Electrochemistry Quiz

Units

• 15
• 16
• 17
• 18

Thermodynamics

• A branch of knowledge essential for understanding natural sciences.
• Focuses on the system level, describing the whole system rather than individual particles.
• Fundamental laws are universal and applicable to various fields.

Chemical reaction - energy change

• Physical and chemical changes involve using or producing energy.
• Combustion of fuels (kerosene, coal, wood, natural gas) produces heat and light.
• Electrical energy comes from chemical reactions in batteries.
• Photosynthesis uses light energy to form glucose.

Types of energy

• Kinetic (motion)
• Potential (stored)
• Mechanical
• Nuclear
• Ionization
• Chemical
• Electrical/Electromagnetic
• Thermal
• Sonic
• Gravitational

Conversion of energy

• Energy can be transferred from one location to another or change forms.
• Total energy is always conserved (law of conservation of energy).
• The sun converts nuclear energy into heat and light energy.
• Our bodies convert chemical energy in food into mechanical energy for movement.
• An electric fan transforms electrical energy into kinetic energy.
• Lightning converts electrical energy into light, heat, and sound.

Definitions important to thermodynamics

• System: The part of the universe under study, separated from the surroundings by a boundary.
• Surroundings: Everything outside the system.
• Universe: System + Surroundings

Types of systems

• Open system: Exchanges both energy and matter with the surroundings.
• Closed system: Exchanges energy but not matter.
• Isolated system: Exchanges neither energy nor matter.

Fundamental terms in thermodynamics

• Energy (E)
• Work (w)
• Heat (q)
• Temperature (T)

Energy

• Capacity of a system to do work or supply heat.
• Total energy equals kinetic energy + potential energy + internal energy.
• Internal energy is the sum of the microscopic energies of all particles in the system.

Internal energy

• Sum of kinetic and potential energies of particles in a system.
• Depends on temperature, particle type, and number of particles.

Kinetic energy

• Energy of motion at the particle level (translational, rotational, vibrational).
• Due to the movement of molecules through space, around their center of mass and back and forth motion of atoms.

Potential energy

• Energy due to particle interactions, including intermolecular forces and chemical bonds.

Law of conservation of energy

• All forms of energy can transform into others but cannot be created or destroyed.
• Energy exchange between the system and surroundings.
• The energy of the universe remains constant.

Units of energy

• SI unit of energy is the joule (J) = 1 kg m² / s².
• Calorie (cal) raises the temperature of 1.00 g of water by 1°C.
• Calorie (nutritional) is a unit of energy provided by food/drinks.

Energy flow to/from a system

• Internal energy (at given state variables) cannot be determined, only changes can be calculated or measured.
• ΔE = E₂ - E₁ (final - initial)
• When a system changes from reactants to products, the internal energy changes: • ΔE = E(products) - E(reactants).

Entropy

• The microscopic randomness, or disorder, of a system; the spread of energy.
• A measure of how dispersed energy is amongst possible energy states in a system.
• Greater disorder = higher entropy.

Enthalpy

• Heat content of a reaction at constant pressure.
• ΔH = qp
• ΔH is a state function and extensive property.
• Positive ΔH = endothermic (heat absorption).
• Negative ΔH = exothermic (heat release).

Heat of reaction

• Total heat evolved or absorbed in a chemical reaction at constant pressure (qp).
• Exothermic: Energy released (q < 0, ΔH negative); system warms.
• Endothermic: Energy absorbed (q > 0, ΔH positive); system cools.

Enthalpies of reaction

• Energy changes resulting from bond breaking and formation.
• ΔHrxn = ΣH(products) - ΣH(reactants).
• Thermochemical equations show reactions with enthalpy changes.
• Coefficients in balanced equations give the number of moles of reactant/product producing a certain ΔH.

Standard enthalpy changes

• Data based on enthalpy changes with all materials in their standard states.
• Standard state = most stable physical form at 0.1 MPa and 298K.

Standard enthalpies of formation (△H°f)

• Change in enthalpy when 1 mole of a substance forms from its elements in standard states.
• Standard enthalpy of formation of elements in their most stable form = 0.

