Acid bases and salts

^ PH

* Ph tells us how acidic or basic a substance is. The acidity or basicity of a solution can be described in terms of the hydrogen ion concentration. 

[ H+] → concentration of hydrogen is represented in [] - square brackets. 

* Acids have Phs lower than 7 and bases higher than

1. [ H+] ( In all acids) is = to the [OH-] (base/alkali) making the solution neutral since none is greater than the other. If the acid is equal to the base you have a neutral solution. 

2. The [H+] (acid) is > - greater than the [OH-] (base) making the solution acidic. When the acid is greater than the base the solution is acidic. 

3. [H+] (acid) is < - less than the [OH-] (base) making the solution basic. When the base is greater than the acid, or the acid is less than the base the solution is basic. 

* [H+] ( The hydrogen ion concentration ) is the acid. [OH-] ( the hydroxide ion concentration) is the base or alkali. 

* The litmus paper tells if the solution is acidic or basic. 

* Bases and acids react with each other to produce salt. 

- Finding the Ph of substances

* The universal indicator ( universal paper) is the test used to find Ph. 

1. Universal indicator/paper: 

A mixture of acid-base indicators that show changes in colour at different Ph values. It is matched on the Ph scale. 

2. Ph meter: 

It translates the concentration of H+ ion into an electric signal that is converted into a digital display or deflection on a meter. The Ph meter would give a more accurate reading because it is able to give decimals. 

* The universal and litmus paper have indicators which is why they can change colours.

- Indicators 

* Liquid indicators exist. 

^ ACIDS

- Properties of acids 

• Corrosive 

• Sour taste 

• Ph less than 7

• change blue litmus to red 

• conduct electricity when in solution (aqueous) 

- Definition of acid

* Acids are substances that produce a greater concentration of hydrogen ions than hydroxide ions when dissolved in water or aqueous solutions. 

- The states acids exist as 

Gases, liquids & solids 

* Gases: 

CO ₂  Carbon dioxide 

SO₂  Sulpher dioxide 

SO₃ Sulphite trioxide 

NO₂  Nitrogen oxide 

* Liquids: 

HCL 

Nitric acid 

Sulphuric acid 

*Solids: 

Citric acid [ from citric fruits eg lime, orange, lemon] C₆ H₈ O₇

Ascorbic acid ( vitamin C) C₆ H₈ O₆

Tartaric acid ( from grapes)

- Types of  Acids 

Mineral acids     

Organic acids 

1. Mineral Acids: 

* Obtained from minerals 

H₂SO₄ Sulphuric acid 

HCL hydrochloric acid 

Carbonic acid H₂CO₃

Nitric Acid HNO₃ 

2. Organic Acids 

* Obtain from plants and animals 

CH₃COOH ethanoic acid 

C₆ H₈ O₇  Cirtric acid 

C6 H8 O6 Ascorbic Acid 

C4H6O6 Tartarc acid 

* A dilute solution for ethanoic acid is vinegar. Vinegar is made from apples - apple cider vinegar. 

- Concentrated & Dilute

* The concentrated solution contains more moles of solute per volume of solution than dilute ones. To go from a dilute to a concentrated solution you heat it. 

{ BASICITY in acid} 

* This is the number of H+ ions produced per molecule of acid when the acid dissolves in water. 

* A hydrogen ion can also be called a proton because after ionization only one proton is left. 

- The Three Types 

• Monobasic or Monoprotic

• Dibasic or Diprotic 

• Tribasic or Tri protic 

* Examples 

• Monobasic or Monoprotic : Hydrochloric acid (HCL), Nitric acid ( HNO₃ ), Acetic acid 

• Dibasic or Diprotic: Sulphuric acid ( H₂SO₄), Carbonic acid ( H₂SO₃) 

• Tribasic or Tri protic: Phosphoric acid ( H₃PO₄) 

- monobasic

* These produce 1 Hydrogen H+ per molecule when in solution. 

- Dibasic 

* Produces 2 Hydrogen H+ molecules when in solution 

Tribasic 

* Produce 3 Hydrogen H+ molecules when in solution. 

