Acids, Bases and Salts

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Acids, Bases and Salts
MyElimu Offline
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Information Acids, Bases and Salts

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Acids

Electrical conductivity
Any solution's ability to conduct electricity is defined by is charges ions in it. As a result, a strong acid will produce more charged ions than a weak one, and so it's solution will be a better electrical conductor than a weak acid. The same goes for strong/weak bases.

Acids in daily life
  • Ethanoic acid – found in vinegar and tomato juice
  • Citric acid – found in citrus foods like lemons, oranges and grapefruit
  • Lactic acid – found in sour milk and yoghurt, and in muscle respiration
  • Tartaric acid – found in grapes
  • Tannic acid – found in tea and ant’s body
  • Formic acid – found in bee stings
  • Hydrochloric acid – found in stomach juices
Common laboratory acids
  • Hydrochloric acid (HCl)
  • Sulphuric acid (H2SO4)
  • Nitric acid (HNO3)
Dilute acid: solution containing small amount of acid dissolved in water
Concentrated acids: solution containing large amount of acid dissolved in water

Properties of acids
    • sour taste
    • hazardous - irritants to skin, causing skin to redden and blister
    • change the color of indicators - turn blue litmus red
    • react with metals to produce hydrogen gas - gas is tested with a burning splint which burns with a 'pop' sound
    • react with carbonates and hydrogencarbonates to produce carbon dioxide - to test this, the gas produced is bubbled into limewater which forms a white precipitate
    • react with metal oxides and hydroxides - reach slowly with warm dilute acid to form salt and water
Storage of acids
  • Acids are stored in claypots, glass or plastic containers as sand, glass and plastic do not react with acids. 
  • If it is stored in metal container, metal would react with acids
Uses of acids
  • Sulphuric Acid 
    • Used in car batteries
    • Manufacture of ammonium sulphate for fertilisers
    • Manufacture of detergents, paints, dyes, artificial fibres & plastics
  • Hydrochloric acid 
    • can remove rust (iron(III) oxide) which dissolves in acids
  • Acids are used in preservation of foods (e.g. ethanoic acid)
Acids and hydrogen ions
  • Acids are covalent compounds and do not behave as acids in the absence of water as water reacts with acids to produce H+ ions, responsible for its acidic properties
    • e.g. Citric acid crystals doesn’t react with metals and doesn’t change colours of indicators; citric acid in water reacts with metals and change turns litmus red.
  • Hydrogen gas is formed by acids as H+(aq) ions are present in acid solutions. This means when a solid/gas acid dissolved in water, they produce H+ ions in it
    • Chemical eqation: HCl(s) ---(water)---> HCl (aq)
    • Ionic Equation: HCl(s) ---(water)---> H+ (aq) + Cl- (aq)
  • However when dissolved in organic solutions, they don’t show acidic properties. When metals react with acids, only the hydrogen ions react with metals, e.g.:
    • Chemical equation: 2Na(s) + 2HCl(aq) → 2NaCl(aq) + H2(g)
    • Ionic equation: 2Na(s) + 2H+(aq) → 2Na+(aq) + H2(g)
  • Basicity of an acid is maximum number of H+ ions produced by a molecule of acid
    • dibasic: can replace two hydrogen atoms
    • tribasic: can replace three hydrogen atoms
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 Acids Reaction with water Basicity
 Hydrochloric acid HCl (aq) → H(aq) + Cl(aq)monobasic
 Nitric acid HNO(aq) → H(aq) + NO3(aq)monobasic
 Ethanoic acidCH3COOH (aq) ⇌ H(aq) + CH3COO(aq)monobasic
 Sulphuric acid H2SO(aq) → 2H(aq) + SO42- (aq)dibasic
 Fizzy drinks
  • Soft drink tablets contains solid acid (e.g. citric acid, C6H8O7) & sodium bicarbonate
