8.1 - Acids, Bases and Salts - IGCSE Chemistry

pH and pH scale

- To obtain an idea how acidic or alkaline a substance is, we use universal indicator.

- This indicator is a mixture of many other indicators. Methyl orange, thymolphthalein and blue and red litmus are all indicators used in titrations.

-The colour shown by this indicator can be matched against a pH scale. -- The pH scale was developed by a Scandinavian Chemist called Soren Sorenson.

- The pH scale runs from below 0 to 14. A pH of less than 7 is an acid and greater than 7 is alkaline. A pH of 7 is neutral

More about pH

- Using UI helps determine how concentrated one solution is compared to another.

-The redder the colour, the more acidic. Highly Acidic solutions contain higher concentration of hydrogen ions.

-In a pH meter, a pH electrode is placed into the solution and a pH reading is given on the digital display.

The Bronsted-Lowry theory

-This theory defines an acid as an H+ ion (or proton donor) and a base as an H+ ion (or proton acceptor). Check out:https://byjus.com/jee/bronsted-lowry-theory/

-Solutions of weak acids are poorer conductors of electricity and have slower reactions with metals, bases and metal carbonates.

-All acids when in aqueous solution produce hydrogen ions, H+(aq).

-The strength of an acid tells you how easily it dissociates (ionises) to produce hydrogen ions.

-The concentration of an acid indicates the proportions of water and acid present in aqueous solution. It is still a strong acid when it is in a dilute solution and a weak acid even if it is concentrated.

Neutralisation

-A neutralisation reaction occurs when an acid reacts with an alkali to produce water.

-When these substances react together in a neutralisation reaction, the H+ ions react with the OH– ions to produce water

-For example, when hydrochloric acid is neutralised a sodium chloride and water are produced:

The net ionic equation of all acid-base neutralisations and is what leads to a neutral solution, since water has a pH of 7:

H+ + OH– ⟶ H2O

Neutralisation is very important in the treatment of soils to raise the pH as some crops cannot tolerate pH levels below 7

This is achieved by adding bases to the soil such as limestone and quicklime

Exam Tip: Not all reactions of acids are neutralisations. For example, when a metal reacts with an acid, although a salt is produced there is no water formed so it does not fit the definition of neutralisation.

Ionic equation

-The reaction between any acid and any alkali in aqueous solution produces water and can be summarised by this ionic equation.

- It shows the ion which causes acidity (H+(aq)) reacting with the ion which causes alkalinity (OH-aq)) to produce neutral water (H20(l)).

- Aqueous solutions of acids contain H ions and aqueous solutions of alkalis contain OH-ions.

8.2 Formation of salts

-All sodium , potassium and ammonium salts are soluble

-All nitrates are soluble

-All chlorides except lead and silver are soluble

-All sulfates except barium, calcium and lead are soluble

-All carbonates and hydroxides are insoluble except sodium ,potassium and ammonium

If the acid being neutralised is HCl, salts called chlorides are formed.

Types of oxides

-Non metal oxides such as sulfur dioxide, SO2 and carbon dioxide are acidic.

-In aqueous solution they produce aqueous hydrogen ions,H+(aq)

-Metal oxides are basic oxides

-If these oxides are soluble they will dissolve in water to produce aqueous hydroxide ions, OH-(aq)

-For example, coppper (II) oxide, Cu0 and calcium oxide,Ca0

- Metal oxides which react with both acids as well as bases to produce salts and water are known as amphoteric oxides. Many metals (such as zinc, tin, lead, aluminium, and beryllium) form amphoteric oxides or hydroxides. Amphoteric oxides also include lead (II) oxide, and zinc (II) oxide, among many others.

8.3 Methods of preparing soluble salts

-Excess metal

-Excess insoluble carbonate

-Excess insoluble base

-An alkali by titration

Acid + metal

Acid + metal -> salt +hydrogen

- Used with less reactive metals such as aluminium, the platinum group metals (platinum, palladium, rhodium, iridium, ruthenium, and osmium), and the coinage metals (copper, silver, and gold).

- It would be very dangerous to use a reactive metal, such as sodium.

-Metals used are MAZIT, magnesium, aluminium ,zinc ,iron and tin.

Acid + carbonate

Acid + carbonate -> salt+ water + C02

- Used with any metal carbonate/ any acid, providing the salt produced is soluble

Acid + alkali (soluble base)

-Titration is used for preparing the salts of very reactive metals such as potassium or sodium

-It would be too dangerous to add the metal directly to the acid. So, use an alkali which contains the reactive metal whose salt we wish to prepare

-Most metal oxides and hydroxides are insoluble bases

-A few metal oxides and hydroxides that do dissolve in water to produce OH- (Aq) ions are known as alkalis or soluble bases

-If the metal oxide or hydroxide does not dissolve in water it is known as an insoluble base

Alkalis are soluble bases. Aqueous solutions of alkalis contain OH-ions.

-A base is a substance which neutralises an acid producing a salt and water as the only products.

