Chapter 1 - Chemical Reactions and Equations
Chemical change: - is a change in which one or more new substances are formed.In a chemical change—
• New substances are formed. • Energy changes are involved.
• There is a change in mass during the reaction. • Permanent change takes place.
Examples –
❖ Cooking of food ❖ Rusting of iron
❖ Heating of Lead nitrate ❖ Souring of milk
❖ Ripening of fruit.
Rusting of iron is a chemical change because ❖ A new substance iron oxide is formed.
❖ The change is permanent; the article has got a rust layer (which may only peal off). ❖ There is an increase in mass when rust forms.
❖ An energy change has taken place (which may not be visible). Chemical changes are also known as chemical reactions.
Exothermic Reaction - A chemical reaction which is accompanied by evolution of heat energy
is known as exothermic reaction. Examples:
The amount of heat (energy) produced is written along with the products. This indicates that heat is given out.
Endothermic Reactions - A chemical reaction which is accompanied by absorption of heat
energy is known as exothermic reaction. Examples:
The amount of heat (energy) produced is written along with the reactants. This indicates that heat is absorbed.
Chemical Formulae - The chemical formula of a substance is the symbolic representation of the
actual number of atoms present in one molecule of that substance.
(a) Formula of one molecule of water is H2O. It shows that one molecule of water is made up
of 2 atoms of hydrogen and one atom of oxygen.
(b) Formula of one molecule of sulphuric acid is H2SO4. It shows that one molecule of
sulphuric acid is made up of 2 atoms of hydrogen, 1 atom of Sulphur and 4 atoms of oxygen.
Chemical Equations - Representation of a chemical change in terms of symbols and formulae of
the reactants and products is known as chemical equation of the reaction. Example
+ + + +
Steps for writing Chemical Equation Step I: Writing skeletal Equation
A word-equation shows change of reactants to products through an arrow placed between them ( ). The arrow may be read as "to yield" or "to form" or "to give" and shows the direction of the reaction.
The reactants are placed on the left hand side (LHS) of the arrow and the products on the right hand side (RHS). The different reactants as well as products are connected by a plus sign (+). A complete chemical equation represents the reactants, products and their physical states symbolically and is a balanced account of a chemical transaction.
Step II: Balancing of Chemical Equation
An equation in which number of atoms of each element is equal on both the sides of the equation is known as balanced chemical equation.
A chemical equation is balanced so that the numbers of atoms of each type involved in a chemical reaction are the same on the reactant and product sides of the equation.
Equations must always be balanced.
Whenever H2O is present on any side, the number of hydrogen on both the sides should be an
even number (2 atoms of hydrogen in water). If there are 4H2O, then 4 × 2 = 8 hydrogen atoms.
On the reactant side, there must be an even number in front of HCl. (What that number is, we will find out later). As a result, the number of chlorine atoms will also be even. But on the product side, the number of chlorine atoms is odd (i.e., KCl = 1, MnCl2= 2, Cl2 = 2. i.e., 1 + 2 +
2 = 5). The only odd number of chlorine atoms is in KCl. Let us change it into the simplest even number possible i.e., 2.
Since number of K atoms in 2 KCl = 2, place 2 in front of KMnO4 to balance K atom.
In 2 KMnO4, there are 2 K, 2 Mn, and 8 O. So add these numbers in front of K, Mn and O, (K is
already done).
If there are 8 H2O on the product side, there should be 16 H (8 x 2) on the reactants side as well.
Now the only unbalanced one is Chlorine. On the left hand side, there are 16 Cl. On the right hand side, firstly, there are 2 Cl in 2 KCl + 4 Cl in 2 MnCl2, making total of 6(2 + 4).
Subsequently 10 more Cl atoms are to be accounted for. So place 5 in front of Cl2 to make it 10
(5 x 2).
This type of balancing the chemical equations is known as the Hit and trial method.
Information conveyed by a chemical equation
1. Names of various reactants and products 2. Formulae of reactants and products
3. Relative number of moles of the reactants and products 4. Relative masses of reactants and products
5. Relative volumes of gaseous reactants and products
Limitation of a Chemical Equation
• It does not mention the state of the substances. Accordingly, the following symbols should be added to make it informative: (s) for solid, (l) for liquid, (aq) if the reactant or product is present as a solution in water, (g) for gas and (vap) for vapour.
Example: CaCO3 (s)+ HCl (aq) CaCl2 (s) + H2O (l) + CO2(g)
• The reaction may or may not be complete. An equation does not reveal this. • It does not give any information regarding the speed of the reaction.
• It does not give the concentration of the substances. In some cases, terms like diluted (dil) and concentrated (conc) may be added.
• It does not give the conditions of temperature, pressure, catalyst, etc. This is overcome by mentioning these above or below the arrow. e.g.
• It does not give any idea about color changes, which has to be mentioned separately.
• It does not give any indication regarding the production or absorption of heat. This is mentioned separately.
C + O2 CO2 + Heat
2C + O2 2CO + Heat
• Some reactions are reversible. They are represented by or
• Chemical reactions that proceed with evolution of heat energy, that is, in which heat is given out along with the product, are called exothermic reactions.
• Chemical reactions that proceed with the absorption of heat energy are called endothermic reactions.
Balancing of simple Chemical Equations - The number of atoms of each element should
remain same before and after the reaction.
Balancing: To make the number of atoms of all the elements equal on both the sides in a skeletal equation. A simple equation is balanced by Hit and trial method.
Steps involved in balancing a Chemical Equation: Step 1: Write the correct skeleton equation.
Step 2: Start with the compound that has the maximum atoms or maximum kinds of atoms and
the atoms present in it are balanced first.
Step 3: Balance elements that appear only once on each side of the arrow first. Then balance
elements that appear more than once on a side.
Step 4: Elementary substances are balanced last of all.
Step 5: If required the whole equation is multiplied by some suitable number in order to make
all the coefficients whole numbers.
Balancing of Ionic Equations:
A balanced ionic equation must satisfy mass as well as charge balance.
Calculations based on Chemical Equations:
We can get a lot of quantitative information from a chemical reaction. Number of moles of reactants and products can be calculated from a chemical reaction.
Types of Reactions
Combination or Synthesis Reactions:
The reactions in which two or more substances combine to form a single new substance. Types of Combination reactions:
1. Combination of two elements to form a compound
• Burning of hydrogen in air or oxygen to produce water.
2. Combination Reactions involving an Element and a Compound • Burning of carbon monoxide in oxygen to form carbon dioxide. 2CO (g) + O2(g) 2CO2 (g)
3. Combination Reactions involving Two Compounds
• Combination of ammonia and hydrogen chloride to produce ammonium chloride.
Decomposition reactions: are opposite to combination reactions. In a decomposition reaction a
compound breaks down into two or more simple substances by the application of heat or electricity.
When a substance decomposes due to heat it is called thermal decomposition, while decomposition due to electricity, is called electrolytic decomposition.
Electrolysis: The decomposition of a substance by passing electric current through it is called
electrolysis.
Photolysis: The decomposition of a compound with light is called photolysis.
1. Mercuric oxide, when heated, undergoes thermal decomposition, to give mercury and oxygen.
2. Similarly, if blue crystals of copper nitrate are heated, they undergo thermal decomposition to give black colored copper oxide, reddish brown fumes of nitrogen dioxide, and a colorless gas of oxygen.
