Chemistry

LEARNING OBJECTIVES                                                                               

After reading this  unit, you will be able to:

• explain the nature of ionic, covalent and metallic bonding.

• state the typical physical properties of ionic compounds, and of covalent compounds.

• explain the physical properties of metals in terms of their structure.

KNOWLEDGE GAIN

In 1985, a new all otrope of carbon Buck minsterfullerene was discovered. It has a cage-like ring structure which resembles a football. It is made of twenty hexagons and twelve pentagons.

1.1 STABILITY OF ATOMS

ACTIVITY 1.1: Showing Stability of Atoms

• Take a glass full of water. Try adding water into it. Are you able to add?

• Now take another glass of water but a quarter (one-fourth) filled.

• Try adding water into it. Now, are you able to add or not?

Perform the two activities in classroom and then discuss your answers among your classmates.
 

In the above activity, you will observe that when the glass was already filled, there was no space to add more water into it. Thus, the water in the glass remained stable. A noble gas has a fully filled outermost shell just like the glass full of water. It has eight electrons in the outermost shells except helium (2 electrons).

When atoms or the elements combine to form molecules, a force of attraction is developed between the atoms (or ions) which holds them together. The force which links the atoms (or ions) in a compound is called a chemical bond (or just “bond”). A bond is formed so that each atom acquires a stable electronic configuration similar to that of a noble gas.

The atoms combine with one another to achieve the inert gas electron arrangement and become more stable. So, when atoms combine to form compounds, they do so in such a way that each atom gets 8 electrons in its outermost shell or 2 electrons in the outermost n shell.

An atom can achieve the inert gas electron arrangement in three ways:

• By losing one or more electrons (to another atom). Atoms with 1, 2 or 3 electrons in the outermost shell lose electrons to achieve stability.

• By gaining one or more electrons (from another atom).  Atoms with five, six or seven electrons in the outermost shell gain three, two or one electron respectively to achieve stability.

• By sharing one or more electrons (with another atom).  Atoms with four to seven electrons in outermost shell may achieve stability by sharing them with each other.

EXERCISES 1.1

1. What do you mean by a chemical bond?

2. When a bond is formed, each atom acquires a stable configuration similar to _______ 

3. Generally, metals lose electrons to achieve inert gas electron arrangement. (True or False)

4. Which of the following is not a noble gas?  (a) Helium (b) Neon  (c) Hydrogen (d) Argon

5. Among, phosphorus, sulphur, and calcium; which element achieves stability by losing electron.

1.2 FORMATION OF IONS FROM ATOMS

ACTIVITY 1.2: Illustrating Formation of Ion

Divide the class into two groups. Half of the students hold positive plank cards and another half hold negative plank cards. Positive plank cards are protons and negative plank cards are electrons. Now perform the following and analyze:

Students with 5 positive and 5 negative plank cards are grouped together. Their total charge being neutral in the group.

• Now, one electron is removed from the group. 4 students are left holding negative plank cards.  Can you tell the net charge now of this group?

• Add one electron to the neutral group.
 6 students are now holding negative plank cards. 

Can you now tell what is the charge of this group?

• Similarly, perform the above activity with 7 students and analyze the charge.

An atom contains electrons, protons and neutrons.

Protons carry positive charges, electrons carry negative charges and neutrons carry no charges. Every atom contains an equal number of “positively charged protons” and “negatively charged electrons”.

Thus, an atom is electrically neutral.
                                   

An ion is formed when an atom loses or gains one or more electrons. The atom may be of a metal or a non-metal.

A metal readily loses its outermost electron or electrons to form a positive ion or cation. The number of positive charges carried on a cation is equal to the number of electron(s) lost by the metal atom. Examples are given in Table 1.3.

                                    

Metal ions carry positive charges because the number of positively charged protons in the nucleus becomes greater than the number of negatively charged electrons surrounding it. For example, in a sodium atom there are 11 protons in the nucleus and 11 electrons surrounding it. Loss of one electron to form a sodium ion means that there are 11 protons but only 10 electrons. There is a net charge of 1⁺. This charge is written as a superscript at the right of the symbol of the element (Figure 1.1).

                                      Figure 1.1: Formation of a sodium ion.

Hydrogen atoms can also lose an electron to form an ion with one positive charge. Some non-metals readily gain one or more electrons into their outermost shell to form a negative ion or anion. The number of negative charges an anion carries is equal to the number of electron(s) gained by the non-metal atom. Examples are given in Table 1.4.
     

