Introduction to Thermodynamics

• Thermodynamics is the study of energy and its transformations.
• Thermochemistry is a subdiscipline about chemical reactions and energy changes.

Kinetic Energy

• Energy of motion present in all particles and measurable as temperature.
• A rolling bowling ball and the heat from a hot stove are examples.

Potential Energy

• Stored energy, like that in chemical bonds due to atomic arrangements.
• A stretched rubber band and food are examples.

Units of Energy

• Energy is measured in joules (J), kilojoules (kJ), calories (cal), and nutritional calories (Cal or kcal).
• Conversions: 4184 J = 4.184 kJ = 1000 cal = 1 Cal = 1 kcal

System and Surroundings

• System: The specific part of the universe under study.
• Surroundings: Everything else in the universe outside the system.

Types of Systems

• Open System: Allows exchange of both matter and heat with surroundings.
• A cup of coffee is an example.
• Closed System: Allows heat exchange but not matter exchange.
• A sealed container of boiling water is an example.
• Isolated System: Prevents both matter and heat exchange.
• A perfectly insulated thermos is a close approximation.
• A closed system in chemistry allows energy exchange but prevents matter exchange.

Work and Heat

• Work (W): Done when an unbalanced force moves an object over a distance, W = Fd.
• Inflating balloons is an example.
• Heat (q): Energy transferred from a hotter object to a colder one.
• People in a polar plunge is an example.

First Law of Thermodynamics

• Energy changes form, but the total amount of energy remains constant.
• Hydroelectric dams and nuclear power plants illustrates this.

Internal Energy (ΔE)

• The total kinetic and potential energy of a system's components, impossible to measure directly.
• Calculated as the difference between final and initial energy (ΔE = Efinal – Einitial).
• ΔE = q + w describes changes based on heat transfer (q) and work done (w).
• Positive q: system gains heat.
• Negative q: system loses heat.
• Positive w: work done on the system.
• Negative w: system does work

Endothermic Processes

• Require energy input which they absorb heat.
• Melting, boiling, and sublimation, and a flowers blooming are examples.

Exothermic Processes

• Release energy, giving off heat.
• Freezing, condensation, and deposition.
• A fire perfectly demonstrate this process.

Enthalpy (H)

• At constant pressure, enthalpy is the sum of a system's internal energy (E) and the product of its pressure (P) and volume (V): H = E + PV.
• Enthalpy change (ΔH) at constant pressure equals the heat transferred (qp).
• Positive ΔH: system gained heat (endothermic).
• Negative ΔH: system lost heat (exothermic).
• Enthalpy is an extensive property, depending on the amount of material.
• ΔHrxn = Hproducts - Hreactants.

Calorimetry

• Measures heat flow using a calorimeter.
• Heat capacity: the heat needed to raise an object's temperature by 1 K (or 1°C).
• Molar heat capacity: the heat needed to raise the temperature of 1 mole of a substance by 1 K (q = nCmΔT).
• Specific heat (capacity): the heat needed to raise the temperature of 1 gram of a substance by 1 K (q = mcΔT).
• Molar heat capacity is the product of molar mass and specific heat.
• Heat gained or lost is calculated using q = mCpΔT (same state) or q = mCx (changes state).
• Cx is the heat of fusion (solid to liquid) or the heat of vaporization (liquid to gas).

Heating Curve

• Shows temperature changes with added heat.
• Flat portions: phase transitions (melting, boiling).
• Temperature remains constant during phase transitions.
• Sloped portions: temperature changes within a single phase.
• Equations q = mcΔT and q = nΔH are shown.

Calorimeters

• Coffee-cup calorimeter: constant pressure.
• qabsorbed = -qreleased (heat released by the reaction is absorbed by the solution).
• Specific heat capacity of dilute aqueous solutions is approximately the same as that of water (4.18 J/g·K).
• Bomb calorimeter: constant-volume calorimetry.
• The heat released (q) by the combustion reaction is absorbed by the calorimeter and the water surrounding the bomb.
• Heat capacity of the calorimeter (Ccal) must be known to calculate the heat released using the equation q = Ccal x ΔT, where ΔT is the temperature change.

Hess's Law

• Ttotal enthalpy change for a reaction is independent of the pathway taken.
• Hess's Law states that the enthalpy change of a physical or chemical process depends only on the initial (reactants) and final (products) states.
• Calculating energy change can be done by algebraically combining chemical equations or by using standard enthalpies of formation.
• Enthalpy of formation (ΔHf) is the enthalpy change associated with forming one mole of a compound from its constituent elements in their standard states. Standard state is at 25°C (298 K), and most stable form.

Standard Enthalpy of Formation (ΔH°f)

• ΔH°f for the most stable form of any element in its standard state is zero.
• ΔHf values are for one mole of substance, so the units are typically kJ/mol.
• The standard enthalpy of a reaction (ΔH°rxn) is the change in enthalpy of a reaction when all substances are in their standard states (at 25°C).
• ΔH°rxn = ΣnΔH°f(products) - ΣmΔH°f(reactants)

Fuel Value

• Energy released when 1 gram of a material is combusted.
• Measured using calorimetry.

Energy in the Body

• The human body utilizes glucose (C₆H₁₂O₆) as its primary energy source.
• Insulin facilitates the movement of glucose from the bloodstream into cells.
• Carbohydrates: Provide 4 kcal/g of energy. Rapidly broken down into glucose.
• Fats: Provide 9 kcal/g of energy. Digested more slowly and easily stored.
• Proteins: Provide 4 kcal/g of energy. Contain nitrogen, converted to urea [(NH₂)₂CO] during digestion.

Fossil Fuels

• Coal, petroleum, and natural gas, derived from decayed ancient organisms.
• Nonrenewable resources with slow formation rates.
• Coal gasification produces combustible gases.

Nuclear Energy

• Derived from nuclear fission or fusion.
• Nonrenewable but produces significant energy.
• Presents a challenge of hazardous radioactive waste disposal.

Renewable Energy Sources

• Examples include solar and wind energy.
• Geothermal energy harnesses heat from the Earth's interior.
• Hydroelectric energy uses the power of moving water.
• Biomass energy utilizes organic matter.
• Solar heating is used to generate carbon monoxide (CO) and hydrogen (H₂).
• Solar energy has its limitations: it's a dilute energy source, storing it for later use is challenging, and its availability fluctuates.