ACIDS, bases and salts Flashcards
Properties of acids
Acids have pH values of below 7, have a sour taste and are corrosive.
In acidic conditions, blue litmus paper turns red and methyl orange indicator
turns red.
Acids are substances that can neutralise a base, forming a salt and water.
When acids react, they will lose electrons to form positively charged hydrogen ions (H+).
The presence of H+ ions is what makes a solution acidic.
Acids and metals
Only metals above hydrogen in the reactivity series will react with dilute acids.
When acids react with metals they form a salt and hydrogen gas:
Acids with Bases (Alkalis)
Metal oxides and metal hydroxides can act as bases.
When they react with acid, a neutralisation reaction occurs.
Acids and bases will react together in a neutralisation reaction and produce a salt and water:
Examples of Reaction Between Acids and Bases:
Acids with Metal Carbonates
Acids will react with metal carbonates to form the corresponding metal salt, carbon dioxide and water:
Acid + Metal Carbonate → Salt + Carbon Dioxide + Water
Examples of Reaction Between Acids and Bases:
Properties of bases
Bases have pH values of above 7.
A base which is water soluble is referred to as an alkali.
In basic (alkaline) conditions red litmus paper turns blue and methyl orange indicator turns yellow.
Bases are substances which can neutralise an acid, forming a salt and water.
Bases are usually oxides or hydroxides of metals.
When alkalis react, they gain electrons to form negative hydroxide ions (OH–).
The presence of the OH– ions is what makes the aqueous solution an alkali.
Bases and acids
When they react with an acid, a neutralisation reaction occurs.
Acids and bases react together in a neutralisation reaction and produce a salt and water:
Alkalis and ammonium salts
Ammonium salts undergo decomposition
when warmed with an alkali.
Even though ammonia is itself a weak base, it is very volatile and can easily by displaced from the salt by another alkali.
A salt, water and ammonia are produced.
Example:
NH4Cl + NaOH → NaCl + H2O + NH3
This reaction is used as a chemical test to confirm the presence of the ammonium ion (NH4+).
Alkali is added to the substance with gentle warming followed by the test for ammonia gas using damp red litmus paper.
The litmus paper will turn from red to blue if ammonia is present.
The pH scale
The pH scale is a numerical scale which is used to show how acidic or alkaline a solution is.
It goes from 1 – 14 (extremely acidic substances can have values of below 1).
All acids have pH values of below 7, all alkalis have pH values of above 7.
The lower the pH then the more acidic the solution is.
The higher the pH then the more alkaline the solution is.
A solution of pH 7 is described as being neutral e.g. water.
Universal indicator
Universal indicator is a mixture of different indicators which is used to measure the pH.
A drop is added to the solution and the colour is matched with a colour chart which indicates the pH which matches specific colours.
The pH scale with the Universal Indicator colours which can be used to determine the pH of a solution
The importance of pH and soil acidity
Soil pH is analysed to indicate the acidity or alkalinity of soil.
Most plants favour a pH value of between 5.5 and 8.
Changes in soil which cause a pH to be outside this range adversely affect plant processes resulting in reduced growth and crop yield.
Soils may become acid from acid rain, overuse of fertilisers which contain ammonium salts or by the excessive breakdown of organic matter by bacteria.
Crushed or powdered limestone (calcium carbonate) or lime (calcium oxide) or slaked lime (calcium hydroxide) is added to neutralise the excess acidity in the soil.
The addition process must be carefully monitored though, as if added in excess, further damage could be done if the pH goes too high.
Extended Subject Content
Proton transfer
The earlier definition of an acid and a base can be extended.
In terms of proton transfer, we can further define each substance in how they interact with protons.
Acids
Acids are proton donors as they ionize in solution producing protons, H+ ions.
These H+ ions make the aqueous solution acidic.
Bases (Alkalis)
Bases (alkalis) are proton acceptors as they ionize in solution producing OH– ions which can accept protons.
These OH– ions make the aqueous solution alkaline.
Strong acids and bases
Acids and alkalis can be either strong or weak, depending on how many ions they produce when dissolved in water.
Strong acids and bases ionize completely in water, producing solutions of very low pH for an acid or very high pH for a base.
Strong acids include HCl and H2SO4 and strong bases include the Group I hydroxides.
