Reaction Rate and Collision Theory

Controlling the Rate

• Chemists control reaction rates for economic viability and safety.
• Reaction rate is proportional to reactant concentration; as reactants deplete, the rate slows.
• Reaction rate is calculable by monitoring volume, concentration, or mass changes.
• Units for relative rate are s-1; units for average rate depend on the measurable quantity (e.g., g s-1 for mass change, mol L-1 s-1 for concentration change).

Average Rate of Reaction

• Average rate is calculated by the change in mass over time, measured in grams per second (g s-1).
• average rate = Δquantity / Δtime

Relative Rate of Reaction

• Relative rate is the rate at a specific point in time.
• relative rate = 1/t

Collision Theory

• For a chemical reaction, reactant molecules must collide with enough energy, known as activation energy (EA).
• Reactant molecules need sufficient energy to overcome repulsive forces and break atomic bonds.
• The activated complex is a high-energy, unstable arrangement of atoms formed as an intermediate stage.
• Successful collisions require correct collision geometry to form the activated complex, and energy is released as new bonds form.

Factors Altering Reaction Rate

Temperature

• Increased temperature gives particles more energy, leading to more frequent collisions and a faster reaction rate.
• A 10°C rise can double the reaction rate due to a significant increase in molecules with kinetic energy greater than activation energy (EK > EA).

Concentration

• Higher concentration of reactants increases particle collisions, thus increasing reaction rate.

Particle Size

• Decreasing particle size increases surface area, which increases collision opportunities.

Catalysts

• Catalysts provide a reaction surface and hold reactant molecules at favorable angles creating successful collisions and lowering activation energy which increases likelihood

Activation Energy

• Activation energy is the minimum energy for a successful reaction.
• High activation energy results in slow reactions; low activation energy yields fast reactions.

Catalysts

• Catalysts lower the temperature needed for a reaction, making industrial processes more economical.
• Catalysts in industry:
  • Iron for ammonia (Haber Process)
  • Platinum for nitric acid (Ostwald Process)
  • Rhodium and Platinum in catalytic converters
  • Nickel for margarine (hardening vegetable oil)
  • Vanadium (V) Oxide for sulfuric acid (contact process)
• Catalysts offer alternative, lower-energy pathways.
• Catalysts do not affect reactants or products, so the overall enthalpy change (ΔH) remains the same.

Potential Energy Diagrams

• Potential energy diagrams calculate enthalpy change and activation energy.

Enthalpy

• Chemical reactions involve energy changes, measured as enthalpy (H). -ΔH is the overall enthalpy change for a reaction.

Exothermic Reactions

• Exothermic reactions release heat energy; products have less energy than reactants, and ΔH is negative.
• EA is the activation energy to start the reaction which releases ΔH which is the enthalpy change.

Endothermic Reactions

• Endothermic reactions absorb heat energy; products have more enthalpy than reactants, and ΔH is positive.

Activated Complex

• The activated complex is a high-energy intermediate shown on potential energy diagrams.

Concentration of Solutions

• Solutions are formed when solutes dissolve in solvents.
• Concentration of a solution is measured in moles per litre (mol L-1).

Periodicity

Covalent Radius

• On the periodic table, covalent radius decreases across a period (left to right).
  • Atoms gain electrons and protons, increasing attraction and decreasing size.
• Covalent radius increases down a group.
  • Atoms add energy levels, shielding outer electrons and decreasing attraction.

Ionization Energy

• Ionization energy increases across a period due to increasing nuclear charge. More energy is required to remove electrons.
• Ionization energy decreases down a group because outer electrons are farther from the nucleus and more easily removed.
• First Ionization Energy of Magnesium: Mg(g) → Mg+(g) + e- (744 kJ mol-1)
• Second Ionization Energy of Magnesium: Mg+(g) → Mg2+(g) + e- (1460 kJ mol-1)
• Third Ionization Energy of Magnesium: Mg2+(g) → Mg3+(g) + e- (7750 kJ mol-1)

Electronegativity

• Electronegativity increases across a period as nuclear charge increases and covalent radius decreases, resulting in stronger attraction of bonding electrons.
• Electronegativity decreases down a group as atoms increase in size and energy levels, which lowers the attraction for bonding electrons due to shielding effect.
• Fluorine is highly electronegative, attracting bonding electrons strongly.

Bonding in Elements

Metallic Bonding

• Metallic bonding occurs between metal atoms where outer electrons are delocalized (free to move).
• Produces an electrostatic attraction between metal ions, and allows electric conductivity.

Covalent Molecules

• Discrete covalent molecules are small groups of atoms with strong covalent bonds and weak intermolecular forces.
• A covalent bond is a shared pair of electrostatically attracted electrons.
• Diatomic elements are discrete covalent molecules.
• Other molecules are phosphorous (P4), sulfur (S8), and fullerenes (C60).

Covalent Network

• Covalent networks are three-dimensional arrangements of atoms with strong covalent bonds.
• Exhibit high melting points.
• Examples are diamond and graphite.

Monatomic Elements

• Noble gases are monatomic elements
• Exhibit low melting and boiling points.

Summary of Bonding

• A table describes the properties of each type of bond
• Metallic lattices (metals) have high density, melting point, and conductivity.
• Covalent networks have very high density and melting points but are non-conductive (except graphite).
• Covalent molecules and monatomic elements have lower densities, melting points, and are non-conductive.

