Chemistry colloq 2 Flashcards
Equilibrium
state in a reversible reaction where the rate of the forward reaction is equal to the rate of the reverse reaction, resulting in no net change in the concentrations of the reactants and products over time.
2 types of equilibriums:
Homogeneous equilibrium: All reactants and products are in the same
state (e.g all gaseous)
Heterogeneous equilibrium: Reactants and products are in different
states
- If ΔG∘ < 0, the reaction is spontaneous under standard conditions, and
the equilibrium will favor WHAT? - If ΔG∘ = 0, the reaction is non-spontaneous under standard conditions and the equilibrium will favor WHAT?
- If ΔG∘ > 0, the reaction is non-spontaneous under standard conditions,
and the equilibrium will favor the WHAT? (small Keq)
- Products
- Reactants
- Reactants
Le Chatelier’s principle:
“If a dynamic equilibrium is disturbed by changing the conditions, the
position of equilibrium will shift to counteract the change”
A+B ⇋ C+D
4 things that can be added to chemical equilibrium
Concentration
Pressure
Temperature
Catalyst
How change of concentration affects equilibrium
A+B ⇋ C+D
Increase in Reactants:
Adding more of a reactant causes the
equilibrium to shift to the right (toward the products) to consume the
added reactant.
Decrease in Reactants: Removing a reactant causes the
equilibrium to shift to the left (toward the reactants) to replace the
lost reactant.
AND VICE VERSA WITH PRODUCTS
Similarly, changes in the concentration of products will have the
opposite
How change of pressure affects equilibrium
A+B ⇋ C+D
If pressure is increased by reducing the volume of the
system, the equilibrium shifts toward the side with fewer gas
molecules to reduce pressure.
If pressure is decreased (by increasing volume), the
equilibrium shifts toward the side with more gas molecules.
How change of temperature affects equilibrium
A+B ⇋ C+D
- For an exothermic reaction (releases heat): Increasing temperature shifts the equilibrium to the left (toward reactants) as the system attempts to absorb the added heat.
- For an endothermic reaction (absorbs heat): Increasing
temperature shifts the equilibrium to the right (toward products).
Addition of a Catalyst
A catalyst speeds up the rate at which equilibrium is reached but does not shift the position of the equilibrium itself. It lowers the activation energy for both the forward and reverse reactions equally
Dissociation Degree (α)
The degree of dissociation (α) refers to the how much of a substance
has dissociated into ions or smaller molecules when it dissolves in a
solvent
Dissociation Degree (α)
numbers. Explainish
Equilibrium Constant (Keq)
The equilibrium constant (Keq) describes the ratio of the concentrations of products to reactants at equilibrium for a reversible chemical reaction.
aA + bB ⇌ cC + dD
Solubility product constant (Ksp)
The solubility product constant (Ksp) is the equilibrium constant for the dissolution of nearly insoluble salts in a solution.
Ksp= [A+] × [B−]
[A+] is the concentration of the cation (A) in the solution at equilibrium.
[B−] is the concentration of the anion (B) in the solution at equilibrium.
Ostwald’s dilution law and
Dissociation equilibrium constant
(Edis) OR (Kdis)
Ostwald’s dilution law is used to calculate the dissociation equilibrium
constant (Edis)
The dissociation equilibrium constant (Edis) describes the equilibrium
of the dissociation of a weak electrolyte in a solution.
Common Ion Effect
The solubility of a salt decreases when a solution** already contains** one of the ions present in the salt, due to the shift in equilibrium to counteract the
change.
Complex Compounds
- Structure of Complex Compound
- Central Metal Ion:
- Ligands:
- consist of a central metal ion or atom bound to surrounding ligands, which can be atoms, ions, or molecules.
Central metal
Definition: The metal ion or atom at the core of the complex.
Characteristics: Typically transition metals with varying oxidation
states.
Coordination Number: The number of ligand atoms directly bonded to
the central metal.
Ligands:
Definition: Molecules or ions that coordinate to the central metal.
Types: Monodentate (bind through a single site), bidentate (bind
through two sites), polydentate (bind through multiple sites).
Common Ligands: Water (H₂O), ammonia (NH₃), cyanide (CN⁻),
ethylenediamine (en), oxalate (C₂O₄
²⁻).
Stability of Complex Compounds
NOT IMPORTANT
Dissociation types:
Differences
Primary Dissociation
Definition: The initial breaking of ionic bonds, often releasing simple
ions or molecules.
Examples: Dissociation of [Co(NH₃)₆]Cl₃ into [Co(NH₃)₆]³⁺ and Cl⁻
ions.
Contributing Factors: Solubility in water, ionic strength of the medium,
and temperature.
———————————-
Secondary Dissociation
Definition: The breaking of covalent/coordination bonds, leading to
further disintegration of the complex.
Examples: Decomposition of [Co(NH₃)₆]³⁺ into Co²⁺ and ammonia
molecules.