* Monobasic acids can only form 1 salt which is normal salt. Dibasic acid can form 2 salts. Tribasic acid can from 3 salts. 

- Strong & Weak Acids 

* Weak Acids: 

A weak acid is only partially ionized when dissolved in water. 

* Strong Acids: 

A strong acid is fully ionized (dissociated) when dissolved in water. 

Examples of Strong Acid: 

• HCl - hydrochloric acid.

• HNO₃ - nitric acid.

• H₂SO₄ - sulfuric acid.

Examples of Weak Acids: 

• Formic acid  HCOOH

• Acetic acid  CH₃COOH

• Benzoic acid  C₆H₅COOH

• Oxalic acid C₂H₂O₄

• Hydrofluoric acid  HF

• Nitrous acid HNO₂

• Sulfurous acid  H₂SO₃

• Phosphoric acid (chemical formula: H₃PO₄)

- Reactions of Acids

1. Acids & Carbonates 

Acid + Carbonate → salt + water + carbon dioxide 

* In reaction A, all of the sodium carbonate will dissolve once enough H₂SO₄ is added to it,  however, in creation B all of the calcium carbonate will not dissolve because a layer of calcium sulphate will form around the calcium carbonate preventing more H₂SO₄  from reaching the calcium carbonate hence stopping the reaction. 

2. Acid & Metal 

Acid + metal → salt + Hydrogen gas 

* All metals except Cu, Ag and Au will displace the H+ from the acid. 

Reactivity series: 

3. Acid & Metal Oxide 

* Acid + Metal oxide → salt + water

4. Acid & Alkali 

* Acid + Alkali → salt + water 

* All alkalis are soluble except CaH₂

^ BASES 

* Bases are proton acceptors.

* It is an oxide or hydroxide of metal. 

- Examples of Bases 

* Alkalis are soluble bases.  

Examples: KOH & NaOH (NH₄OH is weak) 

- Properties of bases

• Soapy slippery feel 

• Bitter 

• Ph greater than 7 

• Change red litmus paper to blue 

• Conduct electricity when dissolved in water 

• corrosive

- Strong & Weak Alkalis 

* Strong Alkalis are fully ionized in water. Weak Alkalis are not fully ionized in water. 

* Once there is OH- it means it's basic or Alkali but it does not make it strong or weak. The 

majority don't dissolve so they're weak. 

Examples of strong bases: 

• Potassium hydroxide KOH 

• Sodium hydroxide NaOH 

• Barium hydroxide Ba(OH)₂

• Strontium hydroxide Sr(OH)₂

• Calcium hydroxide Ca(OH)₂

• Lithium hydroxide LiOH 

• Rubidium hydroxide RbOH 

Examples of weak bases:

• Al(OH)₃  aluminium hydroxide.

• Pb(OH)₂  lead hydroxide.

• Fe(OH)₃  iron hydroxide.

• Cu(OH)₂ copper hydroxide.

• Zn(OH)₂ zinc hydroxide.

• NH₃ Ammonium 

- Amphoteric Oxides 

* They can react with both acids and strong Alkalis to form salt and water. 

amphoteric oxide or hydroxide + acid → salt + water 

amphoteric oxide or hydroxide + strong alkali → salt + water 

* Alkalis can react with other alkalis

Examples of Amphoteric oxides: 

- Acidic Oxide 

* Acidic oxide reacts with water and produces an acid, it is the oxide of non-metals. 

Acidic oxide + water → acid 

- Neutral Oxide

* Neutral oxide show neither basic nor acidic properties and they do not form salts when reacted with acids or bases. 

Examples: 

Carbon Monoxide CO

Nitrus Oxide N₂O 

Nitric Oxide NO 

^ SALTS 

What is a salt?

*  A salt is a compound formed when some or all of the hydrogen ions in an acid are replaced by metal ammonium ions. 

* Ionic compounds is another name for salts. 

* Zinc hydroxide, and sulphuric acid have positive & negative charges but are not salts. Has to be a metal or ammonium ions. 

