Acids and bases

• Common acids


  • Sulfuric acid- H2SO4
  • Hydrochloric acid- HCl
  • Nitric acid- HNO3
  • Ethanoic acid- CH3COOH
• Acid


  Proton donor
• Acids


  • Release H+ ions (protons) when added to water
  • General formula for reaction of acids and water: HA + H2O <-> H3O+ + A-
• Common bases


  • Sodium hydroxide- NaOH
  • Ammonia- NH3
  • Copper (II) oxide- CuO
  • Potassium Carbonate- K2CO3
• Base


  Proton acceptor
• Bases


  • Metal hydroxides, ammonia, metal oxides, and metal carbonates are all bases (metal hydroxides and aqueous ammonia are alkalis)
  • General formula for reaction of bases and water: B + H2O <-> BH+ + OH
  • Bases accept protons from water
• Alkalis


  Soluble bases
• Strong acids and bases

  • Strong acids dissociate/ionise completely (lose all their H+ ions), whereas weak acids only partially dissociate/ionise
  • Strong bases fully ionise/dissociate (they lose all their 'OH'es), whereas weak bases only partially ionise
  • In full dissociation/ionisation, the concentration of the ion in mol dm^-3 will be the same as it was for the whole acid/base (high H+/OH- concentration)
  • In partial dissociation, the forward reaction does not happen as readily as the backwards reaction (low H+/OH- concentration, lower than the concentration of the overall acid/base)
• Reactions of acids and bases

  1. Acid (aq) + alkali (aq) -> salt (aq) + water (l)

  2. Aqueous ammonia (NH3) becomes NH4OH in water

  3. Ionic equation for acid alkali reaction: H+ (aq) + OH- (aq) -> H2O (l)
• Acids and metal oxides

  Acid + metal oxide -> salt + water
• Acids and metal carbonates

  Acid + metal carbonate -> salt + water + carbon dioxide
• Acids and metals


  Acid + metal -> salt + hydrogen
• Salt


  • A substance formed when the H+ of an acid is replaced by a metal ion or an ammonium ion
  • Salt name = first part of the name of the acid + 'ate'
• Solution


  Formed when a solute is dissolved in a solvent
• Concentration


  • The amount of solute in moles dissolved per 1 dm^3 of solution
  • Units are mol dm^-3
  • Concentration equation: C= mol/dm^3
  • Concentration= no. of moles (of solid)/volume (of liquid)
  • Volume must always be in decimetres cubed (ml=cm, 1dm = 1000cm, litres and dm^3 are actually the same thing)
• Strong solution


  Concentrated (high no. moles/dm^3)
• Weak solution


  Dilute (low no. moles/dm^3)
• Sometimes we are asked to express the concentration in g dm^-3
• In this situation we use the mass in grams of the solute rather than the no. of moles to calculate concentration
• When mass of solute is increased, mass of solution must also be increased for the solution to stay the same concentration
• Dilution


  Adding more water to a solution
• 1M


  1 molar = 1 mol dm^-3
• Standard solution


  We know its concentration
• The brown line on a volumetric flask caps the volume to 250 cm^3
• To do the dilution calculation

  1. Work out the number of moles originally in the amount of standard solution we diluted using n = c x v

  2. Work out the total volume of the solution we now have, using the no. of moles in the original solution to calculate concentration
• Always convert into the units that you are asked for
• To do a reacting masses calculation involving solutions (titration/neutralisation reaction)

  1. Identify the known and unknown substances

  2. Work out the moles of the known (c x v)

  3. Work out the moles of the unknown based on the ratio of acid to alkali in the equation

  4. Work out the concentration of the unknown using the volume and number of moles (n/v)
• If asked for in grams per decimetre cubed, we multiply our answer by the Mr of the substance, as our answer acts as the number of moles for if the substance was a solid and not dissolved
• With complex questions, visualise what's going on at each point to avoid getting lost, and write out the balanced equation if needed and not provided
• How to make a standard solution (known volume and concentration)

