Acids

Definition

• Arrhenius Acid → An acid is a substance in which its molecules dissociates in water to produce $\ce{H+(aq)}$ ions (or $\ce{H3O+(aq)}$ ions)
  • In aqueous solutions $\ce{H+}$ ions bond to water molecules to form $\ce{H3O+}$ ions (hydronium or hydroxonium ions), both referred to as solvated $\ce{H+}$ ions
• Bronsted-Lowry acid –> $\ce{H+}$ (proton) donor #smtp
• Lewis acid –> electron pair acceptor #smtp

Basicity of acids

• Monobasic acids –> 1 molecule of the acid yields 1 $\ce{H+}$ ion when dissociated
• Dibasic acids –> 1 molecule of the acid yields 2 $\ce{H+}$ ions when dissociated
• Tribasic acids –> 1 molecule of the acid yields 3 $\ce{H+}$ ions when dissociated

Strength

• The strength of an acid depends on the extent of its dissociation in water
  • A strong acid is one in which its molecules dissociate completely in water to give a high concentration of hydrogen ions ($\ce{H+}$)
  • A weak acid is one in which its molecules dissociate partially in water to give a low concentration of hydrogen ions ($\ce{H+}$)

concentration

• A concentrated acid solution that contains a large amount of the dissolved acid
• A dilute acid solution that contains a small amount of the dissolved acid
• The concentration of $\ce{H+}$ ions in an aqueous solution of an acid is measured by its pH
  • $pH=-\log_{10}[\ce{H+}]$

Properties

• Acids have a sour taste
• Acids dissolve in water to form solutions that conduct electricity
  • Prescence of mobile ions. E.g. in dilute $\ce{HCl}$, $\ce{HCl}$ dissociate completely in water to form mobile $\ce{H+}$ ions and $\ce{Cl-}$ ions, which act as charge carriers to conduct electricity
• Acids turn blue litmus paper red
• Acids have a pH less than 7

Reactions

Acid-metal

• Acids react with reactive metals (above hydrogen in the reactivity series) to form hydrogen gas and salt $\ce{\text{metal}+\text{acid}->\text{salt} + \text{hydrogen}}$

reactivity series

$$$$

\begin{array}{ll} K & \text{Potassium} \ Na & \text{Sodium} \ Ba & \text{Barium} \ Ca & \text{Calcium} \ Mg & \text{Magnesium} \ Al & \text{Aluminium} \ \mathbf{C} & \text{\textbf{Carbon}} \ Zn & \text{Zinc} \ Fe & \text{Iron} \ Sn & \text{Tin} \ Pb & \text{Lead} \ \mathbf{H} & \text{\textbf{Hydrogen}} \ Cu & \text{Copper} \ Hg & \text{Mecury} \ Ag & \text{Silver} \ Au & \text{Gold} \ Pt & \text{Platinum} \end{array}

**Hydrogen** and **Carbon** are not metals and are included for comparison Oxides of metals below Carbon can be extracted using carbon; e.g. $$\ce{2ZnO + C -> Cu +H2O}$$ Oxides of metals below Hydrogen can be extracted using hydrogen gas $$\ce{CuO + H2 -> Cu + H2O}$$ ^185da6

Acid-Carbonate

• Acids react with all carbonates to form a salt, carbon dioxide and water
  • A salt is the ionic compound formed when hydrogen ions from an acid are displaced by metal or ammonium ions The anion of salt comes from an acid $\ce{\text{carbonate} + \text{acid} -> \text{salt} + \text{water} + \text{carbon dioxide}}$

Acid-Oxide/Hydroxide

• Acids react with metal oxides and hydroxides to form a salt and water $\ce{\text{metal oxide} + \text{acid} -> \text{salt} + \text{water}}$ $\ce{\text{metal hydroxide} + \text{acid} -> \text{salt} + \text{water}}$

Neutralisation

• Neutralisation is the combination of hydrogen ions from an acid with the hydroxide ions from an alkali or base to form water molecules and salt as the only products $\ce{H+(aq) + OH-(aq) -> H2O(l)}$

