversible reactions what is actually happening is that the reactions are typically always in forward and reverse at the same time. This is called a dynamic equilibrium (it s called dynamic because it s changing all the time and it s called an equilibrium because it s in balance). As reactant particles collide and react (in single or double replacement reactions) or in some cases decompose the reverse reaction is also happening. This may be an entirely different way of thinking about chemistry and you need to take a minute and read and reread this paragraph over and again until it really sinks in. Let s look at the reaction of hydrogen and nitrogen to make ammonia that we described earlier as the Haber process. 3H 2 (g) + N 2 (g) 2NH 3 (g) While we often write this reaction as simply moving forward the serve reaction also takes place in the Haber reaction chamber: 2NH 3 (g) 3H 2 (g) + N 2 (g) Another way of showing this forward and backward reaction occurring all at the same time I with a double arrow: 3H 2 (g) + N 2 (g) 2NH 3 (g) A dynamic equilibrium can be thought of as a balancing act. The when the reaction reaches equilibrium, the rate of the forward reaction is occurring as fast as the rate of the reverse reaction. As a result the relative amounts of reactants and products remain relatively unchanged at least until something comes along to disturb the equilibrium. Once again you need to take a minute and read and reread this paragraph over and again until it really sinks in. Illustrate that equilibrium as it applies to Reversible chemical reactions In the last section we left you with the thought that in a dynamic equilibrium the relative amounts of reactants and products remain relatively unchanged at least until something comes along to disturb the equilibrium. Lot s of things can change what we might call the balance point of the equilibrium. Actually, the balance point is the equilibrium. In our Haber reaction... 3H 2 (g) + N 2 (g) 2NH 3 (g) there are 4 moles of gas on the left side and 2 moles of gas on the right. Because the right side has fewer moles than the left side, if you put the whole reaction system under increased pressure, the chemicals on the right side are favored. In other words, increasing the pressure will increase the amount of ammonia (the right side) that will be formed.

14 Unit 10 Text Chemistry I CP 14 Solubility Have you ever put to much sugar in a glass of iced tea and couldn t get all of it to dissolve? It looks like the sugar is just sitting on the bottom but, in fact, there is a dynamic equilibrium here, too. Solid sugar is constantly dissolving while at the same time dissolved sugar is recrystallizing. This too is a dynamic equilibrium. If a solid is dissolved in a liquid but too much solid is placed in the liquid such that the saturation point is exceeded, a dynamic equilibrium system is established. This kind of dynamic equilibrium keeps the relative amount of dissolved solid constant (unchanging). Phase change When liquids freeze, we often think about the processes as particles linking together and staying that way. But during the freezing process there are particles that link together and some that link and then let go, creating a dynamic equilibrium especially along the interface between solid and liquid phases. An interface is kind of dividing surface. In this case the surface of the solid is the interface between the solid and liquid. The same thing happens between the liquid and gaseous states and the solid and gaseous state during sublimation. Recall that at any given temperature, there are particles that are moving faster and some moving slower than the average. So it stands to reason that as a substance is changing phase, such as when changing from solid to liquid, some particles will be moving faster than the average and others will be moving slower than the average. Example: Thinking about what can happen along the solid to liquid transition temperature (that s 0.00C for pure water so long as the pressure is at 1.00 atm). Some particles along the solid-liquid interface that are part of the solid phase will wiggle