Energy and Thermochemistry

Energy and Thermochemistry

• Energy is the capacity to perform work or transfer heat.
• Thermodynamics involves studying energy and its various forms.
• Thermochemistry is the study of chemical reactions and the energy changes, mainly heat, associated with them.

Nature of Chemical Energy

• Electrostatic potential energy in charged particles constitutes the most significant form of potential energy: Eel = kQ1Q2/d.
• The unit of energy is the Joule (J): 1 J = 1 kg⋅m²/s².
• Electrostatic attraction occurs between oppositely charged ions.
• Energy is released when chemical bonds form (Eel < 0), and consumed when chemical bonds break (Eel > 0).

First Law of Thermodynamics

• Energy can change forms but cannot be created or destroyed.
• Chemical energy is converted to heat to warm homes.
• Green plants convert sunlight into chemical energy.

Definitions: System and Surroundings

• The system is the specific portion of the universe being studied, such as hydrogen and oxygen molecules in a reaction.
• The surroundings include everything else outside the system, like the cylinder, piston, and beyond.

Types of Systems

• Open system: Exchanges both heat and mass with its surroundings.
• Closed system: Exchanges only heat but not mass with its surroundings.
• Isolated system: Exchanges neither heat nor mass with its surroundings.

Internal Energy

• The change in internal energy (ΔE) is the difference between the final and initial energies of a system: ΔE = Efinal - Einitial.
• Delta E is a state function, meaning it only depends on the initial and final states and not the path taken.

Thermodynamic Quantities

• Thermodynamic quantities have three parts: a number, a unit, and a sign.
• A positive ΔE means the system gains energy from the surroundings.
• A negative ΔE means the system loses energy to the surroundings.
• Internal energy (E) is the sum of kinetic and potential energies of a system's components, but only its changes (ΔE) are generally known.

Internal Energy Changes

• If ΔE < 0 and Efinal < Einitial, the system releases energy to the surroundings.
• If ΔE > 0 and Efinal > Einitial, the system absorbs energy from the surroundings.

Heat and Work

• Energy exchanged between a system and its surroundings occurs as either heat (q) or work (w).
• The change in internal energy is given by ΔE = q + w.
• A system gains heat (q > 0) if heat is added, and loses heat (q < 0) if heat is removed.
• Work done on the system is positive (w > 0), while work done by the system is negative (w < 0).

Sign Conventions

• A plus sign (+) for 'q' means the system gains heat.
• A minus sign (-) for 'q' means the system loses heat.
• A plus sign (+) for 'w' means work is done on the system.
• A minus sign (-) for 'w' means work is done by the system.
• A plus sign (+) for 'ΔE' means the system has a net gain of energy.
• A minus sign (-) for 'ΔE' means the system has a net loss of energy.
• Endothermic processes absorb heat from the surroundings, causing a temperature drop, and exothermic release heat.

State Functions

• The internal energy of a system is independent of the path by which the system achieves that state.
• Internal energy is a state function as it depends only on the present state, therefore ΔE depends only on Einitial and Efinal.
• Heat (q) and work (w) are not state functions; they are path-dependent.

Enthalpy

• At constant pressure, the heat flow is measured by enthalpy (H), defined as H = E + PV.
• The change in enthalpy (ΔH) at constant pressure is ΔH = ΔE + PΔV.
• The change in enthalpy is equal to the heat gained or lost (ΔH=q).

Enthalpy Changes

• A process is endothermic when ΔH is positive.
• A process is exothermic when ΔH is negative.

Pressure-Volume Work

• Mechanical work linked to volume changes in gases is typically the work done in chemical or physical changes
• Pressure volume work is measured as w = -PΔV using reactions in a vessel fitted with a piston.
• The work done by the system is negative.

Enthalpies of Reaction

• The enthalpy of reaction (ΔHrxn) is the difference between the enthalpy of the products and the enthalpy of the reactants: ΔHrxn = Hproducts - Hreactants.