  • When tablet is added to water, citric acid ionises and the H+ produced reacts with sodium bicarbonate to produce carbon dioxide gas, making them fizz
Strong and weak acids
  • Strong Acids - acid that completely ionises in water. 
    • Their reactions are irreversible. 
    • E.g. H2SO4, HNO3, HCl
    • H2SO4 (aq) → 2H(aq) + SO42- (aq)
      • In above H2SO4 has completely been ionized in water, forming 3 kinds of particles:
        • H+ ions
        • SO42- ions
        • H2O molecules
    • Strong acids react more vigorously with metals than weak acids – hydrogen gas bubbles are produced rapidly
  • Weak acids - acids that partially ionise in water. 
    • The remaining molecules remain unchanged as acids. 
    • Their reactions are reversible. 
    • E.g. CH3COOH, H2CO3, H3PO4
    • H3PO4 (aq) ⇌ 3H(aq) + PO42- (aq)
    • Weak acids react slowly with metals than strong acids – hydrogen gas bubbles are produced slowly.
Concentration vs Strength#000000 solid;margin:15px;width:95%;">
 CONCENTRATION STRENGTH
 Is the amount of solute (acids or alkalis) dissolved in 1 dm3 of a solution Is how much ions can be disassociated into from acid or alkali
 It can be diluted by adding more water to solution or concentrated by adding more solute to solution The strength cannot be changed
  • Comparing 10 mol/dm3 and 0.1 mol/dm3 of hydrochloric acids and 10 mol/dm3 and 0.1 mol/dm3 of ethanoic acids
    • 10 mol/dm3 of ethanoic acid solution is a concentrated solution of weak acid
    • 0.1 mol/dm3 of ethanoic acid solution is a dilute solution of weak acid
    • 10 mol/dm3 of hydrochloric acid solution is a concentrated solution of strong acid
    • 0.1 mol/dm3 of hydrochloric acid solution is a dilute solution of strong acid
Bases and Alkalis
  • Bases are oxides or hydroxides of metals
  • Alkalis are bases which are soluble in water
  • All alkalis produces hydroxide ions (OH-) when dissolved in water. 
  • Hydroxide ions give the properties of alkalis. 
  • They don’t behave as acids in absence of water.
  • Alkalis are therefore substances that produce hydroxide ions, OH(aq), in water.
Laboratory Alkalis
  • Sodium Hydroxide, NaOH
  • Aqueous Ammonia, NH4OH
  • Calcium Hydroxide, Ca(OH)2
Properties of Alkalis
  • have a slippery feel
  • hazardous
  • Dilute alkalis are irritants
  • Concentrated alkalis are corrosive and burn skin (caustic(i.e. burning) alkalis)
  • change the colour of indicators: turn common indicator litmus – red litmus to blue
  • react with acids
    • The reaction is called neutralisation
  • react with ammonium compounds
    • They react with heated solid ammonium compounds to produce ammonia gas
    • (NH4)2SO4 (s) + Ca(OH)(aq) → CaSO4 (aq) + 2NH3 (g) + 2H2O (l)
  • react with solutions of metal ions
    • Barium sulphate, BaSO(aq), contains Ba2+ (aq) ions
    • Ca(OH)(aq) + BaSO(aq) → Ba(OH)(s) + CaSO4 (aq)
    • The solid formed is precipitate – the reaction is called precipitate reaction
Strong and weak alkalis
  • Strong Alkalis: base that completely ionises in water to form OH(aq) ions. 
    • Their reactions are irreversible. 
    • E.g. NaOH, KOH, Ca(OH)2
    • Ca(OH)(s) → Ca2+ (aq) + 2OH(aq)
  • Weak Alkalis: base that partially ionise in water. 
    • The remaining molecules remain unchanged as base. 
    • Their reactions are reversible. 
    • E.g. NH3
    • NH(g) + H2O (l) ⇌ NH4(aq) + OH(aq)
Uses of Alkalis
  • Alkalis neutralise acids in teeth (toothpaste) and stomach (indigestion)
  • Soap and detergents contain weak alkalis to dissolve grease
  • Floor and oven cleaners contain NaOH (strong alkalis)
  • Ammonia (mild alkalis) is used in liquids to remove dirt and grease from glass
pH and IndicatorsIndicators are substances that has different colours in acidic and alkaline solutions

