Class 1: Collision Theory Flashcards
Describe what interactions are involved in a physical vs. a chemical change.
A physical change breaks the IMF— Solid → Liquid → Gas, ect.
A chemical change breaks the covalent or ionic bonds— N2 + 2H2 → 2 NH2
Physical Change:
- No new substances formed
- Molecules retain their composition and structure
- Only interactions holding molecules/atoms together are disrupted
- Intermolecular forces broken/re-formed
- Ionic, dipole-dipole, hydrogen bonding, van der Waals
- Energy required to overcome these forces
- Examples: Melting, boiling, sublimation
Chemical Change:
- New substances with different properties are formed
- Molecules break apart and re-form with new bonds
- Intramolecular bonds are broken and new ones created
- Interactions between atoms within molecules are disrupted
- Electron sharing/transfer between atoms
- Energy absorbed/released as chemical bonds break/form
- Examples: Combustion, photosynthesis, acid-base reactions
Key Differences:
- Physical - No bonds broken within molecules
- Chemical - Bonds within molecules broken/re-formed
- Physical - Intermolecular forces disrupted
- Chemical - Intramolecular bonds disrupted
- Physical - Composition stays the same
- Chemical - Composition changes to new substance(s)
Write reactions to illustrate changes of state.
- Physical changes of state:
- Melting (solid → liquid)
- Freezing (liquid → solid)
- Vaporization/Evaporation (liquid → gas)
- Condensation (gas → liquid)
- Sublimation (solid → gas)
- Deposition (gas → solid)
- Chemical changes of state:
- Combustion (reactants → products + heat/light)
- Decomposition (reactant → multiple products)
- Single displacement (element + compound → new compounds)
- Double displacement (two compounds exchange ions/atoms)
- Represent with balanced chemical equations
- Reactants on left, products on right
- Coefficients balance atoms/molecules
- State symbols: (s) solid, (l) liquid, (g) gas, (aq) aqueous
- Examples:
- H2O(s) → H2O(l) (melting)
- 2Na(s) + 2HCl(aq) → 2NaCl(aq) + H2(g) (single displacement)
- 2NaHCO3(s) → Na2CO3(s) + H2O(l) + CO2(g) (decomposition)
Define temperature and how it relates to kinetic energy.
Temperature is a measure of the average kinetic energy of the molecules (gas or in solution)
Breaking bonds requires an input of energy from somewhere else
Conservation of energy means that (during an ISOLATED collision):
When bonds are broken, the atoms PE goes up so their KE must decrease (takes up energy)—KE→ PE
When bonds are formed, the atoms PE goes down so their KE must go up (releases energy)—PE→ KE
If a molecules KE is not high enough, it cannot overcome the activation energy of the reaction
Temperature:
- Measure of average kinetic energy of particles in a substance
- Kinetic energy from vibrational/rotational/translational motion
- Higher temperature = higher average kinetic energy
Kinetic Energy (KE):
- Energy of motion possessed by particles
- KE = 1/2 mv^2 (m = mass, v = velocity)
- Higher velocities = higher kinetic energies
Temperature and KE Relationship:
- As temperature increases, particles gain more KE
- Higher temperatures = faster average particle velocities
- Increased particle velocities = increased collision rates/frequency
- More collisions transfer KE, heating the substance
- At higher temps, particles have more KE to overcome intermolecular forces
Key Points:
- Temperature is macroscopic; kinetic energy is microscopic
- Both increase with faster, more energetic particle motion
- Higher temps drive faster motion/more particle collisions
- Higher KE allows particles to escape intermolecular attractions
Define activation energy and explain how changing temperature affects particle collisions and the rate of reaction.