3. When water taken in an electrolytic cell, is acidified with a small quantity of sulphuric acid and a direct current passed through it undergoes electrolytic decomposition to yield hydrogen and oxygen.
2𝐻2𝑂(𝐼)𝐸𝑙𝑒𝑐𝑡𝑟𝑜𝑙𝑦𝑡𝑖𝑐 𝐷𝑒𝑐𝑜𝑚𝑝𝑜𝑠𝑖𝑡𝑖𝑜𝑛→ 2𝐻2(𝑔)+ 𝑂2(𝑔) Water Hydrogen Oxygen
4. If an electric current is passed through molten lead bromide, it decomposes to give lead and bromine.
Activity Series: The activity series of metals is a list of metals in the order of their decreasing
chemical activity.
A substitution or displacement reaction is a chemical change in which atoms of one element
replace the atoms of another element from the molecules of a compound. Elements which are higher in the Activity Series displace those elements which are placed below them. More electro positive elements displace lesser electropositive elements. Conversely, higher electro negative elements will displace lesser electro negative elements. For e.g.,
The iodine so liberated, dissolves in the chloroform, giving it a purple color.
Double Displacement Reactions / Metathesis reactions: The reactions in which two
Types of Double displacement reactions: 1. Precipitation Reaction
2. Neutralization Reactions
1. Precipitation
In all the above reactions a white substance, which is insoluble in water, is formed. This insoluble substance formed is known as a precipitate. A reaction that produces a precipitate is called a precipitation reaction.
2. Neutralization
Neutralization is a type of double displacement reaction, in which, the reactants are a base and an acid, and the products are salt and water. The positive charge of the hydrogen ion of the acid, and the negative charge of hydroxyl ions or oxide ions of the base, lose their electrical charge, and become covalent molecule of water.
A neutralization reaction is basically a reaction between H+ and OH- ions i.e.,
Oxidation - Reduction Reactions/Redox Reactions Classical Concept of Oxidation and Reduction
"Oxidation is a reaction in which oxygen is added or hydrogen is removed from a substance."
Removal of Hydrogen - Hydrogen is removed from hydro iodic acid to liberate free iodine.
"Reduction is a reaction in which oxygen is removed from a substance or hydrogen is added to a substance."
Removal of Oxygen - Oxygen is removed from copper oxide to form copper metal.
Addition of Hydrogen - Hydrogen adds to chlorine to form hydrogen chloride gas.
Oxidation and reduction reactions may occur simultaneously, these reactions as "redox" reactions.
Example of Redox Reaction
Example of Non-redox Reaction
Redox reactions are the reactions in which oxidation and reduction takes place simultaneously.
Oxidizing and Reducing Agents
Oxidizing Agent: A substance that brings about oxidation. Reducing Agents: A substance that brings about reduction. Another definition of Oxidation and Reduction
Oxidation: The process of addition of electronegative element or radical or removal of electropositive element or radical.
Reduction: The process of addition of electropositive element or radical or removal of electronegative element or radical.
Electronic Concept of Oxidation and Reduction
Reduction: The process in which there is gain of electrons. The Effects of Oxidation Reactions in Every Day Life Corrosion
Many metals are chemically active elements and get easily affected by substances like moisture, air, acids, etc. One must have observed iron articles that are shiny when new, but get coated with a reddish brown powder when left for some time. This process is commonly known as rusting of iron. The problem with iron (as well as many other metals) is that oxidation takes place and the oxide formed does not firmly adhere to the surface of the metal causing it to flake off easily. This eventually causes structural weakness and disintegration of the metal.
Hence metal is attacked by substances around it, it is said to corrode and this process is called corrosion. Corrosion causes deterioration of essential properties in a material.
What happens to copper vessels or artifacts when exposed to air and water? They slowly get tarnished by acquiring a thin green oxide layer. Similarly, silver quickly acquires a thin black oxide coating in moist air. The heaviest metal lead also tarnishes in moist weather. The black coating on silver and the green coating on copper are examples of corrosion in which the oxides formed strongly bond to the surface of the metal, preventing the surface from further exposure to oxygen and consequently slowing down corrosion.
Rancidity
Have you ever tasted or smelt the fat/oil containing food materials left for a long time? This unpleasant change in the flavor and odour of a food is called rancidity. The most important cause of rancidity is the deterioration in fats and fatty foods because of oxidation process. When an oxygen atom replaces hydrogen atom in the fatty acid molecule it destabilizes the molecule. Factors which accelerate fat oxidation include, salt, light, water, bacteria, moulds trace metals (iron, zinc, etc.).
Usually substances which retard fat oxidation or rancidity are called antioxidants (such as BHT, BHA, vitamin E, and vitamin C, and spices such as sage and rosemary). These are added to foods containing fats and oil to prevent such spoiling. Keeping food in air tight containers or air tight wrapping also helps to slow down oxidation.
Some high fat foods such as potato chips are packaged in materials that protect them from light and oxygen and the containers are flooded with nitrogen to further exclude oxygen. At times, to avoid the presence of oxygen altogether, vacuum packaging is used in some processed foodstuff.
Chapter 2 - Acids, Bases and Salts
IntroductionElectrolytes: The substances which when dissolved in water conduct electricity.
Acids, bases and salts are three main categories of chemical compounds. The sour taste of many fruits and vegetables, lemon for instance, is due to various types of acids present in them. The digestive fluids of most animals and humans also contain acids.
The word 'acid' is derived from a Latin word, which means "sour". The acids we use in the laboratory are stronger acids like hydrochloric acid and sulphuric acid. Strong acids are corrosive and can burn your skin. Bases on the other hand are the chemical opposites of acids. They are bitter in taste and soapy to touch. Sea water and detergents are some examples of substances that are basic. Many bases are oxide or hydroxide compounds of metals. Strong bases can also burn ones skin.
• Acids present in plant materials & animals are called Organic acids. Some naturally occurring acids,
Vinegar Acetic acid
Sour milk (curd) Lactic acid
Oranges Citric acid
Lemons Citric acid
Tamarind Tartaric acid
Ant sting Formic acid
Apples Malic acid
Tomatoes Oxalic acid.
Inorganic or Minerals acids are derived from minerals occurring in nature.
Some common acids that are found in laboratories are Hydrochloric acid (HCl),
Sulphuric acid (H2SO4) and
Nitric acid (HNO3).
Some of the lesser used acids are Acetic acid (CH3COOH),
Hydrofluoric acid (HF), Hydrofluoric acid is a highly corrosive acid and is used to etch glass. Carbonic acid (H2CO3).
General properties of Acids:
• Tastes sour
• Changes the colour of litmus from blue to red. • Conducts electricity.
General properties of Bases
• Have a soapy feel, • May also burn the skin
• Common examples are soaps & detergents.
• Commonly found bases in laboratories and in our daily life are: Caustic soda, NaOH; Caustic potash, KOH; Milk of magnesia, Mg(OH)2; Liquor ammonia, NH3; Washing powder, Tooth
paste.
ACIDS
Concentrated & Dilute acids - A concentrated acid is one which contains the minimum amount
of water in it. A dilute acid is obtained by mixing the concentrated acid with water.
Addition of Acids or Bases to Water - The process of dissolving an acid or a base in water is a
highly exothermic one. As this reaction generates lot of heat care must be taken while mixing concentrated acids with water, specially nitric acid or sulphuric acid with water. As a rule:
Always add Acid to Water and Never the Other Way! The acid must be added slowly to water
with constant stirring. If one mixes the other way by adding water to a concentrated acid, the heat generated causes the mixture to splash out and cause burns. The glass container may also break due to excessive local heating and cause damages! Mixing an acid or base with water results in dilution. It decreases the concentration of ions (H3O+/OH-) per unit volume thereby dissipating the
heat effect easily.
