Chemistry Redox and Gas Laws Quiz

Chapter 7: Gases

• Gases are composed of small molecules in constant, random motion.
• The volume of the gas particle is insignificant compared to the volume of the gas.
• Forces between the gas particles are negligible, except when the particles collide.
• Molecular collisions are perfectly elastic; no energy is lost during collisions.
• The average kinetic energy of gas molecules is directly proportional to the absolute temperature of the gas.
• Gas pressure is the sum of forces exerted by gas particles impacting the surface, divided by the area of the surface.
• Units of pressure include atmospheres (atm), millimeters of mercury (mmHg), bars (bar), pascals (Pa), and torr.
• Standard conversions between pressure units are available.

Section 7.1 Gas Pressure

• A gas particle striking a surface exerts a force against the surface.
• Gas pressure is determined by the sum of forces exerted by gas particles impacting the surface, divided by the surface area.

Section 7.2 Boyle's Law

• Boyle's Law: The volume of a gas at constant temperature is inversely proportional to its pressure.
• As volume decreases, gas pressure increases, and vice versa.

Section 7.2 Boyle's Law - Calculations

• If you change the pressure of a gas sample from P1 to P2 at constant temperature, the volume will change from V1 to V2: P₁V₁ = P₂V₂

Section 7.3 Charles's Law

• Charles's Law: At constant pressure, the volume of a gas is directly proportional to its Kelvin temperature.
• As temperature increases, volume increases, and vice versa

Section 7.3 Charles's Law - Calculations

• If you change the temperature of a gas sample from T1 to T2 at constant pressure, the volume will change from V1 to V2: V₁/T₁ = V₂/T₂.
• Temperatures must be in Kelvin

Section 7.4 The Combined Gas Law

• The combined gas law merges Boyle's law and Charles's Law.
• The combined gas law equation is P₁V₁/T₁ = P₂V₂/T₂.
• Temperatures must be in Kelvin.

Section 7.5 Avogadro's Law

• Avogadro's Law: At constant temperature and pressure, the volume of a gas is directly proportional to the number of moles of gas. (V ∝ n).

Section 7.6 The Ideal Gas Law

• The ideal gas law combines Boyle's, Charles', and Avogadro's Laws.
• Ideal gas law equation: PV = nRT.
• R: Ideal gas constant (with units given)

Section 7.7 Dalton's Law of Partial Pressures

• Dalton's Law of Partial Pressures: The total pressure of a mixture of gases is equal to the sum of the individual pressures of the gases.
• Ptotal = P₁ + P₂ +...+ Pn, where Pi is the partial pressure of the ith gas.

Section 7.8 Molar Mass and Density of Gas Law Calculations

• If the gas identity is known, molar mass can be used to convert mass to moles, allowing for applying the ideal gas law.
• If the gas identity is unknown, using ideal gas law, mass of the gas and other pertinent data is used to determine molar mass
• Gas density (given in g/L) provides mass, P, V, and T of a gas sample.

Section 7.9 Gases in Chemical Reactions

• Gay-Lussac's law of combining volumes: the volume ratio of gases in a chemical reaction is equal to the mole ratio in the balanced equation if all gases are at the same temperature and pressure.

Section 7.10 Kinetic Molecular Theory of Gases

• Postulates of the Kinetic Molecular Theory of Gases
• Gases composed of small, constantly randomly moving molecules.
• Volume of molecules insignificant compared to total volume of gas.
• Intermolecular forces insignificant except during collision.
• Collisions are perfectly elastic (no energy loss).
• Average kinetic energy is directly proportional to absolute temperature.

Section 7.11 Movement of Gas Particles

• The root-mean-square speed (v_rms) of a gas is proportional to the square root of the absolute temperature (T) divided by the molar mass (M).
• v_rms = √(3RT/M)

Section 7.12 Real Gases

• Ideal gas behavior assumptions fail at high pressures and low temperatures. Ideal gas law and calculations are not applicable.
• The van der Waals equation is a more accurate model for real gases as it accounts for intermolecular forces and the finite volume of gas particles.

Chapter 19: Electrochemistry

• Redox reactions involve electron transfer.
• Oxidation is loss of electrons (LEO).
• Reduction is gain of electrons (GER).
• Oxidation number increases in oxidation, decreases in reduction.

Section 19.1 Redox Reactions

• A redox reaction is a chemical reaction that involves the movement of electrons from one reactant to another.

Section 19.2 Balancing Redox Reactions

• Redox reactions involve balancing mass and charge in half-reactions.
• Balancing in acidic/basic solutions require specific steps.

Section 19.4 Voltaic Cells

• Voltaic cells involve spontaneous redox reactions to generate electricity.
• Consist of an anode (oxidation) and a cathode (reduction).

