Book 1 Chapter 11 Flashcards
(27 cards)
Titration is an analytical technique used to determine the unknown concentration of a substance by reacting it with a solution of known concentration. It typically involves the gradual addition of a standard solution (titrant) from a burette to a known volume of the analyte until the reaction reaches the equivalence point. An appropriate indicator is often used to signal when the reaction is complete, usually by a sharp colour change. The volume of titrant used allows for the calculation of the unknown concentration using stoichiometry. Titration is widely used in acid-base, redox, and complexometric analyses.
Mass change refers to the difference in mass observed during a chemical reaction, often due to the release of a gas or the formation of a solid. This technique can be used to monitor reaction rates or determine the amount of a product or reactant. A balance is used to record the mass at regular time intervals. The rate of mass loss can give insight into the speed of a gas-evolving reaction. It is especially useful for reactions such as decomposition or acid-carbonate reactions.
A substance is said to be titrated when it is subjected to the process of titration, usually to determine its exact concentration. During the experiment, the titrated solution reacts with a standard reagent until a clear endpoint is observed. The term implies that a controlled reaction has taken place between the analyte and titrant. The precision of the titration depends on accurate measurement and clear endpoint detection. A titrated sample provides data for quantitative chemical analysis.
The noun titrate refers to the substance or sample that is being analysed through the titration process. It is the chemical solution with an unknown concentration, to which a titrant is added until the reaction is complete. The titrate reacts with the titrant in a fixed molar ratio based on a balanced chemical equation. Accurate determination of the titrate’s concentration relies on proper technique and clear endpoint identification. In practice, the titrate is often placed in a conical flask during titration.
Colorimetry is a quantitative technique used to determine the concentration of coloured substances in solution by measuring light absorbance. A colorimeter passes light of a specific wavelength through a sample and detects how much light is absorbed. The absorbance is directly related to concentration via the Beer-Lambert Law. This technique is particularly useful for analysing reaction rates involving coloured compounds or metal ions. It provides rapid and accurate measurements for both qualitative and kinetic investigations.
This method measures the amount of gas produced during a chemical reaction, often using a gas syringe or an inverted measuring cylinder. The volume collected over time can be used to calculate the reaction rate or determine the amount of reactant consumed. It is particularly suitable for reactions that produce gases like CO₂, H₂, or O₂. Accurate timing and temperature control are important to ensure reliable data. It is a direct and visual method of tracking progress in gas-generating reactions.
a technique refers to a standardised and reliable method or procedure used in experimental or analytical tasks. Examples include titration, filtration, distillation, or colorimetry, each with specific steps to ensure accurate and repeatable results. Good technique involves attention to detail, appropriate use of equipment, and minimisation of errors. Mastery of techniques is essential for drawing valid conclusions from practical work. Techniques are assessed for precision, accuracy, and safety.
The initial-rate method involves measuring the rate of a reaction immediately after it begins, before significant changes in concentration occur. This approach ensures that the measured rate reflects only the starting concentrations of the reactants. It typically involves repeating the experiment with varying reactant concentrations to determine orders of reaction. A plot of rate versus concentration can then be used to deduce kinetic relationships. This method is particularly useful for deriving a rate equation experimentally.
To investigate a reaction rate means to observe and measure how quickly a chemical reaction proceeds under controlled conditions. This involves monitoring changes such as mass, volume, colour, or concentration over time. Techniques such as titration, colorimetry, or gas collection may be used. Variables like temperature, concentration, and surface area are often altered to study their effect on the rate. The data collected helps derive conclusions about the kinetic behaviour of the reaction.
Continuous Monitoring Method
The continuous monitoring method tracks a reaction’s progress by recording changes (e.g., colour, mass, volume, pH) at regular intervals throughout the reaction. It provides a complete dataset from start to finish, allowing for analysis of rate changes over time. Graphs such as concentration–time or volume–time can be plotted to examine reaction dynamics. This method allows for half-life determination and is especially useful for first-order kinetics. It is widely applied in school experiments and research settings.
The rate of reaction measures how quickly the concentration of a reactant decreases or a product increases over time, usually expressed in mol dm⁻³ s⁻¹. It depends on factors such as concentration, temperature, surface area, and catalysts. For example, the rate of reaction increases when the concentration of hydrochloric acid increases in its reaction with magnesium.
Examples:
* The reaction between sodium thiosulfate and hydrochloric acid becomes faster as temperature rises.
* The burning of methane gas is a rapid reaction.
* Rusting of iron is a slow reaction rate.
* Enzymes increase the rate of biochemical reactions.
* Increasing surface area of a solid reactant speeds up its reaction rate.
A rate equation shows the mathematical relationship between the rate of reaction and the concentration of reactants raised to their orders. It is written as rate = k[A]^m[B]^n, where k is the rate constant and m, n are the orders of reaction. For example, rate = k[NO]^2[O₂] means the reaction is second order in NO and first order in O₂.
Examples:
* If doubling [A] doubles the rate, rate = k[A] (first order).
* Rate = k means zero order (rate independent of concentration).
* Rate = k[B]^2 means second order with respect to B.
* Experimental data is required to find the rate equation.
* Rate equations do not always match the balanced chemical equation.