What Happens to an Acid in a Water Solution? Acids - Since all acids contain hydrogen ions,
the more hydrogen ions they contain, the stronger the acids are. A good definition of an acid is a
compound that produces H+ ions when it dissolved in water. Hydrogen ions cannot exist alone,
but they exist after combining with water molecules. H+ ions in association with a water molecule form H3O+ ions or hydronium ion.
H+ + H
2O H3O+
For example, when hydrogen chloride gas is dissolved in water, the hydrogen chloride molecules immediately dissociate or split into hydrogen ions and chloride ions. The solution becomes a very strong acid solution called hydrochloric acid.
The separation of H+ ion from HCl molecules cannot occur in the absence of water. Thus hydrogen ions must always be shown as H+(aq) or (H3O+).
The strength of an acid depends on the concentration of the hydronium ions (H3O+) present in a
solution. We know that greater the number of hydronium ions present, greater is the strength of acid. However, some acids do not dissociate to any appreciable extent in water such as carbonic acid. Therefore, these acids will have a low concentration of hydronium ions
How Strong are Acid Solutions? Strong Acid - An acid, which dissociates completely or almost
completely in water, is classified as a strong acid. An aqueous solution of a strong acid contains only ions along with water.
It must be noted that in these acids all the hydrogen ions (H+) combine with water molecule and exist as hydronium ions (H3O+). Examples of strong acids are: Hydrochloric acid, Sulphuric acid,
Nitric acid etc.
H2SO4 (aq) 2H+(aq) + SO42- (aq)
Weak Acid – An acid that dissociates only partially when dissolved in water, is classified as a
weak acid. An Aqueous solution of a weak acid contains ions and molecules. Examples are: acetic acid, formic acid, carbonic acid etc.
Reaction of Acids with Metals
(a) All metals above hydrogen in the metal reactivity series generally react with dilute acids to form their respective salt and liberate hydrogen.
(b) Very active metals like potassium, sodium and calcium also react similarly, but tend to explode when combining with acids.
(c) Nitric acid (of various concentrations) usually exhibits oxidizing property, rather than acidic properties. Metals such as magnesium combine with extremely dilute (1%) nitric acid to liberate hydrogen.
(d) Reaction of Metal Carbonates and Metal Hydrogen Carbonates with Acids?
Acids react with carbonates and hydrogen carbonates (bicarbonates) to form their respective salt, water and carbon dioxide.
Carbonate/Bicarbonate + Acid Salt + Water + Carbon dioxide
Neutralization
The reaction between the hydrogen ions of an acid and the hydroxyl ions of a base is called neutralization. In general, a neutralization reaction can be written as:
Acid + Base Salt + Water Examples:
1.
2.
The acidic property of an acid is due to the presence of hydrogen ions (H+) while that of a base or alkali, is due to the presence of hydroxyl (OH-) ions in them. When an acid and base (alkali)
combine, the positively charged hydrogen ion of the acid combines with the negatively charged hydroxyl ion of the base to form a molecule of water. Hence, the water molecule formed does not have any charge because the positive and negative charges of the hydrogen ions and hydroxyl ions get neutralized.
Neutralization can be viewed as a reaction in which an acid combine with a base, neutralizing the positively charged hydrogen ion and the negatively charged hydroxyl ion, to form a molecule of water and the respective salt.
Reaction of Metallic Oxides with Acids Action with Basic Oxides
Oxides that react with an acid to form salt and water are called basic oxides. These oxides get neutralized when they react with acids.
Basic oxide + Acid Salt + Water
1.
2.
3.
Action with Basic Hydroxides
Acids undergo neutralization reaction with basic hydroxides to form salt and water. Basic hydroxide + Acid Salt + Water
2.
3.
Reaction of Non-metallic Salts with Base
Calcium hydroxide, which is a base, reacts with carbon dioxide to produce salt and water. Since this is similar to the reaction between a base and an acid, we can conclude that nonmetallic oxides are acidic in nature.
This reaction occurs during white washing.
Bases
The oxides and hydroxides of metals are called bases. Examples of bases- sodium hydroxide,
magnesium oxide, calcium oxide, copper oxide, potassium hydroxide, magnesium hydroxide etc. Some bases are water soluble and dissolve in water to produce hydroxyl ions. A base that is soluble in water is an alkali. For example, when sodium hydroxide is dissolved in water it readily dissociates to produce a lot of hydroxide ions.
All alkalis are bases that dissociate in water to yield hydroxyl ion (OH-) as the only negative
ions. Sodium hydroxide, potassium hydroxide, calcium hydroxide and ammonium hydroxide are
the common alkalis.
NaOH(aq) Na+(aq) + OH-(aq) KOH(aq) K+(aq) + OH- (aq)
Ca(OH)2(aq) Ca2+(aq) + 2OH-(aq)
NH4OH(aq) NH4+(aq) + OH-(aq)
Strong Base /Alkali - The strength of a base depends on the concentration of the hydroxyl ions
when it is dissolved in water. A base that dissociates completely or almost completely in water to give a high concentration of hydroxyl ions is classified as a strong base. The greater the number
of hydroxyl ions the base produces, the stronger is the base. NaOH, KOH, & LiOH are strong alkalis.
Example:
Weak Base / Alkali - A base that dissociates in water only partially to give a low concentration
of hydroxyl ions is known as a weak base. Calcium hydroxide & ammonium hydroxide are weak alkalis.
Example:
Reactions of Bases/alkalis
Neutralization Reaction – Already done
Action of Alkalis/Base with Ammonium Salts
Alkalis combine with ammonium salts to liberate ammonia. Alkali + Ammonium salt Salt + Water + Ammonia
pH scale: The pH of a solution is defined as the negative logarithm of hydrogen ion concentration
in moles per litre.
pH = - log [H+(aq)]
The pH scale is a continuous scale and the value of pH varies between 0 to 14.
The pH of pure or neutral water is 7. Solutions having pH less than 7 are acidic in nature and the solutions with pH more than 7 are basic in nature.
Indicators
Acids and bases can be better distinguished with the help of indicators. Indicators are substances that undergo a change of color with a change of acidic, neutral or basic medium.
Litmus, a purple dye extracted from the lichen plant, is commonly used as an indicator in laboratories. Acids change the color of litmus solution to red, and bases change the color of litmus solution to blue.
Turmeric is another common household indicator. A stain of turmeric based food spill on a white cloth becomes reddish-brown when soap is scrubbed on it. Soap is basic in nature and changes the color of the turmeric stain. It turns yellow again when the cloth is washed with plenty of water. Other indicators—
• Red cabbage extract gives red color in acidic solutions & yellow color in basic solutions. • Onion has a characteristic smell. In basic solutions like NaOH, there is no smell. Acids
however, do not destroy the smell of onions.
• Vanilla extract has a pleasant smell in acidic solutions, in basic solutions there is no smell.
The common indicators used and the color changes observed are mentioned below: Indicator Acid Alkali
Litmus Red Blue
Methyl orange Pink Yellow
Phenolphthalein Colorless Deep pink
Methyl red Yellow Red
Universal Indicator: It is a mixture of indicators which give a gradual change of various colors
over a wide range of pH.
