Introduction to Thermochemistry. Thermochemistry Unit. Definition. Terminology. Terminology. Terminology 07/04/2016. Chemistry 30

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1 Thermochemistry Unit Introduction to Thermochemistry Chemistry 30 Definition Thermochemistry is the branch of chemistry concerned with the heat produced and used in chemical reactions. Most of thermochemistry deals with determining quantities of heat by measurement and calculation. This is important because quantities of heat cannot always be measured directly. Terminology System: the part of the universe being studied can be tiny (one atom) or big (the Earth) Surroundings: the part of the universe outside of the system We are concerned with the transfer of energy and matter between system and surroundings. Terminology Terminology Open system: freely exchanges both matter and energy with its surroundings Closed system: can exchange only energy with its surroundings Isolated system: no interaction with surroundings 1

2 Energy Energy is the capacity to do work. In thermochemistry, work is done mostly in the form of thermal energy. Kinetic energy is the energy associated with motion. Potential energy is stored energy associated with forces of attraction or repulsion, like in chemical bonds. Energy Thermal energy is the energy associated with random molecular motion. In general, thermal energy of a system is proportional to temperature, but they are not the same thing. Measured in Joules (1 J = 1 kg m 2 /s 2 ) Conservation of Energy In interactions between a system and its surroundings, energy cannot be created or destroyed the total energy will always remain constant. 0 This means, in general, that all energy transferred out of the system goes in to the surroundings, or vice versa. Heat and Temperature Particle Theory Review Temperature Temperature is a measure of the average molecular kinetic energy of a substance. Temperature is measured in degrees Celsius ( C), degrees Fahrenheit ( F) or Kelvin (K) Kelvin is like Celsius, but does not use negative numbers, which can be useful in some calculations. To convert from C to K: 273 2

3 Heat Heat is energy transferred between a system and its surroundings due to temperature difference. It always moves from an area of high temperature to an area of low temperature. Heat flows until the average molecular kinetic energy between system and surroundings are the same. Heat Heat is represented by the variable q. Note: a system cannot contain heat it has internal energy which can be transferred in the form of heat. Heat versus Temperature Quantity of Heat Heat (q) Measure of energy content flowing into or out of a system Energy that is transferred during a temperature change Temperature (T) Measure of random motion of particles in a substance Indication of the direction heat energy will flow In a system, the quantity of heat needed to change the temperature of the substance depends on: How much the temperature will change The quantity of the substance The nature of the substance Units of Heat The quantity of heat needed to increase the temperature of 1.0 g of water by 1.0 C is called the calorie, abbreviated cal. The Calorie (note the capital C) used for energy content of food is actually 1 kcal. Joules (J) are also commonly used for quantities of heat. (1 cal = J) Specific Heat Capacity 3

4 Heat Capacity Heat capacity is a physical property of a substance. Molar heat capacity is the amount of heat needed to increase the temperature of one mole of a substance by one degree (J/mol C or J/mol K) Specific heat capacity is the same, but for one gram of a substance (J/g C or J/g K) Specific Heat Capacity Every substance has a specific heat capacity associated with it. Due to more complex structure at the molecular level, compounds generally have higher specific heats than elements. Heat capacity relates not only to how easy it is for a substance to heat up, but also how long it takes to cool. (lake effect) Specific Heat Capacity q = heat energy transferred or absorbed from the system (J) m = mass of substance (g) c = specific heat capacity of substance (J/g C ) T = temperature change of the substance, ( C) Example 1: Specific Heat Capacity It takes 132 J of energy to heat a sample of iron from 12.3 C to 33.9 C. What is the mass of the iron sample? Example 2: Specific Heat Capacity ml of ethanol is cooled from 60.0 C to 43.7 C. The density of ethanol is g/ml. How much energy is transferred? Example 3: Specific Heat Capacity A 3.50 g sample of pure gold releases 33.1 J of heat. Its initial temperature was 78.2 C. What is the final temperature of the sample? 4

5 Example 4: Specific Heat Capacity A g sample of iron is heated to C. The sample of iron is then submerged in a calorimeter full of water at 10.6 C. The final temperature of both substances is 24.2 C. What is the volume of water used? Calorimetry Calorimetry Calorimetry is the science associated with determining the changes in energy of a system by measuring the heat exchanged with the surroundings. A calorimeter is a well-insulated vessel (isolated system). Calorimetry Inside, a heated or cooled object (or chemical reaction) the system is submerged in water the surroundings. The temperature change of the water can be used to determine the heat released by the system. The calorimeter needs to prevent heat loss to the surroundings so the temperature change within it can be accurately recorded. Calorimetry This is a coffee-cup calorimeter. It works well because Styrofoam is an excellent insulator. Scientists that are doing more sophisticated experiments may use fancier equipment, but this device is very accurate for experiments at the high school level! Enthalpy and Thermochemical Equations 5