Order refers to the power to which the concentration of a reactant is raised in the rate equation, indicating how concentration affects reaction rate. It must be determined experimentally, not from the balanced equation. For instance, a reaction first order in A means the rate doubles when [A] doubles.
Examples:
* First order: rate doubles as concentration doubles.
* Zero order: rate remains constant regardless of concentration.
* Second order: rate quadruples when concentration doubles.
* Half order (less common) means rate increases by the square root of concentration.
* Fractional orders can occur in complex reactions.
Order with Respect to a Substance in a Rate Equation
This is the exponent indicating how the concentration of a specific reactant affects the rate. For example, in rate = k[A]^2[B]^0, the order with respect to A is 2 and with respect to B is 0. It reflects the sensitivity of the rate to that particular reactant.
Examples:
* If rate doubles when [A] doubles, order in A is 1.
* If rate remains unchanged when [B] changes, order in B is 0.
* Order in C could be 3 if rate increases eightfold when [C] doubles.
* Negative orders mean rate decreases as concentration increases.
* The order can be fractional, e.g., 0.5.
Overall Order of a Reaction
The overall order is the sum of the orders with respect to each reactant in the rate equation. For example, if rate = k[A]^1[B]^2, the overall order is 3. It indicates how the overall concentration changes affect the rate.
Examples:
* Zero overall order means rate does not depend on concentration.
* First order overall means rate proportional to concentration.
* Second order overall could come from one reactant squared or two first orders added.
* A reaction with rate = k[A][B] has overall order 2.
* Overall order helps determine units for the rate constant.
The rate constant (k) is a proportionality constant linking the rate and concentrations in the rate equation. It depends on temperature and catalysts but not on reactant concentration. For example, increasing temperature generally increases k.
Examples:
* For a first order reaction, k has units s⁻¹.
* For a second order reaction, k has units mol⁻¹ dm³ s⁻¹.
* Catalysts increase k without being consumed.
* A higher activation energy means a lower k at a given temperature.
* The Arrhenius equation relates k to temperature and activation energy.
Half-Life
Half-life is the time taken for the concentration of a reactant to reduce to half its initial value. In first-order reactions, half-life is constant regardless of concentration. For example, radioactive decay is a first-order process with a fixed half-life.
Examples:
* The half-life of a drug in the bloodstream can indicate dosing intervals.
* For zero-order reactions, half-life decreases as concentration decreases.
* Half-life can be calculated from a concentration-time graph.
* Successive half-lives for first-order reactions are equal.
* Half-life can be used to estimate the rate constant.
The rate-determining step is the slowest step in a multi-step reaction mechanism that limits the overall rate. Only reactants involved in this step appear in the rate equation. For example, in the reaction mechanism for NO₂ + CO, the slow step determines the rate law.
Examples:
* A fast initial step followed by a slow step means the slow step controls the rate.
* The rate law is derived from the slowest step’s reactants.
* Intermediate species may be formed in fast steps.
* Catalysts can change the rate-determining step.
* Complex mechanisms often have more than one step but only one slow step.
Activation energy is the minimum energy that colliding particles must have for a reaction to occur. It determines how temperature affects the rate of reaction. Lower activation energy means a faster reaction at a given temperature.
Examples:
* Enzymes lower activation energy in biological reactions.
* Catalysts reduce activation energy without being consumed.
* The peak in an energy profile diagram represents activation energy.
* High activation energy leads to slower reactions at room temperature.
* The Arrhenius equation shows the effect of activation energy on rate constant.
A heterogeneous catalyst exists in a different phase from the reactants, usually a solid catalyst with gaseous or liquid reactants. It works by providing a surface for reactants to adsorb and react. For example, iron catalyses the Haber process where gases react on a solid surface.
Examples:
* Platinum in catalytic converters for car exhausts.
* Vanadium(V) oxide in the Contact process for sulfuric acid.
* Nickel in hydrogenation of alkenes.
* Solid catalysts can be poisoned by impurities.
* Surface area of the catalyst affects its activity.
Homogeneous Catalyst
A homogeneous catalyst is in the same phase as the reactants, often a liquid solution. It forms intermediate species with reactants and is regenerated at the end of the reaction. For example, Fe²⁺ ions catalyse the reaction between S₂O₈²⁻ and I⁻ in aqueous solution.
Examples:
* Acid catalysts like H⁺ ions in esterification reactions.
* Enzymes in biological systems.
* Transition metal ions catalysing redox reactions in solution.
* Homogeneous catalysts may be sensitive to pH changes.
* These catalysts provide an alternative reaction pathway.
Mechanism
A mechanism is the step-by-step sequence of elementary reactions by which an overall chemical change occurs. It explains how reactants transform into products at the molecular level. For example, the mechanism of the reaction between NO and O₃ involves intermediate species like NO₂.
Species
A species is any chemical entity involved in a reaction, including atoms, ions, molecules, or intermediates. For example, in the reaction between H₂ and Cl₂, the species are H₂, Cl₂, and HCl. The behaviour of different species determines the reaction pathway.
Rate-Determining Step
The rate-determining step is the slowest step in a multi-step reaction that controls the overall reaction rate. Only the reactants involved in this step typically appear in the rate equation. It acts as a bottleneck limiting how fast products form.