Non-metal ions carry negative charges because the number of negatively charged electrons surrounding the nucleus becomes greater than the number of positively charged protons in it. For example, in a chlorine atom there are 17 protons in the nucleus and 17 electrons surrounding it. Gain of one electron to form a chloride ion means that there are 18 electrons and only 17 protons. There is a net charge of –1. This charge is written as a superscript at the right of the symbol of the element (Figure 1.2).

                 

Notice that when a non-metal forms an anion, the name changes slightly; chlorine forms a chloride ion, oxygen forms an oxide ion. Several common radicals exist as negative ions including nitrate (NO^{– 3}), carbonate
 (CO₂^{– 3} ) and phosphate (PO₃ ^{4 –}).

EXERCISE 1.2

 1. Define ion.

2. Which of the following is an anion?  (a) Cl^–                            (b) Na⁺ 

                                                             (c) Mg²⁺                         (d) Al³⁺

3. Why do metal ions carry positive charges?

4. The number of negative charges an anion carries is equal to the number of electrons gained by the non-metal atom. (True or False)

5. Give two examples of each:  (i) anion (ii) cation

1.3 IONIC BONDING

The compounds which are made up of ions are known as ionic compounds. In an ionic compound, the positively charged ions (cations) and negatively charged ions (anions) are held together by the strong electrostatic forces of attraction. The forces which hold the ions together in an ionic compound are known as ionic bonds or electrovalent bonds. Since an ionic bond consists of an equal number of positive ions and negative ions, the overall charge on an ionic compound is zero. For example, sodium chloride (NaCl) is an ionic compound which is made up of equal number of positively charged sodium ions (Na⁺) and negatively charged chloride ions (Cl^–). Some of the common ionic compounds, their formulae and the ions present in them are given in Table 1.5.

 

 

Ionic compounds are made up of a metal and a non-metal (except ammonium chloride which is an ionic compound made up of only non-metals). So, whenever a bond involves a metal and a non-metal, we call it ionic bond.

EXERCISE 1.3

 1. Give two examples of ionic compounds. Write their chemical formulae.

2. The overall charge on an ionic compound is zero. (True or False)

3. Name the ions present in calcium nitrate.

4. Ionic compounds are made up of a ______ and a ______ .

5. Give an example of an ionic compound made up of only non-metals.

1.4 FORMATION OF IONIC BOND

An ionic bond changes the electronic configurations of the atoms. Metal atoms lose their outermost electron(s), forming cations. Non-metal atoms gain electron(s) to fill their outermost shell, forming anions. The electrostatic force of attraction between the oppositely charged ions holds the ions together. For example,

(a) When a hot sodium atom is placed in chlorine gas, a reaction takes place

Resulting in formation of sodium chloride

                                  

 (b) When a magnesium atom comes in contact with chlorine gas, it forms magnesium chloride.

 1. With the help of dot and cross, show the formation of CaCl₂.

1.5 PROPERTIES OF IONIC COMPOUNDS

ACTIVITY 1.3: Illustrating Physical Properties of Ionic Compounds

• Take a sample of sodium chloride or any other salt from the science laboratory.

• What is the physical state of this salt?

• Take a small amount of a sample on a metal spatula and heat directly on the flame as shown in figure (a).

                                               

(a) Testing melting point of sodium chloride

• What did you observe? Did the sample impart any color to the flame? Does this compound melt

• Try to dissolve the sample in water, petrol and kerosene. Is it soluble?

• Make a circuit as shown in figure (b) and insert the electrodes into a solution of salt. What did you observe?

                                    

               (b) Testing electrical conductivity of salt solution
              • What is your inference about the nature of this compound?

You may have observed the following general properties for ionic compounds:

• Ionic compounds are usually crystalline solids.

• Ionic compounds have high melting and high boiling points.

The temperature at which a solid melts into liquids is called the melting point of the solid. The temperature does not change during melting.

Boiling point is the temperature at which a liquid changes into a gas. The temperature of a liquid remains the same once boiling has started.

• Ionic compounds are usually soluble in water but insoluble in organic solvents like petrol and kerosene.

• Ionic compounds conduct electricity when dissolved in water or when melted. When we dissolve the ionic solid in water or melt it, the crystal structure is broken down to form ions. These ions help in conducting electricity.