If the base is soluble, the term alkali can be used nut there are several bases which are insoluble. It is also a substance which accepts a hydrgoen ion.

In general, most metal oxides and hydroxides are bases.

Acid + insoluble base

Acid + base -> salt + water

- Used to prepare a salt of an unreactive metal

-The acid is neutralised using the particular metal oxide

Reaction of a base with ammonium salts

- Small quantities of ammonia gas, NH3 can be produced by heating any ammonium salt such as ammonium chloride with a base such as calcium hydroxide.

-Calcium hydroxide + ammonium chloride -> Calcium Chloride + water + ammonia

- The ammonia produced can be detected as being formed by its pungent odour and by turning damp red litmus blue.

8.4 Preparing Insoluble Salts

-An insoluble salt such as barium sulfate can be made by precipitation.

-In this case, solutions of the 2 chosen soluble salts are mixed.

- To produce barium sulfate, barium chloride and sodium sulfate can be used.

- The barium sulfate precipitate can be filtered off, washed with distilled water and dried.

-Reaction : Barium chloride + sodium sulfate -> Barium sulfate sodium chloride

-Soluble salt + Soluble salt -> Insoluble salt + Soluble salt

8.5 Testing for different salts

Testing for a sulfate

-Sulfate ions in solution, SO 42-, are detected using barium chloride solution.

-The test solution is acidified using a few drops of dilute hydrochloric acid, and then a few drops of barium chloride solution are added.

-A white precipitate of barium sulfate forms if sulfate ions are present.

Testing for a chloride, a bromide or an iodide

-Chlorine, bromine and iodine are halogens. Their ions are called halide ions, eg chloride, Cl –.

Testing for halide ions

-chloride ions give a white precipitate of silver chloride.

-bromide ions give a cream precipitate of silver bromide.

-iodide ions give a yellow precipitate of silver iodide.

Testing for a carbonate

Carbonate ions , CO 32- can be detected whether in a solid compound or in solution. An acid , such as dilute hydrochloric acid, is added to the test compound. Carbon dioxide gas bubbles if carbonate ions are present. Limewater is used to confirm that the gas is carbon dioxide.

Testing for a nitrate

-By adding aqueous sodium hydroxide and then aluminium foil and warming gently, nitrates are reduced to ammonia.

-The ammonia can be identified using damp indicator paper, which turns blue

-In the reaction, the nitrate ion is reduced, as oxygen is removed from the nitrogen atom, and it gains hydrogen to form ammonia,NH3.

-The gain of hydrogen is also a definition of reduction

8.6 Water of crystallisation

-Some salts, such as sodium chloride, copper carbonate and sodium nitrate, crystallise in their anhydrous forms (without water).

-However, many salts produce hydrates when they crystallise from solution

-A hydrate is a salt which incorporates water into its crystal structure

-This water is referred to as water of crystallisation

-A hydrate substance, or hydrate, is one that is chemically combined with water.

-An anhydrous substance is one containing no water

- When many hydrates are heated, the water of crystallisation is driven away.

-For example, if crystals of copper (II) sulfate hydrate (blue) are heated strongly, they lose their water of crystallisation.

-Anhydrous copper (II0 sulfate remains as a white powder:

Copper (II) sulfate pentahydrate -> anhydrous copper (II) sulfate + water

-When water is added to anhydrous copper (II) sulfate, the reverse process occurs.

It turns blue and the pentahydrate is produced

-This is an extremely exothermic process

-Water of crystallisation is the water molecules present in crystals,e.g. CuSO4, 5H2O amd CoCl2, 6H2O

-Because the colour change only takes place in the presence of water, the reaction is used to test for the presence of water

These processes give a simple example of a reversible reaction

Titration

A titration is an experiment where a volume of a solution of known concentration is added to a volume of another solution in order to determine its concentration. Many titrations are acid-base neutralization reactions, though other types of titrations can also be performed.

In order to perform an acid-base titration, the chemist must have a way to visually detect that the neutralization reaction has occurred. An indicator is a substance that has a distinctly different color when in an acidic or basic solution. A commonly used indicator for strong acid-strong base titrations is phenolphthalein. Solutions in which a few drops of phenolphthalein have been added turn from colorless to brilliant pink as the solution turns from acidic to basic. The steps in a titration reaction are outlined below.

A measured volume of an acid of unknown concentration is added to an Erlenmeyer flask.

Several drops of an indicator are added to the acid and mixed by swirling the flask.

A burette is filled with the base solution of known molarity.

The stopcock of the burette is opened and base is slowly added to the acid while the flask is constantly swirled to insure mixing. The stopcock is closed at the exact point at which the indicator just changes color.

Phenolphthalein turns pink in basic solutions

Figure 1. Phenolphthalein in basic solution.

The standard solution is the solution in a titration whose concentration is known. In the titration described above the base solution is the standard solution. It is very important in a titration to add the solution from the burette slowly so that the point at which the indicator changes color can be found accurately.

The end point of a titration is the point at which the indicator changes color. When phenolphthalein is the indicator, the end point will be signified by a faint pink color.