fast enough to shake loose and become liquid. Others particles in the liquid phase will be moving slowly enough that they will collide with another slow moving particle or a collection of solid particles and become part of the solid. This solid-particles-becomingliquid-while-liquid-particles-become-solid is a dynamic equilibrium. As was said earlier, in a dynamic equilibrium you can disturb the equilibrium. Well, if you disturb the equilibrium, which means to change the conditions affecting the equilibrium, this will change the equilibrium point. Well, if you change the pressure, it will change the equilibrium point for a solid to liquid phase change. For most substances, an increase in pressure will result in more particles joining the solid phase. In other words, particles with greater momentum will be come part of a collection of solid particles if the pressure is increased. As you will find out later, water is a major exception to this rule. Understand and apply Le Châtelier s Principle in reference to the following stresses La Châtelier s Principle describes what happens when you disturb the equilibrium. The principle states that if a stress or force is applied to a system at equilibrium, the system responds by changing the equilibrium (the balancing point) in a direction which tends to eliminate the effect of the stress. Note: Le Chatelier's Principle is named for the discoverer of the principal: Henry Louis Le Chatelier. It is sometimes called the Le Chatelier-Braun principle, because Karl Ferdinand Braun discovered it independently. A change in concentration

15 Unit 10 Text Chemistry I CP 15 Recall the concept of solubility. Solubility When so much solute has been added to a solvent that the saturation point has been passed, some of the solute will precipitate due to stirring or simply will not dissolve. Recall that when this happens a dynamic equilibrium is established in which particles of the solid are constantly dissolving and dissolved particles are constantly becoming part of the solid. When the rate of particles dissolving equals the rate of dissolving particles becoming solid, a dynamic equilibrium exists. The maximum amount of solute that can be dissolved in a solute at a given temperature and pressure is the solubility. Now let s look at the concept called the common ion effect. Let s use the example of silver chloride (AgCl). Silver chloride is listed in your solubility rules as insoluble, but as we said earlier, everything is soluble. Substances that are soluble a lot are considered soluble and those that are soluble very little are considered insoluble. AgCl a slightly soluble salt that in solution dissociates into the ions Ag + and Cl, the equilibrium state being represented by the equation: AgCl(s) Ag + (aq) + Cl (aq). According to Le Châtelier's principle, when a stress is placed on a system in equilibrium, the system responds by reducing that stress and reestablishing a new equilibrium point. If another solute containing one of those 2 ions (Ag + or Cl ) is added, the equilibrium will be disturbed. If we add a substance with a common ion we disturb the equilibrium. For example, if sodium chloride, NaCl, is added to the solution, it supplies Cl ions and the balance between the AgCl solid on the left and the Ag + and Cl ions on the right is disturbed. This change in the amount or concentration of chloride ions (Cl ) is a stress on the dynamic equilibrium. The solubility equilibrium of the solution will be shifted to remove more Cl from the solution, so that at the new equilibrium point there will be fewer Ag + ions in solution and more AgCl precipitated out as a solid. The explanation above can be hard to understand when you 1 st read it, so go back and read it one sentence at a time until it make sense to you. A change in temperature Let s take a look one of particular dissolution equilibrium to examine how the Le Chatelier's Principle works with a change in temperature. Recall that salts are called soluble when a lot of it will dissolve in water and Figure 5. The solubility curves for selected salts and gases from the test references.

