IB Chemistry › Reactivity 1 › Flashcards
FC

Reactivity 1.1 Flashcards

What Drives Chemical Reactions. Select a sub-topic below

R1.1 Enthalpy & Calorimetry (SL): read all 10 questions and answers
  1. [SL] What is the difference between heat and temperature?
    Temperature measures the average kinetic energy of particles. Heat is the transfer of thermal energy between system and surroundings due to a temperature difference.
  2. [SL] What is the sign of ΔH for an exothermic reaction?
    Negative (\(\Delta H < 0\)). Energy is released to the surroundings and the products have less enthalpy than the reactants.
  3. [SL] Why does the temperature of surroundings decrease during an endothermic reaction?
    The reacting chemical system absorbs heat energy from the surroundings and converts it into chemical potential energy, reducing the average kinetic energy of surrounding particles.
  4. [SL] Define system and surroundings in thermodynamics.
    The system consists of the reacting chemicals. The surroundings are everything else: the solvent, the container, and the air.
  5. [SL] State the calorimetry formula and define c.
    \(q = mc\Delta T\). \(c\) is the specific heat capacity: the energy required to raise the temperature of 1 g of a substance by 1 K. For water, \(c = 4.18\text{ J g}^{-1}\text{K}^{-1}\).
  6. [SL] How do you convert q to the molar enthalpy change?
    \(\Delta H = -\frac{q}{n}\), where \(n\) is the moles of limiting reactant. The negative sign corrects for the thermochemical convention (exo = negative).
  7. [SL] Why do we extrapolate the cooling curve in calorimetry?
    To find the maximum theoretical temperature change and compensate for heat loss to the surroundings during the experiment.
  8. [SL] On an energy profile diagram, what does \(E_{\text{a}}\) represent?
    Activation energy: the minimum energy required for a successful collision, shown as the vertical rise from reactants to the transition state peak.
  9. [SL] Why are products more stable in an exothermic reaction?
    Because the products sit at a lower potential energy than the reactants, meaning they have released energy and are more energetically stable.
  10. [SL] Why is bond breaking endothermic?
    Energy must be absorbed to overcome the electrostatic attractions holding atoms together in the bond.
R1.1 Standard Enthalpies & Definitions (HL): read all 7 questions and answers
  1. [SL] What are standard conditions for ΔH?
    Pressure of 100 kPa, concentration of 1 mol dm⁻³ for solutions, all substances in standard states, typically at 298 K (25 degrees C).
  2. [SL] Define standard enthalpy of formation (DeltaHf).
    The enthalpy change when exactly one mole of a compound is formed from its elements in their standard states under standard conditions. DeltaHf of any element in its standard state = 0.
  3. [SL] Define standard enthalpy of combustion (DeltaHc).
    The enthalpy change when one mole of a substance is completely burned in excess oxygen under standard conditions. Always negative.
  4. [SL] How is DeltaH estimated using bond enthalpies?
    DeltaH = sum of bond enthalpies broken minus sum of bond enthalpies formed. If more energy is released making bonds than absorbed breaking them, the reaction is exothermic.
  5. [SL] Why do experimental calorimetry values differ from literature values?
    Due to heat loss to surroundings, incomplete combustion, evaporation of water, and heat absorbed by the calorimeter itself.
  6. [HL] What is a Born-Haber cycle?
    An application of Hess Law used to show the energy changes involved in forming an ionic compound from its elements. Primarily used to calculate lattice enthalpy.
  7. [HL] Define lattice enthalpy.
    The enthalpy change when one mole of an ionic compound is separated into its constituent gaseous ions under standard conditions. Defined as endothermic (DeltaH_lat > 0).
R1.2 Hess Law & Bond Enthalpies (SL): read all 6 questions and answers
  1. [SL] State Hess Law.
    The total enthalpy change for a reaction is independent of the route taken, provided the initial and final conditions are the same.
  2. [SL] What is a mean bond enthalpy?
    The average energy required to break one mole of a particular type of bond in the gaseous state, averaged across many different compounds.
  3. [SL] Why are bond enthalpy calculations only approximate?
    Because mean bond enthalpies are averaged across many compounds. The actual bond energy varies depending on the molecular environment and neighbouring atoms.
  4. [SL] State the bond enthalpy equation for DeltaH.
    DeltaH = sum of bonds broken minus sum of bonds formed. Breaking bonds = endothermic (+), making bonds = exothermic (-).
