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1. A first-order reaction A → products is studied. If doubling [A] from 0.10 M to 0.20 M increases the observed rate constant k from 0.12 s⁻¹ to 0.12 s⁻¹ (unchanged), what does this imply about the reaction order with respect to A?
2.How is the order of a reaction with respect to a reactant determined?
3.How do solvent properties and ionic strength affect rate?
4.Why do mixing or stirring affect reaction rate?
5.How does pressure influence rate for gas-phase reactions?
6.How does surface area impact the rate of a heterogeneous reaction?
7. What is the role of a catalyst in reaction rate?
8.How does reactant concentration influence rate?
9.How does temperature affect reaction rate?
10.What is the reaction rate?
1 1: What is the purpose of the electrolysis of potassium iodide?
12: What setup is typically used for the electrolysis of potassium iodide?
13: What are the expected products at the anode during the electrolysis of potassium iodide?
14: What are the expected products at the cathode during the electrolysis of potassium iodide?
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1) If doubling [A] from 0.10 → 0.20 M leaves the observed rate constant k unchanged (k = 0.12 s⁻¹ in both cases), this implies the reaction is first order in A. For a first‑order reaction k is a true constant (rate = k[A]), so k does not depend on [A].
2) How reaction order is determined:
- Initial rates method: measure initial rate at different reactant concentrations and fit rate ∝ [reactant]^n to find n.
- Integrated rate laws: test which plot is linear — zero order ([A] vs t), first order (ln[A] vs t), second order (1/[A] vs t).
- Half‑life behavior: for first order t1/2 is independent of [A]0; for other orders it depends on [A]0.
- Isolation method: hold other reactants in large excess so the rate depends effectively on one species and determine its exponent.
- For ionic reactions use activity/ionic‑strength corrections when needed.
3) Solvent properties and ionic strength effect on rate:
- Solvent polarity/proticity can stabilize reactants or the transition state differently, changing activation energy and rate (e.g., SN1 vs SN2 behaviors differ in protic vs aprotic solvents).
- Viscosity controls diffusion; high viscosity can slow diffusion‑limited reactions.
- Ionic strength alters activity coefficients of ions; according to Debye–Hückel/Brønsted–Bjerrum relationships, ionic strength can speed up or slow down ionic reactions depending on the charges of reactants (like‑charged ions often react faster as I increases; opposite charges often slower).
4) Why mixing/stirring affect rate:
- Stirring reduces concentration gradients and boundary layers, improving mass transport to/from reactive sites.
- In diffusion‑limited or heterogeneous systems mixing increases contact between reactants/phases and often increases the observed rate (up to the intrinsic chemical‑kinetic limit).
5) Pressure influence for gas‑phase reactions:
- Increasing total pressure raises gas concentrations (ideal gas law), increasing collision frequency and hence rate for reactions involving gaseous reactants.
- Pressure also affects rates through the activation volume (ΔV‡): reactions with negative ΔV‡ are accelerated by higher pressure; those with positive ΔV‡ are slowed.
6) Surface area impact on heterogeneous reactions:
- Greater surface area (e.g., finely divided solid) provides more active sites for adsorption and reaction, so rate generally increases roughly proportionally to available surface area, all else equal.
- If reaction is diffusion/adsorption limited, increasing surface area or improving dispersion reduces those limitations.
7) Role of a catalyst:
- Provides an alternative reaction pathway with a lower activation energy, increasing the rate without being consumed.
- Catalysts can change selectivity and mechanism (homogeneous vs heterogeneous vs enzymatic).
- Catalysts do not change the equilibrium constant; they only speed attainment of equilibrium.
8) How reactant concentration influences rate:
- Rate generally depends on concentration according to the rate law (rate ∝ [A]^n). For positive n, increasing concentration increases rate by increasing collision frequency; for zero order rate is independent of concentration.
- In enzyme or surface‑saturated systems, rates can plateau (e.g., Michaelis–Menten behavior).
9) How temperature affects rate:
- Rate constants typically increase with temperature according to the Arrhenius equation k = A e^(−Ea/RT). Higher T increases fraction of molecules with energy ≥ Ea, often producing large increases in rate.
- Rule of thumb: many reaction rates roughly double for each 10 °C rise (depends on Ea).
10) What is the reaction rate?
- The reaction rate is the change in concentration of a reactant or product per unit time (e.g., −d[A]/dt for reactant A). It can be an instantaneous rate (d[A]/dt at a time) or an average over an interval; units are concentration/time (e.g., M s⁻¹). For a balanced reaction aA + bB → … the rate is often expressed as −(1/a) d[A]/dt = −(1/b) d[B]/dt = (1/… ) d[product]/dt.
11) Purpose of the electrolysis of potassium iodide:
- Typically done to oxidize iodide to iodine and demonstrate electrolysis/redox chemistry, to produce iodine (I2) and hydrogen, and/or to study electrode reactions and product separation (and to generate KOH locally at the cathode).
12) Typical setup for electrolysis of KI:
- Acell containing aqueous KI solution, two inert electrodes (graphite or platinum), a DC power source (battery or power supply) with electrodes connected to positive (anode) and negative (cathode).
- Optionally use a porous barrier or divided cell if you want to keep anodic and cathodic products separate.
- Optional starch indicator can detect iodine formation. Provide ventilation and PPE (iodine vapors are hazardous).
