9-10

• What is a Titration?

  Titration: a solution of known concentration (the titrant) is added to a solution of the substance being studied (the analyte).
• What is the solution in the beaker and burrette

  • An analytical procedure in which a solute in a solution of known concentration reacts with a known stoichiometry with a substance whose concentration is to be determined.
  •   The solution in the beaker is the analyte and the solution in the burette is the titrant
• What are the types of Titration and explain

  • Acid/ base:

  Reacting analyte and titrant in which one is an acid and one is a base that exploits the acid/base neutralisation reaction

   

  • Precipitation

  A reaction of analyte and titrant  until Precipitation

   

  • Complexometric
  • A metal with a complexing agent as the titrant that forms a noticeable and specific colour changes at the end point
  • Redox
  •  analyte and titrant Exploits oxidation
• Titration requirements


  • We must know the equation for the reaction
  • The reaction must be rapid and complete
  • The point at which the reactants combine exactly is the equivalence point (the point at which the indicator changes colour is the end point)
  • We must be able to accurately measure the amount of each reactant
• pH indicator


  A substance that changes colour when a certain pH range has been achieved
• Indicators


  • They are typically weak acids
  • We use different indicators depending on the target pH
• Typical titration


  1. A solution of known concentration – the titrant

  2. A solution of unknown concentration – the analyte

  3. A means of measuring the end point – an indicator (or a pH meter)
• Acid-base titrations – Strong acids: Strong acid – strong base titration: • The initial pH is the pH of the strong acid. • The acidic region: • adding OH– to a solution of H3O+ concentration of H3O+ decreases. • The equivalence point: • point reaction stoichiometry is satisfied. • The alkaline region: • adding excess OH–, no chemical reaction occurring
• Weak Acid – Strong Base Titration:

  Initial pH:

  • Higher than strong acids.

  Acidic Region:

  • pH rises quickly as base is added.

  Equivalence Point:

  • Higher pH than with strong acids.
  • Beyond this point, pH is controlled by the strong base, similar to strong acid titration.
• ph scale: the range goes from 0 - 14, with 7 being neutral. pHs of less than 7 indicate acidity, whereas a pH of greater than 7 indicates a base.
• Basic Calculations • Volume of known solution required (mean titre volume, in L) n = c x V • Moles of known solution (in mol) • Moles of unknown solution (in mol) (x) by molar ratio • Concentration of unknown solution (in mol L-1 or M) c = n/V
• We must know the concentration of at least one solution, and the reaction equation of the solutions • We must have some way of measuring when a reaction is finished • Most important equations to remember: n=cV and molar ratios
• Lets remember:.... Arrhenius acid/base: acids produce H3O+ ions and bases produce OH in water • Brønsted–Lowry acid/base: acids are H+ donors and bases are H+ acceptors • Lewis acid/bases: acids are electron acceptors and bases are electron donors
• Properties of pure water

  Water can act as both an acid or a base • In pure water, water molecules can collide and form conjugate acids and bases during autoionization.
• What is the equilibrium position for this reaction where it is a strong base?

  • Stronger base means that the products going back to the reactants
  • To the left
  • Push away
• What is kw?


  -Kw is called the ion product of water (also called the water dissociation constant)

  • We don’t include pure water cause the concentration does not change, everything is dissolved in water.
  • SO cancel wate
• Properties of pure water

  • Kw is always 1.0 x 10-14 @ 25oC

  • What does this mean?

   

   

  Kw is always 1.0 x 10-14 @ 25oC. In pure water nothing else to interferes amount of hydroxide is going to equal  hydronium

   

   

   

  The equation for the ionization of water applies to any aqueous solution whether acidic, basic or neutral
• • Any of the concentration multiply by hydroxide and hydronium  is going to equal this


  (i.e. [H3O+ ] x [OH- ] = Kw = 10-14 @ 25oC). Kw is always 1.0 x 10^-14 @ 25oC
• One small caveat though

  Just remember: The auto-ionization of water is endothermic • Kw only = 1 x 10-14 at 25oC • If T increases, Kw also increases (not much, but still) • [H3O+ ]=[OH- ] should still hold true for water, regardless
• Logarithms


  if a number equals a base to a power, then the log of that number to the base equals the power”
• Logarithms
example:log10 100 = 2


  100 = 10^2
• ‘p’ means to take the negative log. Calculate the pH of a 0.010 M NaOH solution Similar process, but here use the relationship:

  pH + pOH = 14 ( remember this)
• Pure water will contain a small number of H3O+ and OHions • Kw is called the ion product of water • [H3O+ ] x [OH- ] = Kw • The pH scale ranges from 0 to 14 • How to calculate pH, pOH, [H3O+ ], [OH- ] • pH + pOH = 14
• Equilibrium & acid-base reactions

  Strong acids/bases completely dissociate (100%) – commonly mineral acids (HCl, HNO3 , HI, etc.) and alkali bases (NaOH, KOH etc.) • Weak acids/bases partially dissociate (5-10%) – commonly organic acids (CH3COOH, HCOOH, HCN etc.) and bases (NH3 , NH4OH), but also HF
• Equilibrium & acid-base reactions • Consider the following reaction:

  • HCl is a strong acid, it dissociates completely • The position of its equilibrium lies to the right • The reaction is written with a unidirectional arrow

  
• Equilibrium & acid-base reactions • Consider the following reaction:

  CH3COOH is a weak acid, it partially dissociates acid base C. base C. acid • The reaction is written with a reversible arrow • The position of its equilibrium lies to the left

  
• Predicting the equilibrium position:

  1. Identify the two acids in the equilibrium

   

  2. Determine which acid is the stronger acid and which acid is the weaker acid

   

  3. Identify the stronger base and the weaker base (Remember: the stronger the acid the weaker the conjugate base and vice versa)

   

   4. The stronger acid and stronger base react to give the weaker acid and weaker base
• Acid ionisation constants


  • Acids vary in the extent to which they dissociate • The dissociation (ionisation) of weak acids in water are all in equilibria
• Acid ionisation constants 


  • This equilibrium constant (Ka ), the acid dissociation (ionisation) constant, tells us quantitatively just how strong any weak acid is
• Relationship between pKa and Ka

  • pKa can be used (like pH) to indicate how strong/weak acid dissociation is
• whats the formula for pka

  pKa = -log Ka
• what does it mean when it's low pka?

  Low pKa (high Ka) means product formation is favoured, and vice versa
• pKa, Ka and pH


  pH and pKa are not the same • pH changes with concentration, pKa is always the same at room temperature for the same acid • BUT pH = pKa at half-equivalence point of your titration – this is most useful when we’re talking about acid-base buffering systems
• what do we express a weak acid

  The strength of a weak acid is expressed by Ka
• What is a conjugate acid-base pair


  Any pair of molecules or ions that can be interconverted by the transfer of a proton is called a conjugate acid-base pair
• The pH of Blood

  The pH of human blood is 7.4. The acceptable pH range is 7.35 to 7.45
• how to maintain ph?

  Buffer
• What is a buffer?

  • A buffer solution contains a weak acid and its conjugate base
• what happens if we add a strong acid?


  The addition of strong acid to the solution in eq 1 won’t cause a great change in the pH
• A buffer • Two substances are needed:

  An acid that can react with added OH-. A base that can react with added H3O+. The acid and base must not react with each other