Common indicators:
  • Litmus
  • Methyl orange
  • Phenolphtalein
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 Indicator Colour in acidscolour changes at pH  Colour in alkalis
 Phenolphtalein Colourless 9 Pink
 Methyl orange Red 4Yellow
 Litmus Red 7 Blue
 Screened methyl orange Red 4 Green
 Bromothymol blue Yellow 7 Blue
 pH ScaleA measure of acidity or alkalinity of a solution is known as pH
  • pH 7 is neutral – in pure water
  • solutions of less than pH 7 are acidic. 
    • The solutions contain hydrogen ions. 
      • The lower the pH, the more acidic the solution is and more hydrogen ions it contains.
  • solutions of more than pH 7 are alkaline. 
    • The solution contains hydroxide ions. 
      • The higher the pH, the more alkaline the solution and more hydroxide ions it contains.
Measuring pH of a Solution
  • Universal indicators
    • It can be in paper or solution form. 
    • Universal paper can be dipped into a solution then pH found is matched with the colour chart. 
    • It gives approximate pH value.
  • pH meter
    • A hand-held pH probe is dipped into solution and meter will show the pH digitally or by a scale. 
    • Measures pH water in lakes, water, and streams accurately
  • pH sensor and computer
    • A probe is dipped into solution and will be sent to computer through interface used to measure pH of solution. 
    • The pH reading is displayed on computer screen.
Ionic EquationsIonic equation is equation involving ions in aqueous solution, showing formation and changes of ions during the reaction

Rule to make ionic equations:
  • Only formulae of ions that change is included; ions don’t change = omitted
  • Only aqueous solutions are written as ions; liquids, solids and gases written in full
Reaction between Metals and Acids Eg. reaction of sodium with hydrochloric acid

2Na (s) + 2HCl (aq) → 2NaCl (aq) + H2 (g)

Its ionic equation is written as:
2Na (s) + 2H(aq) + 2Cl(aq) → 2Na(aq) + 2Cl(aq) + H(g)

Since 2 Cl(aq) ions don’t change, they’re not involved in reaction. 
As ionic equation is used to show changes in reactions, we omit Cl(aq) ions. 

So we’re left with:
2Na (s) + 2H(aq) → 2Na(aq) + H(g)Reaction between soluble ionic compounds and acids e.g. Reaction of sodium hydrogencarbonate with hydrochloric acid

NaHCO(aq) + HCl (aq) → NaCl (aq) + CO(g) + H2O (l)

Its ionic equation is:
Na(aq) + H(aq) + CO32- (aq) + H(aq) + Cl(aq) → Na(aq) + Cl(aq) + CO(g) + H2O (l)

Since Na(aq) and Cl(aq) ions don’t change, we omit them, leaving:
H(aq) + CO32- (aq) + H(aq) → CO(g) + H2O (l)
CO32- (aq) + 2H(aq) → CO


Code:
(g) + H2O (l)Reaction between insoluble ionic compounds and acids e.g. Reaction between iron(II) oxide and sulphuric acid

FeO (s) + H2SO(aq) → FeSO


Code:
(aq) + H2O (g)

Its ionic equation is:
FeO(s) + 2H+ (aq) + SO42- (aq) → Fe2+ (aq) + SO42- (aq) + H2O (g)

Note: FeO is written in full as it is solid (although it is an ionic compound)

Since SO42-  (aq) ions don’t change, we omit SO42- ions, leaving:
FeO (s) + 2H+ (aq) → Fe2+ (aq) + H2O (g)