EXERCISE 1.5

 1. Why do ionic compounds conduct electricity when dissolved in water?

2. Ionic compounds are insoluble in

                    (a) kerosene                     (b) petrol

                    (c) both (a) and (b)            (d) neither (a) nor (b)

3. Ionic compounds have low melting and boiling points. (True or False)

4. Ionic compounds are usually ______ solids

1.6 COVALENT BONDING

The chemical bond formed by sharing of electrons between two atoms is known as a covalent bond. The compounds containing covalent bonds are known as covalent compounds. A covalent bond is formed when both the reacting atoms need electrons to achieve the inert gas electron arrangement. Now, the non-metals have usually 5, 6 or 7 electrons in the outermost shells of their atoms. So, all the non-metal atoms need electrons to achieve the inert gas structure. They get these electrons by mutual sharing. Thus, whenever a non-metal combines with another non-metal, covalent bond is formed.

       

        

      

1.7 FORMATION OF COVALENT BOND

Covalent bonding between atoms of different elements. 

(i) Carbon atom shares four electrons to form methane.

                        

 (ii) As in water molecule, 2 hydrogen atoms share electrons with oxygen atom.

                              

             

1.8 PROPERTIES OF COVALENT COMPOUNDS

ACTIVITY 1.4: Illustrating Physical Properties of Covalent Compounds

Let us test some covalent compounds in different ways:

• Take sample of cooking oil. Try to dissolve it in water and ethanol. Does it dissolve?

• Have you ever observed a burning candle wax? If not, take a candle wax and observe it burning. How much time does it take to melt down?

• Take a pan and add water to it. Let it boil. Do you know the boiling point of water?

• Now add two electrodes to the water pan making a circuit. What did you observe? What would have happened if you would have added NaCl salt in the pan?

• What can you now say about these covalent compounds?
You have observed the following properties of covalent compounds:

• Covalent compounds are usually liquids, gases or solids. For example, alcohol, benzene, water and cooking oil are liquids. Methane, ethane and chlorine are gases. Glucose, urea, and wax are solid covalent compounds.

• Covalent compounds have usually low melting points and  low boiling points.

• Covalent compounds are usually insoluble in water, but they are soluble in organic solvents. Some of the covalent compounds like glucose, sugar and urea are soluble in water.

• Covalent compounds do not conduct electricity because they do not contain ions.

ACTIVITY 1.5: Detecting an Ionic Bond or Covalent Bond

• Take the sample such as common salt (NaCl) provided.

• Try to dissolve it in water.

• If it dissolves, chances are it is likely to be an ionic compound. But, you already know some covalent compounds like glucose, urea and sugar are soluble in water.

• Now, perform electrical conductivity test.

• If the NaCl sample dissolves in water, arrange a circuit with two electrodes and a bulb.

• Figure  out whether the bulb glows or not. According to your observation conclude the bond present in the sample.

• Make a report on the properties of ionic and covalent compounds.

                                        

  

  

ACTIVITY 1.6: Identifying Ionic and Covalent Compounds

 Choose the ionic as well as covalent compounds from the bubbles and make a table in your exercise notebook.

                                   

EXERCISE 1.8

 1. Some covalent compounds are solid. (True or False)

2. Most covalent compounds are ______ in water but ______ in organic solvents.

3. Name two covalent compounds which are soluble in water.

4. Why most covalent compounds do not conduct electricity?

5. Melting and boiling points of covalent compounds are

(a) high  (b) low  (c) between 500°C and 1000°C  (d) cannot be determined.


1.9 GIANT COVALENT STRUCTURES

Diamond, graphite and silicon dioxide have giant covalent structures.

1.9.1 Diamond and its Properties

Diamond is a colorless transparent substance having extraordinary brilliance. Diamond is quite heavy. Diamond is extremely hard. It is the hardest natural substance known. Diamond does not conduct electricity. Diamond burns on strong heating to form carbon dioxide. It has a very high melting point. If we burn diamond in oxygen, then only carbon dioxide gas is formed and nothing is left behind. This shows that diamond is made up of carbon only. Since diamond is made up of carbon atoms only, its symbol is taken to be C.

                                 

1.9.2 Graphite and its Properties

Graphite is a greyish-black opaque substance. Graphite is lighter than diamond. Graphite is soft and slippery to touch. Graphite conducts electricity. Graphite burns on strong heating to form carbon dioxide. Like diamond, graphite also has very high melting point. If we burn graphite in oxygen, then only carbon dioxide gas is formed and nothing is left behind. This shows that graphite is made up of carbon only. Since graphite is made up of carbon atoms only, its symbol is taken to be C.

                              

1.9.3 Silicon Dioxide and its Properties

Silicon dioxide (also known as Silica) has a giant covalent structure. Each silicon atom is covalently bonded to four oxygen atoms. Each oxygen atom is covalently bonded to two silicon atoms. This means that, overall, the ratio is two oxygen atoms to each silicon atom, giving the formula SiO₂. Silicon dioxide is very hard. It has a very high melting point (1,610°C) and boiling point (2,230°C). It  is insoluble in water, and does not conduct electricity.