Uses of Acids

• Sulfuric acid is used in
  • making detergents
  • making fertilisers
  • car batteries
• hydrochloric acid is used
  • in leather processing
  • for cleaning metals
• Ethanoic acid is used
  • to preserve food in vinegar
  • in making adhesives such as glue

Bases and Alkalis

• A base is a substance that dissociates in water to produce $\ce{OH-}$ ions
  • Most bases are metal oxides or metal hydroxides
• An alkali is a base which is soluble in water

Strength

• A strong base/alkali is a substance that dissociates completely in water to give a high concentration of hydroxide ions ($\ce{OH-}$)
• A weak base/alkali is a substance that dissociates partially in water to give a low concentration of hydroxide ions ($\ce{OH-}$)

Acidity of base

• The acidity of a base is the number of ionisable hydroxide ions that can be produced by a formula unit of the base
• Monoacidic base: 1 formula unit of the alkali yields 1 $\ce{OH-}$ ion when dissociated
• Diacidic base: 1 formula unit of the alkali yields 2 $\ce{OH-}$ ions when dissociated

Properties

• Alkalis have a pH more than 7
• Alkalis have a bitter taste and soapy, slippery feel
• Alkalis dissolve in water to form solutions that conduct electricity
  • Presence of mobile ions. E.g. in aqueous $\ce{NaOH}$, $\ce{NaOH}$ dissociates completely to form mobile $\ce{Na+}$ and $\ce{OH-}$ ions, which act as charge carriers
• All alkalis produce hydroxide ions when dissolved in water

Ammonia

$\ce{NaOH(aq) -> Na+(aq) + OH-(aq)}$ Ammonium hydroxide/oxide does not exist and is known as ammonia solution

Reactions

Precipitation

• Alkalis when mixed with a salt solution, can produce an insoluble metal hydroxide and a soluble metal salt through a precipitation reaction; e.g. $\ce{2NaOH(aq) + CuSO4(aq) -> Na2SO4aq) + Cu(OH)2(s)}$
  • A precipitation reaction is reaction in an aqueous solution where two soluble ionic compounds combine to form an insoluble solid; e.g. $\ce{AgNO3​(aq) + NaCl(aq) -> AgCl(s) + NaNO3​(aq)}$

Alkali-Ammonium salt

• Alkalis react with ammonium salts on warming to displace ammonia gas $\ce{\text{ammonium salt} + \text{alkali} ->[heat] \text{salt} + \text{water} + \text{ammonia}}$ $\ce{NH4Cl(s) + NaOH(aq) ->[heat] NaCl(aq) + H2O(l) + NH3(g)}$

Uses of bases and alkali

• Ammonia solution is used
  • In window cleaning solutions
  • In fertilizers
• Calcium oxide (Quicklime) is used
  • In neutralizing acidic soil
  • To make iron, concrete and cement
• Magnesium hydroxide is used:
  • In toothpaste to neutralize acid on teeth
  • In antacids, to relieve gastric problems and indigestion
• Sodium hydroxide is used
  • In making soaps and detergents
  • In industrial-cleaning detergents

pH scale

• It measures the degree of alkalinity or acidity of an aqueous solution
• the pH of pure water at $25^\circ C$ is 7
• Decreasing pH is increasing acidity
• Increasing pH is increase alkalinity

pH and pOH

• [$\ce{H+}$] > [$\ce{OH^-}$] ≥ acidic
• [$\ce{OH^-}$] > [$\ce{H+}$] ≥ basic
• $pH = -\log([\ce{H+}])$
• $pOH = -\log([\ce{OH-}])$
• $pH + pOH = pK_{w}\quad \text{(14 at 25C)}$

pH of water at varying temperatures

pH of water may be 6.5, although water is neutral \ce{H2O(l) ≤> OH-(aq) + H+(aq)}

at $25^\circ C$, amount of $\ce{H^+}$ ions and $\ce{OH^-}$ ions are $10^{-7}moldm^{-3}$ for both

pH is only based on scale at $25^\circ C$

if temperature is increased to $80^\circ C$ more water will ionise to form ions, increasing the concentration of $\ce{H+}$ ions