The potential energy barrier to a reaction
This energy is required because we need to partially break or at least deform one bond first before forming a new one
Kinetic energy is key to make the reaction happen
If Ea > KEavg then the reaction will occur slowly, only a few fast moving will react
If Ea < KEavg then the fraction of molecules with enough energy to react is large so the reaction will occur quickly
The Ea of the forward reaction + change in H is the Ea of the reverse reaction
If there is NO PARTICLES with enough KE to get over the Ea barrier, no spontaneous reaction will occur [e.g. when the total energy of the system is lower than the activation energy]
Here are concise notes on activation energy, and how changing temperature affects particle collisions and reaction rates:
Activation Energy:
- Minimum energy required to start a chemical reaction
- Energy needed to break bonds in reactant molecules
- Barrier that must be overcome for reaction to occur
Effect of Temperature on Particle Collisions:
- Higher temperature = faster moving particles with more kinetic energy
- More particle collisions occur per unit time
- More particles have enough energy to overcome activation energy barrier
Effect on Reaction Rate:
- At higher temperatures, more particles have KE > activation energy
- More effective collisions occur that can initiate the reaction
- Rate of reaction increases exponentially with temperature
- Lower temperatures result in fewer effective collisions and slower rates
Key Points:
- Activation energy is the energy barrier to reaction
- Higher temps provide more KE to particles
- More KE allows more particles to overcome activation energy
- More effective collisions occur at higher temps
- Reaction rate increases exponentially with temperature increase
Describe how particle orientation affects whether or not a reaction will occur.
Even if you have enough energy for a reaction, if the molecules do not collide in a position that geometrically promotes the weakening of the breaking bond and the forming of the new bond then it will not react.
They need to react in a way that makes sense/the way that they would be bonded
Here are concise notes on how particle orientation affects whether or not a chemical reaction will occur:
Particle Orientation:
- Describes the relative positioning and alignment of particles during collisions
- Particles must collide with proper orientation for reaction to occur
Reaction Requirements:
- Effective collisions require sufficient energy (KE > activation energy)
- Particles must also collide with correct orientations
- Proper alignment allows formation of new bonds/rearrangement
Orientation Effects:
- If particles collide with improper orientation, no reaction occurs
- Particles must line up so bonds can break and reform properly
- Correct orientation allows reactive sites to interact
- More orientations possible = higher probability of reaction
Increasing Odds:
- Higher temperatures increase particle motion and number of collisions
- More collisions mean more chances for proper orientation
- Catalysts provide alternative reaction pathways with easier orientations
- Overall, proper orientation is critical for an effective, reactive collision
Key Points:
- Orientation is the alignment of particles during collision
- Particles must collide with proper orientation for reaction
- Incorrect alignments prevent bond making/breaking
- Higher temps and catalysts increase odds of proper orientations
Explain how some particles can react even when the average kinetic energy of the system is less than the activation energy for the reaction.
The kinetic energy is the AVERAGE of all of the molecular energies—some will have a higher kinetic energy above the average
If the total energy is too small, then the reaction will not occur without the addition of energy into the system
Here are concise notes explaining how some particles can react even when the average kinetic energy of the system is less than the activation energy:
Kinetic Energy Distribution:
- At any temperature, particles in a system have a range of kinetic energies
- Though most have average KE, some have much higher or lower KE
- KE is distributed among particles according to Maxwell-Boltzmann distribution
Particle Collisions:
- For a reaction to occur, colliding particles must have KE ≥ Activation Energy
- Even if average KE < Activation Energy, reaction can still happen
- This is due to the high-KE tail of the Maxwell-Boltzmann distribution
High Energy Particles:
- A small fraction of particles will have KE»_space; Average KE
- Some of these high-KE particles have enough energy to overcome Ea barrier
- They can collide effectively and initiate the reaction
Increasing Temperature:
- As temperature rises, the high-KE tail of the distribution grows larger
- More particles populate the high-KE region above the Ea
- This increases the number of effective, reactive collisions
- So reactions can occur even when average KE < Ea
Key Points:
- Particle KEs are distributed, not all have the average value
- High-KE tail means some particles have KE > Activation Energy
- These high-KE collisions allow reactions when average KE < Ea
- Raising temperature increases high-KE particle population
Describe how concentration influences rates within a gas-phase reaction
When you increase the concentration of either one, the increase in concentration is proportional with the increase in collisions
EX: 2x the amount of C will just about double the amount of collisions
Here are concise notes describing how concentration influences rates within a gas-phase reaction:
Concentration and Collision Theory:
- For a reaction to occur, particles must collide
- Higher concentrations mean more particles per unit volume
- More particles leads to more frequent collisions
Effect on Collision Frequency:
- In gases, particles move randomly and collide
- Higher gas concentrations increase number of particles in a given space
- More particles per volume leads to a higher collision frequency
Rates and Concentration:
- Reaction rates depend on number of effective, reactive collisions
- With higher concentrations, there are more total collisions
- This increases the number of reactive, product-forming collisions
- So higher concentrations directly increase the reaction rate
Concentration Dependence:
- For many reactions, rate is directly proportional to concentrations
- Rate = k[A]x[B]y (k is rate constant, x and y are reaction orders)
- Doubling a reactant concentration can double or triple the rate
Exceptions:
- Zero-order reactions have rates independent of concentrations
- But for most reactions, increasing concentrations increases rate
Key Points:
- Higher concentrations mean more particles/volume
- More particles leads to more total collisions
- This increases reactive, product-forming collisions
- So higher gas concentrations directly increase reaction rates
Interpret and draw thermal energy distribution graphs.