Approximate pH Values of Some Common Substances
Substance pH Value Hydrochloric acid 1.0 Sulphuric acid 1.2 Gastric juice 2.0 Rain water 6.2 Lemon 2.3 Milk 6.5
Vinegar (Acetic acid) 2.8
Pure water 7.0
Soft drink 3.0
Apple 3.1
Grape 3.1
Ammonium hydroxide 11.1
Tomato 4.2
Sodium hydroxide 13.0
Importance of pH in our daily life:
(a) pH and Plants: Proper pH of soil is required for healthy growth of plants. It should not be too acidic or too basic.
(b) pH in the digestive system: Human body secrets hydrochloric acid which aids in digestion.
Hyperacidity: The Condition of excess acid in the stomach. Hyperacidity can be cured by
taking ant-acid tablets or suspensions.
(c) pH and tooth decay: Tooth enamel which is the hardest substance in our body is corroded when the pH of the mouth is below 5.5. Cleaning of teeth using toothpaste helps in preventing tooth decay. Toothpastes are basic in nature, therefore neutralize the excess acid in the mouth and thus prevent tooth decay.
Salts & pH of Salts
Salts are obtained by treating an acid with a base. Salts consist of both positive ions or 'cations', and negative ions or 'anions'. The cations are called basic radicals and are mostly obtained from metallic ions (ammonium ion being one exception), while the anions are called acidic radicals and are obtained from acids.
Salt is a compound, which on dissociation in water yields positive ions other than a hydrogen ion or hydronium ion, and a negative ion other than hydroxyl ion.
Family of Salts: Salts can be classified into the following types:
Normal or Neutral Salts - A salt that is formed by the complete replacement of the replaceable hydrogen ions of an acid by a metal ion or ammonium ion is called a normal salt.
Examples: NaCl, Na2SO4, Na3PO4, NH4Cl, K2CO3 etc.
A neutral salt arises due to the neutralization reaction. Here, salts of strong acid and strong base combine to form such salts that show a neutral pH of 7.
Sodium Chloride
Sodium chloride is the commonly available salt and so is called common salt. Seawater is the main source of sodium chloride. Seawater contains about 3.5% of soluble salts, the most common of which is sodium chloride (2.7 to 2.9%). Saline water of inland lakes is also a good source of this salt. Sodium chloride is also found as rock salt.
Common salt is generally obtained by evaporation of seawater. Crude sodium chloride is obtained by crystallization of 'brine' that contains sodium sulphate, calcium sulphate, calcium chloride and magnesium chloride as impurities.
Pure sodium chloride is obtained from the crude salt by dissolving it in minimum amount of water and filtering it to remove insoluble impurities. The solution is then saturated with hydrogen chloride gas, when crystals of pure sodium chloride separate out. Calcium and magnesium chlorides, being more soluble than sodium chloride, remain in solution.
Properties
• Sodium chloride is a white crystalline solid having a density of 2.17 g/ml. • It melts at 1080 K (807°C) and boils at 1713 K (1440°C).
• It is soluble in water and its solubility is 36 g per 100 g of water at 273 K. (0°C). The solubility in water remains constant with temperature.
• Pure sodium chloride is non-hygroscopic, but behaves as hygroscopic due to the impurities of CaCl2 and MgCl2 in it.
• Solid Sodium chloride does not conduct electricity at room temperature but molten sodium chloride is a very good ionic conductor.
Uses
• As table salt, an essential constituent of our food. • In the manufacture of Na2CO3, NaOH, Cl2, etc.
• For salting out soap, and organic dyes. • In freezing mixtures.
• In tanning and textile industries.
• As a preservative for fish, meat, butter etc.
Sodium Carbonate (Na2CO3)
Sodium carbonate exists as anhydrous (Na2CO3) and also as hydrated salt. The dehydrated salt
(Na2CO3.10H2O) is known as washing soda while the anhydrous salt is called soda ash. Manufacture of Sodium Carbonate
Sodium carbonate is usually made by the Ammonia-soda process or Solvay process. The raw materials for this process are common salt, ammonia and limestone (for supplying CO2 and
When carbon dioxide is passed into a concentrated solution of brine saturated with ammonia, ammonium bicarbonate is produced. The ammonium bicarbonate then reacts with common salt forming sodium bicarbonate.
Sodium bicarbonate being slightly soluble (in presence of sodium ions) gets precipitated. Precipitated sodium bicarbonate is removed by filtration and changed into sodium carbonate by heating.
Steps in the Solvay process
Step 1. Ammoniacal brine reacts with carbon dioxide to produce sodium hydrogen carbonate.
NaCl + NH3 + H2O + CO2 NH4Cl + NaHCO3 .
Step 2. Sodium hydrogen carbonate is heated to get sodium carbonate. 2 NaHCO3 Na2CO3 + H2O + CO2.
Step 3. Sodium carbonate is recrystallized by dissolving in water to get washing soda. Na2CO3 + 10 H2O Na2CO3 .10H2O
Limestone is heated to obtain CO2. CaCO3 CaO + CO2.
The quicklime is dissolved in water to obtain slaked lime which is made to react with ammonium chloride to obtain ammonia which is used in step 1.
CaO + H2O Ca (OH)2
Ca(OH)2 + NH4Cl CaCl2 + 2 NH3 + 2H2O Properties
Sodium carbonate is a white crystalline solid, which can exist as anhydrous salt (Na2CO3),
monohydrate salt (Na2CO3.H2O), heptahydrate salt (Na2CO3.7H2O) and decahydrate
(Na2CO3.10H2O - washing soda). Sodium carbonate is readily soluble in water. On heating, the
decahydrate salt gradually loses water to, finally give anhydrous salt (Na2CO3 - soda ash).
Uses
• For washing purposes in laundries.
• For the manufacture of other sodium compounds like sodium silicates, sodium hydroxide, borax, hypo etc.
• As a household cleansing agent. • In paper and soap/detergent industries. • For the softening of water.
• A mixture of NaCO3 and KCO3 is used as a fusion mixture.
• In textile industry and petroleum refining.
Sodium Hydrogen Carbonate, (NaHCO3)
Sodium hydrogen carbonate is also known as sodium bicarbonate or baking soda because it decomposes on heating to generate bubbles of carbon dioxide (leaving pores in cakes or pastries and making them light and fluffy).
Preparation
NaHCO3 is made by saturating a solution of sodium carbonate with carbon dioxide. The white
crystalline powder of sodium hydrogen carbonate being less soluble gets separated.
On an industrial scale, sodium hydrogen carbonate (NaHCO3) is obtained as an intermediate
product in Solvay process for the manufacture of sodium carbonate.
Properties
• Sodium hydrogen carbonate is a white crystalline solid having a density of about 2.2 g/ml. • It has alkaline taste and is sparingly soluble in water. The solubility of sodium hydrogen
carbonate increases with the rise of temperature.
Uses
• As a component of baking powder. • In fire extinguishers.
• In medicines as a mild antiseptic for skin diseases and to neutralize the acidity of stomach. • As a reagent in laboratory.
Sodium Hydroxide (NaOH)
Sodium hydroxide is commonly called caustic soda because of its corrosive action on animal and vegetable tissues. Large quantity of sodium hydroxide is prepared by electrolytic process called the 'Chlor-alkali process'. Here, chlorine gas is given of at the anode and hydrogen gas at the cathode. Sodium hydroxide solution is formed near the cathode.
