Unit B Review Thermochemical Changes The study of

Unit B Review Thermochemical Changes

The study of energy changes by a chemical system during a chemical reaction is called thermochemistry. Calorimetry is the technological process of measuring energy changes of an isolated system called a calorimeter. Recall that an isolated system does not exchange matter or energy with its outside environment. No calorimeter is 100% sealed and insulated, so they only approximate an isolated system.

Analyzing Energy Changes Heat refers to the form of energy that is transferred from an object at a higher temperature to an object at a lower temperature. Thermal energy is the total kinetic energy of the entities of a substance. Q = quantity of thermal energy (J) Q = mcΔt m = mass (g) c = specific heat capacity (J/g·°C) Δt = temperature change (°C) The S. I. unit for energy is the joule (J).

The specific heat capacity of a substance is the quantity of energy required to raise one gram of a substance by one degree Celcius. The change in temperature of the water is used to determine the quantity of heat energy released or absorbed by the chemical system.

Heat Transfer and Enthalpy Change Kinetic energy (energy of motion) of a chemical system includes: • moving electrons within atoms • the vibration of atoms connected by chemical bonds • the rotation of molecules • the translation of molecules (moving from one place to another) The temperature of a chemical system is a measure of the average kinetic energy of the entities that make it up. So a change in temperature means a change in kinetic energy.

Potential energy (stored energy in chemical bonds) includes: • covalent and/or ionic bonds between the entities (intramolecular) • intermolecular forces between entities hydrogen bonding covalent bonding

The enthalpy (H) of a system is the sum of the kinetic and potential energy within it. An enthalpy change, ΔH, is the difference between the enthalpy of the products and the enthalpy of the reactants for a system under constant pressure. ΔH = Hproducts – Hreactants WE CANNOT MEASURE ENTHALPY DIRECTLY! We can calculate the quantity of heat that is released or absorbed by the surroundings of a chemical system by measuring a change in temperature of the surroundings. ΔH (system) = Q (calorimeter)

Zn(s) + 2 HCl(aq) → H 2(g) + Zn. Cl 2(aq) The change in potential energy of the chemical system equals the change in kinetic energy of the surroundings.

For an exothermic reaction (chemical system releases heat), ΔH (the enthalpy change) is negative. For an endothermic reaction (chemical system absorbs heat), ΔH (the enthalpy change) is positive.

Molar Enthalpies and Calorimetry Enthalpy of reaction (or enthalpy change of reaction) refers to the energy change for a whole chemical system when reactants change to products. Δr H = enthalpy of reaction (k. J) Δr H = n Δr H m n = chemical amount (mol) Δr Hm = molar enthalpy of reaction (k. J/mol) Molar enthalpy of reaction is the enthalpy change in a chemical system per mole of a specific chemical in a system at constant pressure. In a calorimeter, the change in enthalpy of the chemical system is equal to the change in thermal energy of the calorimeter. ΔH = Q n Δr Hm = mcΔt

Method 1: Molar Enthalpies of Reaction, Δr. Hm When reactants and products are in their standard state, they are at a pressure of 100 k. Pa, an aqueous concentration of 1. 0 mol/L. and liquids and solids are in their pure state. To communicate a molar enthalpy, both the substance and the reaction must be specified.

Formation Reaction Δf Hm° = – 239. 2 k. J/mol CH 3 OH When 1 mol of methanol is formed from its elements when they are in their standard states at SATP, 239. 2 k. J of energy is released. Combustion Reaction Δc Hm° = – 725. 9 k. J/mol CH 3 OH The complete combustion of 1 mol of methanol releases 725. 9 k. J of energy. Note that the above reactions are balanced for one mole of the compound.

Method 2: Enthalpy Changes, Δr. H Write an enthalpy change (Δr H) beside the chemical equation. CO(g) + 2 H 2(g) → CH 3 OH(l) Δr H = – 725. 9 k. J The enthalpy change is not a molar value, so does not require the “m” subscript and is not in k. J/mol. Δc H° = – 98. 9 k. J Δc H° = – 197. 8 k. J When 2 moles of sulfur dioxide are burned, twice as much heat energy is released as when 1 mole of sulfur dioxide is burned.

Method 3: Energy Terms in Balanced Equations For endothermic reactions, the energy is listed along with the reactants + energy → products For exothermic reactions, the energy is listed along with the products. reactants → products + energy

Method 4: Chemical Potential Energy Diagrams During an exothermic reaction, the enthalpy of the system decreases. Heat flows out of the system and into the surroundings and we observe a temperature increase.

Method 4: Chemical Potential Energy Diagrams During an endothermic reaction, the enthalpy of the system increases. Heat flows into the system from the surroundings and we observe a temperature decrease.

Not all chemical energy changes can be studied conveniently using simple calorimetry. Methods used to study these reactions are based on the principle that net changes in all properties of a system are independent of the way the system changes from the initial state to the final state.