Indicators

Oxides

• Metal Oxides
• Non-metal oxides
  • acidic oxides

Metal Oxides

Basic Oxides

• Most are insoluble
• Solids at room temperature

limewater

limewater is formed by dissolving $\ce{CaO}$ in water

$\ce{CaO(s) + H2O(l) -> Ca(OH)2(aq)}$

Reactions

• react with acids to form a salt and water; e.g. $\ce{Cao(s) + 2HNO3(aq) -> Ca(NO3)2(aq) + H2O(l)}$

Amphoteric Oxides

• All are insoluble in water and produce a pH of 7 when tested

Reactions

• Amphoteric oxides react with both acids and bases to form salts and water

Example

$\ce{ZnO(s) + 2NaOH(aq) -> Na2ZnO2(aq) + H2O(l)}$ $\ce{PbO(s) + 2NaOH(aq) -> Na2PbO2(aq) + H2O(l)}$ $\ce{Al2O3(s) + 2NaOH(aq) -> 2NaAlO2(aq) + H2O(l)}$ $\ce{ZnO(s) + 2HCl(aq) -> ZnCl2(aq) + H2O(l)}$ $\ce{PbO(s) + 2HNO3(aq) -> Pb(NO3)2(aq) + H2O(l)}$ $\ce{Al2O3(s) + 6HCl(aq) -> 2AlCl3(aq) + 3H2O(l)}$

Non-metal Oxides

Acidic Oxides

• Most dissolve in water to form acids

Example

Acidic Oxide	Formula	Physical state	Acid produced in water - Name	Acid produced in water - Formula
Carbon dioxide	$\ce{CO2}$	Gas	Carbonic acid	$\ce{H2CO3}$
Sulfur dioxide	$\ce{SO2}$	Gas	Sulfurous acid	$\ce{H2SO3}$
Sulfur trioxide	$\ce{SO3}$	Gas	Sulfuric acid	$\ce{H2SO4}$
Phosphorus(V) oxide	$\ce{P4O10}$	Solid	Phosphoric acid	$\ce{H3PO4}$

Reactions

• Acidic oxides react with alkalis to form a salt and water; e.g. $\ce{SO4(g) + 2NaOH(aq) -> Na2SO4(aq) + H2O(l)}$

Salts

• A salt is a compound produced alongside water when an acid reacts with a base, specifically when a metallic or ammonium ion replaces one or more hydrogen ions in the acid
  • e.g. sulfuric acid forms sulfates, hydrochloric acid forms chlorides, and nitric acid forms nitrates.
• Types of Salts:
  • Normal Salts: Formed through the full replacement of the hydrogen ions in an acid.
  • Acid Salts: Formed through the partial replacement of hydrogen ions, which can occur in polyprotic(diabasic, tribasic) acids like $H_2CO_3$, $H_2SO_4$, and $H_3PO_4$

Water of crystallisation

• Salts which combine with water to form crystals are known as hydrated salts
  • Water molecules in the crystal are known as water of crystallisation
• Salts which do not contain water molecules are known as anhydrous salts

Example

Name of Salt	Anhydrous salt	Hydrated salt
Sodium carbonate	$\ce{Na2CO3}$	$\ce{Na2CO3.10H2O}$
Copper (II) sulfate	$\ce{CuSO4}$	$\ce{CuSO4.5H2O}$
Zinc sulfate	$\ce{ZnSO4}$	$\ce{ZnSO4.7H2O}$
Magnesium sulfate	$\ce{MgSO4}$	$\ce{MgSO4.7H2O}$

• When a hydrated salt is heated, its water of crystallisation is driven off and the anhydrous salt is formed $\ce{CuSO4*5H2O(s) ->[heat] CuSO4(s) + 5H2O(l)}$

Uses of salts

Salt	Use
Ammonium chloride	In torch batteries
Ammonium nitrate	In fertilisers
Calcium carbonate	Extraction of iron, making cement, glass making
Calcium chloride	In the extraction of sodium, drying agent (anhydrous)
Calcium fluorite	Added to toothpaste to make teeth and bones stronger
Calcium sulfate	For making plaster boards, plaster casts for injured limbs
Iron(II) sulfate	In 'iron' tablets
Lead(II) chromate	Bright yellow substance in "no parking lines"
Magnesium sulfate	In medicines
Potassium nitrate	In fertiliser and explosives manufacture
Silver bromide	In film photography
Sodium carbonate	Glass making, softening water, making modern washing powders
Sodium chloride	Making hydrochloric acid, for food flavouring, hospital saline, in the Solvay process for the manufacture of sodium carbonate
Sodium stearate	In some soaps
Tin(II) fluoride	Additive to toothpaste