6 Enthalpy Enthalpy, H, is the total heat content of a system at constant pressure. This quantity is unmeasurable; however, we are more interested in knowing about energy changes, which is measurable. The change in heat of a system is called the enthalpy of reaction, H rxn. Enthalpy of Reaction H rxn is the difference between the energy contained by the reactants and the energy contained by the products. It shows whether energy has been transferred into or out of the system. H rxn = H products -H reactants Thermochemical Equations Thermochemical equations are chemical equations that include the following: All reactants and products, properly balanced States of all reactants and products Energy change of the reaction (enthalpy, H) Thermochemical Equations 4 Fe (s) + 3 O 2 (g) 2 Fe 2 O 3 (s) H = kj C 6 H 12 O 6 (s) + O 2 (g) 6 CO 2 (g) + H 2 O (l) H comb = kj Thermochemical Equations A thermochemical equation tells the enthalpy associated with the reaction based on the coefficients in the balanced equation. For example: 2 H 2 S (g) + 3 O 2 (g) 2 SO 2 (g) + 2 H 2 O (g) H = kj In this reaction, the enthalpy is based on 2 moles of H 2 S, 3 moles of O 2, etc. Example 1: Thermochemical Equations Aqueous sodium bicarbonate solution reacts with hydrochloric acid to produce aqueous sodium chloride, water and carbon dioxide gas. The reaction absorbs 11.8 kj/mol of heat at a constant pressure for each mole of sodium bicarbonate. Write the thermochemical equation for the reaction. 6

7 Example 2: Thermochemical Equations Gasoline contains ethanol (C 2 H 5 OH) which completely combusts to produce 1235 kj/mol of heat energy. a. Write the thermochemical equation. b. How much heat is released when g of ethanol completely combusts? Endothermic and Exothermic Reactions Endo- and Exothermic In an endothermic reaction: Energy (heat) is absorbed by the system from the surroundings. H rxn is positive (H products > H reactants ) In an exothermic reaction: Energy (heat) is released from the system to the surroundings. H rxn is negative (H reactants > H products ) Enthalpy Diagrams Enthalpy diagrams show the reactants and products and how energy is transferred into or out of the system: Mechanism of Enthalpy Energy is absorbed into the system to break chemical bonds. Energy is released from the system when new bonds form. In an exothermic reaction: 2H 2 (g) + O 2 (g) 2H 2 O (l) + energy The bonds for water are more stable than the bonds for the reactants, so energy is released. Mechanism of Enthalpy In an endothermic reaction: CaCO 3 (s) + energy CO 2 (g) + CaO (s) The bonds for calcium carbonate are more stable than the bonds for the product, so energy needs to be absorbed for the reactant to break apart. 7

8 Heat of Reaction/Solution Heat of Reaction When a chemical reaction takes place in an aqueous environment (or a solid is dissolved into water), the temperature change of the water can be used to determine the amount of heat released or absorbed by the reaction. Using the specific heat capacity of water and the law of conservation of energy, the experimental value for heat of reaction or heat of solution can be determined. Heat of Reaction/Solution To find ΔH from q, the equation is: Δ Where q water is the heat absorbed or released by the water in kj and n LR is the number of moles of the limiting reactant. Example 1: Heat of Reaction 75.0 ml of 10.0 M hydrochloric acid is mixed with 75.0 ml of 10.0 M sodium hydroxide in a calorimeter. Both solutions start at 25 C. The resulting mixture has a temperature of 35 C. a. How much heat is produced by the reaction? b. Write the reaction as a thermochemical equation, with the enthalpy in kj/mol. Example 2: Heat of Solution 7.98 g of solid magnesium sulfate is dissolved in 50.0 ml (50.0 g) of water. What will be the temperature change of the water? Assumptions for Heat of Reaction Assume that any solution has the same properties of water: Density is 1.00 g/ml Specific heat capacity is J/g C Assume that any heat released by the reaction is absorbed by the water and that no heat loss occurs. (Law of Conservation of Energy) Assume that the limiting reactant is entirely consumed or that the dissolving solid is completely dissolved. 8

9 Phase Changes Phase Changes and Latent Heat Phase changes also have thermochemical equations associated with them. The reactant and product are the same, but the state changes. There is also an energy change for phase changes. Phase Changes Phase Changes of Water Endothermic phase changes: H 2 O (l) H 2 O (g) H vap = 40.7 kj H 2 O (s) H 2 O (l) H fus = 6.01 kj Exothermic phase changes: H 2 O (g) H 2 O (l) H 2 O (l) H 2 O (s) H cond = kj H solid = kj Latent Heat Latent Heat Latent heat occurs when thermal energy is being absorbed into or released from a system but no temperature change occurs. This happens during a phase change, right at the melting and boiling points. All of the energy is used to break or form bonds between the molecules, instead of increasing or decreasing the kinetic energy of the molecules. This causes a temperature plateau in a heating or cooling curve. 9