Properties
• Sodium hydroxide is a white deliquescent solid having melting point at 591 K (318°C). • It is stable towards heat.
• It is highly soluble in water and considerable amount of heat is evolved due to the formation of a number of hydrates e.g., NaOH.H2O, NaOH.2H2O. It is also soluble in alcohol.
• Aqueous solution of sodium hydroxide is strongly alkaline due to its complete dissociation into Na+ and OH-.
• Solution of sodium hydroxide is soapy to touch. It has a bitter taste. When a concentrated solution of sodium hydroxide comes in contact with skin, it breaks down the skin and flesh to a pasty mass.
Uses
• In the manufacture of soap, paper, viscose rayon (artificial silk), organic dyestuffs, and many other chemicals.
• In the refining of petroleum and vegetable oils.
• In the purification of bauxite for the extraction of aluminum.
• As a cleansing agent and in washing powder for machines, metal sheets etc. It is too caustic to be used in washing clothes or hands.
• For mercerizing cotton. • As a reagent in the laboratory. • In reclaiming rubber.
• In the preparation of soda lime.
Plaster of Paris, [CaSO4. ½ H2O]
Calcium sulphate with half a molecule of water per molecule of the salt (hemi-hydrate) is called plaster of paris (plaster of paris).
Water of Crystallization
When crystals of certain salts are formed, they do so with a definite number of molecules of water, chemically combined in a definite proportion. Water of crystallization is the number of water molecules, chemically combined in a definite molecular proportion, with the salt in its crystalline state. This water is responsible for the geometric shape and color of the crystals.
Remember
A substance containing water of crystallization is called a hydrous substance or a hydrate. This water can be expelled, by heating, and then the salt is said to have become anhydrous.
Preparation
Plaster of Paris is prepared by heating gypsum (CaSO4.2H2O) at 120°C in rotary kilns, where it
gets partially dehydrated.
The temperature should be kept below 140°C otherwise further dehydration will take place and the setting property of the plaster will be partially reduced.
Properties
It is a white powder. When mixed with water (1/3 of its mass), it evolves heat and quickly sets to a hard porous mass within 5 to 15 minutes. During setting, a slight expansion (about 1%) in volume occurs so that it fills the mould completely and takes a sharp impression. The process of setting occurs as follows:
The first step is called the setting stage, and the second, the hardening stage. The setting of plaster of Paris is catalyzed by sodium chloride, while it is reduced by borax, or alum.
Uses
• In surgery for setting broken or fractured bones.
• For making casts for statues, in dentistry, for surgical instruments, and toys etc. • In making black board chalks, and statues.
Chapter 3 - Metals and Non-metals
IntroductionThere are 118 elements present in the periodic table, 92 of which are naturally occurring. Metals and non-metals are characterized by distinctly different physical and chemical properties. At present about 80 metals are known to us.
At room temperature, over half of the non-metals are gases, except bromine, which is a liquid. The most abundant non-metal in the earth's crust is oxygen, which constitutes about 50% of the earth's crust and along with nitrogen it forms the main constituents of air.
The next abundant nonmetal is silicon which constitutes about 26% of the earth’s crust. Oxygen and silicon are the two major constituents of earth. Hydrogen and oxygen are the two major constituents of the oceans.
Position of Metals and Non-metals in the Periodic Table
Metals occupy the groups on the left of the periodic table. Group IA consists of highly reactive metals called the alkali metals, while group II A elements are called alkaline earth metals. Elements between group IIA and IIIA are all called transition metals.
The non-metals are elements (with the exception of hydrogen) that are found to the right on the Periodic Table i.e., groups IVA, VA, VIA &VIIA. The non-metallic character of these elements increases from top to the bottom of the group. For example, in group VA the first and second members are non-metals, the third and fourth are metalloids and the last member is a metal. The metalloids are a group of elements which have properties similar to both the metals and non-metals. These metalloids are: Boron, silicon, germanium, arsenic, antimony, tellurium and astatine. The non-metals are elements found to the right of these metalloids, including the element, hydrogen.
Group V A
Non Metals Nitrogen, Phosphorous, Metalloids Arsenic, Antimony
Metal Bismuth
Physical Properties of Metals
Physical State - Metals are solids at room temperature with the exception of mercury and gallium,
which are liquids at room temperature.
Lustre - Metals have the quality of reflecting light from its surface and can be polished e.g., gold,
Malleability - Metals have the ability to withstand hammering and can be made into thin sheets
known as foils. Except Zinc which is brittle.
Ductility - Metals can be drawn into wires. Except Zinc which is brittle.
Hardness - All metals are hard except sodium and potassium, which are soft and can be cut with a
knife.
Conduction - Metals are good conductors because they have free electrons. Silver and copper are
the two best conductors of heat and electricity. Lead is the poorest conductor of heat. Bismuth, mercury and iron are also poor conductors
Density - Metals have high density and are very heavy. Iridium and osmium have the highest
densities whereas lithium has the lowest density.
Melting and Boiling Point - Metals have high melting and boiling point. Tungsten has the highest
melting point where as silver has low boiling point. Sodium and potassium have low melting points.
Alloy Formation - Metals form homogeneous mixture with each other called an alloy. Example-
Brass is an alloy of copper and zinc.
Sonorous - Metals are sonorous i.e. they produce sound when hit with some solid object. Physical Properties of Non-metals
Physical State - Most of the non-metals exist in two of the three states of matter at room
temperature: gases (oxygen) and solids (iodine, carbon, sulphur). These have no metallic lustre, (except iodine) and do not reflect light. (Except carbon in the form of diamond).
Nature - Non-metals are very brittle, and cannot be rolled into wires or pounded into sheets.
Except- diamond is the hardest substance known.
Conduction - They are poor conductors of heat and electricity. (Except graphite conducts heat, both
graphite & gas carbon conduct electricity.)
Electronegative Character - Non-metals have a tendency to gain or share electrons with other
atoms. They are electronegative in character.
Reactivity - They generally form acidic or neutral oxides with oxygen.
Melting and Boiling Points – Non-metals have low melting and boiling points. Comparative Properties of Metals and Non-Metals
Property Metals Non-metals
Property Metals Non-metals
which is a liquid at room temperature. Gallium and Caesium melt below 30oC.
So if room temperature is around 30oC, they may also be in liquid state
Bromine is the only liquid. Solids – iodine, carbon, sulphur.
Density They usually have high density, except
for sodium, potassium, calcium etc.
Their densities are usually low.
Melting point They usually have a high melting point except mercury, cesium, gallium, tin, lead.
Their melting points are low.
Boiling point Their boiling points are usually high. Their boiling points are low.
Hardness They are usually hard, except mercury,
sodium, calcium, potassium, lead etc.
They are usually not hard. But the
exception is the non-metal
diamond, the hardest substance. Malleability They can be beaten into thin sheets. They are generally brittle.
Ductility They can be drawn into thin wires,
except sodium, potassium, calcium etc.
They cannot be drawn into thin wires.
Conduction of heat They are good conductors of heat. They are poor conductors of heat. (exception- carbon in the form of graphite)
Conduction of
electricity
They are good conductors of electricity. They are non-conductors, except for carbon in the form of graphite and the gas carbon.
Lustre Newly cut metals have high lustre.
Some get tarnished immediately.
Usually not lustrous, except iodine and diamond - the most lustrous of all the substances.
Alloy formation They form alloys. Generally, they do not form
alloys. However, carbon,
phosphorus, sulphur etc. can be present in some alloys.