Weak acids and bases
Weak acids and bases partially ionize in water and produce pH values which are closer to the middle of the pH scale.
Weak acids include organic acids such as ethanoic acid, CH3COOH and weak bases include aqueous ammonia.
For both weak acids and bases, there is usually an equilibrium set-up between the molecules and their ions once they have been added to water.
Example of a weak acid: propanoic acid
In both cases the equilibrium lies to the … indicating a / base molecules, with a concentration of ions in solution.
- left,
- high concentration of intact acid
- low
Effect of concentration on strong and weak acids and alkalis
A concentrated solution of either an acid or a base is one that contains a high number of acid or base molecules per dm3 of solution.
It does not necessarily mean that the acid or base is strong though, as it may be made from a weak acid or base which does not dissociate
For example a dilute solution of HCl will be more acidic than a concentrated solution of ethanoic acid, since most of the HCl molecules dissociate but very few of the CH3COOH do.
In acid-base chemistry, the terms strong and weak refer to the…
If referring to concentration when answering a question, then the words … should be used.
- ability to dissociate and produce H+/OH– ions.
- concentrated or dilute
Acid and basic oxides general
Acidic and basic oxides have different properties and values of pH.
The difference in their pH stems from whether they are bonded to a metal or a nonmetal element.
The metallic character of the element influences the acidic or alkaline behaviour of the molecule.
Acidic oxides
Acidic oxides are formed when a nonmetal element combines with oxygen.
They react with bases to form a salt and water.
When dissolved in water they produce an acidic solution with a low pH.
Common examples include SO2 and SiO2.
Basic oxides
Basic oxides are formed when a metal element combines with oxygen.
They react with acids to form a salt and water.
When dissolved in water they produce a basic solution with a high pH.
Common examples include NaOH, KOH and Ca(OH)2.
Neutral oxides
Some oxides do not react with either acids or bases and thus are said to be neutral.
Examples include N2O, NO and CO.
Amphoteric oxides
Amphoteric oxides are a curious group of oxides that can behave as both acidic and basic, depending on whether the other reactant is an acid or a base.
In both cases a salt and water is formed.
Two most common amphoteric oxides are zinc oxide and aluminum oxide.
The hydroxides of both of these elements also behave amphoterically.
Example of aluminium oxide behaving as a base:
Salts
A salt is a compound that is formed when the hydrogen atom in an acid is replaced by a metal.
For example if we replace the H in HCl with a potassium atom, then the salt potassium chloride is formed, KCl.
Salts are an important branch of chemistry due to the varied and important uses of this class of compounds.
These uses include fertilisers, batteries, cleaning products, healthcare products and fungicides.
Naming salts
The name of a salt has two parts.
The first part comes from the metal, metal oxide or metal carbonate used in the reaction.
The second part comes from the acid.
The name of the salt can be determined by looking at the reactants
For example hydrochloric acid always produces salts that end in chloride and contain the chloride ion, Cl–.
Other examples:
Sodium hydroxide reacts with hydrochloric acid to produce sodium chloride.
Zinc oxide reacts with sulfuric acid to produce zinc sulfate.
Preparing salts
Some salts can be extracted by mining but others need to be prepared in the laboratory.
There are two key ideas to consider when preparing salts:
Is the salt being formed soluble or insoluble in water?
Is there water of crystallisation present in the salt crystals?
Preparing soluble salts
Method A: adding acid to a solid metal, base or carbonate.
Method:
Add dilute acid into a beaker and heat using a bunsen burner flame.
Add the insoluble metal, base or carbonate, a little at a time, to the warm dilute acid and stir until the base is in excess (i.e. until the base stops disappearing and a suspension of the base forms in the acid).
Filter the mixture into an evaporating basin to remove the excess base.
Heat the solution to evaporate water and to make the solution saturated.
Check the solution is saturated by dipping a cold, glass rod into the solution and seeing if crystals form on the end.
Leave the filtrate in a warm place to dry and crystallize.
Decant excess solution and allow crystals to dry.
Preparation of Pure, Hydrated Copper (II) Sulfate Crystals using Method A
Method:
Add dilute sulfuric acid into a beaker and heat using a bunsen burner flame.
Add copper (II) oxide (insoluble base), a little at a time to the warm dilute sulfuric acid and stir until the copper (II) oxide is in excess (stops disappearing).