  1. Weigh out a known mass of solid (recording the mass you used if the mass was not perfect)

  2. Transfer the solid to a beaker, including the washings

  3. Dissolve the solid in a small amount of distilled water

  4. Stir with a stirring rod, which would then be rinsed so that the washings could be included

  5. Transfer the solution into a volumetric flask of the required volume (usually 250 cm^3) rinsing the beaker to include the washings

  6. Make the solution up to the mark, making sure that the base of the meniscus sits on the mark

  7. Put the stopper on the volumetric flask and invert/shake to ensure that the solute was fully dissolved

  8. Label the flask with the name of the solution inside, its concentration, and the date it was produced
• How to do an acid-base titration

  1. Titrations are used to determine volumes of acid and alkali needed to make a neutral salt solution

  2. Titrations are also used to determine concentration of an acid or alkali if the concentration of the other reactant is known
• How to carry out a titration

  1. Wash burette using the solution that is going into the burette (i.e: the acid) and then distilled water

  2. Fill burette to 100cm3 said solution with the meniscus' base on the 100cm3 line using a funnel, draining any excess solution out through the tap into a waste beaker

  3. Use 25cm3 pipette to add 25cm3 of alkali into a conical flask that is placed on a white tile, drawing alkali into the pipette using a pipette filler

  4. Add a few drops of a suitable indicator to the conical flask (eg: phenolphthalein which is pink when alkaline and colourless when acidic)

  5. Add acid from burette to alkali until end-point is reached (as shown by indicator)

  6. The titre (volume of acid needed to exactly neutralise the alkali) is the difference between the first (100cm3) and second readings on the burette)

  7. Repeat the experiment to gain more precise results
• Titration tips


  • Horizontal position of burette tap= closed
  • Have the zero point on the burette at eye level so that you can check that the meniscus is in the correct position
  • Never have the funnel still on/in the burette whilst the titration is occurring as excess solution can drip off it into the burette, altering your results
  • Make sure the jet (bit above the tap before the graduations begin) is full too by overfilling the burette, then draining out the solution to the correct point, thus removing any trapped air
  • The meniscus is a result of the water's surface tension/cohesion,causing water to stick to the sides of the glass and creating a ellipse
  • The resolution of the burette is to 0.01 (results go up in .1s), so results should be recorded to 2 d.p
  • As long as you know where the burette starts, it doesn't have to be perfectly on zero
  • If the meniscus is between two 0.1s, then round it to exactly halfway between (i.e: 1.15 if between 0.1 and 0.2) as the burette does not have a high enough resolution to show you conclusively the exact value, so the second decimal place is always a zero or a five
  • Hold the pipette at the top to avoid snapping it
• What to record in a titration

  • Initial volume in the burette in cm^3
  • Final volume of the burette in cm^3 (the end point)
  • Titre (difference between the two/amount used) in cm^3
  • Mean titre over the course of however many readings you take
• Titration procedure


  • The first reading is used as a trial to get a rough idea of the endpoint (as it's a trial, you shouldn't use this attempts results)
  • The titration is repeated as many times as necessary until you get two readings that are within 0.10 of each other (concordant results)
  • Always quote mean titre to 1 d.p
• Percentage error


  • All equipment used to measure out is not totally accurate
  • Usually stamped on the side of the equipment
  • To calculate percentage error, we do the amount of error divided by the measured amount, then timesed by 100
  • Always times the percentage error by the number of times the apparatus was used
  • When you use multiple apparatus (i.e: a burette and a pipette) you add the percentage error of each together to work out overall percentage error

  
• Theoretical yield


  The amount of product produced if there is 100% conversion
• Reasons most reactions don't reach the theoretical yield

  • The reaction wasn't given enough time to complete
  • Product may be lost in transfer between containers/filtering
  • Reaction may be reversible
• Actual yield


  How much product was actually produced
• Percentage yield


  Actual yield/theoretical yield x 100 (usually in moles, but can be done by mass)