Tenacity These usually have high tensile strength
except sodium, potassium, calcium, lead etc.
These have low tensile strength.
Brittleness They are hard but not brittle, except zinc at room temperature.
They are generally brittle.
Electronic They usually have 1, 2 or 3 electrons in
their valence shell. The greater the
They usually have 4, 5, 6 or 7 electrons in the valence shell. If it
Property Metals Non-metals
configuration number of shells and lesser the number of valence electrons, the greater is the reactivity of the metal.
has 8 electrons, it is called a noble gas. Lesser the number of shells and greater the number of valence electrons, greater is the reactivity of the non-metal.
Ionization They always ionize by losing electrons: They always ionize by gaining
electrons:
Charge of ions Positively charged. Negatively charged.
Type of valency Metals always exhibit electrovalency. Non-metal exhibit both
electrovalency or covalency. Deposition during
electrolysis
They are always deposited at the cathode.
They are always deposited at the anode.
Redox reaction These lose electrons and hence get oxidized.
These gain electrons and hence get reduced.
Redox agents They are reducing agents. They are oxidizing agents.
Nature of oxides They generally form basic oxides, some of which are also amphoteric, such as aluminum oxide, zinc oxide, lead oxide etc.
They generally form acidic
oxides.
Neutral oxides are nitrous oxide, nitric oxide, carbon monoxide water etc.
Hydrides They do form hydrides except some
transition elements.
They do form hydrides, e.g. NH3,
PH3, HCl, HBr, HI, H2S, H2O etc.
Atomicity These are always monatomic. These can be mono, di, tri, or
polyatomic.
Solubility They do not dissolve in solvents except
by chemical action.
They dissolve in solvents and can be re-obtained by evaporation. Example: Sulphur in carbon disulphide.
Action with
chlorine
They produce chlorides, which are electrovalent.
They produce chlorides, which are covalent.
Action with dilute acids
On reaction with dilute acids they give respective salt and hydrogen.
They do not react with dilute acids.
Chemical Properties of Metals Metals are Electropositive Elements
Metals are very reactive. Metals tend to lose electrons easily and form positively charged ions; therefore, metals are called electropositive elements. Sodium metal forms sodium ions Na+. The electropositive nature allows metals to form compounds with other elements easily.
Reaction of Metals with Oxygen
Metals like sodium (Na) and potassium (K) are some of the most reactive metals. Potassium, sodium, lithium, calcium and magnesium react with oxygen and burn in air.
Metals from aluminum to copper in the activity series of metals react slowly when heated in air to form the metal oxides. Aluminum is the fastest and copper is the slowest of them.
• Sodium metal reacts with the oxygen of the air at room temperature to form sodium oxide. Hence, sodium is stored under kerosene to prevent its reaction with oxygen, moisture and carbon dioxide.
• Sodium oxide is a basic oxide which reacts with water to form sodium hydroxide.
• Mg does not react with oxygen at room temperature. On heating, Mg burns in air with intense light and heat to form MgO.
• Zinc metal burns in air only on strong heating to form zinc oxide.
• In moist air, iron is oxidized to give rust.
• On heating in air it burns with a brilliant flame forming triferric tetroxide. • Copper is the least reactive metal and does not burn in air even on heating. However, on prolonged strong heating copper reacts with oxygen and forms copper (II) oxide (CuO) outside and copper (I) oxide (Cu2O) inside.
• Gold and platinum do not react with oxygen in air.
Reaction of Metals with Water
Potassium, sodium, lithium and calcium react with cold water.
• Sodium reacts vigorously with cold water forming sodium hydroxide and hydrogen.
• Metals from magnesium to iron in the activity series of metals, react with steam (but not cold H2O) to form the metal oxide and hydrogen gas.
• Red hot iron reacts with steam to form Iron (II, III) oxide.
Note: The reaction between iron and steam is irreversible. Tin, lead, copper, silver, gold and
platinum do not react with water or steam.
Reaction of Metals with Acids
• Potassium, sodium, lithium and calcium react violently with dilute H2SO4 and dilute HCl,
forming the metal salt (either sulphate or chloride) and hydrogen gas. The reaction is similar to the reaction with water.
• Magnesium, aluminum, zinc, iron, tin and lead react safely with dilute acid. Magnesium is the fastest and lead is the slowest of the six.
Zinc with dilute sulphuric acid is often used for the laboratory preparation of hydrogen. The reaction is slow at room temperature, but its rate can be increased by the addition of a little copper (II) sulphate. Zinc displaces copper metal, which acts as a catalyst.
Metals below hydrogen (copper, silver, gold and platinum), will not react with dilute acid to liberate hydrogen. In general,
• Hydrochloric acid makes a metal chloride. • Sulphuric acid makes a metal sulphate.
• Reactions with nitric acid are more complex, the nitrate is formed but the gas is rarely hydrogen, and more often, an oxide of nitrogen.
Reaction of Metals with Salt Solutions
Reactive metals can displace any metal less reactive than itself, from the oxide, chloride or sulphate of the less reactive metal in solution or their molten state. If metal A displaces metal B
from its solution, it is more reactive than B.
Copper (II) sulphate solution is blue; iron sulphate solution is almost colourless when dilute. During the displacement, the blue solution loses its color and the iron metal is seen to turn pink-brown as the displaced copper becomes deposited on it.
On heating the mixture of magnesium powder and black copper (II) oxide, white magnesium oxide is formed with brown bits of copper:
Adding magnesium to blue copper (II) sulphate solution, the blue color fades as colourless magnesium sulphate is formed and brown bits of copper metal form a precipitate:
Electronic Nature of Metals and Non-metals
The atoms of all elements, except noble gases, have an incomplete outermost shell. . Noble gases have their outermost shell complete and hence they are not reactive or "inert".
Most elements are reactive and try to achieve the stability of the noble or inert gases by electron transfer or by electron sharing. Elements that can donate electrons are called metals. They form positive ions by losing electrons.
The elements that accept electrons are called non-metals. They form negative ions by gaining electrons. Metals have 1 to 3 electrons in the outermost shell of their atom and non-metals have 4 to 8 electrons in the outermost shell.
There are two exceptions to this rule: Hydrogen and helium. Hydrogen is a non-metal having 1 electron in the valence shell and helium too is an inert gas having 2 electrons in the valence shell.
Type of Elements Element Atomic Number Number of Electrons in Shells K L M N
Noble Gases Helium (He) 2 2
Type of Elements Element Atomic Number Number of Electrons in Shells Argon (Ar) 18 2 8 8
Metals Sodium (Na) 11 2 8 1
Magnesium (Mg) 12 2 8 2 Aluminium (Al) 13 2 8 3 Potassium (K) 19 2 8 8 1 Calcium (Ca) 20 2 8 8 2 Non-metals Nitrogen (N) 7 2 5 Oxygen (O) 8 2 6 Fluorine (F) 9 2 7 Phosphorus (P) 15 2 8 5 Sulphur (S) 16 2 8 6 Chlorine (Cl) 17 2 8 7
Important Point
Metals that donate electrons gain positive charge equal to the number of electrons donated. For example, atomic number of aluminum is 13, so the electronic configuration of Al is 2, 8, and 3. Aluminum has 3 electrons in the valence shell; it loses 3 electrons to form Al3+.
Alo – 3e- Al3+ Other examples,
Non-metals gain electrons and hence gain negative charge equal to the number of electrons accepted.