16 Unit 10 Text Chemistry I CP 16 insoluble when only a little will dissolve in water. Note that the terms a lot and a little are qualitative. In other words, they aren t very specific about how much of a substance will dissolve and they are subject to personal opinion (a lot to once person is not much at all to another). In fact, all substance can be dissolved in water, so nothing is completely insoluble. Let s take a look at the case if potassium chlorate (KClO 3 ). In your test references you can see that the solubility of KClO 3 is highly dependent on temperature. If 20.0 g of KClO 3 is added to100.0 g of water at 20.0C, only about half of the KClO 3 will dissolve. The remaining half will precipitate to the bottom on the container if left undisturbed. Take a second and put your finger on the KClO 3 solubility curve in the illustration to the right. Start on the left (at 0.0C and move your finger to the right until you come to 20.0C and stop. Look at how the KClO 3 solubility curve intersects with the vertical line indicating 20C at about the same time as the curve intersects with the horizontal line indicating 10 grams of solute per 100 grams of water. This tells us that 10 grams of KClO 3 (our solute in this example) will dissolve in 100 grams of water when the water is at 20C and the inference is that the remaining 10 grams of KClO 3 in our example will not dissolve. If we change the temperature of the water in our example to 30C, about 3 more grams of KClO 3 will dissolve. Use your finger to follow the KClO 3 solubility curve as I explain what happens as the temperature changes. If we continue to increase the temperature (use your finger to follow the KClO 3 solubility curve) then at about 45C all of that 20.0g of the KClO 3 in our example will dissolve. If we start to cool our solution back to 40C, then about 2 grams of our KClO 3 sample will precipitate. This change in solubility with a change in temperature is an example of Le Chatelier's Principle. The change in temperature is a stress on the solubility system. Recall that this solubility system is in dynamic equilibrium and the change in the stress applied to the system at equilibrium causes the system to respond by changing the equilibrium in a direction which tends to eliminate the effect of the stress. In this example, the stress is eliminated by increasing the solubility of potassium chlorate as the temperature increases and decreasing the solubility as the temperature decreases. A change in pressure A change in pressure primarily affects only the equilibrium solubility of gases in a liquid. Let s use an equation used in the Haber process to produce ammonia for fertilizer as our example. Nitrogen (N 2 ) and hydrogen (H 2 ) exist naturally in the air. They combine together to make ammonia gas (NH 3 ). N 2 (g) + 3H 2 (g) 2 NH 3 (g) Note that there are 4 moles of gas represented on the left side of the equation: 1 mole of N 2 and 3 moles of H 2 (g). Note also that there are 2 moles of gas on the right side of the equation. You may recall from the Ideal Gas Law that pressure and the number of moles of gas in a closed container have an affect on each other. Well, if the pressure is increased that puts a stress on the system. It should be obvious that the system wants to respond in a way that will reduce the increased pressure. The response is to shift the reaction in the direction that has the fewer number of moles of gas. So, increasing the pressure tends to shift the reaction in the direction of the formation of ammonia gas (NH 3 ) because that side of the equation has the fewer number of moles. Reducing the

17 Unit 10 Text Chemistry I CP 17 pressure shifts the equilibrium to the left, favoring the formation of nitrogen and hydrogen gas. Write your learning objective for this section here: Indicator C-5.4 Indicator C-5.4 in the South Carolina Science Standards students should be able to illustrate and interpret heating and cooling curves (including how boiling and melting points can be identified and how boiling points vary with changes in pressure). In physical science, students learned to Explain the process of phase change in terms of temperature, heat transfer, and particle arrangement (PS-3.7) Physical science students explain phase change in terms of the Kinetic Molecular Theory Physical science students explain why temperature vs. time graphs show constant temperature during phase change. In this unit of chemistry students should: Define phase changes in terms of kinetic energy of the particles, heat transfer, and particle orientation and arrangement. melting boiling condensation freezing sublimation Differentiate the processes of evaporation and boiling The difference between heat and temperature Heat is a form of energy and you shouldn t confuse energy with temperature. Temperature is a relative measure of the amount of stored energy there is in a material. It's a bit more complex than that because you have to take into account the mass of the particles, but for this analogy you can think of this a measure of how fast the structural particles are wiggling, twisting, or moving. To understand what is going on it helps to think about how the structural particles in a solid, liquid, or gas are arranged. The particles in a solid aren't sitting still, but they're not moving around each other either. The energy that is put into a solid material during the melting process goes into the structural particles causing those particles to wiggle and twist faster and faster until the attraction between them that is holding them in place is overcome. The particles in a liquid are touching each other but moving around each other. You might think of this like marbles that you are rolling around in your hand. The energy that is put into a liquid material during the boiling process goes into the structural particles causing those particles to move around each other faster and faster until the attraction that keeps them touching each other is overcome and they go flying apart. The particles in a liquid are touching each other but moving around each other. You might think of this like marbles that you are rolling around in your hand. The energy that is put into a liquid material during the boiling process goes into the structural particles causing those particles to move around each other faster and faster until the attraction that keeps them touching each other is overcome and they go flying apart. The particles in a gas are--for the most part--not touching at all. They bump into each other but bounce off. If energy is removed from a gaseous material during the condensation process, the structural particles keep slowing down until the attraction between the particles takes over causing those particles to stick together but they are not yet locked in place relative to each other. Recall that the particles in a liquid are touching each other but moving around each other. If energy is removed from a liquid material during the freezing process, the structural particles keep slowing down until the attraction between the particles takes over causing those particles to lock in place relative to each other.

