Download Thermochemistry: Energy Flow and Chemical

Thermochemistry: Energy Flow and Chemical Change
6.1 Forms of Energy and Their Interconversion
When energy is transferred from one object to another, it appears as work and/or heat
∆E = Efinal – Einitial = Eproducts – Ereactants
∆ – refers to the final state of the system minus the initial state
Because the total energy must be conserved, a change in the energy of the system is always
accompanied by an opposite change in the energy of the surroundings
Heat – thermal energy; symbol q; is the energy transferred between a system and its surroundings as a
result of a difference in their temperatures only
Work – symbol w; the energy transferred when an object is moved by a force
• ∆E = q + w
• We define the sign of the energy transfer from the system's perspective. Energy coming into the
system is positive. Energy going out from the system is negative
Energy Transfer as Heat Only: w = 0, therefore, ∆E = q
1. Heat flowing out from a system: The systems energy decreases as heat flows out from the
system, so the final energy of the system is less thanits intial energy. Heat was lost by the
system, so q is negative, and therefore ∆E is negative
2. Heat flowing into a system: energy is transferred into the system, so the final energy of the
system is higher than its initial energy. Heat was gained by the system, so q is positive, and
therefore ∆E is positive
Energy Transfer as Work Only: q = 0, therefore, ∆E = w
1. Work done by a system; some of the internal energy of the system is doing work on the
surroundings; Energy is lost by the system as work, so w is negative and ∆E is negative.
Pressure-Volume Work – the type of work in which a volume changes against an external
pressure
2. Work done on a system; the system gains energy because work is done on the system by the
surroundings; w is positive, so ∆E is positive
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Law of Conservation of Energy (First Law of Thermodynamics) – the total energy of the universe
is constant
• the energy of the system plus the energy of the surroundings remains constant: energy is
conserved
• ∆Euniverse = ∆Esystem + ∆Esurroundings = 0
Joule (J) – 1J = 1 kg·m²/s² – both heat and work are expressed in Joules
• The work done on a mass is the force (F) times the distance (d) that the mass moves
• w=Fxd
• A force changes the velocity of (accelerates) a mass
• Velocity has units of m/s, so acceleration (a) has units of m/s²
• Force therefore, has units of mass (kg) times acceleration
• F = m x a in units of kg·m²
• w = F x d in units of (kg·m/s²) x m = kg·m²/s² = J
calorie (cal) – defined originally as the quantity of energy needed to raise the temperature of 1 g of
water by 1º C; the calorie is now defined in terms of the J – 1 cal = 4.184 J or 1 J = .2390 cal
• The nutritional Calorie is actually a kilocalorie (kcal)
• Btu – British Thermal Unit = 1055 J
State Function – a property dependent only on the current state of the system (its composition,
volume, pressure, and temperature), not on the path the system took to reach that state; the current state
depends only on the difference between the final and initial states
• For a given change, ∆E (sum of q and w) is constant, even though q and w can vary
6.2 Enthalpy: Heats of Reaction and Chemical Change
We find the quantity of PV work done by multiplying the external pressure by the change in
volume of the gas – w = –P∆V – the result is negative as the work is being done on the
surroundings
Enthalpy (H) – the internal energy plus the product of the pressure and volume: H = E + PV
Change in Enthalpy (∆H) – change in internal energy plus the product of the constant pressure and
the change in volume: ∆H = ∆E + P∆V
• at constant pressure we denote q as qP and solve for it: ∆E = q + w = q + (–P∆V) = q – P∆V
• qP = ∆E + P∆V = ∆H
• the change in enthalpy equals the heat gained or lost at constant pressure
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Comparing ∆E and ∆H:
• Reactions that do not involve gases: Because liquids and solid undergo very small volume
changes, ∆V = 0; thus P∆V = 0 and ∆H ≈ ∆E
• Reactions in which the amount (mol) of gas does not change: when the total amount of gaseous
reactants equals the total amount of gaseous products, ∆V = 0, so P∆V = 0 and ∆H = ∆E
• Reactions in which the amount (mol) of as does change: in these cases P∆V ≠ 0. However, qP is
usually so much larger than P∆V that ∆H is very close to ∆E; ∆H ≈ ∆E
Heat of reaction, ∆Hrxn – enthalpy change of a reaction; refers to Hfinal – Hinitial
Exothermic Process – releases heat and results in a decrease in the enthalpy of the system
• Hfinal < Hinitial and ∆H < 0
Endothermic Process – absorbs heat and results in an increase in the enthalpy of the system
• Hfinal > Hinitial and ∆H > 0
Heat of Formation, ∆Hf – when 1 mol of a compound is produced from its elements
Heat of Fusion, ∆Hfus – when 1 mol of a substance melts
Heat of Vaporization, ∆Hvap – when 1 mol of a substance vaporizes
6.3 Calorimetry: Laboratory Measurement of Heats of Reaction
• The quantity of heat absorbed by an object is proportional to its temperature change.
