Equilibrium Flashcards
(124 cards)
1
Q
Equilibrium state
A
An equilibrium state is a system condition where there is no change over time, and any change in one direction is balanced by a change in the opposite direction
2
Q
Dynamic Equilibrium
A
Dynamic equilibrium is a state where opposing forces or reactions reach a balance, resulting in stability over time.
3
Q
Equilibrium Mixture
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Mixture of reactants and products in the equilibrium state
4
Q
Ionic Equilibrium
A
Ionic equilibrium is the state of balance between ions and unionized molecules in a solution of weak electrolytes
5
Q
Equilibrium for solid-liquid
A
- Rate of melting = Rate of freezing
- Both processes occur simultaneously
- Both processes occur at the same rate
6
Q
Normal Freezing Point / Normal Melting Point
A
Temperature at which the solid and liquid phases are at equilibrium at atmospheric pressure
7
Q
Liquid Vapour Equilibrium
A
Rate of evaporation = Rate of condensation
* IMPORTANT: NEEDS TO HAPPEN IN A CLOSED VESSEL OTHERWISE REVERSE PROCESS WILL NOT HAPPEN
8
Q
Equilibrium vapour pressure / Vapour pressure of liquid
A
Pressure exerted by vapours in equilibrium with the liquid at a particular temperature
9
Q
Normal Boiling Point
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Temperature at which the liquid and vapours are at equilibrium
10
Q
Relation between boiling point and atmospheric pressure
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Inversely proportional
At high altitude = BP decreases
11
Q
Solid Vapour equilibrium
A
Can be explained by iodine
After sometime vessel gets filled up with violet vapour and intensity of colour increases with time
12
Q
Solids in Liquids
A
Sugar solution
13
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Saturated Solution
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When no more of a solute can be dissolved in it at a given temperature
14
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Gases in liquids
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A soda water bottle opened (giving CO2)
15
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State Henry’s Law
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Mass of a gas dissolved in a given mass of a solvent at any temperature is proportional to the pressure of the gas above the solvent
16
Q
Tell the processes and conclusions attained:
1. Liquid to vapour (H2O liq. to H2O gas)
2. Solid to liquid (H2O solid to H2O liq)
3. Solute (s) to Solute (solution)
4. Gas (g) to Gas (aq)
A
- pH2O constant at given temperature
- Melting point is fixed at constant pressure
- Concentration of solute in a solution is constant at a given temperature
- Gas (aq) / Gas (g) is constant at a given temp
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Q
Equilibrium equation
A
kc (equilibrium constant) = [C][D] / [A][B]
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Equilibrium equation for general equation
A
aA + bB –><– cC + dD
Kc = [C]^c [D]^d / [A]^a [B]^b
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Q
Equilibrium constant unit
A
mol L^-1
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Q
Equilibrium Law / Law of Chemical Equilibrium
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Products of concentrations of reaction products raised to the respective stoichiometric coefficient in the balanced chemical equation divided by the product of concentrations of the reactants raised to their individual stoichiometric coefficients has a constant value
21
Q
Homogenous Equilibrium
A
All reactants and products are in the same phase
22
Q
Relation betwen pressure and conc. at constant temperature
A
Proportional
p = [gas]RT
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Q
Difference between Kc and Kp
A
The main difference between Kc and Kp is that Kc is the equilibrium constant expressed in terms of concentration, while Kp is expressed in terms of pressure
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Q
General equation with Kp
A
For aA + bB -><- cC + dD
Kp = Kc (RT)^delta n
Where delta n = no. of moles of gaseous products - no. of moles of gaseous reactants
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Bar into Pa
1 bar = 10^5 Pa
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Heterogeneous Equilibrium
System having more than one phase
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Concentration of pure liquid or solid
Taken as constant (independent of amount present)
Gaseous and aq. will vary
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Important features of eq. constant
1. Expression for eq. const. applicable only when conc. of react. and prod. attained constant value
2. Independent of initial concentrations of react and prod
3. Temperature dependent
4. K of forward = 1 / K of backward
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What happens when Kc > 10^3
* Products predominate
* Reaction proceeds nearly to completion
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If Kc < 10^-3
* Reactants predominate over products
* Reaction rarely proceeds
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Kc between 10^-3 and 10^3
Appreciable concentrations of both reactants and products present
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Reaction quotient (Q)
The reaction quotient, or Q, is a measurement of the relative amounts of products and reactants in a chemical reaction at a specific time
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What happens when:
1) Qc > Kc
2) Qc < Kc
3) Qc = Kc
1. Reaction will proceed in direction of reactants (reverse reaction)
2. Will proceed in direction of products (forward reaction)
3. At equilibrium
