Thermochemistry

CHAPTER 6: THERMOCHEMISTRY

THERMOCHEMISTRY & DEFINITIONS

• Thermochemistry: The study of relationships between chemistry and energy.

  • Focus on the exchange of energy during chemical and physical changes.

  • Quantification of energy relationships in chemical reactions.

NATURE OF ENERGY

• Energy: Capacity to do work.

  • Work: Defined as force acting through a distance.

• Types of energy include:

  • Kinetic Energy: Energy of moving objects.

  • Potential Energy: Energy stored due to position or state.

  • Thermal Energy: Related to temperature and heat.

  • Chemical Energy: Stored in the bonds between atoms.

CONSERVATION OF ENERGY

• Law of Conservation of Energy: Energy cannot be created or destroyed, only transformed from one form to another.

SYSTEM AND SURROUNDINGS

• System: The specific part of the universe under study (e.g. chemicals in a beaker).

• Surroundings: Everything outside the system that can exchange energy with it (e.g. solutions, air).

• Types of Systems:

  • Open System: Exchanges both energy and matter with surroundings.

  • Closed System: Exchanges energy but not matter.

  • Isolated System: No exchange of energy or matter.

UNITS OF ENERGY

• SI Unit of Energy: Joules (J) defined as $1 ext{ kg m}^2/ ext{s}^2$.

• Calorie (cal): Energy required to raise 1 g of water by 1°C.

• Kilowatt-hour (kWh): Another large unit, often used in home energy costs.

ENERGY CONVERSION FACTORS

1. $1 ext{ cal} = 4.184 ext{ J}$

2. $1 ext{ Cal} ext{ (kilocalorie)} = 1000 ext{ cal} = 4184 ext{ J}$

3. $1 ext{ kWh} = 3.60 imes 10^6 ext{ J}$

INTERNAL ENERGY

• Internal Energy (E): Sum of kinetic and potential energies of all particles in a system.

  • $E = KE + PE$

  • Kinetic Energy: Associated with motion.

  • Potential Energy: Associated with chemical bonding and molecular attraction.

• State Function: Internal energy is a state function, independent of the process to reach that state.

CHANGES IN INTERNAL ENERGY

• Change in internal energy when a system shifts between states:

  • $riangle E = E<em>{ ext{products}} - E</em>{ ext{reactants}}$

• Internal energy can be understood from exothermic and endothermic reactions:

  • Exothermic: Reactants > Products, riangle E < 0 (energy flows to surroundings).

  • Endothermic: Reactants < Products, riangle E > 0 (energy flows into the system).

FIRST LAW OF THERMODYNAMICS

• Changes in internal energy ($riangle E$) are more crucial than absolute values.

• Internal energy change can be quantified via heat (q) and work (w):

  • $riangle E = q + w$

HEAT

• Heat: Energy transfer due to temperature differences.

  • If heat is lost, q < 0.

  • If heat is gained, q > 0.

WORK

• Work: Energy exchange resulting from moving an object through a distance.

  • $w = ext{Force} imes ext{Distance}$.

  • If work is done by the system, w < 0; if done on the system, w > 0.

SIGN CONVENTIONS FOR q, w, AND $riangle E$

EXAMPLES OF ENERGY EXCHANGE

• Example of identifying heat/work signs:

  • Ice melting in a drink loses heat (-).

  • Metal cylinder rolling uphill gains work (+).

  • Condensing steam releases heat (-).

QUANTIFYING HEAT

• Heat Transfer: Defined by the temperature difference between two substances:

  • Heat always flows from high to low temperature.

• Quantification: $q = m imes C_s imes riangle T$

  • Where $C_s$ is specific heat capacity, $m$ is mass, and $riangle T$ is the temperature change.

SPECIFIC HEAT CAPACITY

• Specific Heat Capacity (C_s): Amount of heat to raise the temperature of 1 g of substance by 1°C.

  • Units: $ext{J/g°C}$.

• Molar Heat Capacity: Heat required to raise 1 mole by 1°C.

  • Units: $ext{J/mol°C}$.

CALCULATING HEAT EXAMPLES

1. Warm 1.5 kg of sand from 25.0°C to 100.0°C:

   • $q = (1.5 ext{ kg})(0.84 ext{ kJ/kg°C})(75.0 ext{ °C}) = 94.5 ext{ kJ}$.

2. Final temperature of aluminum in a calorimeter:

   • Given heat absorbed: 725 J, mass: 55.0 g, and specific heat: 0.903 J/g°C, find final temperature.

HEAT TRANSFER

• Upon combining two substances at different temperatures, heat transfers until thermal equilibrium is achieved.

• The relationship with heat transfers $(q<em>metal = -q</em>water)$ allows for the calculation of unknowns (mass, temperature changes).

PRESSURE-VOLUME WORK

• Pressure-volume work occurs when a gas expands or contracts against external pressure:

  • $w = -P riangle V$

  • Expansion (more moles produced) exhibits negative work, contraction (less moles produced) demonstrates positive work.

CONSTANT-VOLUME CALORIMETRY

• Measured heat exchange via temperature changes at constant volume conditions.

• A calorimeter measures heat and changes in internal energy through the relationship:

  • $q<em>{rxn} = -q</em>{cal}$

  • Heat absorbed by calorimeter equals heat released by the reaction.

ENTHALPY

• Enthalpy (H): Sum of internal energy and the product of pressure and volume:

  • $H = E + PV$

• Change in enthalpy is important under constant pressure conditions:

  • $riangle H = q_p$

• Endothermic reactions absorb heat (positive $riangle H$).

• Exothermic reactions release heat (negative $riangle H$).

THERMOCHEMICAL EQUATIONS

• Represents the enthalpy change for a reaction ($riangle H_{rxn}$).

• Enthalpy of formation relates to the heat absorbed or released according to stoichiometric amounts.

CALORIMETRY EXAMPLES

• Calorimetry techniques adapt to experimental needs, measuring relevant changes in heat due to reactions.

ENTHALPIES OF FORMATION

• Defined standard states for measurements.

• Standard enthalpy of formation ($riangle H_f°$) for a compound formed from its elements under standard conditions is defined as zero for elements in their standard state.

• Calculating reactions through their thermodynamic cycles can yield valuable thermodynamic properties.

• Example Calculations:

  • Finding $riangle H_{rxn}$ based on enthalpy changes in multi-step processes by applying Hess's Law.

End of Notes.