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Acids and Basis

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Acids and Basis
« on: September 08, 2014, 10:29:58 AM »
Acids, Bases and Salts  All substances are acidic, neutral or basic (alkaline). How  acidic or basic a substance is shown by its pH. There are several other ways by which we could find out whether a substance is acidic, neutral or basic.
 pH Scale: This is a scale that runs from 0 to 14. Substances with a pH  below 7 are acidic. Substances with pH above 7 are basic. And those with pH 7 are neutral.
Indicators:  Indicators are substances that identify acidity or  alkalinity of substances. They cannot be used in solid form.
  Universal Indicator:  This is a substance that changes color when added to another substance depending on its pH. The indicator and the substance should be in  aqueous form.
 Litmus Paper or Solution:  This indicator is present in two colors: red and blue. We  use blue litmus if we want test a substance for acidity. We use red litmus if  we want to test a substance for alkalinity. Its results are:
 
  • Acids: Turns blue litmus paper/ solution red,
  • Bases: Turns red litmus paper/ solution blue,
  • Neutral: if it is used as paper the color doesn’t change. If  it is used as solution it turns purple.
  Note: use damp litmus paper if testing gases.
 Phenolphthalein:  This is an indicator that is used to test for alkalinity  because it is colorless if used with an acidic or neutral substance and it is pink if it is used with a basic substance.
 Methyl Orange:  This indicator gives fire colors: Red with acids, yellow  with neutrals and orange with bases.
 
 Acids:  Acids are substances made of a hydrogen ion and non-metal  ions.  They have the following  properties:
 
  • They dissolve in water producing a hydrogen ion H+,
  • They have a sour taste,
  • Strong ones are corrosive,
  • Their pH is less than 7.
All acids must be in aqueous form to be called an acid. For  example Hydrochloric acid is hydrogen chloride gas dissolved in water. The most common acids are:
 
  • Hydrochloric acid HCl,
  • Sulphuric Acid H2SO4,
  • Nitric Acid HNO3,
  • Cirtric Acid,
  • Carbonic Acid H2CO3.
Strength of Acids:  One of the most important properties of acids is that it  gives hydrogen ion H+ when dissolved in water. This is why the amount of H+ ions the acid can give when dissolved in water is what determines its strength.  This is called ionization or dissociation. The more ionized the acid is the stronger  it is, the lower its pH. The more H+ ions given when the acid is dissolved in  water the more ionized the acid is.
 
Strong Acids:[/t][/t]
  • Have pH’s: 0,1,2,3
  • They are fully ionized
  • When dissolved in water, they give large amounts of H+ ions
  • Examples:
  • Hydrochloric Acid
  • Sulfuric Acid
  • Nitric Acid
  • [/t][/l][/l]
    Weak Acids:
    [/t]
       
    • Have pH’s: 4,5,6
    • They are partially ionized
    • When dissolved in water, they give small amounts of H+ ions
    • Examples:
    • Ethanoic acid (CH3COOH)
    • Citric Acid
    • Carbonic Acid
    • [/l]   Hydrochloric acid is a strong acid. When it is dissolved in water  all HCl molecules are ionized into H+ and Cl-  ions. It is fully ionized.
      Ethanoic acid has the formula CH3COOH. It is a weak acid. When  it is dissolved in water, only some of the CH3COOH molecules are ionized into  CH3COO-  and H+ ions. It is partially  ionized.
      Note: Acids with pH 3 or 4 can be considered moderate in  strength.
      Solutions of strong acids are better conductors of  electricity than solutions of weak acids. This is because they contain much  more free mobile ions to carry the charge.
      Concentrated acids are not necessarily strong. The concentration  of an acid only means the amount of molecules of the acid dissolved in water.  Concentrated acids have a large amount of acid molecules dissolved in water.  Dilute acids have a small amount of acid molecules dissolved in water.  Concentration is not related to strength of the acids. Strong acids are still  strong even if they are diluted. And weak acids are still weak even if they are  concentrated.
       