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k value in terms of g
k = e ^ (-G eq / RT)
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What happens when G < 0
-G/RT > 0;
K > 1;
spontaneous reaction;
proceeds in a forward direction
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What happens when G > 0
-G/RT < 0;
K < 1;
non - spontaneous reaction;
proceeds in a backward direction
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Le Chatelier's Principle
Change in any of the factors that determine the equilibrium conditions of a system will cause the system to change in such a manner so as to reduce or to counteract the effect of the change
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G value in terms of k
G = -RT ln Kc
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Concentration change (Le Chaterlier's Principle)
When concentration of any of the reactants or products in a reaction at equilibrium is changed, the composition of the equilibrium mixture changes so as to minimize the effect of concentration changes
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Increase in concentration of reactant
* According to Le Chatelier's principle, such a change in the sytem will have to counteract the change
* As concentration of reactants is higher, Qc is now much lower
* Qc < Kc means the reaction will move in the forward direction
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Increase in concentration of product
* According to Le Chatelier's principle, such a change in the sytem will have to counteract the change
* As concentration of products is higher, Qc is now much higher
* Qc > Kc means the reaction will move in the backward direction
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Decrease in concentration of product (Removal)
* According to Le Chatelier's principle, such a change in the sytem will have to counteract the change
* As concentration of reactants is higher, Qc is now much lower
* Qc < Kc means the reaction will move in the forward direction
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Decrease in concentration of reactant (Removal)
* According to Le Chatelier's principle, such a change in the sytem will have to counteract the change
* As concentration of products is higher, Qc is now much higher
* Qc > Kc means the reaction will move in the backward direction
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Effect of pressure change:
N2(g) + 3H2(g) --><-- 2NH3(g)
* An increase in pressure would cause a decrease in volume (P and V inversely related)
* Hence, there will be a large number of moles of gaseous substance per unit volume
* Hence, this can be undone by moving in a direction where pressure decreases or produces lesser number of moles per unit volume {Also moles and P are proportional if you want to write in short}
* Here, the number of moles in products (2 < 4) hence, equilibrium will shift towards the direction of products -- aka the forward direction
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Pressure change (le chatelier) on solids and liquids
Can be ignored because the volume of a solid / liquid is nearly independent of pressure
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Effect of pressure change
(C + CO2(g) --> 2CO(g))
When pressure is increased, the reaction goes in the reverse direction because the number of moles of gas increases in the forward direction
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Pressure change for no change in number of molecules
Pressure has no impact on equilibrium
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Decrease in pressure
Causes a shift in the direction of higher number of moles
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Explain using Kc and Qc for increase in pressure of:
CO(g) + 3H2 (g) --><-- CH4 (g) + H2O (g)
When pressure is doubled:
Qc = (2[CH4])(2[H2O])/(2[CO])(2[H2])^3
= Kc / 4
Since Qc < Kc: reaction will move in forward direction
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# ****
Effect of change of temperature on QC and Kc
For temperature, Kc value changes and not the Qc value
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Equilibrium constant with temperature
* According to Le - Chatelier's Principle, if the temperature of the system at equilibrium is increased, the equilibrium will shift in the direction in which the added heat is absorbed.
* Eq. constant for exothermic reaction (- delta H) decreases as temp increases
* Eq. constant for endothermic reaction (+delta H) increases as temp increases
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Using example of N2 + 3H2 --><-- 2NH3 explain effect of change of temperature
* We know the forward reaction is exothermic while the backward reaction is endothermic
* Hence, when heat is added to the system, it would change in the direction of the backward reaction
* The backward reaction being endothermic, reaction will be favoured by increasing temperature as it tends to undo the effect of added heat
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Eq. constant with decrease in temp
* A decrease in temp for exothermic reaction favours forward reaction
* A decrease in temp for endothermic reaction favours backward reaction
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Change of solubility with temp
* Solubility increases with increase in temp (endothermic reaction)
* Solubility decreases with increase in temp (exothermic reaction)
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Effect of volume change on equilibrium
PV = constant
Increase in Volume = Decrease in P
hence it will be the exact opposite
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What does a catalyst do?