Heat Capacity – the quantity of heat required to change its temperature by 1 K.
• Heat Capacity = q/∆T [in units of J/K]
Specific Heat Capacity (c) – the quantity of heat required to change the temperature of 1 gram of a
substance by 1 K.
• Specific Heat Capacity = q/(mass x ∆T) [in units of J/g-K]
• q = c x mass x ∆T
• If we know the c of the substance we're using, we can measure its mass and temperature change
and calculate the heat absorbed or released.
Molar Heat Capacity is the quantity of heat required to change the temperature 1 mole of a substance
by 1K.
• Molar Heat Capacity (C) = q/(moles x ∆T) [in units of J/mol-K]
Calorimeter – a device used to measure the heat released (or absorbed) by a physical or chemical
process.
• The calorimeter acts as the surroundings that changes temperature when heat is transferred to or
from the system.
“coffee cup calorimeter” – used to measure the heat transferred in processes open to the atmosphere.
• The heat lost by the system is equal in magnitude but opposite in sign to the heat gained by the
surroundings. We assume in a coffee cup calorimeter, that all the heat is gained by the water, but
some must be gained by the stirrer, thermometer, among other variables.
Bomb calorimeter – designed to measure very precisely the heat released in a combustion reaction
• Constant Volume Calorimetry allows for more precise work; the heat capacity of the entire
colorimeter must be known
6.5 Hess's Law of Heat Summation
Hess's Law of Heat Summation: the enthalpy change of an overall process is the sum of the enthalpy
changes of its individual steps.
Calculating an unknown ∆H involves three steps:
1. Identify the target equation, the step whose ∆H is unknown, and note the number of moles of
each reactant and product
2. Manipulate the equations with known ∆H values so that the target numbers of moles of
reactants and products are on the correct sides. Remember to:
▪ Change the sign of ∆H when you reverse an equation
▪ Multiply numbers of moles and ∆H by the same factor
3. Add the manipulated equations to obtain the target equation. All substances except those in the
target equation must cancel. Add their ∆H values to obtain the unknown ∆H
6.6 Standard Heats of Reaction (∆Hºrxn)
An element in its standard state is assigned a ∆Hºf of 0
Most compounds have a negative ∆Hºf – that is most compounds have exothermic formation
reactions under standard conditions: heat is given off when the compound forms.
Determining ∆Hºf from ∆Hºf values of Reactants and Products:
1. Each reactant decomposes to its elements. This is the reverse of the formation reaction for each
standard enthalpy change is -∆Hºf
2. Each product forms from its elements. This step is the formation reaction for each product, so
each standard enthalpy change is ∆Hºf
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the standard heat of reaction is the sum of the standard heats of formation of the products minus
the sum of the standard heats of formation of the reactants
∆Hºrxn = ∑m∆Hºf(products) – ∑n∆Hºf(reactants)
m and n are the amounts (mol) of the products and reactants indicated by the coefficients from
the balanced equation