Contributing Factors: Acid-base reactions, redox reactions, thermal
stability
Acids, Bases, and Salts:
Acids: Substances that donate H⁺ ions; categorized into strong (e.g., HCl, HNO₃) and weak acids (e.g., CH₃COOH). Strong acids fully ionize in water, while weak acids partially ionize.
Bases: Substances that accept H⁺ ions or donate OH⁻ ions; include strong bases (e.g., NaOH) that fully dissociate and weak bases (e.g., NH₃) that partially dissociate.
Salts: Formed by the neutralization reaction between acids and bases, with types (neutral, acidic, basic) determined by the strength of the acid and base from which they
derive.
pH and pOH
pH measures acidity (pH < 7), neutrality (pH = 7), or alkalinity (pH > 7) of solutions.
Formulas for calculating pH, pOH, and dissociation constants (Ka, Kb, pKa, pKb) were introduced.
Brønsted-Lowry Theory:
Defines acids as proton donors and bases as proton acceptors. Discusses conjugate acid-base pairs formed during reactions.
pH effects on the Body:
NOT VERY IMPORTANT
Strong acids and bases can cause tissue damage. Conditions like acidosis (pH < 7.35) and alkalosis (pH > 7.45) disrupt physiological balance, with significant impacts on the nervous, cardiovascular, and respiratory systems.
Salt Hydrolysis
Explains how salts influence solution pH depending on the acid-base strengths of their components (neutral, acidic, or basic salts).
Colligative Properties:
Properties like
- vapor pressure lowering
-boiling point elevation
- freezing point depression
- osmotic pressure
depend on solute particle number, not identity.
Osmosis and Osmotic Pressure:
Describes solvent movement across membranes, driven by solute concentration differences.
Provides formulas for calculating osmotic pressure.
Solution Tonicity
Defines
- isotonic (equal solute concentration)
- hypertonic (higher solute concentration outside cells)
- hypotonic solutions (lower solute concentration outside cells)
and
their medical applications.
General Principles of Potentiometric Titration:
not very important
Potentiometric titration tracks the potential change of a solution during a chemical reaction, using electrodes to monitor the equivalence
point without the need for visual indicators.
Equivalence Point (Reaction End-Point):
The equivalence point is where the titrant and analyte react in exact stoichiometric proportions, marked by a sharp potential change on the titration curve
Buffer Region
The buffer region is the part of the curve where the solution resists pH or potential changes,
caused by the coexistence of a weak acid/base and its conjugate pair.
Half-Neutralization Point:
The half-neutralization point occurs when half the analyte is neutralized, and the pH (or potential) at this point corresponds to the pKa (or pKb) of the analyte.
Using Titration Graph to Recognize Electrolytes
The shape and features of the titration curve, such as initial pH, buffer regions, and equivalence point behavior, help identify whether the substances are strong/weak acids,
bases, or redox agents.
What is a buffer system?
A buffer system is a solution that resists changes in pH when:
small amounts of acid (H⁺ ions) or base (OH⁻ ions) are added.
The 4 main elements of a buffer
1. Weak Acid and its Conjugate Base (Salt):
- Example: Acetic acid (CH3COOH) and acetate (CH3COO−).
- These components react with added bases or acids to stabilize pH.
2. Weak Base and its Conjugate Acid (Salt):
- Example: Ammonia (NH3) and ammonium ion (NH4+).
- These components function similarly but in systems where bases are more prevalent.
3. Weak bivalent acid and its acidic salt (The weak acid can
donate two protons (H⁺)
- Example: H2CO3 (weak bivalent acid) and HCO3- (acidic salt/conjugate base)
- These components function similarly but in systems where
bases are more prevalent.
4. Two salts of the same polyvalent acid (differing in 1 hydrogen ion)
- The salt that contains greater number of H + acts as the acid in a buffer system
Principles of Maintaining pH in
Buffer Systems
𝐻𝐴 ⇔ 𝐻+ + 𝐴−
Le Chatelier’s Principle:
When an external change (like addition of H+ or OH−) disturbs the
equilibrium of the buffer system, the system adjusts by shifting the
equilibrium to minimize the change. (see explanation above)
DO YOU RECOGNISE??
Example
Buffer capacity (β)
3 components
The ability of a buffer solution to resist
changes in pH when small amounts of acid or base are added. It
measures the “strength” of a buffer in maintaining a stable pH. Unit is
mol/L.
* Concentration of buffer components: Higher concentrations of
the weak acid (HA) and its conjugate base (A−) result in a higher
buffer capacity.
*Ratio of acid to base: Buffer capacity is maximum when the
concentrations of the weak acid and its conjugate base are equal
(pH = pKa).
* Volume of the buffer: Larger volumes have greater buffer capacity
because they can neutralize more added acid/base without
significant pH changes.
Calculations buffer capacity
Effective buffer pH region
- range
Buffer system in human body
pH in human blood
Applications in medicine