For example,
The Reactivity Series of Metals
Although most metals are usually electropositive in nature and lose electrons in a chemical reaction they do not react with the same vigour or speed. Metals display different reactions towards different substances. The greater the ease with which an element loses its electrons and acquires a positive charge, the greater is its reactivity. Further, the greater the number of shells and lesser the number of valence electrons, the greater is the reactivity of the metal. The activity series of metals, arranges all metals in order of their decreasing chemical activity. As we go down the activity series from potassium to gold the ease with which a metal loses electrons, and forms positive ions in solutions, decreases.
The most active metal, potassium, is at the top of the list and the least reactive metal, gold, is at the bottom of the list. Although hydrogen is a non-metal it is included in the activity series due to the fact that it behaves like a metal in most chemical reactions i.e., the hydrogen ion has a positive charge [H+] like other metals.
Element Symbol Group Number Potassium K IA Sodium Na IA Lithium Li IA Calcium Ca IIA Magnesium Mg IIA aluminium Al IIIA Carbon C IVA Zinc Zn IIB Iron Fe VIII Tin Sn IVA Lead Pb IVA Hydrogen H IA Copper Cu IB Silver Ag IB
Element Symbol Group Number
Gold Au IB
Platinum Pt VIII
• The higher the metal in the series, the more reactive it is i.e., its reaction is fast and more exothermic.
• This also implies that the reverse reaction becomes more difficult i.e., the more reactive a metal, the more difficult it is to extract from its ore. The metal is also more susceptible to corrosion with oxygen and water.
• The reactivity series can be established by observation of the reaction of metals with water, oxygen or acids.
• Within the general reactivity or activity series there are some periodic table trends:
Metals Reactivity and reactions
Potassium K Very reactive, very rapid with cold water forming the alkali potassium
hydroxide and hydrogen gas (which is ignited).
2K+2H2O(l) 2KOH(aq) +H2(g)
Sodium Na Fast reaction with cold water forming the alkali sodium hydroxide and
hydrogen gas.
2Na(s) + 2H2O(l) 2NaOH(aq) + H2(g)
The reaction of sodium with water-the sodium melts to a silvery ball and fizzes as it spins over the water. The rapid exothermic reaction produces a colourless gas that gives a squeaky pop! with a lit splint-hydrogen. Universal indicator will turn from green to purple/violet-the strong alkali sodium hydroxide is formed. The sodium floats because it is less dense than water.
Calcium Ca Quite reactive with cold water forming the moderately soluble alkali calcium hydroxide and hydrogen gas.
Ca(s) +2H2O(l) Ca(OH)2(aq/s) + H2(g)
Very reactive with dilute hydrochloric acid forming the colourless soluble salt calcium chloride and hydrogen gas.
Ca(s) +2HCl(g) CaCl2(aq) +H2(g)
Not very reactive with dilute sulphuric acid because the colourless calcium sulphate formed is not very soluble and coats the metal inhibiting the reaction.
Metals Reactivity and reactions
Magnesium Mg Slow reaction with water forming the slightly soluble base magnesium
oxide and hydrogen gas.
With steam, the reaction is faster with heated magnesium and a white powder magnesium oxide is formed along with hydrogen.
Magnesium will burn with a bright white flame in steam, if previously ignited in air.
Mg(s) + H2O(g) MgO(s) + H2(g)
In fact it will even burn in carbon dioxide forming black specks of carbon!
2Mg(s) +CO2(g) 2MgO(s) + C(s)
Very reactive with dilute hydrochloric acid forming the colourless soluble salt, magnesium chloride and hydrogen gas.
Mg(s) +2HCl(aq) MgCl2(aq) +H2(g)
Very reactive with dilute sulphuric acid forming colourless soluble magnesium sulphate and hydrogen.
Mg(s) +H2SO4(aq) MgSO4(aq) + H2(g)
Aluminum Al Aluminum has no reaction with water or steam due to a protective aluminum oxide layer of Al2O3. Slow reaction with dilute hydrochloric
acid to form a colourless soluble salt aluminum chloride and hydrogen gas.
2Al(s) +6HCl(aq) 2AlCl3(aq) + 3H2(g)
The reaction with dilute sulphuric acid is extremely slow to form colourless aluminum sulphate and hydrogen.
2Al(s +3H2SO4(aq) Al2(SO4)3(aq) + 3H2(g) (Carbon C,a
non-metal)
Elements higher than carbon i.e aluminum and the more reactive metals must be extracted by electrolysis (or displacing it with an even more reactive metal). Metals below it, i.e., zinc or a less reactive can be extracted by reducing the hot metal oxide with carbon.
Zinc Zn No reaction with cold water. When the metal is heated in steam zinc, oxide and hydrogen are formed.
Zn(s) + H2O(g) ZnO(s) + H2(g)
Quite reactive with dilute hydrochloric acid forming the colourless soluble salt zinc chloride and hydrogen gas. Zn(s) +2HCl(aq) ZnCl2(aq)
+ H2(g)
Quite reactive with dilute sulphuric acid forming the colourless soluble salt zinc sulphate and hydrogen gas.
Metals Reactivity and reactions
Zn(s) + H2SO4(g) ZnSO4(s) + H2(g)
(this reaction is catalyzed by adding a trace of copper sulphate solution) Zinc can be extracted by reducing the hot metal oxide on heating with carbon
2ZnO(s) + C(s) 2Zn(s) + CO2(g)
A zinc coating (galvanizing) is used to protect iron from rusting.
Iron Fe No reaction with cold water (rusting is a joint reaction with oxygen). When the metal is heated in steam an iron oxide (unusual formula) and hydrogen are formed. This oxide is 'technically' Iron (III, II) oxide! 3Fe(s) + 4H2O(g) Fe3O4(s) +4H2(g)
Slow reaction with dilute hydrochloric acid forming the soluble pale green salt Iron (II) chloride and hydrogen gas.
Fe(s) + 2HCl(aq) FeCl2(aq) + H2(g)
Slow reaction with dilute sulphuric acid forming the soluble pale green salt Iron (II) sulphate and hydrogen gas.
Fe(s) + 2H2SO4(g) FeSO4(s) + H2(g)
Iron can be extracted by reducing the hot metal oxide on heating with carbon monoxide formed from carbon in the blast furnace e.g.,
Fe2O3(s) +3CO(g) 2Fe(s) +3CO2(g)
Fe3O4(s) +4CO(g) 3Fe(s) +4CO2(g) (Hydrogen H
non-metal)
None of the metals below hydrogen can react with acids to form hydrogen gas. They are least easily corroded metals and partly accounts for their value and uses in jewellery, electrical contacts etc.
Copper Cu No reaction with cold water or when heated in steam. No reaction with dilute hydrochloric acid or dilute sulphuric acid. Copper can be extracted by reducing the hot black metal oxide on heating with carbon. 2CuO(s) + C(s) 2Cu(s) +CO2(g)
Silver Ag No reaction with cold water or when heated in steam. No reaction with dilute hydrochloric acid or dilute sulphuric acid. Silver can be extracted by reduction but can be found 'native' as the element.
Gold Au No reaction with cold water or when heated in steam. No reaction with dilute hydrochloric acid or dilute sulphuric acid. Gold can be readily extracted from its ores easily by reduction but it is usually found 'native'. Pure gold is 24 carat.
