Chemical Energetics Flashcards

Key vocabulary and concepts related to chemical energetics

Pressure

Force per Area = N m* = Pa (pascal)

Boyle's Law

For a fixed amount of gas at a constant temperature, the gas volume is inversely proportional to the gas pressure.

Charles's Law

The volume of a fixed amount of gas at constant pressure is directly proportional to the Kelvin (absolute) temperature.

Avogadro's Law

Equal volumes of different gases, at the same temperature and pressure, contain equal numbers of molecules.

Ideal Gas Constant (R)

The constant R is the same for all ideal gases, but absolute value depends on the type of problem and the units needed in the solution.

Normal Temperature and Pressure (NTP)

0 °C (273.15 K) and 1 atm (101.33 kPa)

Density (d)

The mass occupied per unit volume

Dalton's Law of Partial Pressures

The total pressure exerted by a mixture of gases is the sum of all the partial pressures of these gases.

Mole fraction

The number of moles of an individual constituent in a gas mixture divided by the total number of moles in the mixture.

Vapour Pressure

The partial pressure of water vapour.

Effusion

The flow of gas molecules at low pressures through tiny pores or pinholes in a container

Graham's Law

The rate of effusion of a gas is inversely proportional to the square root of its molar mass.

Thermodynamics

The study of the energy changes involved in physical and chemical processes

Thermochemistry

The branch of thermodynamics that investigates the heat flow that occurs during chemical reactions.

System

The part of the universe that one wants to study.

Surroundings

The remaining parts of the universe that can interact with a system under study.

Open System

A system which can exchange both matter and energy with its surroundings.

Closed System

A system which can exchange energy but not matter with its surroundings.

Isolated System

A system which exchanges neither matter nor energy with its surroundings.

Work

The product of a force, acting on an object, and the distance, that the object moves in response to the force.

Heat Energy (Thermal Energy)

The energy transferred due to a temperature difference between the system and the surroundings.

Kinetic Energy

The energy associated with motion.

Potential Energy

Stored energy or energy a body possesses due to its position.

Heat Capacity

A measure of how many joules are required to change the temperature of a substance or object by one degree.

Specific Heat Capacity

The heat needed to warm one mass unit (g or kg) of a substance by one degree.

Temperature Change

The difference between final and initial temperatures.

Endothermic

A process when a system absorbs energy from the surroundings; the change in energy is positive.

Exothermic

A process when a system releases energy to the surroundings; the change in energy is negative.

Thermal Equilibrium

The state when a warmer object has transferred sufficient heat to the cooler object, and both objects reach the same final temperature.

Intensive Property

A property that has the same value regardless of the sample size.

Extensive Property

A physical property whose value increases with the sample size.

State Function

A property whose value depends only on the current state of the system but not on how that state was reached.

First Law of Thermodynamics

The total energy of an isolated system is conserved.

Internal Energy (E)

The sum of all kinetic and potential energies of all the atoms, ions, and molecules in the system.

Enthalpy (H)

A thermodynamic state function defined by H = E + PV, where E is the internal energy of the system, P is the pressure, and V is the volume of the system.

Latent Heat

The energy change associated with physical processes that involve a change of state (phase).

Standard Conditions

1 atm (101.33 kPa) and 25 °C (298.15 K), unless otherwise stated

Calorimetry

An experimental technique used to study the heat flow of physical and chemical processes.

Standard Enthalpy of Formation

The enthalpy change for the reaction in which 1 mole of a substance under standard conditions (1 atm and 298.15 K) is formed from the constituent elements in their reference states under standard conditions.

Standard State

The state of a substance under standard conditions.

Reference State

The most stable form of an element under standard conditions.

Allotropes

Two or more different forms of a chemical element

Hess's Law

The enthalpy change for a reaction is the sum of the enthalpy changes for the individual steps of the reaction.

Bond Enthalpy (BE)

The energy required to break one mole of a specific type of bond between two atoms, provided the reactants and products are both gases.

Total Bond Enthalpy (TBE)

The energy required to break all of the bonds in one mole of a gaseous compound into gaseous atoms.

Average Bond Enthalpy

Determined by dividing TBE of a molecule that contains only one type of bond by total number of those bonds.

Spontaneous Process

One that occurs in the absence of any ongoing outside intervention.

Entropy (S)

A measure of the randomness, or the disorder, of a system.

Second Law of Thermodynamics

Every spontaneous process increases the entropy of the universe.

Standard Molar Entropy

Absolute entropies of all elements and compounds are greater than zero because entropy increases as temperature increases

Gibbs Free Energy (G)

Amount of energy that is available for free- to enable spontaneous change to occur at constant temperature and pressure.

Gibbs equation

ΔG = ΔH - TΔS

Hess’s law

ΔG = ΣΔGp(products)- ΣΔGf(reactants)

Specific Rate Constant

k

Henderson Hasselbalch Equation

pH = pKa + log ([conjugate base]/[weak acid])