Instantaneous Reaction Rates and Reaction Mechanisms

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Collision Theory Noun

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[kuh-li-zhun thee-uh-ree]

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Collision Theory

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A principle stating that reacting particles must collide with sufficient energy and correct orientation in order to form products.

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Example: This image shows that for a chemical reaction to occur, molecules must collide with the correct orientation, a key part of Collision Theory.

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Activated Complex Noun

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[ak-tuh-vay-tid kom-pleks]

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Activated Complex

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A temporary, unstable arrangement of atoms at the activation energy peak where old bonds break and new bonds form.

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Example: This graph shows that for a chemical reaction to happen, reactants must gain enough energy to form a high-energy, unstable 'activated complex' at the peak.

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Activation Energy Noun

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[ak-tuh-vay-shun en-er-jee]

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Activation Energy

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The minimum amount of energy that reacting particles must possess to form the activated complex and start a reaction.

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Example: This diagram shows that reactants need a boost of energy, called activation energy, to get over an 'energy hill' and turn into products.

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Average Reaction Rate Noun

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[av-er-ij ree-ak-shun rayt]

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Average Reaction Rate

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The change in concentration of a reactant or product over a specific time interval, expressed in mol/(L·s).

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Example: This image shows a chemical reaction over time. The average reaction rate is found by measuring how fast products, like the gas bubbles, are made.

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Instantaneous Reaction Rate Noun

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[in-stun-tay-nee-us ree-ak-shun rayt]

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Instantaneous Reaction Rate

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The rate of a reaction at a specific moment in time, determined by the slope of a concentration-versus-time curve.

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Example: This graph shows that the instantaneous reaction rate at a specific moment (t=500s) is the slope of the tangent line to the concentration-versus-time curve.

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Rate Law Noun

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[rayt law]

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Rate Law

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A mathematical equation that relates the rate of a chemical reaction to the concentrations of the reactants involved.

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Example: This table shows that as the concentration of a substance (reactant A) increases, the speed (rate) of the chemical reaction also increases.

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Reaction Order Noun

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[ree-ak-shun or-der]

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Reaction Order

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The exponent for a reactant's concentration in the rate law, which defines how the rate is affected by it.

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Example: This table shows that for a zero-order reaction, changing the amount (concentration) of a substance does not change how fast the reaction happens (the rate).

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Specific Rate Constant (k) Noun

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[spuh-sif-ik rayt kon-stunt]

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Specific Rate Constant (k)

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A numerical value in a rate law relating the reaction rate and reactant concentrations at a given temperature.

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Example: This image shows mathematical formulas used to determine the units for the rate constant, 'k'. The specific units for 'k' depend on the overall order of the chemical reaction.

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Method of Initial Rates Noun

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[meth-ud uv i-nish-ul rayts]

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Method of Initial Rates

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An experimental procedure for determining reaction order and the rate law by comparing initial rates at different starting concentrations.

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Example: This table shows how changing the amount of starting chemicals (A and B) in different experiments (trials) affects how fast a reaction happens (Initial Rate).

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Reaction Mechanism Noun

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[ree-ak-shun mek-uh-niz-um]

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Reaction Mechanism

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The complete sequence of elementary steps that make up a complex chemical reaction from reactants to final products.

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Example: This diagram shows how two hydrogen iodide molecules collide and rearrange to form a temporary transition state before becoming new hydrogen and iodine molecules, illustrating the step-by-step path of a reaction.