E.g. Reaction between calcium carbonate and hydrochloric acid

CaCO(s) + 2HCl (aq) → CaCl2 (aq) + CO(g) + H2O (l)

Its ionic equation is:
CaCO3 (s) + 2H(aq) + 2Cl- (aq) → Ca2+ (aq) + 2Cl- (aq) + CO(g) + H2O (l)

Since 2 Cl- (aq) ions don’t change, we omit Cl- ions, leaving:
CaCO(s) + 2H(aq) → Ca2+ (aq) + CO(g) + H2O (l)Reaction producing precipitate  E.g. Reaction between calcium hydroxide and barium sulphate

Ca(OH)2 (aq) + BaSO4 (aq) → Ba(OH)(s) + CaSO(aq)

Its ionic equation is written as:
Ca2+ (aq) + 2OH- (aq) + Ba2+ (aq) + SO42- (aq) → Ba(OH)


Code:
(s) + Ca2+ (aq) + SO42- (aq)

Since Ca2+ (aq) and SO42- (aq) ions don’t change, we omit them, leaving:
Ba2+ (aq) + 2OH(aq) → Ba(OH)(s)Displacement reactions  E.g. Reactions between magnesium with zinc sulphate

Mg (s) + ZnSO4 (aq) → MgSO4 (aq) + Zn (s)

Its ionic equation is written as:
Mg (s) + Zn2+ (aq) + SO42- (aq) → Mg2+ (aq) + SO42- (aq) + Zn (s)

Since SO42-  (aq) ions don’t change, we omit them, leaving:
Mg (s) + Zn2+ (aq) → Mg2+ (aq) + Zn (s)Neutralization
  • Neutralization is the reaction between acid and base to form salt and water only.
  • From ionic equation, we know that the reaction only involves H+ ions from acids with OH- ions from alkali to form water .
E.g. NaOH + H2SO4 forms Na2SO


Code:
4
 + H2O


H2SO4 (aq) + NaOH (aq) -->  Na2SO


Code:
(aq) + H2O (g)

Ionic equation is:
H(aq) + OH- (aq)→ H2O (g)

Plants don’t grow well in acidic soil. Quicklime (calcium hydroxide) is added to neutralise the acidity of soil according to equation:
Acid (aq) + Ca(OH)(aq) --> Ca(acid anion) (aq) + H2O (g)Reaction between Base and Ammonium Salts  E.g. Reaction between NaOH and NH4OH

NaOH (aq) + NH4Cl (aq) --> NaCl (aq) + NH3 (g) + H2O (g)

Ionic equation:
NH4(aq) + OH-(aq) → NH3 (g) + H2O (g)Oxides#000000 solid;margin:15px;width:95%;">[td]Oxides that don’t react with either acids/alkalis, hence do not form salts, e.g. H


Code:
2
O, CO, NO[/td]
 Acidic oxideOxides of non-metals, usually gases which reacts with water to produce acids, e.g. CO2, NO3, P4O10, SO2
 Basic oxideOxides of metals, usually solid which reacts with water to produce alkalis, e.g. CaO, K2O, BaO
 Amphoteric oxideOxides of transition metals, usually solid, which reacts with acids/alkalis to form salt and water, e.g. Al2O3, FeO, PbO
 Neutral oxide
 Preparation of Salts#000000 solid;margin:15px;width:95%;">[td] PbCl(soluble in hot water), 
AgCl, HgCl2[/td]
Soluble Insoluble 
 All Nitrates -
 All sulphates except --> BaSO4, CaSO4, PbSO4
 All Chlorides except -->
 Potassium, Sodium, Ammonium salts -
 K2CO3, Na2CO3, (NH4)2CO3 All other carbonates
 K2O, Na2O All other oxides
 [Image: Screen%20shot%202012-03-03%20at%20PM%2009.41.16.png]
 

 

 
02-27-2014 05:57 PM
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