Presentation on theme: "Principles of Chemical Reactivity"— Presentation transcript:
1 Principles of Chemical Reactivity Chapter 5Principles of Chemical Reactivity
2 Basic Principles Thermodynamics: The science of heat and work Energy: the capacity to do work-chemical, mechanical, thermal, electrical, radiant, sound, nuclear-affects matter by raising its temperature, eventually causing a state change-All physical changes and chemical changes involve energyPotential Energy:energy that results from an object’s position-gravitational, chemical, electrostaticKinetic Energy:energy of motion17
3 Basic Principles Law of Energy Conservation: Temperature vs. Heat: Energy can neither be created nor destroyed-a.k.a. The first law of thermodynamics-The total energy of the universe is constantTemperature vs. Heat:Temperature is the measure of an object’s heat energyHeat ≠ temperature
4 The Measurement of Heat Thermal Energydepends on temperature and the amount (mass or volume) of the object-More thermal energy a substances has the greater the motion its atoms/molecules have-Total thermal energy of an object is the sum of the individual energies of all atoms/molecules/ions that make up that objectSI unit: Joule (J)1 calorie = JEnglish unit = BTU
5 Converting Calories to Joules Convert 60.1 cal to joules19
6 Basic Principles System: object or collection of objects being studied In lab, the system is the chemicals inside the beakerSurroundings: everything outside of the system that can exchange energy with the systemThe surroundings are outside the beakerUniverse: system plus surroundingsExothermic: heat transferred from the system to the surroundingsEndothermic: heat transferred from the surroundings to the system
7 Specific Heat Capacity (C) amount of heat required to raise the temperature of 1gram of a substance by 1 kelvinSI Units: Specific heat capacity = J /g.KSpecific heat of water = 4.184Jg.K
8 Heat Transfer Heat transfer equation used to calculate amounts of heat (q) in a substanceKJJg.Kgq1 + q2 + q3 … = 0 or qsystem + qsurroundings = 0
9 Heat TransferCalculate the amount of heat to raise the temperature of 200 g of water from 10.0 oC to 55.0 oC
10 Heat TransferCalculate the amount of heat energy (in joules) needed to raise the temperature of 7.40 g of water from 29.0°C to 46.0°C22
11 Heat TransferA 1.6 g sample of metal that appears to be gold requires 5.8 J to raise the temperature from 23°C to 41°C. Is the metal pure gold?Jg.KSpecific heat of gold is 0.13Therefore the metal cannot be pure gold.23
12 Changes of Stateoccurs when enough energy is put into a substance to over come molecular interactionsSolid-liquid:molecules in a solid when heated move about vigorously enough to break solid-solid molecular interactions to become a liquidLiquid-gas:molecules in a liquid when heated move about more vigorously enough to break liquid-liquid molecular interactions to become a gasNote: This happens in reverse by removing heat energy
13 Energy and Changes of State Heat of fusion:heat needed to convert a substance from a solid to a liquid (at its melting/freezing point)333 J/g for waterHeat of vaporization:heat needed to convert a substance from a liquid to a gas (at its boiling/condensation point)2256 J/g for waterExample: Calculate the amount of heat involved to convert g of ice at -50.0°C to steam at 200.0°C.
14 The First Law of Thermodynamics This law can be stated as, “The combined amount of energy in the universe is constant”Also called-The Law of Conservation of Energy:Energy is neither created nor destroyed in chemical reactions and physical changes.
15 The First Law of Thermodynamics There are two basic ideas for thermodynamic systems:Chemical systems tend toward a state of minimum potential energySome examples of this include:-H2O flows downhill-Objects fall when droppedEpotential = mg(h)
16 State Function The value of a state function is independent of pathway -An analogy to a state function is the energy required to climb a mountain taking two different paths:E1 = energy at the bottom of the mountainE1 = mgh1E2 = energy at the top of the mountainE2 = mgh2E = E2-E1 = mgh2 – mgh1 = mg(h)Examples of state functions:Temperature, Pressure, Volume, Energy, Entropy, and enthalpyExamples of non-state functions:Number of moles, heat, work
17 The First Law of Thermodynamics Chemical systems tend toward a state of maximum disorderCommon examples of this are:- A mirror shatters when dropped and does not reform- It is easy to scramble an egg and difficult to unscramble it- Food dye when dropped into water disperses
18 The First Law of Thermodynamics Thermodynamic questions:Will these substances react when they are mixed under specified conditions?If they do react, what energy changes and transfers are associated with their reaction?If a reaction occurs, to what extent does it occur?
19 Changes in Internal Energy (E or DU) all of the energy contained within a substanceall forms of energy such as kinetic, potential, gravitational, electromagnetic, etc.First Law of Thermodynamics states that the change in internal energy, E, is determined by the heat flow (q) and the work (w)E = q + wbook: DU = q + w
20 Changes in Internal Energy (E) DE = Eproducts – EreactantsDE = q + w at constant pressure:w = -P x DVq > 0 if heat is absorbed by the systemq < 0 if heat is absorbed by the surroundingsw > 0 if surroundings do work on the systemw < 0 if system does work on the surroundings
21 Changes in Internal Energy (E) If 1200 joules of heat are added to a system, and the system does 800 joules of work on the surroundings, what is the:energy change for the system, Esys?energy change of the surroundings, Esurr?