       Bases:  Bases are substances made of hydroxide OH- ions and a metal.  Bases can be made of:
       
      • Metal hydroxide (metal ion & OH- ion)
      • Metal oxides
      • Metal carbonates (metal ion & CO32-)
      • Metal hydrogen carbonate (Bicarbonate)
      • Ammonium hydroxide (NH4OH)
      • Ammonium Carbonate ((NH4)2CO3)
      Properties of bases:
       
      • Bitter taste
      • Soapy feel
      • Have pH’s above 7
      • Strong ones are corrosive
      Some bases are water soluble and some bases are water  insoluble. Water soluble bases are also called alkalis.
      Like acids, alkalis' strength is determined by its ability  to be ionized into metal and hydroxide OH-  ions. Completely ionized alkalis are the  strongest and partially ionized alkalis are the weakest. Ammonium hydroxide is  one of the strongest alkalis while weak alkalis include the hydroxides of  sodium, potassium and magnesium.
       
       Types of Oxides:
      Basic Oxides
      [/t][/t][/t]
         
      • They are metal oxides
      • They react with acids forming a salt and water
      • They are solids
      • They are insoluble in water except group 1 metal oxides.
      • They react with an acid forming salt and water
      • Examples: Na2O, CaO and CuO
      • [/l]
        Amphoteric Oxides
        [/t]
           
        • These are oxides of Aluminum, Zinc & Lead
        • They act as an acid when reacting with an alkali & vice versa
        • Their element’s hydroxides are amphoteric too
        • They produce salt and water when reacting with an acid or an alkali.
        • [/l]
          Acidic Oxides
          [/t]
             
          • They are all non-metal oxides except non-metal monoxides
          • They are gases
          • They react with an alkali to form salt and water
          • Note:  metal monoxides are neutral oxides
          • Examples: CO2, NO2, SO2 (acidic oxides) & CO, NO,
             H2O (neutral oxides)
          • [/l]   
             Salts:  A salt is a neutral ionic compound. Salts are one of the  products of a reaction between an acid and a base. Salts are formed in  reactions I n which the H+ ion from the acid is replaced by any other metal  ion. Some salts are soluble in water and some are insoluble.
             
            Soluble Salts:[/t][/t]
            [/t]
               
            • All Nitrates
            • All halides EXCEPT AgCl and PbCl2
            • All sulfates EXCEPT CaSO4, BASO4,  PbSO4
            • All group 1 metals salts
            • All ammonium salts
            • [/l]
              Insoluble Salts:
              [/t]
                 
              • Silver and lead chlorides  (AgCl & PbCl2)
              • Calcium, barium and lead sulphates (CaSO4, BASO4,  PbSO4)
              • All carbonates EXCEPT group 1 metals and ammonium carbonates
              • [/l]   
                 Preparing Soluble Salts:  Displacement Method (Excess Metal Method):
                Metal + Acid → Salt + Hydrogen
                Note: this type of method is suitable to for making salts of  moderately reactive metals because highly reactive metals like K, Na and Ca  will cause an explosion. This method is used with the MAZIT (Magnesium,  Aluminum, Zinc, Iron and Tin) metals only.
                Example: set up an experiment to obtain magnesium chloride  salt.
                Mg + 2HCl → MgCl2 + H2
                • Add 100 cm3 of dilute hydrochloric acid to a beaker
                • Add excess mass of powdered magnesium
                • When the reaction is done, filter the mixture to get rid of  excess magnesium (residue)
                • The filtrate is magnesium chloride solution
                • To obtain magnesium chloride powder, evaporate the solution  till dryness
                • To obtain magnesium chloride crystals, heat the solution  while continuously dipping a glass rod in the solution
                • When you observe crystals starting to form on the glass rod,  turn heat off and leave the mixture to cool down slowly
                • When the crystals are obtained, dry them between two filter  papers
                Observations of this type of reactions:
                • Bubbles of colorless gas evolve (hydrogen). To test approach  a lighted splint if hydrogen is present it makes a pop sound
                • The temperature rises (exothermic reaction)
                • The metal disappears
                You know the reaction is over when:
                 
                • No more gas evolves
                • No more magnesium can dissolve
                • The temperature stops rising
                • The solution becomes neutral
                Proton Donor and Acceptor Theory:
                 When an acid and a base react, water is formed. The acid  gives away an H+ ion and the base accepts it to form water by bonding it with  the OH- ion. A hydrogen ion is also called a proton this is why an acid can be  called Proton Donor and a base can be called Proton Acceptor.
                 
                 Neutralization Method:
                Acis + Base → Salt + Water
                Note: This method is used to make salts of metals below  hydrogen in the reactivity series. If the base is a metal oxide or metal  hydroxide, the products will be salt and water only. If the base is a metal carbonate, the products will be salt, water and carbon dioxide.
                 Type 1:
                Acid + Metal Oxide → Salt + Water
                To obtain copper sulfate salt given copper oxide and  sulfuric acid:
                CuO + H2SO4 → CuSO4 + H2O
                • Add 100 cm3 of sulfuric acid to a beaker
                • Add excess mass of Copper oxide
                • When the reaction is over, filter the excess copper oxide  off
                • The filtrate is a copper sulfate solution, to obtain copper  sulfate powder evaporate the solution till dryness
                • To obtain copper sulfate crystals, heat the solution white  continuously dipping a glass rod in it
                • When you observe crystals starting to form on the glass rod,  turn heat of and leave the mixture to cool down slowly
                • When you obtain the crystals dry them between two filter  papers
                Observations of this reaction:
                 
                • The amount of copper oxide decreases
                • The solution changes color from colorless to blue
                • The temperature rises
                • You know the reaction is over when
                • No more copper oxide dissolves
                • The temperature stops rising
                • The solution become neutral
                Type 2:
                Acid + Metal Hydroxide → Salt + Water
                to obtain sodium chloride crystals given sodium hydroxide  and hydrochloric acid:
                HCl + NaOH → NaCl + H2O
                • Add 100 cm3of dilute hydrochloric acid to a beaker
                • Add excess mass of sodium hydroxide
                • When the reaction is over, filter the excess sodium  hydroxide off
                • The filtrate is sodium chloride solution, to obtain sodium  chloride powder, evaporate the solution till dryness
                • To obtain sodium chloride crystals, hear the solution while  continuously dipping a glass rod in it
                • When crystals start to form on the glass rod, turn heat off  and leave the mixture to cool down slowly
                • When the crystals are obtained, dry them between two filter  papers
                Observations:
                 
                • Sodium hydroxide starts disappearing
                • Temperature rises
                You know the reaction is over when:
                 
                • The temperature stops rising
                • No more sodium hydroxide can dissolve
                • The pH of the solution becomes neutral
                Type 3:
                Acid + Metal Carbonate → Salt + Water + Carbon Dioxide
                To obtain copper sulfate salt given copper carbonate and  sulfuric acid:
                CuCO3 + H2SO4 → CuSO4 + H2O + CO2
                • Add 100 cm3 of dilute sulfuric acid to a beaker
                • Add excess mass of copper carbonate
                • When the reaction is over, filter excess copper carbonate  off
                • The filtrate is a copper sulfate solution, to obtain copper  sulfate powder evaporate the solution till dryness
                • To obtain copper sulfate crystals, heat the solution white  continuously dipping a glass rod in it
                • When you observe crystals starting to form on the glass rod,  turn heat of and leave the mixture to cool down slowly
                • When you obtain the crystals dry them between two filter  papers
                  Observations:
                 
                • Bubbles of colorless gas (carbon dioxide) evolve, test by approaching lighted splint, if the CO2  is present the flame will be put off
                • Green Copper carbonate starts to disappear
                • The temperature rises
                • The solution turns blue
                You know the reaction is finished when:
                 
                • No more bubbles are evolving
                • The temperature stops rising
                • No more copper carbonate can dissolve
                • The pH of the solution becomes neutral
                Titration Method: This is a method to make a neutralization reaction between a base and an acid producing a salt without any excess. In this method, the  experiment is preformed twice, the first time is to find the amounts of  reactants to use, and the second experiment is the actual one.
                1st Experiment:
                 
                • Add 50 cm3 of sodium hydroxide using a pipette to be  accurate to flask
                • Add 5 drops of phenolphthalein indicator to the sodium  hydroxide. The solution turns pink indicating presence of a base
                • Fill a burette to zero mark with hydrochloric acid
                • Add drops of the acid to conical flask
                • The pink color of the solution becomes lighter
                • When the solution turns colorless, stop adding the acid (End point: is the point at which every base molecule is  neutralized by an acid molecule)
                • Record the amount of hydrochloric acid used and repeat the  experiment without using the indicator
                • After the 2nd experiment, you will have a sodium chloride solution. Evaporate it till dryness to obtain powdered sodium chloride or crystalize  it to obtain sodium chloride crystals
                 
