9 Aqueous Solutions

  • 1,839 views
Uploaded on

 

More in: Business , Technology
  • Full Name Full Name Comment goes here.
    Are you sure you want to
    Your message goes here
    Be the first to comment
    Be the first to like this
No Downloads

Views

Total Views
1,839
On Slideshare
0
From Embeds
0
Number of Embeds
0

Actions

Shares
Downloads
32
Comments
0
Likes
0

Embeds 0

No embeds

Report content

Flagged as inappropriate Flag as inappropriate
Flag as inappropriate

Select your reason for flagging this presentation as inappropriate.

Cancel
    No notes for slide

Transcript

  • 1. 9 Aqueous Solutions
  • 2. Water as a Polar Solvent
    • The water molecule is polar. That is it has a slightly positive end and a slightly negative end.
     +  + H H  -  + O
  • 3.
    • Oxygen attracts electrons and becomes slightly negatively charged  the hydrogen atoms become slightly positively charged
    The shape of the molecule is important to its polarity. It must have a positive and negative end. H H  -  +  + O
  • 4.
    • When a polar substance (e.g. an ionic compound) is dissolved in water the ions are attracted to the water molecules
    Na + Cl - Na + Na + Cl - Cl - Cl - H H  -  +  + O
  • 5.
    • Water molecules separate, surround and disperse the ions into the liquid.
    • Some ionic compounds are only slightly (sparingly) soluble in water. e.g. silver chloride. This is because the electrostatic attractions within the compound are greater than the attraction between the ions and the water molecules
    • All ionic compounds will dissolve to a certain extent even though we call them insoluble
    • Solubility of NaCl in water at 20 o C = 365 g/L
    • Solubility of AgCl in water at 20 o C = 0.009 g/L
  • 6.
    • Ethanol – a polar molecule
    C C H H H H H H O  + Polar substances will dissolve in ethanol which has a polar and a non-polar part. Water and ethanol are therefore miscible. Note the ethanol and water do not dissociate into ions!  -  +  +  - H H - O
  • 7.
    • Acids
    • Acids are an important group of covalent compounds that dissolve in water.
    • e.g. HCl H 2 SO 4 HNO 3
    • These all interact so strongly with water that the molecules dissociate into ions
    • HCl  H + + Cl -
    • In fact we do not get a lone H +. The H +. Is attracted to the water molecule to give H 3 O +
  • 8.
    • So we should really write the equation as follows
    HCl + H 2 O  H 3 O + + Cl - H + .. H H  -  +  + O
  • 9. Writing Ionic equations
    • A molecular equation
    • AgNO 3(aq) + NaCl (aq)  AgCl (s) + NaNO 3(aq)‏
    • Total Ionic equation
    • Ag + (aq) + NO 3 - (aq + Na + (aq) + Cl - (aq)  AgCl (s) + NaNO 3(aq)‏
  • 10.
    • Sodium nitrate is soluble in water so no reaction occurs between sodium and the nitrate ions. The ions remain dissociated in solution as solvated ions. These ions are called spectator ions.
    • Silver chloride is insoluble so reaction occurs between silver ions and chloride ions
    •  the net ionic equation shows the actual chemical change taking place
    • Ag + (aq) + Cl - (aq)  AgCl (s)
  • 11. Solubility
    • See page 139 for solubility of compounds in water
    • In order to predict whether a precipitate occurs we need to know the solubility of that compound in water
  • 12.
    • Write ionic equations including state symbols for the following reactions
    • Barium nitrate with sodium carbonate
    • Sodium chloride with lead (II) nitrate
    • Aluminium nitrate with sodium phosphate
  • 13. Precipitation Reactions
    • In precipitation reactions two soluble compounds react to form an insoluble product, a precipitate. Precipitates form because the electrostatic attraction between the ions outweighs the tendency for the ions to become solvated (surrounded by the solvent molecules) and move randomly through the solution.
    • You can predict whether a precipitate occurs by considering the solubility of the products.
  • 14.
    • In the reaction between barium nitrate and sodium carbonate an insoluble precipitate of barium carbonate forms.