Filter the mixture into an evaporating basin to remove the excess copper (II) oxide.
Leave the filtrate in a warm place to dry and crystallize.
Decant excess solution.
Blot crystals dry.
Equation Of Reaction:
Copper (II) Oxide + Dilute Sulfuric Acid → Copper (II) Sulphate + Water
CuO (s) H2SO4 (aq) CuSO4 (s) H2O (l)
Method B: reacting a dilute acid and alkali.
Method:
Use a pipette to measure the alkali into a conical flask and add a few drops of indicator (phenolphthalein or methyl orange).
Add the acid into the burette and note the starting volume.
Add the acid very slowly from the burette to the conical flask until the indicator changes to appropriate colour.
Note and record the final volume of acid in burette and calculate the volume of acid added (starting volume of acid – final volume of acid).
Add this same volume of acid into the same volume of alkali without the indicator.
Heat to partially evaporate, leaving a saturated solution.
Leave to crystallise decant excess solution and allow crystals to dry.
Preparing Insoluble Salts
Insoluble salts can be prepared using a precipitation reaction.
The solid salt obtained is the precipitate, thus in order to successfully use this method the solid salt being formed must be insoluble in water.
Using Two Soluble Reactants
Method:
Dissolve soluble salts in water and mix together using a stirring rod in a beaker.
Filter to remove precipitate from mixture.
Wash filtrate with water to remove traces of other solutions.
Leave in an oven to dry.
Preparation Of Pure, Dry Lead (II) Sulfate Crystals using a precipitation reaction
Soluble Salt 1 = Lead (II) Nitrate Soluble Salt 2 = Potassium Sulfate
Method:
Dissolve Lead (II) Nitrate and Potassium Sulfate in water and mix together using a stirring rod in a beaker.
Filter to remove precipitate from mixture.
Wash filtrate with water to remove traces of potassium nitrate solution.
Leave in an oven to dry.
Equation of Reaction:
Lead (II) Nitrate + Potassium Sulfate → Lead (II) Sulfate + Potassium Nitrate
Pb(NO3)2 (s) K2SO4 (s) PbSO4 (s) 2KNO3 (s)
Selecting a Method of Preparation
When deciding the method of preparation, if is important to first know whether the salt being produced is soluble or insoluble.
If it is soluble than it can be prepared using either method (A or B) for preparing a soluble salt.
If it is insoluble then it must be prepared using by precipitation.
Test for aqueous cations
Metal cations in aqueous solution can be identified by the colour of the precipitate they form on addition of sodium hydroxide and ammonia.
If only a small amount of NaOH is used then normally the metal hydroxide precipitates.
In excess NaOH some of the precipitates may dissolve.
A few drops of NaOH is added at first and any colour changes or precipitates formed are noted.
Then the NaOH is added in excess and the reaction is observed again.
The steps are repeated for the test using ammonia solution.
Analysing results
The table below contains the results for each of the cations included in the syllabus.
If a precipitate is formed from either NaOH or aqueous ammonia then the hydroxide is insoluble in water.
Zinc for example reacts as such:
ZnCl2(aq) + 2NaOH(aq) →Zn(OH)2 + H2O(l)
Ca2+ ions can be distinguished from Zn2+ and Al3+ as CaOH calcium hydroxide precipitate does not dissolve in excess NaOH but both zinc hydroxide and aluminium hydroxide do.
Zn2+ ions can be distinguished from Al3+ ions as ZnOH dissolves in excess aqueous ammonia but Al(OH)3 does not.
Most transition metals produce hydroxides with distinctive colours.
Be sure to distinguish between the term “colourless” and “clear”.
A solution that loses its colour has become…
A clear solution is one that…
Solutions can be clear and have colour e.g….
- colourless.
- you can see through such as water.
- dilute copper sulphate.
Tests for cations
The flame test
is used to identify the metal cations by the colour of the flame they produce.
A small sample of the compound is placed on an unreactive metal wire such as nichrome or platinum.
The colour of the flame is observed and used to identify the metal.
When it comes to qualitative inorganic analysis, always remember that there will be:
a test for the metal cation part of the molecule
and another test for the anion part.
Tests for gases
It is easy to confuse the tests for hydrogen and oxygen. Try to remember that:
a ligHted splint has an H for Hydrogen, while a glOwing splint has an O for Oxygen.