  5. [SL] Why must substances be in the gaseous state for bond enthalpy calculations?
    Bond enthalpies only apply to bonds in gaseous molecules. If substances are liquids or solids, additional energy for vaporisation must be considered.
  6. [SL] How do you use a Hess Law energy cycle?
    Construct a cycle with alternative routes from reactants to products. The enthalpy change for the unknown route equals the sum of enthalpy changes for the known routes.
R1.2 Energy Cycles & Born-Haber (HL): read all 8 questions and answers
  1. [HL] How is DeltaH calculated using standard enthalpies of formation?
    DeltaH = sum of DeltaHf(products) minus sum of DeltaHf(reactants).
  2. [HL] How is DeltaH calculated using standard enthalpies of combustion?
    DeltaH = sum of DeltaHc(reactants) minus sum of DeltaHc(products). Note the reversed subtraction compared to formation data.
  3. [HL] List the steps in a Born-Haber cycle for NaCl.
    1. Atomisation of Na(s), 2. Ionisation of Na(g), 3. Atomisation of Cl2(g), 4. Electron affinity of Cl(g), 5. Lattice enthalpy (formation of NaCl lattice).
  4. [HL] Define electron affinity.
    The enthalpy change when one mole of electrons is added to one mole of gaseous atoms. First electron affinity is usually exothermic (negative). Second EA is always endothermic (positive) due to electron-electron repulsion.
  5. [HL] Define enthalpy of atomisation.
    The enthalpy change when one mole of gaseous atoms is formed from the element in its standard state. Always endothermic (positive).
  6. [HL] How do you calculate lattice enthalpy from a Born-Haber cycle?
    By applying Hess Law around the cycle: DeltaHf = DeltaH(atomisation) + DeltaH(ionisation) + DeltaH(electron affinity) - DeltaH_lat(dissociation). Rearrange to find lattice enthalpy.
  7. [HL] Why might experimental and theoretical lattice enthalpies differ?
    Differences indicate covalent character in the ionic bond. If the experimental value is larger than the theoretical (purely ionic) value, the compound has significant covalent character.
  8. [HL] What factors affect the magnitude of lattice enthalpy?
    1. Ionic charge: higher charges give stronger attraction. 2. Ionic radius: smaller ions give shorter interionic distances and stronger attraction.
R1.3 Combustion & Environment (SL): read all 8 questions and answers
  1. [SL] What are the products of complete vs incomplete combustion?
    Complete: CO2 + H2O (in excess oxygen, releases more energy). Incomplete: CO or C (soot) + H2O (in limited oxygen, releases less energy).
  2. [SL] Why is carbon monoxide dangerous?
    CO is colourless and odourless. It binds with very high affinity to haemoglobin, displacing oxygen and preventing oxygen transport, which can lead to unconsciousness and death.
  3. [SL] What are the environmental problems with burning coal?
    Produces CO2 (greenhouse gas), SO2 (acid rain), and particulates (global dimming and respiratory issues).
  4. [SL] Explain the greenhouse effect.
    Shortwave solar radiation passes through the atmosphere. The Earth absorbs it and re-emits longwave IR radiation. Greenhouse gases (CO2, CH4, H2O) absorb IR and re-radiate it, trapping heat.
  5. [SL] Why are biofuels considered carbon neutral?
    The CO2 released when biofuels burn equals the CO2 absorbed by the plants during photosynthesis. However, this ignores energy used in production and transport.
  6. [SL] State two disadvantages of biofuels.
    1. Require large areas of land, competing with food production and causing deforestation. 2. Lower specific energy than fossil fuels, so more fuel is needed.
  7. [SL] How does a hydrogen fuel cell work?
    Converts chemical energy directly into electrical energy via redox reactions at electrodes. Anode: H2 to 2H+ + 2e-. Cathode: O2 + 4H+ + 4e- to 2H2O. Only product is water.
  8. [SL] What is the overall equation for a hydrogen fuel cell?
    2H2(g) + O2(g) to 2H2O(l). The only product is water, making it much cleaner than combustion engines.
R1.3 Fuel Calculations & Comparisons (SL): read all 7 questions and answers
  1. [SL] Why do longer-chain hydrocarbons undergo more incomplete combustion?
    Higher carbon percentage and stronger London forces make them less volatile. Less vaporisation means less optimal mixing with air, resulting in more incomplete combustion.
  2. [SL] Define specific energy.
    The amount of energy stored or released per unit mass of fuel, measured in kJ/g or MJ/kg.