Enthalpies of reaction under standard conditions

• Can be calculated from standard enthalpies of formation values (a table).
• To find standard enthalpy change for a given reaction sum up the enthalpies of formation of the products and subtract from this value the sum of the enthalpy changes of the reactants.

Hess's law (law of heat summation)

• Total enthalpy change of a reaction is independent of the reaction pathway.
• The heat exchange is independent of the pathway.

Changing of enthalpy at phase transfer

• Finding total enthalpy change for a solid to gas conversion can be broken down to finding changes for solid to liquid and then liquid to gas.
• ΔH°sub = ΔH°melt + ΔH°vap

Spontaneity

• Whether a reaction occurs or not when substances are brought together.
• Spontaneous processes: Physical or chemical changes that occur without external intervention.
• Temperature affects spontaneity.

Entropy and Disorder

• Order is less probable than disorder.
• Work is generally required to produce order out of disorder; therefore, energy must be used to produce a highly ordered state.

Entropy is a state function

• Its value depends on the exact state of the system.
• Its change (ΔS) does not depend on the path, only on the initial and final states.
• ΔS = S₂ - S₁
• ΔS = q/T

Processes where entropy increases (ΔS > 0)

• As energy is transferred between systems or from one form to another, it becomes dispersed from localized to more spread out states. – solid < liquid < gas.

Standard molar entropy (S°)

• Entropy gained when 1 mole of a substance changes from a perfect crystal at 0K to 0.1 MPa and 298K.

Entropy change for a reaction

• Increase in entropy: Molecule breaking down into smaller molecules; increase in the moles of gases; solid changing to liquid or liquid to gas.
• ΔS°rxn = ΣnS°(products) - ΣmS°(reactants)

Predicting spontaneity of reactions

• Reactions that have a large and negative ΔH tend to be spontaneous.
• Spontaneous reactions lead to an increase in the entropy of the universe.
• Spontaneity depends on enthalpy, entropy, and temperature.

Gibbs free energy (G)

• Unifies the first and second laws of thermodynamics for constant temp and pressure.
• G = H – TS
• ΔG = ΔH – TΔS (change in Gibbs free energy)
• G is a state function, extensive property, maximum work done at constant temp and pressure.

Gibbs free energy (G) and spontaneity

• Negative ΔG = spontaneous reaction in the forward direction (exergonic process).
• Positive ΔG = non-spontaneous reaction in the forward direction (endergonic process).
• Zero ΔG = system is at equilibrium.

Standard Gibbs free energy

• Free energy change when reactants in standard states are converted to products in standard states.
• ΔG° = ΔH° - TΔS°

Standard Gibbs free energy of formation (△G°f)

• Free energy change when 1 mol of a substance forms from its elements in their reference states at 0.1 MPa and 298K.

Gibbs free energy and temperature

• When enthalpy change (△H) is negative and entropy change (△S) is positive, the reaction is exergonic at all temperatures.
• When △H is positive and △S is negative, the reaction is endergonic at all temperatures.
• When △H is positive and △S is positive, the reaction is spontaneous at high temperatures and nonspontaneous at low temperatures.
• When △H is negative and △S is negative, the reaction is spontaneous at low temperatures and nonspontaneous at high temperatures.

Free energy and eqilibrium

• Q<K → ΔG<0, reactants
• Q=K → ΔG=0 equilibrium q>K → ΔG >0, products

△G° and the chemical equilibrium

• To find △G, use the equation △G = △G° + RT In Q
• At equilibrium △G=0

Electrochemistry

• Deals with chemical reactions that produce or use electricity.
• Important for mobile phones, cars, fuel cells, corrosion, and materials preparation.
• Topics of study: oxidation, reduction, redox reactions, oxidation numbers, electrochemical cells, electrode potentials (cell notation and types of electrodes), standard reduction potentials, Nernst equation, concentration cells, glass electrodes, redox reactions in living cells, cellular respiration, standard redox potentials in living systems.

Test your knowledge on key concepts in electrochemistry, including energy transformations in photosynthesis and the workings of galvanic and electrolytic cells. Explore redox reactions, electrode potentials, and different types of chemical reactions in this comprehensive quiz.