Section 19.5 Cell Potential

• Cell potential (E_cell) is the electromotive force (EMF), measured as voltage.
• Standard reduction potentials (E°_red) are measured against standard hydrogen electrode (SHE) and assigned values.
• The standard cell potential (E°_cell) is determined by E°_cathode – E°_anode.

Section 19.6 Free Energy and Cell Potential

• Gibbs free energy (∆G) relates to cell potential (E_cell) through ∆G = −nFE_cell.
• F is the Faraday constant (96,485 C/mol e−).
• ∆G < 0 indicates spontaneity while ∆G > 0 indicates nonspontaneity.

Section 19.7 The Nernst Equation and Concentration Cells

• The Nernst equation calculates the cell potential (E_cell) under non-standard conditions.
• The Nernst equation equation gives E_cell = E°_cell− RT/nF lnQ.
• Q is the reaction quotient

Section 19.8 Voltaic Cell Applications: Batteries, Fuel Cells, and Corrosion

• Voltaic cells are used in several contexts
• Batteries use chemical reactions to provide a source of electricity.
• Fuel cells utilize chemical energy as a source of electricity, converting fuels into electricity directly.
• Corrosion is the oxidation of metal structures.

Section 19.9 Electrolytic Cells and Applications of Electrolysis

• Electrolytic cells use electricity to facilitate nonspontaneous chemical reactions (redox) to occur.
• Electrolysis involves calculations based on moles of electrons (or Faraday's constant).
• The total charge(q) is determined by current(I) times time(t).

Chapter 18: Chemical Thermodynamics

• Chemical thermodynamics deals with the relationship between heat, work, and forms of energy.

Section 18.1 Entropy and Spontaneity

• Entropy (S) measures the degree of disorder or randomness in a system.
• Systems tend towards spontaneous change toward increased entropy. Processes that decrease entropy can still be spontaneous overall when considering the entropy of the surroundings.

Section 18.2 Entropy Changes

• The third law of thermodynamics states that the entropy of a pure, perfectly ordered crystal substance is zero at absolute zero (0 K)
• Standard changes in entropy (∆S°_rxn) can be determined using standard molar entropy data, by calculating the initial state vs final state entropy.

Section 18.3 Entropy and Temperature

• The overall change in entropy (∆S_univ) is the sum of the entropy change of the system (∆S_sys) and the entropy change of the surroundings (∆S_surr)
• For a process to be spontaneous, ∆S_univ must be positive
• ∆S_surr is related to the enthalpy change (∆H) of a reaction by the equation ∆S_surr = –∆H/T. Where T is temperature (in Kelvin).

Section 18.4 Gibbs Free Energy

• Gibbs free energy (G) is a state function that can be used to determine the spontaneity at a given temperature.
• G=H-TS
• ∆G° can calculated from ∆H° and ∆S°.

Section 18.5 Free-Energy Changes and Temperature

• For process to be spontaneous overall ∆G must be negative.
• The temperature at which spontaneous change begins can be calculated using the equation ∆G = ∆H − T∆S.

Section 18.6 Gibbs Free Energy and Equilibrium

• At equilibrium, the Gibbs free energy change (∆G) is zero
• For a chemical reaction at nonstandard conditions, ∆G = ∆G° + RTlnQ where Q is the reaction quotient

Chapter 17: Aqueous Equilibrium

• Understanding acid-base titrations.
• Understanding the concepts of buffer solutions and calculations.

Section 17.1 Introduction to Buffer Solutions

• A buffer solution consists of a weak acid and its conjugate base.
• Buffer solutions resist changes in pH when a strong acid or base is added.

Section 17.2 The Henderson-Hasselbalch Equation

• The Henderson-Hasselbalch equation relates pH to pKa and the concentrations of a weak acid and its conjugate base.

Section 17.3 Titrations of Strong Acids and Strong Bases

• Titration curve shows how pH changes with added strong acid or strong base .
• Calculations related to titration curves are associated with the equivalence point.

Section 17.4 Titrations of Weak Acids and Weak Bases

• A titration of a weak acid/base involves calculating changes in pH during the addition of a strong base/acid.
• Equilibrium constant calculations are involved in determining the equivalence point.

Section 17.5 Indicators in Acid-Base Titrations

• Acid-base indicators are weak acids/bases that change color at distinct pH values to indicate the equivalence point.
• Indicators have a specific pH range over which a transition from one color to another is observed.

Section 17.6 The Solubility Product Constant, K_sp

• K_sp represents the solubility of an ionic compound.
• Solubility products are related Equilibrium.
• Molar solubility is the number of moles of dissolved solute in 1 L solution

Section 17.7 The Common-Ion Effect and the Effect of pH on Solubility

• Common ion effect – reducing solubility by adding a solute containing an ion from the compound.
• pH affect on solubility – changing the charge of the ionic compound can affect solubility

Section 17.8 Precipitation: Q Versus K_sp

• Q and K_sp are used to determine if a precipitate will form.
• Q represents current conditions while K_sp is for equilibrium.

