Chemistry Chapter: Reaction Rates and Concentration

Chemistry 3202 Introduction

• Instructor email: [email protected]
• Course materials located on Google Classroom
• Seating plan based on last name
• Cell phones, earbuds, and bathroom use guidelines
• Testing and in-class assignments (ICA's)
• Attendance and homework are important
• Course description and formula sheet handouts provided

Unit 1: From Kinetics to Equilibrium - Section 1.1 - Skill #1

• Four factors affecting reaction rate:
  • Surface area
  • Temperature
  • Concentration
  • Catalyst addition

1. Surface Area

• Surface area is the exposed matter of a solid substance.
• Splits of wood have more surface area compared to a whole piece of wood
• A 5.0 g of coal dust has a greater surface area than a 5.0 g lump of coal, reacting faster.
• Rate of reaction is directly proportional to the surface area

2. Temperature

• Temperature, in Kelvin, is proportional to the kinetic energy of particles.
• Doubling the Kelvin temperature doubles the average kinetic energy of particles.
• Higher temperatures lead to more frequent and intense particle collisions, increasing the reaction rate.

3. Concentration of Reactants

• Important for gases and aqueous solutions, not pure solids (e.g., Zn).
• Higher concentration results in more frequent collisions, leading to a faster reaction rate.
• Example: Vinegar on fish and chips reacts more effectively than glacial ethanoic acid.
• Concentration of HCl solution affects the reaction rate with Zn.

4. Catalyst

• A substance that speeds up a chemical reaction without undergoing permanent changes.
• The mass of the catalyst remains constant before and after the reaction.
• Common catalysts include MnO2 in H2O2 decomposition and Fe in NH3 production
• Catalysts provide an alternate pathway with lower activation energy, increasing reaction rates.
• Example: Decomposition of hydrogen peroxide to water and oxygen with KI as catalyst.

Core Lab #1 Rates of Reaction - Section 1.1 - Skill 2

• Measuring reaction speed involves observing/controlling a measurable factor (e.g., gas volume).
• Rate of reaction is calculated as the change in quantity during a specific time period (e.g., ml/s).

Core Lab #1 Cont' Investigating the Rates of Reaction

• Maintaining constant conditions (except for the manipulated variable) is crucial for accurate experiments.
• Manipulable factors in reaction rate experiments include: H2SO4 concentration, surface area of marble chips, and temperature of flask.

1.2 Collision Theory - Kinetic Molecular Theory

• All matter forms atoms, ions, or molecules, constantly moving randomly and colliding with each other/container walls.
• Kinetic energy (Ek) associated with each collision remains constant.
• Temperature of the substance remains consistent
• Pressure exists in gases because of this characteristic.

1.2 Collision Theory – Skill #2

• Pressure: Gas molecule collisions with the container walls cause pressure. Inflating a balloon, fluid pressure increasing with depth, are examples.
• Diffusion: Gas molecules spread out as far as possible. Examples include scents, dyes in water, osmosis in body cells.

1.2 Collision Theory – Skill #3

• Essential Conditions for Reaction: Three factors determine collisions leading to reactions
  • Reactant particles must collide
  • Colliding particles must have proper orientation for reactions
  • Sufficient collision intensity/energy needed for bond breaking / formation. (To overcome activation energy)

1.2 Collision Theory - Example

• Illustrates successful vs. unsuccessful reactions based on collisions:
  • Atom contact
  • Correct orientation of atoms for reactions
  • Minimum energy requirement for effective reactions.

1. Using Observables to Increase Reaction Rate

• Increasing the surface area of a reactant (e.g., breaking up a solid).
• Increasing the concentration of a reactant.
• Adding a catalyst.
• Increasing reaction temperature.

Jeopardy! Question 2

• The correct method for safely removing a stuck lid from a bottle of pickles is to run ice-cold water over the lid.

Jeopardy! Question 3

• Water molecules move faster at higher temperatures (80°C) compared to lower temperatures (15°C).

1.2 Collision Theory – Skill #4: Macroscopic Properties

• Measurable physical properties of chemicals during reactions.
• Variables include mass, pH, electrical conductivity, gas pressure, color change, gas volume, and concentration.
• Units for reaction rates involve quantities per time.

Example #4

• Units for Macrosopic Properties: Mass (grams/sec), Volume (mL/sec), pH (pH/sec), Gas pressure (kPa/sec), Gas volume (mL/sec)
• Cannot be measured quantitatively: Concentration and colour change

Section 1.3 Factors Affecting Reaction Rate (Part II) – Skill #1

• Temperature:
  • Average kinetic energy per particle increases with temperature.
  • Particles collide with more energy and greater frequency
  • Increased chance of successful collisions with correct orientation, leading to faster reaction
• Explanation of milk spoiling faster in a warm place
• Temperature increase means particles collide more often and with more energy this more effective collisions increasing reaction rate.

Section 1.3 Factors Affecting Reaction Rate (Part II) – Skill#2: Concentration

• Concentration is the number of particles per unit volume.
• Higher concentration; greater frequency of collisions and particles are closer this increase chances (and frequency) of successful collisions with enough energy (Ea) and correct orientation - leading to faster reaction rate
• Example*
• Which reaction could be “sped up” by increasing the concentration? Combustion of CH4 (g), A-B neutralization. NOT a solid.

