1.7 Equilibria + acid Flashcards
Reversible reaction?
one that can go in either direction depending on the conditions
Dynamic equilibrium?
when the forward and reverse reactions occur at the same rate
Go to completion?
where the reactants change completely to form products
Water in CuSO4.5H2O?
water of crystallisation
When Copper sulfate is heated. Water of crystallisation is given off as steam, leaving a white powder called anhydrous Copper (II)
Process is represented by?
CuSO4.5H2O —- CuSO4+5H2O
Haber process?
formation of ammonia from nitrogen and hydrogen is known as a indrustrial process e.g
N2(g) + 3H2(g)—– 2NH3(g)
Equilibrium?
term to denote balance
2 types everyday = static + dynamic equilibrium
Example of dynamic equilibrium?
dissolving an ionic compound in water
when Copper sulphate crystals = added to water, crystals begin to dissolve and solution turns blue
more Copper sulphate added, deeper the blue colour
when no more dissolves and copper sulphate crystals remain in the solution, solution = saturated and the intensity of the blue colour remains constant
This point, the solution = in equilibrium with the undissolved solid
The concentration of the saturated solution remains the same, the copper (II) sulphate is still dissolving but as it does so Copper (II) sulphate is recrystalising from solution at the same rate
Features of equilibrium?
closed system
dynamic at a molecular level
forward and reverse reactions occur at the same rate
macroscopic properties remain constant
Position of equilibrium def?
proportion of products to reactants n an equilibrium mixture
Le Chatelier’s principle?
states that if a system of equilibrium is subjected to a change the equilibrium tends to shift to minimise the effect of the chamge
Effect of concentration change?
dissolving hydrated copper (II) sulphate crystals in water gives a blue solution because the ion (Cu(H2O)6)2+ is formed
when concentrated HCl is added to a solution of Copper(II) sulphate, the equilibrium is established
(Cu(H2O)6)2+ (aq) + 4Cl - (aq) —– (CuCl4)2-(aq) + 6H2O(l)
Adding moer Hcl adds chloride ions so system tries to minimise the effect by decreasing the concentration of chloride ions so position of equilibrium moves to the right, forming more CuCl42- ions, making the solution yellow - green
Effect of pressure change?
pressure virtually has no effect on chemistry of solids and liquids
Pressure of gas depends on number of molecules ina given volume of gas
Greater the number of molecules, the greater the number of collisions per unit time, therefore the greater the pressure of the gas
E.g 2NO2(g)—–N2O4(g)
2 moles on LHS
and 1 mole on RHS
if the total pressure is increased, the equilibrium will shift to minimise the increase
Pressure will decrease if the equilibrium system contains fewer gas molecules
Position of equilibrium moves to the right which increases the yield of N2O4 and colour becomes lighter
Effect of temperature change?
If the delta H value is negative, its exothermic
If delta change is positive, the reaction is endothermic
for example
2NO2(g) —- N2O4(g)
enthalpy = -24
so if negative, forward reaction is exothermic and if the temp is increased, the system will try to minimise this increase
system opposes the change by taking in heat so the position of equilibrium moves in the endothermic direction, so the equilibrium moves to the left
decreasing the yield of NO2 - making the mixture appear brown
Decreasing the temp?
shifts the equilibrium to the right, favouring the exothermic direction, increasing the yield of N2O4 so the mixture appears a lighter colour
Effect of temp change?
An endothermic reaction absorbs heat from the surroundings whereas an exothermic reaction gives out heat to the surroundings
e.g for a reversible reaction the forward reaction is exothermic and the backward reaction is endothermic
Enthalpy change of forward reaction = the same magnitude as but opposite sign to backward reaction
again
2NO2(g) —- N2O4(g)
since enthalpy change = negative forward reaction = exothermic
if temp = increased, the system will try and minimise this increase
System opposes the change by taking in heat so position of equilibrium moves in the endothermic reaction.
Therefore, equilibrium moves to the left, decreasing the yield of N2O4 and increasing the yield of NO2 making mixture brown
Decreasing the temp?
shifts the equilibrium to the right, favouring the endothermic direction, increasing the yield of N2O4 so mixture appears a much lighter colour
Effects of catalysts?
catalysts speed up a chemical reaction by lowering the activation energy of the reaction
In a reversible reaction, a catalyst will increase the rate of the forward and backward reaction but to the same extent. Therefore, a catalyst does not affect the position of equilibrium but equilibrium is reached faster
Equilibrium constant?
equilibrium position chnges when the temp, pressure and concentrations change
equilibrium position may be described by combining the equilibrium concentrations to give a value for an equilibrium constant
given by symbol Kc
General for equilibrium?
aA+bB—–cC+dD
Equation for Kc?
(C)^c (D)^d /
(A)^a (B)^b
products = top line
reactants = bottom line
concentrations = raised to powers
unit of Kc can depend on equilibrium
Value greater than Kc?
more products than reactants in equilibrium mixture
Position of equilibrium lies to the right
greater the value of Kc , the further the equilibrium lies to the right
Value less than 1 for kC?
more reactants than products in the equilibrium mixture i.e. position of equilibrium lies to the left
the smaller the value of Kc, equilibrium lies to the left
Acid?
proton donor
Base?
proton acceptor
Common acids?