13) Expected products at the anode (oxidation):
- I− is oxidized preferentially: 2 I− → I2 + 2 e−, so molecular iodine (I2) appears (often as brown solution or forms I3− with excess I−).
- If I− is depleted or potential is high, water oxidation to O2 can occur: 2 H2O → O2 + 4 H+ + 4 e−.
14) Expected products at the cathode (reduction):
- Water is reduced in aqueous KI: 2 H2O + 2 e− → H2 + 2 OH−, so hydrogen gas evolves and hydroxide forms near the cathode (producing KOH in bulk solution).
- Metallic K is not produced in aqueous electrolysis (it would require much more negative potential and occurs only in molten salts).
2) How reaction order is determined:
- Initial rates method: measure initial rate at different reactant concentrations and fit rate ∝ [reactant]^n to find n.
- Integrated rate laws: test which plot is linear — zero order ([A] vs t), first order (ln[A] vs t), second order (1/[A] vs t).
- Half‑life behavior: for first order t1/2 is independent of [A]0; for other orders it depends on [A]0.
- Isolation method: hold other reactants in large excess so the rate depends effectively on one species and determine its exponent.
- For ionic reactions use activity/ionic‑strength corrections when needed.
3) Solvent properties and ionic strength effect on rate:
- Solvent polarity/proticity can stabilize reactants or the transition state differently, changing activation energy and rate (e.g., SN1 vs SN2 behaviors differ in protic vs aprotic solvents).
- Viscosity controls diffusion; high viscosity can slow diffusion‑limited reactions.
- Ionic strength alters activity coefficients of ions; according to Debye–Hückel/Brønsted–Bjerrum relationships, ionic strength can speed up or slow down ionic reactions depending on the charges of reactants (like‑charged ions often react faster as I increases; opposite charges often slower).
4) Why mixing/stirring affect rate:
- Stirring reduces concentration gradients and boundary layers, improving mass transport to/from reactive sites.
- In diffusion‑limited or heterogeneous systems mixing increases contact between reactants/phases and often increases the observed rate (up to the intrinsic chemical‑kinetic limit).
5) Pressure influence for gas‑phase reactions:
- Increasing total pressure raises gas concentrations (ideal gas law), increasing collision frequency and hence rate for reactions involving gaseous reactants.
- Pressure also affects rates through the activation volume (ΔV‡): reactions with negative ΔV‡ are accelerated by higher pressure; those with positive ΔV‡ are slowed.
6) Surface area impact on heterogeneous reactions:
- Greater surface area (e.g., finely divided solid) provides more active sites for adsorption and reaction, so rate generally increases roughly proportionally to available surface area, all else equal.
- If reaction is diffusion/adsorption limited, increasing surface area or improving dispersion reduces those limitations.
7) Role of a catalyst:
- Provides an alternative reaction pathway with a lower activation energy, increasing the rate without being consumed.
- Catalysts can change selectivity and mechanism (homogeneous vs heterogeneous vs enzymatic).
- Catalysts do not change the equilibrium constant; they only speed attainment of equilibrium.
8) How reactant concentration influences rate:
- Rate generally depends on concentration according to the rate law (rate ∝ [A]^n). For positive n, increasing concentration increases rate by increasing collision frequency; for zero order rate is independent of concentration.
- In enzyme or surface‑saturated systems, rates can plateau (e.g., Michaelis–Menten behavior).
9) How temperature affects rate:
- Rate constants typically increase with temperature according to the Arrhenius equation k = A e^(−Ea/RT). Higher T increases fraction of molecules with energy ≥ Ea, often producing large increases in rate.
- Rule of thumb: many reaction rates roughly double for each 10 °C rise (depends on Ea).
10) What is the reaction rate?
- The reaction rate is the change in concentration of a reactant or product per unit time (e.g., −d[A]/dt for reactant A). It can be an instantaneous rate (d[A]/dt at a time) or an average over an interval; units are concentration/time (e.g., M s⁻¹). For a balanced reaction aA + bB → … the rate is often expressed as −(1/a) d[A]/dt = −(1/b) d[B]/dt = (1/… ) d[product]/dt.
11) Purpose of the electrolysis of potassium iodide:
- Typically done to oxidize iodide to iodine and demonstrate electrolysis/redox chemistry, to produce iodine (I2) and hydrogen, and/or to study electrode reactions and product separation (and to generate KOH locally at the cathode).
12) Typical setup for electrolysis of KI:
- Acell containing aqueous KI solution, two inert electrodes (graphite or platinum), a DC power source (battery or power supply) with electrodes connected to positive (anode) and negative (cathode).
- Optionally use a porous barrier or divided cell if you want to keep anodic and cathodic products separate.
- Optional starch indicator can detect iodine formation. Provide ventilation and PPE (iodine vapors are hazardous).
13) Expected products at the anode (oxidation):
- I− is oxidized preferentially: 2 I− → I2 + 2 e−, so molecular iodine (I2) appears (often as brown solution or forms I3− with excess I−).
- If I− is depleted or potential is high, water oxidation to O2 can occur: 2 H2O → O2 + 4 H+ + 4 e−.
14) Expected products at the cathode (reduction):
- Water is reduced in aqueous KI: 2 H2O + 2 e− → H2 + 2 OH−, so hydrogen gas evolves and hydroxide forms near the cathode (producing KOH in bulk solution).
- Metallic K is not produced in aqueous electrolysis (it would require much more negative potential and occurs only in molten salts).
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