* Increases the rate of chemical reaction by making available a new low energy pathway for the conversion of reactants to products
* Increases the rate of forward and backward direction that pass through the same transition rate and does not affect equilibrium
* Lowers the activation energy of forward and backward reaction by the same amount
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When is a catalyst of no use?
When K is an exceedingly small value
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Effect of addition of inert gas (constant pressure)
At constant pressure, the volume increases
Opposite of increase in pressure and favours the direction in which there is an increase in the number of moles of gases
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Effect of addition of an inert gas at constant volume
At constant volume, total pressure will increase
But, no change in concentrations of reactants and products --> hence no effect on equilibrium
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Conductance of electricity increases with increase in
concentration of common salt
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Electrolytes
An electrolyte is a substance that conducts electricity through the movement of ions, but not through the movement of electrons.
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Strong electrolytes vs Weak Electrolytes
Strong: On dissolution with water are ionized almost completely
Weak: Partially dissolved
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Why does NaCl only conduct in aqueous state?
* NaCl is held together with strong electrostatic forces of attraction
* When dissolved in water, a dielectric constant of water cuts down forces of attraction between the ions
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Ionic Equilibrium
Equilibrium established between the unionized molecules and the ions in solution of weak electrolytes
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Dissociation
Process of separation of ions already present in the solid state of solute when dissolved in water or a solvent
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Ionization
Process in which a neutral molecule which does not contain ions splits into charged ions when dissolved in water or a solvent
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Arrhenius concepts of bases
Substances that dissociates in aqeous solutions to give OH- ions
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Arrhenius concept of acids
Substances that dissociate in water to give H+ ions
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Relation between electrostatic forces and dielectric constant
Inversely proportional
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Strong acids vs Weak acids
Strong: Almost completely ionized in aqueous solution
Weak: Weakly ionized in aqueous solution
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Strong base vs Weak Base
Strong base: Almost completely ionized in aqueous solution
Weak base: Slightly ionized in aqueous solution
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Explain Hydronium ions
* A bare proton of H+ is very reactive and cannot exist freely in aqueous solutions
* Bonds to the oxygen atom of a solvent water molecule to give trigonal pyramidal hydronium ion
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Bronsted - Lowry Theory
* Acid is a substance that is capable of donating a hydrogen ion H+ and bases are capable of accepting a hydrogen ion H+
* Acids are proton donors
* Bases are proton acceptors
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Amphoteric
Substances that behave both as an acid and a base
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Conjugate Acid - Base
* A base forms a conjugate acid (as it gives it an extra proton)
* An acid forms a conjugate base (as it removes an extra proton)
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Relative strengths of conjugate acid base
Acid depends upon the tendency to donate a proton and strength of a base depends upon its tendency to accept a proton
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Lewis acids and Lewis bases
Lewis Acid = Substance which can accept a pair of electrons
Lewis Base = Substance which can donate a pair of electrons
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Lewis Acids
* Molecules in which central atom has an incomplete octet
* Simple cations
* Molecules with central atom having an empty d - orbital
* Molecules in which multiple atoms of dissimilar electronegativities are joined by multiple bonds
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Relation between base and conjugate acid
Strong acid = Weak conjugate base
Strong base = Weak conjugate acid
Weak acid = Strong conjugate base
Weak base = Strong conjugate acid
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Lewis Bases
* Neutral species having at least one lone pair of electrons
* Negatively charged species of anions
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Ionic Product of Water
Dissociation constant is represented by the product of [OH-] * [H+] ions
Represented as Kw
Found to be 10^-14 M^2
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pH
The negative logarithm of the H3O+ ion concentration in moles per litre
pH = -log aH+ = - log {[H+] / mol L^-1}
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Degree of Dissociation
Fraction of the total number of molecules of an electrolyte which ionizes into ions
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Acid dissociation constant (Ka)
An acid dissociation constant is a quantitative measure of the strength of an acid in solution
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High value of Ka vs Low value of Ka
Higher value = Strong acid
As it dissociates more
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Ka value for weak acids
Ka = c (alpha)^2
/
(1 - alpha)
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NH3 conjugate acid and conjugate base
Conjugate acid: NH4 +
Conjugate base: NH2-
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Base ionization constant
Equilibrium constant for base ionization (Kb)
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[H2O]Keq value
10^-14 at 298K
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Acid Base Neutral
Neutral: [H+] = [OH-] = 10 ^ - 7
Acid: [H+] > [OH-] (> 10^-7)
Base: [H+] < [OH-] (< 10^-7)
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Ionic product of water
The autoionization of liquid water produces OH− and H3O+ ions. The equilibrium constant for this reaction is called the ion-product constant of liquid water (Kw) and is defined as Kw=[H3O+][OH−]. At 25 °C, Kw is 1.01×10−14; hence pH+pOH=pKw=14.00.
