Ch 10 Acids & Bases

Acids & Bases Chapter 10

April 2005

Examples of some uses of acids  In    

 In   

food e.g. Citric acid & ascorbic acid – health (anti-oxidants) Ethanoic acid – vinegar for taste & preservative Carbonic acid – fizzy drinks Lactic acid – in milk

industry e.g. Hydrochloric acid – cleaning, etching Sulphuric acid – batteries, fertilisers, detergents… Nitric acid – fertilisers, cleaning

Defining an acid Some common acids  HCl hydrochloric acid  HNO nitric acid 3 

H2SO4

Ions produced in solution H+ Cl− H+ NO3−

sulphuric acid

2H+

SO42−

2H+

CO32−

Less common acids (weak ones)  H CO carbonic acid 2 3 

CH3COOH

ethanoic acid

+ − H CH COO 3 An acid is a substance that produces H ions in solution +

Pure acids Pure acids are small covalent molecules, e.g. HCl, H2SO4

H

Cl

H2O

+

H

+

Cl

−

Acids behave as acids only when they are dissolved in water – formation of H+ ions in aqueous solutions

Strong vs. weak acids

split up

Strong acids completely dissociate in water, e.g. HCl, HNO3, H2SO4

H

Cl

100%

H2O

+

H

+

Cl

−

100%

*Nearly all acid molecules dissociate to form ions.

Strong vs. weak acids Weak acids only partially dissociate in water, e.g. carbonic acid H2CO3 & vinegar CH3COOH CH3COOH

100%

→

CH3COO− +

H+

0.5%

* In 1000 molecules of vinegar, only about 4 molecules dissociate into ions. The rest 996 molecules remain as molecules! * This is the reason why we can consume them!

Concentrated = Strong?  E.g.

consider concentrated ethanoic acid (vinegar)

 Concentrated

means lots of acid molecules but little water → few H+ ions

 So 

concentrated ≠ strong E.g. concentrated vinegar

 And 

strong ≠ concentrated

E.g. dilute HCl

Questions: Which conducts electricty better? Concentrated or dilute sulphuric acid? Why? Citric acid, a white solid can be dissolved in both organic solvents as well as inorganic solvents. Why?

Reactions of acids (usually these are used as tests for acids)  Acids

+ metals  Acids + carbonates  Acids + bases hydroxides

metal oxides

1st Acid + metal hydrochloric acid + magnesium → magnesium chloride + hydrogen gas

Acid + metal → salt + H2 2HCl(aq) + Mg(s) → MgCl2(aq) + H2(g) • Effervescence of a gas is observed. • Gas gives a ‘pop’ sound with a lighted/burning splint

1st Acid + metal (cont’d) Only reactive metals react, such as those from Group I, II & III metals and some transition metals. Copper, silver and gold are examples of unreactive metals.

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2nd Acid + carbonate hydrochloric acid + magnesium carbonate → magnesium chloride + carbon dioxide gas + water

Acid + carbonate → salt + CO2 + water 2HCl(aq) + MgCO3(s) → MgCl2(aq) + CO2(g) + H2O(l) • Effervescence of a gas is observed. • Gas gives a white precipitate with limewater.

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3 Acid + base rd

Metal hydroxide Metal oxide

hydrochloric acid + magnesium hydroxide → magnesium chloride + water hydrochloric acid + magnesium oxide → magnesium chloride + water

Acid + base → salt + water 2HCl(aq) + Mg(OH)2(aq) → MgCl2(aq) + 2H2O(l) 2HCl(aq) + MgO(s) → MgCl2(aq) + H2O(l)

Neutralisation reaction

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Metal hydroxide Metal oxide

Bases

Bases

Metal oxide

Metal hydroxide

Insoluble e.g. Cu(OH)2, Fe(OH)3

Soluble e.g. NaOH, KOH, Ca(OH)2

alkalis

water

Soluble e.g. Na2O, K2O, CaO

Insoluble e.g. MgO, CuO, PbO

Common alkalis NaOH KOH Ca(OH)2

Na+, OH− K+, OH− Ca2+, 2OH−

Ba(OH)2

Ba2+, 2OH−

Which gas turns damp red litmus to blue? NH3(g) + H2O(l) → NH4+(aq) + OH−(aq) (NH4OH)

OH− ion and alkalis  OH−

hydroxide ion  This ion makes the substance alkaline Acid + Alkali → Salt + Water (Neutralisation) H+(aq) + OH−(aq) → H2O(l) Ionic equation

Shows what actually reacts in a chemical reaction • ions not featured in the ionic equation do not take part in the chemical reaction (spectator ions)

Another reaction of alkalis heat

Alkali + Ammonium salt → Salt + H2O + ammonia that is: heat NaOH(aq) + NH4Cl(s) → NaCl(aq) + H2O(l) + NH3(g)

Neutralisation only involves H+ and OH− ions HCl(aq) + NaOH(aq)