Metals Reactivity and reactions
Platinum Pt No reaction with cold water or when heated in steam. No reaction with dilute hydrochloric acid or dilute sulphuric acid. Like gold, it is a very rare metal. It is used in expensive jewellery, laboratory ware (inert crucible container) as a industrial catalyst, and catalytic converters in car exhausts.
Bonding
The tendency of an atom to take part in chemical combination is determined by the number of valence electrons (electrons in the outermost shell of an atom). The atoms acquire the stable noble gas configuration of having eight electrons in the outermost shell (called octet rule) during chemical combination.
The combination of atoms occurs in two ways: either by electrovalent bonding or covalent bonding. In all chemical reactions, it is the electrons from the outermost shell of an atom that are involved in interacting with other atoms, either by their transfer or by sharing.
Electrovalent Bonding
When an atom donates one, two or three electrons from its valence shell to another atom, which has the ability to accept these electrons, it is known as electrovalency. As a result of electrovalency, both these atoms achieve the structure of an inert gas. When the chemical bond occurs by the transfer of electrons from the atom of an element to the atom or atoms of another it is called Ionic or Electrovalent bond.
Thus, the electrovalency of sodium is 1+, and that of chlorine is 1- in NaCl. Similarly, calcium, magnesium in their chloride exhibits an electrovalency of 2+. There are many elements, which show different electrovalencies in different compounds. This phenomenon is called 'variable electrovalency' e.g., iron exists as Fe2+ and Fe3+ in ferrous sulphate and ferric sulphate respectively.
Formation of Sodium Chloride
During the formation of an ionic bond between the metal sodium and the non-metal chlorine, sodium loses one electron to complete its octet as it has only one electron in its valence shell. It acquires a noble gas configuration of neon (2, 8). While the chlorine atom has seven electrons in its valence shell and gains one electron to complete its octet and also acquires stable electronic configuration of argon.
Formation of Magnesium Chloride
Magnesium, whose atomic number is 12, has 2, 8, 2 configurations. Its valence shell has two electrons. The electronic configuration of chlorine (At. no. 17) is 2, 8, 7. It has seven valence electrons. Since, magnesium has two electrons in excess of the neon configuration (2, 8), and chlorine is one electron short of the argon configuration (2,8,8), hence one atom of magnesium will look for two atoms of chlorine to transfer its two electrons to (one to each) as shown below:
The Mg2+ and the two Cl- so formed, then form ionic bonds between them.
In terms of Lewis dot structure,
Formation of Magnesium Oxide
Mg (at no 12) has configuration 2, 8, 2.
The atom of Mg loses 2 electrons to become stable like Neon (2, 8) Mg – 2e- Mg2+
Oxygen (at no 8) has configuration 2,6.The atom of Oxygen gains 2 electrons to become stable like Ne (2, 8 )
O + 2e- O 2-
Mg2+ + O2- MgO
Formation of Calcium oxide
Ca (at no 20) has configuration 2,8,8,2.
The atom of Ca loses 2 electrons to become stable like Argon (2,8,8) Ca – 2e- Ca2+
Oxygen (at no 8) has configuration 2, 6.
The atom of Oxygen gains 2 electrons to become stable like Ne (2, 8) O + 2e- O 2-
Properties of Electrovalent Compounds
Property Electrovalent or ionic compounds
Structure of charged
ions They consist of oppositely charged molecules.
Physical state and
hardness
The inter-atomic attraction is high, hence they are brittle, hard, crystalline solids.
Melting and boiling
points
Due to strong attraction between the particles, high temperatures are required to melt or boil them.
Solubility They are usually soluble in water, but insoluble in organic solvents.
Passage of electricity
Ionic compounds do not conduct electricity in the solid state because movement of ions in the solid is not possible due to their rigid structure. In the molten form or in aqueous solution form, since the electrostatic forces of attraction between the oppositely charged ions are overcome they allow the flow of electricity, and get decomposed by it.
Rate of reaction Their reaction usually occurs with high speeds.
Dissociation in solution
Since electrovalent compounds are made up of charged ions, they dissociate to give negative and positive ions in solution.
Electrolysis These compounds can undergo electrolysis. The cations get discharged at cathode and anions at anode.
Covalent Bond
A covalent bond is defined 'as the force of attraction arising due to mutual sharing of electrons between the two nonmetallic atoms'. The combining atoms may share one, two or three pairs of electrons. The covalent bond is formed between two similar or dissimilar atoms of nonmetals by a mutual sharing of electrons, which are counted towards the stability of both the participating atoms.
When the two atoms combine by mutual sharing of electrons, then each of the atoms acquires a stable configuration of the nearest noble gas. The compounds formed due to covalent bonding are called covalent compounds. The shared pair of electrons are called Bond Pairs.
Formation of Covalent Bonds The Hydrogen Molecule
The hydrogen atom (Atomic number = 1) has 1 electron in the K shell. It tries to acquire the configuration of He (Atomic number 2). This is possible if the two combining atoms share their valence electron to form one covalent bond between themselves.
H - H ; H: H 1 covalent bond between two 2 H atoms forms H2
The Oxygen Molecule
Oxygen (Atomic number 8) has 6 valence electrons, 2 short of the octet configuration. The two oxygen atoms share two pairs of electrons to form 2 covalent bonds between them.
To form Oxygen molecule O2. Covalency
The number of electrons, which an atom contributes towards mutual sharing during the formation of a chemical bond, is called its covalency in that compound. Thus, the covalency of hydrogen in H2 (H - H, H : H) is one; that of oxygen in O2 is two (OO,O:xx O)and that of
nitrogen in N2 is three(N N,N N) x x x .
Sometimes one or more pairs of electrons in the valence shell of the atom do not take part in bonding, and are known as a lone pairs; they are also called non-bonding pair of electrons.
Example:
Each atom of oxygen has 2 pairs of non bonding electrons.
Multiple Covalent Bonds
The covalent bonds developed due to mutual sharing of more than one pair of electrons, are termed 'multiple covalent bonds'. These are:
Double covalent bond
The bond formed between two atoms due to the sharing of two electron-pairs is called a double covalent bond or simply a double bond. It is denoted by two small horizontal lines (=) drawn between the two atoms, e.g., O = O,
O = C = O etc.
Triple covalent bond
Bond formed due to the sharing of three electron pairs is called a triple covalent bond or simply a triple bond. Three small horizontal lines between the two atoms denote a triple bond e.g.,
NN, and H-CC-H (acetylene).
Formation of Molecules Having Double Bond Formation of oxygen (O2) molecule
Each oxygen atom has six electrons in its valence shell. Thus, it requires 2 more electrons to achieve the nearest noble gas configuration. This is achieved by sharing two pairs of electrons by the two oxygen atoms as shown below:
Formation of carbon dioxide molecule (CO2)
The electronic configurations of carbon and oxygen are, C 2, 4 and O 2, 6
Thus, each carbon atom requires four, and each oxygen atom requires two more electrons to acquire noble gas configurations. This is achieved as follows:
Formation of molecules having triple bond Formation of nitrogen (N2) molecule
Nitrogen atom has five electrons in its valence shell. Thus, it requires three more electrons to acquire a stable configuration of the nearest noble gas (neon). This is done by mutually sharing three pairs of electrons as shown below:
Formation of hydrogen cyanide (HCN) molecule
The carbon atom in HCN, shares one electron-pair with hydrogen, thus forming a single covalent bond with H atom. The C atom shares three electron pairs with N atom to form a triple bond between C and N. The combining of atoms and Lewis structure of HCN molecule is given below:
