6.0-6.1 Book Notes
📘 CLUE 6.0–6.1 — Chemical Reactions & Equilibrium
BIG QUESTION:
How do molecules interact and rearrange to make life happen, and how do we know when a reaction “settles down”?
6.0 — Overview of Reactions
Key idea:
A chemical reaction is atoms rearranging. Bonds break and new ones form.
Example:
[
A + B \longrightarrow C + D
]
Atoms don’t disappear
Electrons shift to form new bonds
Energy is released or absorbed
Important concepts:
1⃣ Reactants & Products
Reactants = starting molecules
Products = molecules formed
2⃣ Direction of reactions
Not all reactions go only forward
Some reactions reach equilibrium (forward and reverse happen at the same rate)
3⃣ Energy matters
Reactions may require activation energy to start
Overall ΔG determines if reaction is spontaneous
4⃣ Coupled reactions
Some reactions that don’t want to happen can occur if paired with highly favorable reactions (like ATP hydrolysis)
Biological connection:
Digestion = breaking food molecules apart
Muscle contraction = coupling ATP hydrolysis to work
Enzymes = lower activation energy so reactions happen fast enough
6.1 — Reaction Rates and Equilibrium
1⃣ Reaction Rates
Rate = how fast reactants turn into products
Factors that affect rate:
Factor | Effect |
|---|---|
Concentration of reactants | More reactants → faster collisions → faster reaction |
Temperature | Higher temp → molecules move faster → more collisions → faster reaction |
Catalysts (enzymes) | Provide a “shortcut” → lower activation energy → faster reaction |
Surface area | More surface → more collisions → faster reaction (applies to solids) |
Key CLUE idea:
Reactions are molecular collisions — if molecules don’t bump into each other in the right orientation, nothing happens.
2⃣ Reversible Reactions
Many reactions are reversible:
[
A + B \rightleftharpoons C + D
]
Forward rate = rate at which A + B → C + D
Reverse rate = rate at which C + D → A + B
Eventually, rates equalize → system at equilibrium
Equilibrium does NOT mean the reaction stops — it means no net change.
3⃣ Equilibrium Constant (K)
K tells you which direction is favored:
[
K = \frac{[Products]}{[Reactants]}
]
K > 1 → more products at equilibrium
K < 1 → more reactants at equilibrium
K = 1 → equal amounts
Temperature can change K (important in biology)
4⃣ Gibbs Free Energy & Equilibrium
CLUE ties this back to Gibbs energy:
[
\Delta G = \Delta G^\circ + RT \ln\frac{[Products]}{[Reactants]}
]
ΔG = free energy at any moment
ΔG° = free energy under standard conditions
At equilibrium: ΔG = 0 → no net change
Intuition:
If ΔG < 0 → reaction moves forward
If ΔG > 0 → reaction moves backward
If ΔG = 0 → reaction is balanced
5⃣ Le Châtelier’s Principle (CLUE version)
When you stress a system at equilibrium, it will shift to relieve that stress
Stresses include:
Adding/removing reactants or products
Changing temperature
Changing pressure (for gases)
Example:
Adding more A → reaction shifts forward → more C forms
Removing D → reaction shifts forward → restores balance
Biology connection:
Cells constantly shift reactions to maintain homeostasis
Metabolism relies on pushing reactions forward or backward
6⃣ Enzymes & Reaction Direction
Enzymes speed up both forward and reverse reactions
They don’t change equilibrium, just help system get there faster
Life relies on enzymes to make reactions fast enough to matter
7⃣ Summary of Key CLUE Takeaways (6.0–6.1)
Reactions = rearranging atoms; energy matters.
Rate = how often molecules collide in the right orientation.
Reversible reactions reach equilibrium → no net change.
K and ΔG tell you which direction is favored.
Le Châtelier’s principle explains how systems respond to stress.
Enzymes make reactions faster without changing ΔG.
Cells control reactions by concentration, temperature, and coupling.
🧠 Memory Tips
Rate → think “how fast the molecules bump into each other”
Equilibrium → think “balance, not stop”
ΔG → “spontaneous or not”
Le Châtelier → “stress → shift to compensate”
Enzymes → “speed limit, not GPS”