Preparation of salts

• Insoluble Salts
  • pirecipitation method
• soluble salts
  • not containing group 1 or ammonium cations (non-SPA salts)
    • filtration method
    • containing group 1 or ammonium cations (SPA salts)
    • titration method

Precipitation method

^39828e

• To prepare an insoluble salt, mix a soluble salt containing its cations and a soluble salt containing its anions
• steps
  1. Add one aqueous solution to the other aqueous solution
  2. Stir thoroughly until no more precipitate forms
  3. Filter the mixture and collect the residue, which contains the precipitate
  4. Discard the filtrate
  5. Wash the residue with deionised water to remove unwanted reactants
  6. Dry the residue between sheets of filter paper

Coating

• When you are forced to start with an insoluble reagent, you have to use two equations
• This is as a direct reaction would produce an insoluble salt which would coat the original insoluble reagent, preventing further reaction

Example

prepare solid copper(II) carbonate using $\ce{CuO}$ as one of the starting reagents

$$$$

\begin{align*} \ce{CuO(s) + 2HNO3(aq) &-> Cu(NO3)2(aq) + H2O(l))} \ \ce{Cu(NO3)2(aq) + Na2CO3(aq) &-> CuCO3(s) + NaNO3(aq)} \end{align*}

$$$$

Filtration Method ^447e2f

• Prepare an acid, containing the anion, and an insoluble reactant:
  • metal (ensure that it is below calcium on the reactivity series)
    • e.g. $\ce{K}$, $\ce{Na}$, $\ce{ba}$, $\ce{Ca}$ cannot be used as they are too reactive
  • insoluble base (ensure that it is completely insoluble)
    • e.g. $\ce{CaO}$ and $\ce{CaOH}$ cannot be used as they are sparingly soluble
  • insoluble carbonate (ensure it exists)
    • e.g. $\ce{Al2(CO3)2}$ does not exist; any ion with charge of $3^+$ is incompatible with large anion such as carbonate as thye are unstable and will decompose
• steps
  • Add excess insoluble reactant to warm dilute acid and stir until no more reacts
  • Filter the mixture to obtain the filtrate
  • Heat the filtrate till the point of saturation
  • Test for saturation by dipping glass rod into solution
  • Cool the filtrate to allow for crystals to form
  • Filter to obtain the crystals as residue
  • Wash with a small amount of cold deionised water
  • Dry crystals between pieces of filter paper

Titration method ^e226f1

• Soluble SPA salts can be prepared using an acid containing the anion and the SPA alkali, containing the cation
• An acid-base indicator such as methyl orange or phenolphthalein must be used to know when neutralisation is complete
• steps (e.g. using dilute $\ce{KOH(aq)}$ and dilute $\ce{H2SO4(aq)}$)
  • Pipette 25.0 cm3 of dilute sulfuric acid into a conical flask
  • Add 2 to 3 drops of an indicator e.g. phenolphthalein or methyl orange into the conical flask
  • Using a burette filled with the aqueous potassium hydroxide, titrate the sulfuric acid with the aqueous potassium hydroxide until the endpoint is reached, where one drop of the alkali causes the indicator to change colour permanently
  • Record the volume of the aqueous potassium hydroxide used
  • Repeat the titration without the indicator by adding the same volume of the aqueous potassium hydroxide from the burette as in the previous titration
  • Transfer the aqueous potassium sulfate solution to a crucible and heat the solution to evaporate the water to form a saturated solution
    • Test for saturation by dipping a glass rod into the solution and if crystals are formed on the cool rod, no further heating would be required
  • Allow the hot saturated solution to cool to room temperature so that the potassium sulfate salt can crystallise
  • Filter the solution to obtain the potassium sulfate crystals and rinse the crystals with a small amount of cold deionised water
  • Dry the potassium sulfate crystals by pressing between a few sheets of filter paper