18 Unit 10 Text Chemistry I CP 18 Differentiate the terms gas and vapor Explain how atmospheric pressure and vapor pressure affect the boiling point of a substance Analyze a phase diagram (temperature vs. pressure) Explain triple point Critical point Analyze a graph of temperature vs. time which illustrates the heating or cooling of a substance over the range of phase change. Explain the shape of the graph in terms of kinetic energy, potential energy, and heat transfer OK, let s look at all these indicators one at a time: Define phase changes in terms of kinetic energy of the particles, heat transfer, and particle orientation and arrangement. Before getting into these categories, let s recall some basics. Recall the temperature is a relative measure of the average amount of random momentum of the particles in a substance. When momentum increases, the kinetic energy increases. Heat is not the same thing as temperature. Heat is energy. Energy can flow from one place to another. A transfer of heat energy from one place to another will decrease the temperature from the 1 st place and increase the temperature in the second place. Structural particles orient themselves in solids in ways that make the solid most stable. Polar molecules orient themselves in so that partially positive ends are oriented toward partially negative ends. Non-polar particles orient themselves so that the space between the particles is minimized. This maximizes the effect of London forces. Melting A solid substance becomes a liquid substance through a process called melting. In melting, heat energy is transferred to the substance, increasing the kinetic energy or particle momentum. The increased particle momentum or kinetic energy shakes apart the structural links between particles. The increased momentum or kinetic energy also make links between particles (due to polar attractions, opposite charge/ionic attractions, or London forces) mostly impossible, keeping the substance in the liquid state. The attractive forces in liquids are strong enough to hold the particles against each other, rubbing around themselves (kind of like the people in a mosh pit) but not strong enough to overcome the kinetic energy or particle momentum and get the particles to link up to become solids again. Boiling A liquid substance becomes a gas through a process called boiling. In boiling, heat energy is transferred to the substance, increasing the kinetic energy or particle momentum. The increased particle momentum or kinetic energy shakes the structural particles completely apart. The increased momentum or kinetic energy also overcomes the attractive forces between particles (due to polar attractions, opposite charge/ionic attractions, or London forces) keeping the substance in the gaseous state.

19 Unit 10 Text Chemistry I CP 19 Condensation Condensation is the opposite of boiling. In condensation a gas becomes a liquid. In condensation, heat energy is transferred out of the substance, decreasing the kinetic energy or particle momentum. The decreased particle momentum or kinetic energy allows the structural particles to start to stick to each other but the particle momentum or kinetic energy is too high to allow the structural particles to link up in a rigid form. The decreased momentum or kinetic energy allows the attractive forces between particles (due to polar attractions, opposite charge/ionic attractions, or London forces) to overcome the particle momentum or kinetic energy that had previously been keeping the structural particles apart. Freezing Freezing is the opposite of melting. In freezing a liquid becomes a solid. In freezing, heat energy is transferred out of the substance, decreasing the kinetic energy or particle momentum. The decreased particle momentum or kinetic energy allows the structural particles to link up with each other in specific directions forming a rigid form we call a solid. The decreased momentum or kinetic energy allows the attractive forces between particles (due to polar attractions, opposite charge/ionic attractions, or London forces) to overcome the particle momentum or kinetic energy that had previously been keeping the structural particles from linking together. Sublimation A solid substance becomes a gas without first becoming a liquid through a process called sublimation. In sublimation, heat energy is transferred to the substance, increasing the kinetic energy or particle momentum. The increased particle momentum or kinetic energy shakes apart the structural links between particles. In sublimation, the attractive forces that had kept structural particles touching and moving around each other (polar attractions, opposite charge/ionic