Predicting Δr. H: Hess’ Law Hess’ law states that the addition of chemical equations yields a net chemical equation whose enthalpy change is the sum of the individual enthalpy changes. Δr. H° = Δ 1 H° + Δ 2 H° + Δ 3 H° + … = ΣΔr. H° Two things to remember: • If a chemical equation is reversed, then the sign of Δr. H changes. • If the coefficients of a chemical equation are altered by multiplying or dividing by a constant factor, then the Δr. H is altered by the same factor.

Use the following given equations and their standard enthalpy changes.

When comparing enthalpy changes formation reactions of different compounds, we must choose a reference energy state. It is convenient to set the enthalpies of elements in their most stable form at SATP to be zero. As an arbitrary convention, for the sake of simplicity, all other enthalpies of compounds are measured relative to that reference energy state. A formation reaction always begins with elements, so any standard enthalpy of formation reactions are measured from the reference energy state of zero.

Thermal stability is the tendency of a compound to resist decomposition when heated. The lower (i. e. more negative) the value of a compound’s standard molar enthalpy of formation, the more stable it is. Δf Hm° = – 280. 7 k. J/mol Δf Hm° = – 577. 6 k. J/mol Sn. O 2 Tin(IV) oxide has a greater thermal stability than tin(II) oxide.

The standard enthalpy change of a reaction is the sum of the standard enthaplies of formation of the products minus the sum of the standard enthalpies of formation of the reactants. Δr. H° = ΣnΔf. PHm° – ΣnΔf. RHm°

Reaction Progress • Why do some chemicals react faster than others, when all other variables are controlled, except for the type of chemicals? e. g. Mg(s) reacts much faster with HCl(aq) than Zn(s). • Why do some reactions need an initial input of external energy to start? e. g. A match is needed to start the combustion of a hydrocarbon.

Collision-Reaction Theory I 2 HI • A chemical sample consists of entities (atom, ions, or molecules) that are in constant random motion at various speeds, rebounding elastically from collisions with each other. • A chemical reaction must involve collisions of reactant entities.

• An effective collision requires sufficient energy. Collisions with the required energy have the potential to react. • An effective collision also requires the correct orientation (positioning) of the colliding entities so that bonds can be broken and new bonds can be formed.

• Ineffective collisions involve entities that rebound elastically from the collision.

Activation Energy of a Reaction Activation energy is the minimum energy that colliding entities must have in order to react. This initial input energy may be in the form of heat, light, or electricity.

CO(g) + NO 2(g) → CO 2(g) + NO(g) Δr. H° = – 224. 9 k. J

CO(g) + NO 2(g) → CO 2(g) + NO(g) The diagram on the right just tells us about the “before and after. ” Δr. H° = – 224. 9 k. J

H 2(g) + I 2(g) → 2 HI(g) Δr. H° = + 53. 0 k. J

Breaking bonds between atoms or ions requires energy, making it an endothermic process. bonded particles + energy → separated particles Bond energy is the energy required to break a chemical bond. The stronger the bond, the greater the energy needed to break it. Forming bonds between atoms or ions releases energy, making it an exothermic process. bonded particles → separated particles + energy Bond energy is also the energy released when a bond is formed. The stronger the bond, the greater the energy released.

Endothermic Reactions 2 H 2 O(l) → 2 H 2(g) + O 2(g) Δr. H° = +571. 6 k. J In any endothermic reaction, the energy required to break bonds (2 O─H bonds in H 2 O) is greater than the energy released when bonds are formed (O═O and H─H). Exothermic Reactions H 2(g) + Cl 2(g) → 2 HCl(g) Δr. H° = – 184. 6 k. J In any exothermic reaction, the energy required to break bonds (H─H and Cl ─Cl) is less than the energy released when bonds are formed (H─Cl).

Empirical Effect of Catalysis deals with the properties and development of catalysts, and the effects of catalysts on chemical reactions. A catalyst is a substance that increases the rate of a chemical reaction without being consumed in the overall process. Chlorophyll is a catalyst for photosynthesis. The inside of a catalytic converter in a car exhaust system is coated with an alloy that acts as a catalyst for the combustion of exhaust gases.

Theoretical Explanation of Catalysis A catalyst lowers the activation energy for a reaction. This results in a larger number of effective collisions between entities, so the reaction rate increases. The rate of the forward reaction and of the reverse reaction increases by the same amount.

Biological catalysts (enzymes) have binding sites specific to the reactants (substrate). The enzyme temporarily binds the substrate molecules and aligns them in the proper orientation.

The uncatalyzed reaction proceeds slowly at room temperature.

The catalyzed reaction has a lower activation energy, thus having an increased rate.

Uses of Catalysts The Oil Industry Catalysts are used in the cracking and reforming of crude oil which reduces the temperature required for these processes, making them more efficient.

Upgrading of Bitumen from Oil Sands

Emissions Control

Enzymes