- Normal Salts 

* Normal salts are formed when all of the H+ ions in an acid are replaced by metal or ammonium ions. 

Monobasic acids form normal salts. 

Examples: 

KCl   Potassium chloride 

NaCl  Sodium chloride 

FeS0₄  Iron sulphate 

Na₂S0₄  Sodium sulphate

FeCl₂  Iron chloride 

- Acid Salts 

* Acid salts are formed when H+ ions in an acid are only partially replaced by a metal or ammonium ions. 

Examples: 

Ammonium Chloride  NH₄Cl 

Ammonium Sulphate  (NH₄)₂SO₄ 

- Preparation of Salts

* When preparing salts you have to determine if the salt is soluble or insoluble. 

Preparation of Soluble Salts 

A) Acid & Metal carbonate 

Acid + metal carbonate →  salt + water + carbon dioxide 

• Measure 15cm³ of an acid and pour it into a beaker 

• Add excess metal carbonate to the acid 

[ This is to use up all of the acids to replace all of theH+ ions. Forms normal salts. You know when to stop adding when it stops dissolving or effervescence stops. ] 

• Filter the suspension 

• Collect the filter & discard the residue 

• Evapourate the filtrate/solution to about 1/2 the volume

[ This is to make the solution more concentrated by removing water. 

• Leave or allow to cool 

[ So crystallization can take place. ] 

• Filter to collect crystals 

• Wash crystals with small amounts of cold distilled water & leave crystals to dry. 

B) Acid + Metal 

Acid + metal → salt + hydrogen gas 

• Measure 15cm³ of an acid and pour it into a beaker 

• Add excess metal  to the acid 

[ This is to use up all of the acids to replace all of theH+ ions. Forms normal salts. You know when to stop adding when it stops dissolving or effervescence stops. ] 

• Filter the suspension 

• Collect the filter & discard the residue 

• Evapourate the filtrate/solution to about 1/2 the volume

[ This is to make the solution more concentrated by removing water. 

• Leave or allow to cool 

[ So crystallization can take place. ] 

• Filter to collect crystals 

• Wash crystals with small amounts of cold distilled water & leave crystals to dry. 

C) Acid + Metal Oxide 

Acid + Metal Oxide  → salt + water 

• Measure 15cm³ of an acid and pour it into a beaker 

• Add excess metal oxide to the acid 

[ This is to use up all of the acids to replace all of theH+ ions. Forms normal salts. You know when to stop adding when it stops dissolving or effervescence stops. ] 

• Filter the suspension 

• Collect the filter & discard the residue 

• Evapourate the filtrate/solution to about 1/2 the volume

[ This is to make the solution more concentrated by removing water. 

• Leave or allow to cool 

[ So crystallization can take place. ] 

• Filter to collect crystals 

• Wash crystals with small amounts of cold distilled water & leave crystals to dry. 

D) Acid + Alkali 

Acid + Alkali → salt + water 

• Measure 15cm³ of Alkali and use a pipette to add the alkali to a conical flask 

• Add a few drops of indicator to the alkali 

• Fill a burette with acid and record the starting volume 

• Slowly add the acid from the burette to the alkali in the conical flask

• Stop adding the acid when the indicator changes the colour of the solution permanently

• Record the volume of acid 

• fill up a burette with acid 

• Pour alkali into a conical flask 

• Slowly add the acid to the alkali until the end point volume is reached

• filter the suspension 

•  Collect the filter & discard the residue 

• Evapourate the filtrate/solution to about 1/2 the volume [ This is to make the solution more concentrated by removing water. 

• Leave or allow to cool [ So crystallization can take place. ] 

• Filter to collect crystals 

• Wash crystals with small amounts of cold distilled water & leave crystals to dry. 

* You can't have the indicator in the salt so you make over the salt without the indicator. The acid & alkali are colourless solutions so you need the indicator to see a noticeable pH change occurring near the end point of acid-base titrations, an indicator can be used to signal the end of a titration.