attractions, or London forces) are completely overcome by particle momentum or kinetic energy keeping the substance in the gaseous state. Differentiate the processes of evaporation and boiling Recall the definition of models. Models in science describe behavior. Models in science explain the behavior in ways that we cannot observe. Models explain why the behavior occurs by attempting to explain how the underlying phenomena work to cause the observed behavior. The underlying phenomena in chemistry are the parts we cannot usually see or otherwise observe. In the accepted model that explains the processes of evaporation and boiling, any sample of a liquid the structural particles are touching but moving about each other. Structural particles can be atoms, molecules, or ions, but in most practical applications this process involves molecules. Some particles are always moving faster than others. Particles moving fast enough near the surface and moving to the surface will escape the surface. If most of the particles in the substance are moving too slowly (with too little momentum or kinetic energy) then the few particles that are escaping are evaporating. If heat is transferred to the liquid and the temperature is increased to the point that most of the particles are moving fast enough to escape (if given the opportunity) then boiling begins to occur. Coincidentally, the combined force of particles escaping the surface of a liquid is called vapor pressure (recall that pressure is the application of force to an area). When the vapor

20 Unit 10 Text Chemistry I CP 20 pressure of particles escaping the surface of a liquid is greater than the atmospheric pressure above the liquid then boiling takes place. Differentiate the terms gas and vapor A vapor is that part of a solid or liquid that has escaped the surface and is in a gaseous state above the solid or liquid. A vapor in its natural state is a solid or liquid at the environmental temperature. A gas in its natural state will still be a gas. It will not condense to a liquid or deposit to a solid at whatever the environmental temperature is. Examples: At room temperature water molecules in the air would be a vapor because most of the water would remain a liquid. Nitrogen (the most common gas in the Earth s atmosphere) would still be in the gaseous state at room temperature. Explain how atmospheric pressure and vapor pressure affect the boiling point of a substance As was stated earlier when the vapor pressure of particles escaping the surface of a liquid is greater than the atmospheric pressure above the liquid, then boiling of the liquid takes place. If the atmospheric pressure above a liquid is increased then the boiling point (boiling temperature) is increased. Such is the case in Death Valley. Death Valley is below sea level and therefore has a higher than average atmospheric pressure. In Death Valley the boiling temperature of pure water is slightly higher than at sea level. At sea level the boiling temperature of pure water is 100C. If the atmospheric pressure above a liquid is decreased then the boiling point (boiling temperature) is increased. Such is the case in Denver, Colorado. Denver is above sea level and therefore has a lower than average atmospheric pressure. In Denver the boiling temperature of pure water is slightly lower than at sea level.

21 Pressure Unit 10 Text Chemistry I CP 21 Analyze a phase diagram (temperature vs. pressure) A typical phase diagram Super- Critical Fluid Critical point Solid (ice) Liquid Gas (steam) Triple Point Temperature Figure 6. This is a generalized phase diagram for most substances. This diagram does not represent a phase diagram for any one substance but represents the things that most phase diagrams have in common. The diagram above shows the phase behavior that most substances have in common. Note that when the pressure is low enough (along the bottom couple of centimeters of figure 6), liquids do not exist. As the temperature increases, substances will sublimate (which means they go directly from solid to gas) and deposit (which mans that you go directly from gas to solid). Unless this is very clear to you, take a second and put your finger on the vertical axis (the axis for pressure) about a centimeter above the horizontal axis. Then move your finger to the right. Notice that you start out in the region labeled solid and then move into the region labeled gas. Note that at median pressures substances go from solid (at low temperatures) to liquid (at median temperatures) to gas (at higher temperatures). Again, unless this is very clear to you, take a second and put your finger on the vertical axis (the axis for pressure) about