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Brown, LeMay Ch 5 AP Chemistry Monta Vista High School

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1 Brown, LeMay Ch 5 AP Chemistry Monta Vista High School
Thermochemistry Brown, LeMay Ch 5 AP Chemistry Monta Vista High School

2 James Prescott Joule (1818-1889)
5.1: Thermochemistry From Greek therme (heat); study of energy changes in chemical reactions Energy: capacity do work or transfer heat Joules (J) or calories (cal); 1 cal = J Kinetic: energy of motion; dependent on mass & velocity Applies to motion of large objects & molecules Linked to thermal energy (object’s T above 0 K) James Prescott Joule ( ) English, studied under John Dalton, took over his father’s job as brewer where he got interested in steam power.

3 Potential: stored in “fields” (gravitational and electrical/magnetic); dependant on position relative to another object Applies to large objects where gravity is overriding force, but not significantly to molecules where gravity is negligible and electrostatic forces dominate Associated with chemical energy; stored in arrangement of atoms or subatomic particles (electrostatic & nuclear forces, bonding between atoms) Link for Kinetic and Potential Energy Conversion

4 Vocabulary System: “isolated” portion of study (typically just the chemicals in a reaction) Surroundings: everything else (container, room, Earth, etc.) Closed system: easiest to study because exchanges energy with surroundings but matter is not exchanged. Force: a push or pull on an object Work: energy transferred to move an object a certain distance against a force: W = (F)(d) Heat: energy transferred from a hotter object to a colder one

5 5.2: Laws of Thermodynamics
“0th Law”: 2 systems are in thermal equilibrium when they are at the same T. Thermal equilibrium is achieved when the random molecular motion of two substances has the same intensity (and therefore the same T.) 1st Law: Energy can be neither created nor destroyed, or, energy is conserved. 2nd and 3rd Laws: discussed in Ch. 19

6 Internal energy, E DE = Efinal - Einitial Includes:
Translational motion Rotational motion of particles through space Internal vibrations of particles. It is difficult to measure all E, so the change in internal energy (DE) is typically measured: DE = Efinal - Einitial DE > 0 Increase in energy of system (gained from surroundings) DE < 0 Decrease in energy of system (lost to surroundings)

7 First Law of Thermodynamics
When a system undergoes a chemical or physical change, the change in internal energy (E) is equal to the heat (q) added or liberated from the system plus the work (w) done on or by the system: DE = q + w `

8 Sign Conventions q > 0 Heat is added to system q < 0 Heat is removed from system (into surroundings) w > 0 Work done to system w < 0 System does work on surroundings

9 Example: Octane and oxygen gases combust within a closed cylinder in an engine. The cylinder gives off 1150 J of heat and a piston is pushed down by 480 J during the reaction. What is the change in internal energy of the system? q is (-) since heat leaves system; w is (-) since work is done by system. Therefore, DE = q + w = (-1150 J) + (-480 J) = J 1630 J has been liberated from the system (C8H18 and O2) added to the surroundings (engine, atmosphere, etc.)

10 Heat & reactions E E

11 State functions Property of a system that is determined by specifying its condition or its “state” The value of a state function depends only on its present state and not on the history of the sample. T & E are state functions. Consider 50 g of water at 25°C: EH2O does not depend on how the water got to be 25°C (whether it was ice that melted or steam that condensed or…) Work (w) and heat (q) are not state functions because the ratio of q and w are dependent on the scenario. Consider the combustion of gasoline in a car engine vs. burning in the open.

12 5.5: Calorimetry q = C DT q = m c DT Measurement of heat flow
Heat capacity, C: amount of heat required to raise T of an object by 1 K q = C DT Specific heat (or specific heat capacity, c): heat capacity of 1 g of a substance q = m c DT Ex: How much energy is required to heat 40.0 g of iron (c = 0.45 J/(g K) from 0.0ºC to 100.0ºC? q = m c DT = (40.0 g)(0.45 J/(g K))(100.0 – 0.0 ºC) = 1800 J

13 Calorimetry is an experimental technique used to measure the heat transferred in a physical or chemical process. The apparatus used in this procedure is called as a “Calorimeter”. There are two types of calorimeters- Constant pressure (coffee cup) and constant volume (bomb calorimeter). We will discuss constant pressure calorimeter in detail. Constant Pressure Calorimeter: The coffee cup calorimeter is an example of this type of calorimeter. The system in this case is the “contents” of the calorimeter and the surroundings are cup and the immediate surroundings.