Section 1.3 Factors Affecting Reaction Rate (Part II) – Skill #3: Pressure

• Pressure in a gas relates directly to the concentration of a gas.
• Decreasing volume increases pressure, increasing the concentration of reacting particles.
• Increasing pressure leads to more frequent collisions and higher chances of successful collisions, leading to a faster reaction rate.
• Example: How to speed up the reaction H₂(g) + Cl₂(g) → 2HCl(g) by changing the container volume?
• Decreasing the volume to increase pressure in the container.

Section 1.3 Factors Affecting Reaction Rate (Part II) – Skill #4: Surface Area

• Surface Area of solids affects reactions.
• Smaller particles (larger surface area) have more exposed surface, increasing collision frequency with other particles and, thus, the reaction rate.
• Example: Checking your car's rusting areas.

Skill 2: Addition of a Reaction-Specific Catalyst

• Catalysts provide an alternative reaction mechanism with lower activation energy.
• This reduces the energy threshold for collisions, thereby increasing the chances (and frequency) of successful collisions, leading to faster reaction rates.

Section 1.4 Potential Energy – Skill #1: Endothermic Reactions

• Potential energy is defined as the energy due to position.
• Reactions that absorb energy are called endothermic.
• The change in energy (enthalpy) is calculated by the difference in energy between the products and reactants. Endothermic reactions have a positive ΔH value.
• Molar Enthalpy (ΔH) – change in energy for a reaction

2 ways to represent Endothermic reactions

• A) Including energy as a reactant.
• B) Writing the enthalpy change separately

Skill #2: Exothermic Reactions

• Potential energy reactions that release energy called exothermic.
• Exothermic reactions have a negative ΔH value.

1.4 Skill #3 Potential Energy Diagrams

• Terminology:
  • Activation energy: The energy required to start the reaction. Ea(forward) or Ea(reverse)
  • Heat of Reaction: (∆H) is the difference of energy between products and reactants
  • Endothermic vs Exothermic (energy absorbed vs released)

Hunt the Letters (Potential Energy Diagrams)

• Identify specific parts (activation energy, enthalpy, activated complex) of a potential energy diagram for a chemical reaction.

Example #4 (Exothermic Graph): Calculate values

• Identify activation energy (Ea forward and Ea reverse), and enthalpy changes (ΔH forward and ΔH reverse) from an exothermic graph.

Extra Practice (Potential Energy Diagram)

• Identify features/symbols (e.g., activation energy) on a given graph.
• Calculate activation energy for a reverse reaction using heat of reaction and activation energy for forward reaction.

Ex 3: More word Problems (Potential Energy Diagrams)

• Sketch an energy profile graph, classify (endo or exo) a reaction based on the energy profile, and determine enthalpy change for a reaction.

Skill #4: Catalysts with Reactions

• Catalysts speed up reactions by providing alternate reaction pathways with lower activation energy.
• Enthalpy (ΔH) for a reaction remains unchanged during the presence of a catalyst.

Example #2: Catalysts

• Identify minimum catalyzed activation energy from a given uncatalyzed reaction and heat of reaction.
• Find maximum possible catalyzed activation energy for the reverse reaction.

Section 1.5 Reaction Mechanisms – Skill #1: Overall Reaction

• Most reaction occur in multiple small reaction steps, called elementary steps.
• Overall reaction is the sum of elementary steps.
• Reaction intermediates are unstable substances formed and consumed during intermediate steps of multiple reactions. They are not part of the overall reaction equation.

Examples for Hess's Law

• Calculate overall enthalpy change for reaction using values from elementary steps.
• Find missing enthalpy change for elementary steps using Hess's law.

Skill #2: Identifying Reaction Intermediates as catalysts

• Identify reaction intermediate from multiple elementary reaction steps, and overall Reaction.

Catalysts

• Catalysts are substances that speed up a chemical reaction without being consumed in the process.
• Catalysts provide alternate reaction pathways that have lower activation energies compared to uncatalyzed reactions.
• Reaction mechanism for catalyzed reaction is different from uncatalyzed.

Example: Catalyst

• Identifies the overall reaction, reaction intermediates, and the catalyst given multiple elementary steps.

Rate Determining Step

• The slowest step in a multi-step reaction mechanism determines the overall reaction rate.
• Each elementary reaction step has its own specific activation energy. (The Ea of the RDS is the highest)

R.D.S Example

• Determine overall reaction, reaction intermediate, catalyst, and rate-determining step for a given set of chemical equations.

1.5 Skill #2 Continued (Level 3 question)

• Determining missing reactants and products/intermediate from the overall reaction and provided elementary steps.
• Identifying the substances and what shouldn´t be in the overall reaction.

Example (Rates & Reaction Mechanisms)

• Determine missing elementary step and its enthalpy from given information of Overall Reaction, R.I, and RDS, if any.
• Determine which substances (reactants or products) should increase concentration, increasing the overall reaction rate.

1.5 Skill #3 Ep diagrams for multi-step reactions

• Relationship between # of elementary steps vs # of activated complexes (peaks) on an energy profile graph.
• Identify a) Rate-determining step, b) Enthalpy change (∆H) for each reaction step based on a given energy profile graph. c) Classify reaction (endothermic or exothermic) based on energy changes.
• Find which step is most likely to be catalyzed in the given energy profile graph.

Style #2 (Energy Profile Graph construction):

• Construct a potential energy diagram from a provided reaction mechanism.
• Determine the reaction intermediate(s), catalyst(s), rate-determining step(s)
• Calculate the missing enthalpy changes (∆H) for each step in the mechanism.

Test your understanding of how concentration affects reaction rates with this quiz. Explore concepts such as the relationship between reactant concentration and the speed of chemical reactions, as well as factors influencing rate changes. Perfect for students studying the principles of chemical kinetics.