Heat of Reaction

• The enthalpy of reaction is also known as heat of reaction.

Enthalpy Guidelines

• Enthalpy is an extensive property, depending on the amount (moles) of reactant.
• ΔH for the reverse reaction has the same magnitude but opposite sign to the forward reaction.
• The enthalpy change depends on the states (phases) of the reactants and products.

Calorimetry

• Calorimetry measures heat flow, since the exact enthalpy of reactants and products cannot be known.
• A calorimeter is the instrument used to measure heat flow.
• Heat capacity is the energy required to raise the temperature of a substance by 1 K (or 1°C).
• Specific heat is the heat capacity per gram of a substance.
• Molar heat capacity is the heat capacity per mole of a substance.

Constant-Pressure Calorimetry

• In a simple calorimeter with aqueous solutions, heat change for the system can be found by measuring the heat change for the water.
• The specific heat for water is 4.184 J/g⋅K, and this value is used for dilute solutions.
• The equation used is qsoln = Cs × msoln × ΔT = -qrxn.

Bomb Calorimetry

• Reactions can be carried out in a sealed "bomb"
• The heat absorbed (or released) by the water approximates the enthalpy change for the reaction: qrxn = -Ccal × ΔT.

Bomb Calorimetry (Constant Volume)

• Bomb calorimeters measure internal energy changes (ΔE), not ΔH, because volume is constant.
• The difference between ΔE and ΔH is very small and can be equated for most reactions.

Hess's Law

• Hess’s law states that if a reaction is carried out in a series of steps, ΔH for the overall reaction equals the sum of the enthalpy changes for the individual steps.
• Since H is a state function, ΔH is the same whether the reaction occurs in one step or a series of steps.

Enthalpies of Formation

• An enthalpy of formation (ΔHf) is the enthalpy change for the reaction where a compound forms from its constituent elements in their elemental forms.
• For example, for ammonia it's nitrogen(g) and hydrogen(g).
• The coefficient of the compound must be one.
• 1/2 N2(g) + 3/2 H2(g) → NH3(g).
• Use fractional coefficients when necessary.

Standard Enthalpies of Formation

• Standard enthalpies of formation (ΔH°f) are measured under standard conditions (25 °C and 1.00 atm pressure).
• Hess’s law is applied to calculate heat of reaction using tabulated values: ΔH = ΣnΔHf, products - ΣmΔH°f, reactants.
• In this equation, 'n' and 'm' are the stoichiometric coefficients.
• Heats of formation are obtained from published values, with ΔH°f for elements in their standard state defined as zero.

Breaking Down ΔH

• Imagine this reaction as occurring in three steps:
  • Decomposition of reactant propane to the elements: C3H8(g) → 3 C(graphite) + 4 H2(g)
  • Formation of the product CO2: 3 C(graphite) + 3 O2(g) → 3 CO2(g)
  • Formation of the product H₂O: 4 H2(g) + 2O2(g) → 4 H2O(l)
• The sum of these equations is the overall equation.

Bond Enthalpies

• The enthalpy associated with breaking one mole of a particular bond in a gaseous substance is bond enthalpy.
• Bond enthalpy is always positive because energy is required to break chemical bonds.
• The greater the bond enthalpy, the stronger the bond.
• Energy is released when a bond forms between gaseous fragments.
• Chemical reactions can be predicted using bond energies.
• Add bond energy values for all bonds made (+).
• Subtract bond energy values for all bonds broken (-).
• The result is an estimate of ΔH.

Energy in Foods

• The energy released when one gram of food is combusted is its fuel value.
• Most of the energy in foods comes from carbohydrates, fats, and proteins.
• Carbohydrates contain 17kJ/g; Fats contain 38 kJ/g.
• Proteins produce 17 kJ/g.

Energy in Fuels

• The majority of the energy consumed comes from fossil fuels.
• Nuclear fission accounts for 8.6% of U.S. energy needs.
• Renewable energy sources account for 9.9% of U.S. energy needs.