Chapter 14: Chemical Kinetics

• Chemical kinetics deals with the speed of chemical reactions.

Section 14.1 Rates of Reactions

• The rate of a chemical reaction is defined as the change in concentration over time.
• Factors that affect reaction rates include concentration, temperature, surface area, and the presence of a catalyst.

Section 14.2 Reaction Rates and Concentration: Rate Laws

• Rate laws describe the relationship between reaction rate and the concentrations of reactants.
• Rate laws often include rate constant and the orders of each reactant (determined experimentally).

Section 14.3 Integrated Rate Laws and Half-Lives

• Integrated rate laws are used to determine concentrations over time for different reaction orders.
• Half-life is the time it takes for half of the reactant to be used up.

Section 14.4 Reaction Rates and Temperature: Activation Energy

• Activation energy (E_a) is the minimum energy needed for a reaction to occur.
• The Arrhenius equation describes the relationship between reaction rate constant (k) and temperature (T) using activation energy.

Section 14.5 Reaction Mechanisms

• Reaction mechanisms detail the series of elementary steps that occur to form products.
• Each step of the mechanism involves a transition state containing intermediate species.

Section 14.6 Catalysis

• Catalysts speed up chemical reactions by lowering the activation energy.
• Catalysts participate in the elementary steps but do not appear in the overall equation for the reaction.

Chapter 13: Solutions

• Solutions consist of a solvent and one or more solutes. Solutions are homogenous mixtures of solute and solvent.

Section 13.1 The Solution Process

• Substances dissolve in one another when the intermolecular forces between solute-solute, solvent-solvent, and solute-solvent become favorable.
• The enthalpy of solution can be determined by summing the enthalpy changes of the three steps involved in solution formation: breaking solute-solute bonds, breaking solvent-solvent bonds, and forming solute-solvent bonds.
• Most ionic compounds dissolve in water via ion-dipole interactions.

Section 13.2 Saturated, Unsaturated, and Supersaturated Solutions

• A saturated solution contains the max. amount of dissolved solute at a given temperature and pressure.
• An unsaturated solution contains less than the max. amount of dissolved solute at a given temperature and pressure.
• A supersaturated solution contains more dissolved solute than is stable at a given temperature, often crystallizes out when disturbed.

Section 13.3 Concentration Units

• The concentration describes the amount of solute in a given amount of solution. This is written in terms of molarity (M), molality (m), mass percent and mole fraction.

Section 13.4 Colligative Properties of Nonelectrolytes

• Colligative properties of a solution depend only on the number of solute particles and not the identity of the solute.
• These properties include vapor pressure lowering, freezing point depression, boiling point elevation, and osmotic pressure

Section 13.5 Colligative Properties of Electrolytes

• Electrolytes dissolve into more than one particle.
• The Van't Hoff factor (i) is multiplied to the molarity (M) in equations for freezing point depression, boiling point elevation, and osmotic pressure

Chapter 12: Liquids and Solids

• Properties of liquids and solids.

Section 12.1 Intermolecular Forces

• Intermolecular forces hold atoms together.
• Types of intermolecular forces include dispersion forces, dipole-dipole forces, hydrogen bonding, and Ion-dipole forces.

Section 12.2 Properties of Liquids

• Viscosity means how resistant a liquid is to flow.
• Surface tension is the resistance of a liquid to spread out or increase in surface area.

Section 12.3 Phase Changes and Heating Curves

• A heating curve plots temperature vs. heat added, used to calculate the amount of heat required for each transformation.

Section 12.4 Vapor Pressure, Boiling Point, and the Clausius-Clapeyron Equation

• Vapor pressure is the partial pressure of a vapor above a liquid.
• Ideal gas law used and calculations associated with it are used to calculate vapor pressure

Section 12.5 Phase Diagrams

• A phase diagram shows the phases of matter and pressure/temperature ranges.
• The critical point is where properties of liquid and gas become indistinguishable.
• The triple point is where all three phases exist in equilibrium with one another.

Section 12.6 Classification of Solids

• Solids can be classified based on their structure: amorphous and crystalline.
• Crystalline solids include ionic, molecular, covalent network, and metallic solids.

Section 12.7 The Unit Cell and the Structure of Crystalline Solids

• Unit cells make up crystalline solids.
• Properties of simple cubic, body-centered cubic and face-centered cubic unit cells including atoms per unit cell, packing efficiency and coordination number are included.

Test your knowledge on key concepts in chemistry, including the ideal gas law, Dalton's Law, and electrochemistry. This quiz covers fundamental principles related to moles, gas mixtures, redox reactions, and more. Perfect for students studying introductory chemistry.