Bonding and Structure

Intramolecular Bonding

• Intramolecular bonding occurs inside molecules, and consists of varying degrees of covalent, polar covalent and ionic types of bonds.

Pure Covalent bonds

• Electrons are shared equally by two atoms with the same electronegativity.
• Diatomic elements are examples.

Polar Covalent bonds

• Electrons are shared unequally
• One atom attracts them more strongly in polar covalent bonds, leading to partial charges (dipoles).

Ionic Bonds

• Electrons are not shared.
• Electrons are transferred.
• Ionic bonds form between metals and nonmetals with large electronegativity differences, creating ions arranged in a 3D lattice.

Bonding Continuum

• Pure covalent, polar covalent, and ionic bonds exist on a continuum determined by electronegativity differences between atoms.

Intermolecular Bonds

• Intermolecular bonds (Van der Waals forces) occur between molecules.

London Dispersion Forces

• London dispersion forces are the weakest intermolecular bond, present in all atoms and molecules.
• These forces arise from momentary uneven electron distribution, leading to temporary dipoles.
• Strength increases with molecule size.

Permanent Dipole Interactions

• Permanent dipole interactions occur between polar molecules.
• Stronger than London dispersion forces.

Hydrogen Bonding

• Hydrogen bonding is the strongest intermolecular bond.
• It occurs between hydrogen atoms bonded to electronegative atoms (N, O, F) and a neighboring electronegative atom.

Properties Relating to Bonding

Bonding Strength

• The relative strengths of bonds are:
• Bonds > Hydrogen bonds > Permanent dipole interactions > London dispersion forces.

Polar and Non-Polar Molecules

• Molecular polarity depends on bond polarity and molecular shape.
• Symmetrical molecules are non-polar.
• Bent molecules like water is polar.
• Tetrahedral molecules like carbon Tetrachloride is non-polar.

Ionic Lattice

• exhibit high melting/boiling points, conduct electricity when molten/dissolved, and undergo electrolysis.

Covalent Network

• exhibit high melting/boiling points and do not conduct electricity.

Covalent Molecules

• exhibit low melting/boiling points and do not conduct electricity.

Properties of Water

• Water has unusual properties due to hydrogen bonding:
  • Ice is less dense than water due to hydrogen bond lattice formation.
  • High surface tension and viscosity.

Summary of Solubilities

• Ionic lattices and polar covalent molecules are soluble in polar solvents like water.
• Non-polar covalent substances are soluble in non-polar solvents such as hexane.
• "Like dissolves like" is a general rule for solubilities.

Glossary

Activated Complex

• An unstable arrangement of atoms at the maximum potential energy barrier during chemical reaction.

Activation Energy

• Is the minimum kinetic energy required for colliding particles before reaction.

Adsorption

• Occurs when molecules become bonded to the surface of a catalyst.

Allotrope

• Are one of two or more existing forms of an element; for example, diamond, fullerene, and graphite are allotropes of carbon.

Bonding Electrons

• Shared pairs of electrons from both atoms forming the covalent bond.

Chemical Bonding

• A term used to describe the mechanism by which atoms are held together.

Chemical Structure

• Describes the way in which atoms, ions, or molecules are arranged.

Collision Theory

• Reactions need to collide in order to occur.

Covalent Bond

• Formed when two atoms share electrons in their outer shell to complete that shell.

Covalent Radius

• Half he distance between the nuclei of two bonded atoms of an element.

Delocalised

• Electrons are free from attachment to any one metal ion of metallically bonded electrons.

Desorption

• Means the bonds between the molecules and the surface break.

Diatomic

• Molecules with two atoms are diatomic.

Dipole

• Concentration of electrical charge (either positive or negative).

Electronegativity

• A measure of the attraction that an atom involved in a bond has for the electrons of the bond.

Enthalpy Change

• The change in heat energy when 1 mole of reactant is converted to product(s) at a certain pressure. ΔH is the unit and kJ mol-1 is the unit.

Fullerenes

• Pure carbon with combined structures in hollow locations (like footballs)

Hydrogen Bonding

• Electrostatic force of attraction between atoms containing hydrogen and another strongly electronegative element

Intermolecular Forces

• Forces which attract molecules together

Intramolecular Forces

• Forces that exist within a molecule

Ionisation Energy

• Energy required to remove one mole of electrons from a gaseous mole

Isoelectronic

• The same arrangement of electrons

Lattice

• Where a regular 3D arrangement of particles exist, where metal ions can exist or solids with positive and negative ions in general

London Dispersion Forces

• Electrostatic attraction derived from temporary dipoles, movement causes fluctuations that arrange dipoles of atoms and molecules

Lone Pairs

• Pairs of electrons in the outer shell that do not connect

Miscible

• How fluids may mix with another fluid

Periodicity

• The order of similarity of properties as atomic number increase

Polar Covalent Bond

• Electrons are not equally distributed, causes a partial charge and an uneven distribution of electronegatvity

Potential Energy Diagram

• Explains the enthalpy of reactants and products, a view of the change during a chemical reaction

Properties

• Physicality and chemical characteristics of an object

Thermochemical Equation

• Determines the change for the reaction, the amount and the products

Viscosity

• Resistance to flow by liquids