22 Changes in Internal Energy (E) E is negative when energy is released by a system-Energy can be written as a product of the process
23 Changes in Internal Energy (E) E is positive when energy is absorbed by a system undergoing a chemical or physical changeEnergy can be written as a reactant of the process
24 Enthalpy Changes for Chemical Reactions Exothermic reactions:release energy in the form of heat to the surroundings (DH < 0)-heat is transferred from a system to the surroundingsEndothermic reactions:gain energy in the form of heat from the surroundings (DH > 0)-heat is transferred from the surroundings to the systemFor example, the combustion of propane:Combustion of butane:
25 Enthalpy Change (DH) Heat content of a substance at constant pressure - Chemistry is commonly done in open beakers on a desk top at atmospheric pressure- Therefore enthalpy change (H) is the change in heat content: H = qp at constant pressureDE = qv at constant volumeIf DE and DH < 0:energy is transferred to the surroundingsIf DE and DH > 0:energy is transferred to the system- Enthalpy and energy differ by the amount of workDE = DH + w and w = -PDV
26 Enthalpy Changes for Chemical Reactions Exothermic reactions generate specific amounts of heatBecause the potential energies of the products are lower than the potential energies of the reactantsEndothermic reactions consume specific amounts of heatPotential energies ofthe reactants are lowerthan the productsDH for the reverse reactionis equal, but has theopposite sign to the forwardreaction
27 Thermochemical Equations balanced chemical reaction with the H value for the reactionH < 0 designates an exothermic reaction:heat is a product, the container feels hotH > 0 designates an endothermic reaction:heat is a reactant, the container feels cold
28 Hess’s LawIf a reaction is the sum of two or more other reactions, DH for the overall process is the sum of the DH for the component reactionsHess’s Law is true because H is a state functionIf we know the following H values:Target:?1.2.
29 Hess’s LawWe can calculate the H for the reaction by properly adding (or subtracting) the H for reactions 1 and 2Notice that the target reaction has FeO and O2 as reactants and Fe2O3 as a productArrange reactions 1 and 2 so that they also have FeO and O2 as reactants and Fe2O3 as a productEach reaction can be doubled, tripled, or multiplied by a half, etc.H values are then doubled, tripled, etc.If a reaction is reversed the H value is changed to the opposite sign4 FeO Fe O21.++1088 kJ2.1
30 Hess’s Law Given the following equations and H values calculate H for the reaction below:
31 Hess’s Law-The + sign of the H value tells us that the reaction is endothermic.-The reverse reaction is exothermic, i.e.
32 Standard Enthalpy of Formation Thermochemical standard state conditionsThe thermochemical standard T = KThe thermochemical standard P = atm- Be careful not to confuse these values with STPThermochemical standard states of matterFor pure substances in their liquid or solid phase the standard state is the pure liquid or solid- For aqueous solutions the standard state is 1.00 M concentrationFor gases the standard state is the gas at 1.00 atm of pressure- For gaseous mixtures the partial pressure must be 1.00 atm
33 Standard Molar Enthalpies of Formation, Hfo The enthalpy change for the formation of one mole of a substance formed directly from its constituent elements in their standard statesThe symbol for standard molar enthalpy of formation is Hfo (KJ/mol)The standard molar enthalpy of formation for MgCl2 is:
34 Standard Molar Enthalpies of Formation (Hfo) Standard molar enthalpies of formation have been determined for many substances and are tabulated in Appendix LStandard molar enthalpies of elements in their most stable forms at K and atm are zero.Example: The standard molar enthalpy of formation for phosphoric acid is kJ/mol. Write the equation for the reaction for which Hof = kJNote: P in standard state is P4 (s)Phosphoric acid in standard state is H3PO4(s)
35 Standard Molar Enthalpies of Formation Calculate the enthalpy change for the reaction of one mole of H2(g) with one mole of F2(g) to form two moles of HF(g) at 25oC and one atmosphere.
36 Standard Molar Enthalpies of Formation Calculate the enthalpy change for the reaction in which 15.0 g of aluminum reacts with oxygen to form Al2O3 at 25oC and one atmosphere.You do it!
37 Hess’s Law Version II For chemical reaction at standard conditions: the standard enthalpy change is the sum of the standard molar enthalpies of formation of the products minus the sum for the reactants- each enthalpy of formation is multiplied by its coefficient in the balanced chemical equationFinal - Initial
38 C3H8(g) + 5O2(g) 3CO2(g) + 4H2O(l) Hess’s LawCalculate the DH°f for the following reaction fromthe data in appendix L:C3H8(g) + 5O2(g) CO2(g) + 4H2O(l)
39 Hess’s Law Given the following information, calculate Hfo for H2S(g) You do it!
40 CalorimetryAn experimental technique that measures the heat transfer during a chemical or physical processConstant pressure calorimetry:A styrofoam coffee-cup calorimeter isused to measure the amount of heatproduced (or absorbed) in a reactionThis is one method to measureqP (called DH) for reactions in solutionqreaction + qsolution = 0Note: Assuming no heat transfer to the surroundings
41 CH3COOH(aq) + NaOH(aq) NaCH3COO(aq) + H2O(l) CalorimetryIf an exothermic reaction is performed in a calorimeter, the heat evolved by the reaction is determined from the temperature rise of the solutionThis requires a two part calculationWhen we add mL of M NaOH at oC to mL of M CH3COOH already in the calorimeter at the same temperature, the resulting temperature is observed to be oC. Determine heat of reaction and then calculate the change in enthalpy (as KJ/mol) for the production of NaCH3COO.CH3COOH(aq) + NaOH(aq) NaCH3COO(aq) + H2O(l)
42 Calorimetry Constant volume calorimetry: Or Bomb calorimetry measures measure the amount of heat produced (or absorbed) in a chemical reaction-this method is used for measuring qv (DE)qreaction + qbomb + qwater = 0