HCl
H2SO4
HNO3
CH3COOH
Equation for HCl disassociating?
HCl(g)—-H+(aq) + Cl-(aq)
Common bases?
MgO
NaOH
CaO
NH3
If base dissolves in water?
called alkali
ion common to all alkalis = OH-
Strong acid?
one that fully disassociates in aqueous solution
Weak acid?
one that partially dissolves in aqueous solution
pH scale measures?
the acidity of a solution which is a measure of the conc of the aqueous hydrogen ion
Ph def?
- log10 (H+)
H+ = concentration of H+ in mol dm-3
negative sign in equation results in pH decreasing as aqueous Hydrogen ion concentration icnreases
Salt?
compound that forms when a metal ion replaces the hydrogen ion in an acid
Neutralisation def ?
where acids + bases react
E.G?
HCl(aq)+NAOH(aq) —- NACl(aq) + H2O
H+(aq)+Cl-(aq)+OH-(aq) —- Na+ + Cl- + H2O(l)
H+(aq) + OH-(aq) —- H2O(l)
if MGO reacts with H2SO4?
MGO(s) + H2SO4(aq) — MGSO4(aq) + H2O(l)
If PbCO3 reacts with 2HNO3(aq) form?
PbCO3(s)+2HNO3(aq) —– Pb(NO3)2(aq) + H2O(l) + CO2(g)
Neutralisation always produces?
a salt
Standard solution?
one for which concentration is accurately known
How to prepare a standard solution?
prepared using a primary standard
primary standard?
typically a reagent which can be weighed easily and which is so pure that its weight is truly representative of the number of moles of substance contained
Features of a primary standard?
high purity
stability ( low reactivity)
low hygroscopicity
(to minimise the weight changes due to humidity)
high molar mass (to minimise weighing areas)
Why can’t sodium hydroxide be used as a primary standard?
it reacts with atmospheric carbon dioxide
2 examples of Primary standards?
Potassium hydrogen phthalate
(usually called KHP for standardisation of aqueous base solutions)
Sodium carbonate for standardisation of aqueous acids
Standard solution is prepared as?
calculate the solid required and accurately weigh this amount into a weighing bottle
transfer all solid into a beaker. wash out the weighing bottle so that all the weighings run into the beaker.
Add water and stir until all the solid dissolves
Pour all the solution carefully through a funnel into a volumetric flask, washing all the solution off the beaker and the glass rod. Add water until just below graduation mark
Add water drop by drop until the graduation mark is reached and mix solution thoroughly
Performing a titration?
what do they involve?
a burette containing 1 solution e.g acid
a conical flask containing the other solution (e.g a base)
A pipette to accurately transfer the solution to the conical flask
an indicator to show when reaction = completed
2 common indicators?
phenolphthalein
(colourless in acid solution and purple in alkaline solution) and methyl orange which is red in acid solution and yellow in alkaline solution)
All titrations follow the same method which is?
pour the acid into a burette, using a funnel, making sure the jet is filled. Remove the funnel and read the burette
use a pipette to add a measured volume into a conical flask
add a few drops of indicator to the solution in the flask
Run the acid from the burette into a solution in the conical flask, swirling the flask
stop when the indicator just changes colour ( this is the end point of the titration)
read the burette again and subtract the volume of the acid used
Repeat the titration again, making sure the acid is added drop by drop near the end-point, until there are at least 2 readings that are within 0.020cm3 of each other and calculate a mean titre
Double titrations?
Since different indicators change colour at different pH values, if a solution contains a mixture of 2 bases which are of different strengths, 1 titration can be performed but in 2 stages, using 2 different indicators, 1 added at each stage, to calculate the concentrations of both bases
Example?
concentrations of sodium hydroxide and sodium carbonate in a mixture can be determined by titrating with HCl using phenolphthalein and methyl orange as indicators
Phenolphthalein changes colour at around pH9
methyl orange changes colour at around pH 4
Sodium oxide is neutralised by acid according to
OH-(aq)+H+(aq)—H2O(l)
sodium carbonate = neutralised by acid 2 reactions occur
CO32-(aq) + H+(aq) —- HCO3-(aq)
HCO3-(aq) + H+(aq) —- H2O(l) + CO2(g)
At pH 9, all OH ions have been neutralised and the carbonate ions have been converted to hydrogen carbonate ions.
At pH4, the hydrogen carbonate ions are converted to water and carbon dioxide
OH- + H+ —- H2O
XO23- + H+ — HCO3-
HCO3-+H+—H2O+CO2
what does the first stage of titration relate to?
relates to the concentration of the hydroxide and carbonate
what does the second stage of titration relate to?
relates to the concentration of the carbonate only
when do we use back titrations?
sometimes it not possible to use standard titration methods
e.g the reaction between a certain substance and titrant can be too slow, there can be a problem with the end point determination or base = insoluble salt
in such we use a back titration
How does it work?
known excess of one reagent reacts with a unknown amount of reagent B at the end of the reaction. The amount of reagent A that remains is found by titration
