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The stronger an acid is, the easier it can fall apart.

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Accurately measure the mass of the solid
Dissolve this solid into the minimum possible amount of water - no more than 150ml.
Transfer the solution into a volumetric flask using a funnel
Use a dropping pipette to add water until you reach the 250ml mark.

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Primary standards have the following properties:

High purity

Stability, i.e low reactivity

Weight changes due to humidity are minimised

High molar mass

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Acid standardisation allows us to find the concentration of an acid using a standard solution, which is prepared from a chemical that is a primary standard.

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H+ (aq) + OH- (aq) --> H₂O (l)
acid + base --> salt + water
HCl(aq) + NaOH(aq) --> NaCl(aq) + H₂O(l)

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An alkali is a solution of a base.
When we neutralise an acid, the H+ (aq) ions are removed from the solution. Alkaline solutions contain OH- (aq) ions, i.e hydroxide ions.
They neutralise acidity.
Bases are H+ (aq) acceptors.

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Hydrogen in acids enters solution H⁺ (aq) ions, i.e protons.
It is the H+ (aq) ions that make a solution acidic.
The releasing of a hydrogen atom into water is called dissociation.
Acids are H+ (aq) donors.

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pH = -1log ₁₀ [H+ (aq)]

[H+ (aq) ] = conc. of H+

[H⁺ ] = 10-pH

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HCl- Hydrochloric acid
H ₂ SO 4 - Sulfuric  Acid
HNO ₃ - Nitric Acid
CH ₃ COOH - Ethanoic Acid
NaOH - Sodium hydroxide

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When the products are favoured, Kc will have large values.
When the reactants are favoured, Kc will have small values.

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A catalyst will speed up the reaction by lowering the activation energy go the reaction. The catalyst will not affection the PoE, but equilibrium is reached faster.

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Generally: aA + bB ⇌ cC + dD
[X]eq : concentration of chemical X at equilibrium
Equilibrium constant, K c - [C]c x [D]d ÷ [A]a x [B]b

Example: H2(g) + I2(g) ⇌ 2HI(g)
Kc = [HI]² ÷ [H₂] x [I₂]

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At dynamic equilibrium the forward & reverse reactions happen at the same rate.
The position of equilibrium(PoE) can lie in favour of the reactants or products.

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If we increase the temperature, the PoE shifts to try to counteract the temperature change. But we need to know whether the forward/reverse reaction are exothermic(giving out heat) or endothermic(taking in heat).
In general:
increasing temp. favours endothermic reactions
decreasing temp. favours exothermic reactions
Example:
2NO₂(g) ⇌ N₂O₄(g)
-> exothermic
<- endothermic
ΔH = negative tells us forward reactions is exo
ΔH = positive tells us forward reaction is endo
ΔH = kJ mol⁻¹

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If we increase pressure, the equilibrium shifts to try to reduce it. Moving the PoE to the side with least gas molecules will lower the pressure.
If we decrease pressure, the equilibrium shifts to try to increase it again. Moving PoE to the side with most molecules will increase the pressure.
Example:
CO₂(g) ⇌ CO₂(aq) (coke bottle opening)
On opening the bottle, pressure decreases & PoE shifts left, the fizz means more gas

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If we increase the concentration of a chemical, the equilibrium will try to minimise the change by using up the chemical. This shift the position of the equilibrium ‘away’ from the chemical in the equation.
Example:
The Haber Process is an example, if the N₂ concentration is increased, the forward reaction taking place happens strongly. The position of equilibrium moves to the right & the amount of ammonia is increased.
If we decrease the concentration of a chemical, the equilibrium will try to minimise the change by making more of the chemical. This shifts the PoE 'towards' the chemical in the equation.

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The position of equilibrium shifts to try to cancel out any changes that you introduce.

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The CuSO₄.5H₂O is hydrated copper sulfate & the CuSo₄ in the products is anhydrous.

Example

CuSO4 .5H₂O ⇌  CuSO₄ + 5H₂O