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How to do pH questions for strong acids or strong bases
* They dissociate completely so final concentration for the acid is always 0
* alpha = 1
* Then just using molarity calculate the [H+] ions conc.
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pH +pOH =
14
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pH for acids with smaller concentrations
Consider the dissocation of the acid and the self ionization of water
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For pH of acids with smaller conc. why do we take 'x' while calculating ionization of water
* Water dissociates less as we add more H+ from the acids due to Le Chatelier's Principle
* The concentration of OH- is x
* The concentration of H+ is 10^-a + x because we have to consider both ways
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Relation between Ka and Kb
[Ka] [Kb] = Kw
pKa + pKb = pKw
-logKa + -logKb = -logKw
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Polybasic / Polyprotic acids
Bronsted acids which can donate more than one proton
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Why is Ka1 > Ka2
* It is more difficult to remove a positively charged proton from a negative ion due to electrostatic forces
* In Ka2 a proton is removed from a negatively charged species; which has / undergoes higher electrostatic forces
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Ka in terms of Ka1 and Ka2
Ka = Ka1 * Ka2
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Relation between acidity in the same group
* In the same group compare the bond strength
* When the bond strength decreases, energy required to break the bond decreases, HA becomes a stronger bond
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Relation between acidity in the same period
* In the same period compare the electronegativity difference
* When the electronegativity difference increases, the polarity increases and hence the strength of acid increases
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Hydrolysis
Interaction of anion or cation of the salt with water to produce an acidic or basic solution
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Why is hydrolysis opposite of neutralization
In Neutralization reaction acid and base react to form water and salt and involves the combination of hydrogen ions and hydroxyl ions while in hydrolysis chemical decomposition takes place of any salt with water giving rise to acid and base.
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Buffer Solution
The solution which resists changes in the hydrogen ion concentration on the addition of small amount of acid or base
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Acidic Buffer
Contains equimolar quantities of a weak acid and its salt with strong base
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Basic Buffer
Containd equimolar quantities of a weak base and its salt with a strong acid
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pH of a buffer solution (acidic buffer)
[H+] = x = c1 / c2 * Ka
c1 = conc. of acid
c2 = conc. of salt
pH = pKa + log10 (salt / acid)
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pH of buffer solution (basic buffer)
pOH = pKb + log(salt / base)
pH = 14 - pOH
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Neutral Buffer
A neutral buffer is a buffer solution that is created by mixing a weak acid and a weak base
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Range of buffer
A buffer has an effective pH range of one pH unit on either side of the pKₐ value for the weak acid. If the pH of a buffer goes out of this range, the buffer will no longer be effective at resisting large changes in pH.
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Equilibrium constant for hydrolysis constant
Kh = Kw / Ka
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How to know the pH when a solution dissolves in water
Strong base weak acid --> pH > 7
Strong acid weak base --> pH < 7
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pH of strong base weak acid
pH = 7 + 1/2 (pKa + log c)
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pH of strong acid weak base
pH = 7 - 1/2 (log c + pKb)
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pH of weak acid weak base
pH = 7 + 1/2 (pKa - pKb)
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pH of strong acid strong base
* Does not undergo hydrolysis
* Salts formed by the neutralisation of strong acid and strong base are neutral in nature as the bonds in the salt solution will not break apart
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Relation between the solubility and lattice enthalpy
* Inversely related
* High lattice enthalpy = Lower Solubility
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Solvation Enthalpy
Enthalpy of solution is the enthalpy change associated with the dissolution of a substance in a solvent at constant pressure resulting in infinite dilution.
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Sparingly Soluble
Materials which have low solubility - or ability to dissolve in a solvent
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Solubility Product
Solubility product refers to the equilibrium constant used to predict the formation and dissolution of precipitates based on factors like temperature, pH, and concentrations of reactants
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When does precipitation occur
When ionic product > Solubility product
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