                 Preparing Insoluble Salts:  Precipitation Method:  A precipitation reaction is a reaction between two soluble  salts. The products of a precipitation reaction are two other salts, one of them is soluble and one is insoluble (precipitate).
                Example: To obtain barium sulfate salt given barium chloride  and sodium sulfate:
                BaCl2 + Na2SO4 → BaSO4 + 2NaCl
                 Ionic Equation: Ba2+ + SO42- → BaSO4
                • Add the two salt solutions in a beaker
                • When the reaction is over, filter and take the residue
                • Wash the residue with distilled water and dry it in the oven
                Observations:
                 
                • Temperature increases
                • An insoluble solid precipitate (Barium sulfate) forms
                You know the reaction is over when:
                 
                • The temperature stops rising
                • No more precipitate is being formed
                 
                 Controlling Soil pH:  If the pH of the soil goes below or above 7, it has to be neutralized using an acid or a base. If the pH of the soil goes below 7, calcium carbonate (lime stone) is used to neutralize it. The pH of the soil can  be measured by taking a sample from the soil, crushing it, dissolving in water  then measuring the pH of the solution.
                 Colors of Salts:
                SaltFormulaSolid In Solution
                Hydrated copper sulfate[/t][/t]
                [/t]
                CuSO4.5H2O
                Blue crystals
                Blue
                Anhydrous copper sulfate
                CuSO4
                White powder
                Blue
                Copper nitrate
                Cu(NO3)2
                Blue crystals
                Blue
                Copper chloride
                CuCl2
                Green
                Green
                Copper carbonate
                CuCO3
                Green
                Insoluble
                Copper oxide
                CuO
                Black
                Insoluble
                Iron(II) salts
                E.g.: FeSO4, Fe(NO3)2
                Pale green crystals
                Pale green
                Iron(III) salts
                E.g.: Fe(NO3)3
                Reddish brown
                Reddish brown
                   
                 Tests for Gases:
                GasFormula Tests
                Ammonia[/t][/t]
                [/t]
                NH3
                Turns damp red litmus paper blue
                Carbon dioxide
                CO2
                Turns limewater milky
                Oxygen
                O2
                Relights a glowing splint
                Hydrogen
                H2
                ‘Pops’ with a lighted splint
                Chlorine
                Cl2
                Bleaches damp litmus paper
                Nitrogen dioxide
                NO2
                Turns damp blue litmus paper red
                Sulfur dioxide
                SO2
                Turns acidified aqueous potassium dichromate(VI) from orange to green
                   
                 Tests for Anions:
                AnionTestResult
                Carbonate (CO32-)Add dilute acidEffervescence,
                 carbon dioxide produced
                Chloride (Cl-)
                 (in solution)
                Acidify with dilute nitric acid, then add
                 aqueous silver nitrate
                White ppt.
                Iodide (I-)
                 (in solution)
                Acidify with dilute nitric acid, then add
                 aqueous silver nitrate
                Yellow ppt.
                Nitrate (NO3-)
                 (in solution)
                Add aqueous sodium hydroxide, then
                 aluminium foil; warm carefully
                Ammonia produced
                Sulfate (SO42-)Acidify, then add aqueous barium nitrateWhite ppt.
                 
                 Tests for aqueous cations:
                CationEffect of aqueous sodium hydroxideEffect of aqueous ammonia
                Aluminium (Al3+)White ppt., soluble in excess giving a
                 colourless solution
                White ppt., insoluble in excess
                Ammonium (NH4+)Ammonia produced on warming
                Calcium (Ca2+)White ppt., insoluble in excessNo ppt. or very slight white ppt.
                Copper (Cu2+)Light blue ppt., insoluble in excessLight blue ppt., soluble in excess,
                 giving a dark blue solution
                Iron(II) (Fe2+)Green ppt., insoluble in excessGreen ppt., insoluble in excess
                Iron(III) (Fe3+)Red-brown ppt., insoluble in excessRed-brown ppt., insoluble in excess
                Zinc (Zn2+)White ppt., soluble in excess,
                 giving a colourless solution
                White ppt., soluble in excess,
                 giving a colourless solution


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