14 qrxn + q solution = 0 During the rxn:
where qrxn is the heat gained or lost in the chemical reaction and qsolutionis the heat gained or lost by solution. Heat exchange in this system (qrxn), is equal to enthalpy change. Assuming no heat transfer takes place between the system and surroundings, qrxn + q solution =


16 Mg (s) + 2HCl (aq)  H2 (g) + MgCl2 (aq)
Sample Problem #1: g of magnesium chips are placed in a coffee-cup calorimeter and ml of 1.00 M HCl is added to it. The reaction that occurs is: Mg (s) + 2HCl (aq)  H2 (g) + MgCl2 (aq) The temperature of the solution increases from 22.2oC (295.4 K) to 44.8 oC (318.0 K). What’s the enthalpy change for the reaction, per mole of Mg? (Assume specific heat capacity of solution is 4.20 J/(g * K) and the density of the HCl solution is 1.00 g/ml.) Ans x 105 J/mol Mg

17 Sample Problem #2: 200. ml of M HCl is mixed with the same amount and molarity of NaOH solution, inside a coffee-cup calorimeter. The temperature of the solutions before mixing was oC, and oC after mixing and letting the reaction occur. Find the molar enthalpy of the neutralization of the acid, assuming the densities of all solutions are 1.00 g/ml and their specific heat capacities are 4.20 J/(g * K).

18 5.3: Enthalpy, H Since most reactions occur in containers open to the air, w is often negligible. If a reaction produces a gas, the gas must do work to expand against the atmosphere. This mechanical work of expansion is called PV (pressure-volume) work. Enthalpy (H): change in the heat content (qp) of a reaction at constant pressure H = E + PV H = E + PV (at constant P) H = (qp + w) + (-w) H = qp

19 Sign conventions H > 0 Heat is gained from surroundings + H in endothermic reaction H < 0 Heat is released to surroundings - H in exothermic reaction

20 5.4: Enthalpy of Reaction (Hrxn)
Also called heat of reaction: Enthalpy is an extensive property (depends on amounts of reactants involved). Ex: CH4 (g) + 2 O2 (g)  CO2 (g) + 2 H2O (l) Hrxn = kJ Combustion of 1 mol CH4 produces 890. kJ … of 2 mol CH4 → (2)(-890. kJ) = kJ What is the H of the combustion of 100. g CH4?

21 Hreaction = - Hreverse reaction
CH4 (g) + 2 O2 (g)  CO2 (g) + 2 H2O (l) H = kJ CO2 (g) + 2 H2O (l)  CH4 (g) + 2 O2 (g) H = kJ

22 CO2 (g) + 2 H2O (g) CH4 (g) + 2 O2 (g) H = 802 kJ
Hrxn depends on states of reactants and products. CO2 (g) + 2 H2O (g) CH4 (g) + 2 O2 (g) H = 802 kJ 2 H2O (l)  2 H2O (g) H = 88 kJ So: CO2 (g) + 2 H2O (l)  CH4 (g) + 2 O2 (g) H = 890. kJ CH4 (g) + 2 O2 (g) CH4(g) + 2 O2(g) 802 kJ 890. kJ CO2 (g) + 2 H2O (g) 2 H2O (g) 88 kJ 2 H2O (l) CO2(g) + 2 H2O(l)

23 Hess’s Law: http://www. media. pearson. com
If a rxn is carried out in a series of steps, Hrxn =  (Hsteps) = H1 + H2 + H3 + … Germain Hess ( ) Ex. What is DHrxn of the combustion of propane? C3H8 (g) + 5 O2 (g)  3 CO2 (g) + 4 H2O (l) 3 C (s) + 4 H2 (g)  C3 H8 (g) H1 = kJ C (s) + O2 (g)  CO2 (g) H2 = kJ H2 (g) + ½ O2 (g)  H2O (l) H3 = kJ C3H8 (g)  3 C (s) + 4 H2 (g) H1 = kJ Born in Switzerland, moved to St Petersburg at age 3 with his family. His father was an artist. Hess was also interested in geology. 3[ ] ( ) 4[ ] ( ) Hrxn = ( ) + 4( ) = kJ

24 5.7: Enthalpy of Formation (Hf)
Formation: a reaction that describes a substance formed from its elements NH4NO3 (s) Standard enthalpy of formation (Hf): forms 1 mole of compound from its elements in their standard state (at 298 K) C2H5OH (l) Hf = kJ Hf of the most stable form of any element equals zero. H2, N2 , O2 , F2 , Cl2 (g) Br2 (l), Hg (l) C (graphite), P4 (s, white), S8 (s), I2 (s) Ex: 2 N2 (g) + 4 H2 (g) + 3 O2 (g)  2 2 C (graphite) + 3 H2 (g) + ½ O2 (g) 

25 The complete (long) way to calculate the ENTHALPY of any reaction is to take apart the reactant molecules piece by piece to elements, and then build the product molecules from all of the pieces, bond by bond.