10 Enthalpy of Phase Changes Specific heat capacity is used to determine how much energy is needed to increase the temperature of a substance in a specific state. Latent heat works in a similar way, and is used to calculate the amount of energy needed to change the state of a substance. Δ Δ Hess Law Hess Law Hess Law is a way to determine the enthalpy of a reaction using other known reactions in the same conditions. Reaction equations can be manipulated and added together to make another reaction. Rules for Hess Law 1. If you add two or more equations to get your target equation, add the H values of each. 2. If you multiply an equation by a number to change the coefficients, multiply H by the same number. 3. If you flip an equation, change the sign of H. Example 1: Hess Law Calculate the energy involved in the oxidation of elemental sulfur to sulfur trioxide from the reactions: (1) S (s) + O 2(g) SO 2(g) H 1 = kj (2) 2SO 3(g) 2SO 2(g) + O 2(g) H 2 = kj We are looking for this equation (target equation): (3) S (s) + 3/2 O 2(g) SO 3(g) H 3 =? Example 2: Hess Law Calculate the enthalpy change for: H (g) + Br (g) HBr (g) (1) H 2 (g) + Br 2 (g) 2 HBr (g) H 1 = -72 kj (2) H 2 (g) 2 H (g) H 2 = 436 kj (3) Br 2 (g) 2 Br (g) H 3 = 224 kj 10

11 Other Tips Work with substances that only appear in one place first (e.g. don t start by trying to balance O 2 if it is in three of the reactions). Start by putting the unique substances on the right side of the target equation. Fraction coefficients: You can leave them for diatomic gases ONLY do not use them for any other elements or compounds. How to Know You ve Got It All coefficients are the same as the target equation. All compounds are the same as the target equation. All states are the same as the target equation. You have followed the rules for Hess Law. Standard Enthalpy of Formation Standard Enthalpy of Formation H f is the change in enthalpy (heat energy) that occurs when 1 mole of a compound is formed from elements in their standard states (25 C and 1 atm, and the state on the periodic table). These values are on a chart in your duotang. For example: ½ N 2 (g) + O 2(g) NO 2 (g) H f = 33.2 kj/mol Formation Equations Example : Formation Equations A formation equation is a special type of thermochemical equation. All of the reactants are elements (e.g. Cu, O 2 ) and the state is what they are on the periodic table. Elements have H f of zero. The product is one mole of a compound. Write the formation equation for potassium chlorate. 11

12 Hess Law with H f Enthalpies of formation equations can be used just like in Hess law to determine the overall enthalpy change of a reaction, but some of the lengthy steps can be skipped. The limitation of this method (even though it s easier) is that it only works for reactions in standard conditions, and that each compound must be in the state given on the table. Calculating Enthalpy with H f As a general rule, for a reaction at standard conditions: = sum of (add them all together) n = coefficient of each species (multiply H f by the coefficent of that compound) H f = the enthalpy of formation for the compound (from the table!) Example: Standard Enthalpies Determine the enthalpy of the following reaction at standard conditions: Cl 2 (g) + 2 HBr (g) 2 HCl (g) + Br 2 (g) Is the reaction endothermic or exothermic? Lewis Diagrams and Bond Enthalpy Lewis Diagrams Lewis Diagrams Lewis diagrams are used to show the structure of covalent molecules. (Not to be confused with VSEPR diagrams, which show three-dimensional structure) Bonds are shown by lines (single, double or triple) and lone electron pairs are shown by two dots. Example: 12

13 Lewis Diagrams Usually, bond angles are important when drawing Lewis diagrams; however, for our purposes (bond enthalpies), the most important part is to draw the bonds correctly. Bonds can be single, double or triple. Rules and Exceptions Most atoms will follow the octet rule (8 valence electrons). Exceptions to the octet rule: Free radicals, which have an unpaired electron seen when there is an odd number of electrons Small atoms might not have a full octet B, Be Large atoms might have more than an octet any elements in Period 3 or higher Hydrogen always has only one single bond and no lone pairs. Steps to Drawing Lewis Diagrams Example: Lewis Diagrams 1. Count number of valence electrons in all atoms. 2. Draw the central atom(s) with its valence electrons the least electronegative (but never H). 3. Draw single bonds between outer atoms and central atom(s). Determine available electrons (total valence (2 # single bonds). 4. Arrange remaining electrons in pairs around outside atoms, except hydrogen, then around the central atom. 5. If necessary to complete the octet on the central atom, move lone pairs from outer atoms to form multiple bonds. F 2 NH 3 HCN C 2 H 5 OH Bond Enthalpy Bond Enthalpy Bond enthalpy (also called bond energy) is the amount of energy it takes to break a bond. With the units kj/mol, the bond enthalpy value for a certain bond (e.g. H-H) is the amount of energy released when one mole of that bond is broken. Bond energy will not always give the same value as enthalpy of formation, but will be close, because these values are approximate. BE = bond enthalpy x = number of that type of bond in the molecule n = coefficient in balanced equation Note that this is opposite from the process for enthalpy of formation. (reactants products) 13

14 Example 1: Bond Enthalpy C 2 H 4 (g) + Cl 2 (g) C 2 H 4 Cl 2 (g) Example 2: Bond Enthalpy CH 4 (g) + 2 O 2 (g) CO 2 (g) + 2 H 2 O (l) *Must be liquid water, since bond enthalpies are only valid at 298K (25 C). 14