  3. [SL] Define energy density.
    The amount of energy stored per unit volume of fuel, measured in MJ L⁻¹ or kJ dm⁻³.
  4. [Enrichment] Why do hydrocarbons have higher specific energy than alcohols?
    Alcohols are already partially oxidised (contain oxygen), so they release less energy per gram when burned compared to hydrocarbons.
  5. [Enrichment] Why does methane have higher specific energy than butane?
    Methane has a higher hydrogen-to-carbon ratio. Hydrogen has a very high enthalpy of combustion relative to its tiny atomic mass.
  6. [HL] (AHL R1.2.4) How do you calculate DeltaH using enthalpies of combustion via Hess Law?
    DeltaH = sum of DeltaHc(reactants) minus sum of DeltaHc(products). Both reactants and products are combusted to common products (CO2 and H2O). Note: Enthalpy of combustion cycles are assessed under AHL R1.2.4.
  7. [SL] Is DeltaHc always negative?
    Yes, combustion is always exothermic. DeltaHf can be positive or negative depending on whether the compound is more or less stable than its elements.
R1.4 Entropy Basics (HL): read all 7 questions and answers
  1. [HL] What is entropy (S)?
    A thermodynamic property that measures the dispersal or distribution of matter and energy within a system. Higher entropy = more disorder.
  2. [HL] How does entropy compare across states of matter?
    S(gas) >> S(liquid) > S(solid). Gases have the most dispersal and freedom of movement.
  3. [HL] What is the entropy of a perfect crystal at 0 K?
    Exactly zero (Third Law of Thermodynamics). Complete order with no thermal motion.
  4. [HL] How do you predict if DeltaS is positive or negative?
    Look at states: more gas moles in products = positive DeltaS. Also positive when a solid dissolves, or when fewer moles become more moles.
  5. [HL] What defines a spontaneous process?
    A process that occurs naturally without a continual input of energy from an external source. Note: spontaneous does NOT mean fast.
  6. [HL] According to the 2nd Law, what must happen for a reaction to be spontaneous?
    The total entropy of the universe (system + surroundings) must increase: DeltaS(total) > 0.
  7. [HL] Why can some endothermic reactions still be spontaneous?
    If the increase in entropy of the system is large enough to compensate for the decrease in entropy of the surroundings caused by heat absorption.
R1.4 Gibbs energy (HL): read all 10 questions and answers
  1. [HL] How is standard entropy change calculated?
    DeltaS = sum of S(products) minus sum of S(reactants), using standard entropy values from the data booklet.
  2. [HL] State the Gibbs energy equation.
    DeltaG = DeltaH minus TDeltaS. G represents the maximum energy available to do useful work.
  3. [HL] How do you use DeltaG to predict spontaneity?
    DeltaG < 0: spontaneous. DeltaG > 0: non-spontaneous. DeltaG = 0: at equilibrium.
  4. [HL] What unit conversion must you make when calculating DeltaG?
    Entropy is in J/(K mol) but enthalpy is in kJ/mol. Divide DeltaS by 1000 to convert to kJ/(K mol) before using the Gibbs equation.
  5. [HL] When is a reaction always spontaneous at all temperatures?
    When DeltaH is negative (exothermic) AND DeltaS is positive (more disorder). Both terms favour spontaneity, making DeltaG always negative.
  6. [HL] When is a reaction never spontaneous?
    When DeltaH is positive (endothermic) AND DeltaS is negative (more order). DeltaG is always positive.
  7. [HL] How does temperature affect spontaneity when both DeltaH and DeltaS are positive?
    Non-spontaneous at low T but becomes spontaneous at high T when TDeltaS exceeds DeltaH. The transition temperature is T = DeltaH / DeltaS.
  8. [HL] How do you find the temperature at which a reaction becomes spontaneous?
    Set DeltaG = 0 and solve: T = DeltaH / DeltaS. Above this T, the reaction is spontaneous (if DeltaH and DeltaS are both positive).
  9. [HL] How is DeltaG related to the equilibrium constant K?
    DeltaG(standard) = -RT ln K. A large negative DeltaG means a large K (products strongly favoured). At equilibrium, DeltaG = 0.
  10. [HL] State the equation linking DeltaG to the reaction quotient Q.
    DeltaG = DeltaG(standard) + RT ln Q. As a reaction approaches equilibrium, Q approaches K and DeltaG approaches zero.
← Back to Reactivity 1Exam Practice →