26 A GREAT SHORTCUT exists however:
To estimate the enthalpy of a reaction ΔRH ° we need only to know the enthalpies of formation of the reactants and the products. (we calculate the large RED arrow by the difference between the known small GREEN arrows)

27 An example: The combustion of liquid benzene :
Firstly, we need to define the overall reaction, and balance it to 1 mole of the molecule we’re interested in : 1 C6H6 (l) + 15/2 O2 (g)  6 CO2 (g) + 3 H2O (l) ΔRH ° = ??? ΔRH ° = ΔfH ° { products } ΔfH ° { reactants } Secondly, we need to look up the ΔfH ° values corresponding to both the reactants and the products, and multiply by the moles:

28 Example Contd: ΔfH ° { products } = 6 ( kJ/mol ) + 3 ( kJ/mol ) = kJ/mol to form the products ΔfH ° { reactants } = ( + 49 kJ/mol ) + 15/2 ( zero ) = kJ/mol to form the reactants Overall reaction ΔRH ° = kJ/mol

29 Enthalpies of combustion ΔCH ° are also commonly tabulated :
( converted into enthalpies per gram, for convenience ) methane ΔCH ° = kJ/g octane ΔCH ° = kJ/g dodecane ΔCH ° = kJ/g methanol ΔCH ° = kJ/g glucose ΔCH ° = kJ/g carbohydrates ΔCH ° ~ kJ/g tristearin ΔCH ° = kJ/g human beings at age 20 need around 10,000 kJ per day of energy from combustions (~12,000 kJ for men, and ~9000 kJ for women).

30 Calculating heat of reaction from Heat of formation
Ex. Combustion of propane: C3H8 (g) + 5 O2 (g)  3 CO2 (g) + 4 H2O (l) Given: Compound Hrxn (kJ/mol) C3H8 (g) CO2 (g) H2O (l) H2O (g) Hrxn = [3( ) + 4( )] – [1( ) + 5(0)] = kJ

31 Hrxn ≈  (Hbonds broken) -  (Hbonds formed)
8.9: Bond enthalpy: Amount of energy required to break a particular bond between two elements in gaseous state. Given in kJ/mol. Remember, breaking a bond always requires energy! Bond enthalpy indicates the “strength” of a bond. Bond enthalpies can be used to figure out Hrxn . Ex: CH4 (g) + Cl2 (g) → CH3Cl (g) + HCl (g) DHrxn = ? 1 C-H & 1 Cl-Cl bond are broken (per mole) 1 C-Cl & 1 H-Cl bond are formed (per mole) Hrxn ≈  (Hbonds broken) -  (Hbonds formed) Note: this is the “opposite” of Hess’ Law where Hrxn = DHproducts – Dhreactants Bond Enthalpy link

32 Ex: CH4 (g) + Cl2 (g) → CH3Cl (g) + HCl (g) DHrxn = ?
Bond Ave DH/mol Bond Ave DH/mol C-H Cl-Cl 242 H-Cl C-Cl 328 C-C C=C 614 Hrxn ≈  (Hbonds broken) -  (Hbonds formed) Hrxn ≈ [(1(413) + 1(242)] – [1(328) + 1(431)] Hrxn ≈ -104 kJ/mol Hrxn = kJ/mol (actual) Note: 2 C-C ≠ 1 C=C 2(348) = 696 kJ ≠ 614 kJ

33 Ex: CH4(g) + Cl2(g) → CH3Cl(g) + HCl(g) DHrxn=?
*CH3(g) + H(g) + 2 Cl(g) Absorb E, break 1 C-H and 1 Cl-Cl bond Release E, form 1 C-Cl and 1 H-Cl bond H CH4(g) + Cl2(g) Breaking a bond decreases stability, therefore releasing energy. Creating a bond increases stability, therefore absorbing energy. The formation of ADP from ATP has a net release of energy, but the breaking of the peptide bond itself does not release energy. CH3Cl (g) + HCl (g) DHrxn Hrxn =  (Hbonds broken) +  (- Hbonds formed) Hrxn =  (Hbonds broken) -  (Hbonds formed)


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