The graph is a probability distribution graph
The greater the area under the curve, the greater the amount of electrons with enough energy to overcome the activation energy barrier
Here are concise notes on interpreting and drawing thermal energy distribution graphs:
Thermal Energy Distribution Graphs:
- Plot fraction/percentage of particles vs. their kinetic energies
- Shows distribution of kinetic energies in a sample at a given temp
Graph Features:
- X-axis: Kinetic Energy
- Y-axis: Fraction/Percentage of Particles
- Peak indicates most probable/average kinetic energy
- Curve is asymmetrical, tails off to higher and lower energies
Interpreting Graphs:
- Higher temperatures shift peak and curve to higher kinetic energies
- Wider curve means greater spread of kinetic energies
- Larger high-energy tail = more particles with energy > Activation Energy
Drawing Graphs:
- Plot relativistic Maxwell-Boltzmann distribution curves
- Higher temps = peak shifts right (higher average KE)
- Wider, flatter curve for higher temps (broader KE distribution)
- Larger area under high-energy tail for higher temps
Key Points:
- Graphs show spread of kinetic energies at a given temperature
- Higher temps = higher average KE and wider energy distribution
- Can estimate reactive fraction from high-energy tail area
- Useful for visualizing effects of temperature on reaction rates
What factors related to the orientations of the colliding molecules and their direction of motion seem
to be important?
- Collision where the particles are moving directly towards each other will have greater
impact and provide more energy to activate the reaction than where they are moving more
nearly in the same direction - A collision that brings the A and C directly in contact will not produce a reaction, since
these atoms do not form a bond
How is it possible for a reaction to occur if the total average (kinetic) energy supplied is less than the
activation energy?
As they collide, some molecules gain more than the average energy at the expense of others;
these extra-energetic molecules are able to react.
Extra Notes:
Postulates of Collision theory
1. The rate of reaction is proportional to the rate of reactant collisions— = means proportional: [reaction rate = # collisions/time]
2. The reacting species must collide in an orientation that allows contact between atoms correctly to bond
3. The collision must occur with adequate energy to permit a mutual penetration of the reacting species valence shells so that the electrons can rearrange to form new bonds
Concentration is key to determining the rate of collisions
Changing temperature and concentration won’t have significant effect on the orientation
Raising the P.E. to break a bond requires an input of energy from somewhere else -ex. kinetic energy
Temperature
is a measure of KINETIC ENERGY at a molecular level
Kelvin is proportional to the amount of KE per atom- includes intramolecular motion and motion of the whole molecule
conservation of energy MEANS that DURING AN ISOLATED COLLISION
when a bond is BROKEN potential energy goes up so Kinetic energy goes down
when a bond is FORMED potential energy goes down so kinetic energy goes up
further collisions will redistribute KE among molecules
If the molecules KE is not high enough , breaking bond is IMPOSSIBLE
breaking bonds more likely at high