Cl−

H+

Na+ OH− H+

Na+

H+

OH−

Na

OH− H+

OH− Na+

OH− H+

OH− Cl−

−

Cl− Na+

H+ OH

Na+

+

Cl−

+

Cl− Na+

H

Na

+

Cl

−

H+

Cl− OH− Cl−

Na+ and Cl− ions remain as ions in solution Same state as when they are in HCl and NaOH

Constructing ionic equations

Sincl2.mov

- must always include state symbols H2SO4(aq) + NaOH(aq) → Na2SO4(aq) + H2O(l) AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq) 2HCl(aq)+ Na2CO3(aq) → 2NaCl(aq)+ CO2(g) + H2O(l) H2SO4(aq) + CuO(s) → CuSO4(aq) + H2O(l) Fe

+ H SO

→ FeSO

+ H

Quick Check Page 175

Use of OH− ions to identify cations Reaction with NaOH

Reaction with NH4OH

Ca2+

White ppt insoluble in excess NaOH

No rxn

Al3+

White ppt soluble in excess NaOH

White ppt insoluble in excess NH4OH

Pb2+

White ppt soluble in excess NaOH

White ppt insoluble in excess NH4OH

Zn2+

White ppt soluble in excess NaOH

White ppt soluble in excess NH4OH

Cu2+

Blue ppt insoluble in excess NaOH Blue ppt soluble in excess NH4OH to give a deep blue solution Dirty-green ppt insoluble in excess Dirty-green ppt insoluble in NaOH excess NH4OH

Fe2+ Fe3+

Reddish-brown ppt insoluble in excess NaOH

Reddish-brown ppt insoluble in excess NH4OH

NH4+

NH3 gas produced on warming that turns damp red litmus to blue

No rxn

Use of OH− ions to identify cations Reaction with NaOH

Reaction with NH4OH

Ca2+

White2+ ppt insoluble in− excess NaOH

No rxn

Al3+

White 3+ppt soluble in−excess NaOH

White ppt insoluble in excess NH43OH

Pb2+

White2+ ppt soluble in excess − NaOH

White ppt insoluble in excess NH4OH 2

Zn2+

White2+ ppt soluble in excess − NaOH

White ppt soluble in excess NH4OH 2

Cu2+

Blue ppt NaOH Blue ppt soluble in excess NH4OH 2+ insoluble in excess − to give2a deep blue solution Dirty-green ppt insoluble 2+ − in excess Dirty-green ppt insoluble in NaOH excess 2 NH4OH

Sincl2.mov

Fe2+

Ca Al

Pb Zn

+ 2OH → Ca(OH)2

+ 3OH → Al(OH)

+ 2OH → Pb(OH)

+ 2OH → Zn(OH)

Cu

+ 2OH → Cu(OH)

Fe

+ 2OH → Fe(OH)

Fe3+

Reddish-brown ppt insoluble in 3+ − excess NaOH

NH4+

NH3 gas No rxn + produced on − warming 4 damp red litmus to blue 3 2 that turns

Fe

NH

Reddish-brown ppt insoluble in excess 3 NH4OH

+ 3OH → Fe(OH)

+ OH → NH + H O

pH scale for Universal indicator tells H+ and OH− ion concentration

Liquid drain cleaner Bleach, oven cleaner Soapy water, detergents Ammonia solution Milk of magnesia Baking soda, toothpaste

7

Sea water, blood Pure water, e.g. distilled water

Urine, saliva, rain water Soft drinks, fizzy drinks, coffee

Tomato juice, acid rain Orange juice, grapefruit Lemon juice, vinegar Hydrochloric acid secreted by stomach

Battery acid, Hydrofluoric acid

10 11 12 13 14 9 8 6 5 4 3 2 1 0

more alkaline more acidic

pH indicators  Expected

to know the colour changes for methyl orange and screened methyl orange screened methyl orange methyl orange

acid

neutral

alkaline

purple

grey

green

red

peach orange

yellow

phenolphthalein colourless colourless

pink

bromothymol blue

blue

yellow

-

Other pH indicators  

Testsol2.mov

Redlit2.mov

Bluelit2.mov

Litmus papers – detects acidity or alkalinity only. pH meters – measures the exact pH value directly and electronically.

Acialkb2.mov

Controlling pH in soils 



Excess acidity in soils causes crops not to grow well (usually grow well only under slightly acidic – neutral pH 6-7 conditions) Caused by acid rain. can be neutralised by

Slaked lime – calcium hydroxide Quick lime – calcium oxide

H2SO4 + Ca(OH)2 → CaSO4 + 2H2O H2SO4 + CaO → CaSO4 + H2O

Extension…. Magnesium oxide and calcium carbonate are both used in indigestion tablets to neutralise the excess acids in our stomach. Which do you think will have less side effects? Why? Sodium hydroxide neutralises HCl more readily, wouldn’t this be more effective then?