  • 15. Acid-Base Reactions
    • An acid is a substance that produces H + ions when dissolved in water
    • H 2 O
    • HX  H + (aq) + X - (aq)‏
    • A base is a substance that produces OH - ions when dissolved in water
    • H 2 O
    • MOH  M + (aq) + OH - (aq)‏
  • 16.
    • Strength of an acid or base
    • Strong acids or bases dissociate completely into ions when they dissolve in water
    • Weak acids or bases dissociate so little that most of their molecules remain intact
    • Strong acids and bases are therefore electrolytes (conduct a current)‏
    • Weak acids and bases are very weak electrolytes
  • 17.
    • For a strong acid or base dissociation into ions is virtually 100% so the conc.
    •  If we dissolve 1mol of solid sodium hydroxide in water we will get 1mol of OH - ions
    • NaOH  Na + + OH -
    • 1mol 1mol 1mol
    • This is not true of a weak acid or base
  • 18.
    • As noted earlier an acid produces H + ions when dissolved in water
    • So for example HCl is a covalent gas and does not behave as an acid. Only when it is dissolved in water does it dissociate into ions.
    • HCl (g) + H 2 O (l)  H 3 O + (aq) + Cl - (aq)‏
  • 19. Writing the net ionic equation for an acid base reaction
    • HCl (aq) + NaOH (aq)  NaCl (aq) + H 2 O (l)‏
    • NaCl will be dissociated but water will be undissociated
    • The net ionic equation is
    • H + (aq) + OH - (aq)  H 2 O(l)‏
  • 20.
    • Strong acids Most inorganic acids e.g.
    • sulphuric, nitric, hydrochloric
    • Strong bases sodium hydroxide, calcium hydroxide
    • Weak acids ALL organic acids e.g. acetic (ethanoic acid)
    • Weak bases ammonia
  • 21. Ammonia
    • Not all bases actually contain OH -. Ammonia NH 3 is a covalent gas that reacts with water to produce a basic solution. It is a weak base i.e. the following reaction does not go to completion but is in equilibrium. Both forward and backward reactions are occuring at the same rate .
    • NH 3 + H 2 O  NH 4 + + OH -
    • Ammonia removes a proton from water to leave OH - ion
  • 22.
    • Because it is a weak base and doesn’t dissociate fully in water the no of moles of OH - produced will be much less than the actual concentration of the ammonia solution.
    • This is also true of weak acids.
  • 23. The pH scale
    • pH is the negative log of H + concentration
    • pH = - log [H + ]
    •  for a 0.1molar solution of HCl (a strong acid)‏
    • pH = -log 0.1 = 1
  • 24.
    • What is the pH of a 0.01 mol solution of HCl?
    -log 0.01 = 2 What is the pH of a 0.1 molar solution of sulphuric acid? H 2 SO 4  2H + + SO 4 2-  [H + ] = 2 x 0.1 = 0.2 mol/l And pH = -log 0.2 = 0.7
  • 25.
    • What is the pH of the following?
    • 0.005 mol/l HNO 3
    • 0.05 mol/l H 3 PO 4
    • [H + ] = 0.005 mol/L
    • pH = -log 0.005
    •  pH = 2.3
    b) [H + ] = 3 x 0.05 mol/L pH = -log 0.15  pH = 0.82
  • 26. Finding pH from [H + ]
    • What is the [H + ] of a solution of HCl with a pH of 3?
    Antilog -3 = 0.001  [H] = 0.001 What is the [HCl]?
    • From the equation HCl  H + + Cl - we can see that the ratio of [H + ] to [HCl] = 1:1
    • [HCl] = 0.001 mol/L
  • 27.
    • What is the [H + ] of a solution of H 2 SO 4 with a pH of 2.4?
    Antilog -2.4 = 0.004 mol/L  [H+] = 0.004 mol/L H 2 SO 4  2H+ + SO 4 2-  [H 2 SO 4 ] = 0.004/2 = 0.002 mol/L What is the [H 2 SO 4 ]?
  • 28. Finding pH from [OH - ]
    • pOH = -log[OH - ]
    • and
    • pH +pOH = 14
    What is the pH of a 0.1 mol/L solution of sodium hydroxide? pOH = -log 0.1 = 1  pH = 14 – 1 = 13
  • 29.
    • Find pH of the following solutions
    • 0.001 mol/l potassium hydroxide
    • 0.05 mol/L calcium hydroxide
    a) pOH = -log 0.001 = 3  pH = 14 -3 = 11 b) [OH - ] = 2 x 0.05  pOH = -log 0.1 = 1  pH = 14 -1 = 13
  • 30.
    • Calculate the pH of of a solution made by dissolving 1.00g of calcium hydroxide in 500ml of water.
    RMM Ca(OH) 2 = 40 + (2 x 16) + (2 x 1) = 74 1.00g = 1.00 = 0.0135mol 74 Moles [OH - ] = 2 x 0.0135 = 0.027 Moles per litre = 2 x 0.027 = 0.054 pOH = -log 0.054 = 1.27 pH = 14 – 1.27 = 12.3