These properties result from the very strong covalent bonds that hold the silicon and oxygen atoms in the giant covalent structure. Silicon dioxide is found as quartz in granite, and is the major compound in sandstone. The sand on a beach is made mostly of silicon dioxide.

                                                   

1.9.4  Uses of Diamond, Graphite and Silicon Dioxide

Uses of Diamond

• Since diamond is extremely hard, it is used for cutting and grinding other hard materials. It is also used for drilling holes in the earth’s rocky layers. Diamond ‘dies’ are used for drawing thin wires like the tungsten filament of an electric bulb. 

• Diamonds are used for making jewelry. The use of diamonds in making jewelry is because of their extraordinary brilliance. Diamond is also used in the tip of glass cutter. A sharp diamond-edged knife called keratome is used by eye surgeons to remove cataract from the eyes.

                  

Uses of Graphite

Due to its softness, powdered graphite is used as a lubricant for fast moving parts of machinery. Graphite can be used as a dry lubricant in the form of graphite powder or mixed with petroleum jelly to form graphite grease. Graphite powder can also be mixed with lubricant oils.
Anode (Zinc Inner Case)
Cathode (Graphite Rod)
Paste of MnO ,
NH Cl, and Carbon 2 4

                                             

                                       

                                 Figure 1.7: Some of the uses of graphite.

• Graphite is a good conductor of electricity due to which graphite is used for making carbon electrodes or graphite electrodes in dry cells and electric arcs. The black colored ‘anode’ of a dry cell is made of graphite. The carbon brushes of electric motors are also made of graphite.

• Graphite is used for making the cores of our pencils called ‘pencil leads’ and black points. Graphite is black in color and quite soft. So, it marks black lines on paper. Due to this property, graphite is used for making pencil leads. For making pencil leads, graphite is usually mixed with clay.

Uses of Silicon Dioxide

• An estimated 95% of silicon dioxide produced is consumed in the construction industry,

   e.g. for the production of Portland cement

• Silica is used primarily in the production of glass for windows, drinking glasses, beverage bottles, and many other uses.

  • The majority of optical fibers for telecommunication are also made from silica.

                                           

                                       Figure 1.8: Some of the uses of silicon dioxide

EXERCISE 1.9

1. Diamond and Graphite are two common all tropes of Carbon. (True or False) 

2. Which of the following is correct? 

(a) Diamond is the hardest substance known. 

(b) Graphite has very low melting point. 

(c) Graphite does not conduct electricity. 

(d) Diamond burns on strong heating to form helium gas.

3. Why are diamonds used for making jewelry?

4. Graphite is used for making 

(a) pencil lead

(b) electrodes 

(c) both (a) and (b)

(d) none of these.

5. Diamond and Graphite have very ______ melting point.

1.10 METALLIC BONDING

The force which binds various metal atoms together is called metallic bond. The metallic bond is neither a covalent bond nor an ionic bond because these bonds are not able to explain properties of metals.

For example, metals are very good conductors of electricity but in solid state. Both ionic and covalent compounds cannot do so with the exception of graphite.

1.11 FORMATION OF METALLIC BOND

Loreutz proposed the theory of electron gas model or electron sea model for metallic bonding.

In this model, the metal is pictured as an array of metal cations in a “sea” of electrons. The atoms in a metallic solid contribute their valence electrons to form a “sea” of electrons that surrounds metallic cations. Delocalized electrons are not held by any specific atom and can move easily throughout the solid. A metallic bond is the attraction between these electrons and the metallic cation.

                                     

EXERCISE 1.10

1. Name the scientist who proposed the theory of electron sea model.

2. Metallic bond is neither a covalent bond nor an electrovalent bond. (True or False)

3. The force which binds various metal atoms together is called ______ .

4. Make a 3D structure of electron sea model.

5. Write a short note on formation of metallic bonding.

1.12 PROPERTIES OF METALLIC BOND

ACTIVITY 1.7: Illustrating the Properties of Metals

• Take samples of iron, copper, aluminum, sodium, carbon and iodine. Note the appearance of each sample.

• Clean the surface of each sample by rubbing them with sand paper and note their appearance again.

• Try to cut these elements with a sharp knife and note your observations.

• Hold a piece of sodium with a pair of tongs.  Caution: Always handle sodium with care. Dry it by pressing between the folds of a filter paper.

• Put it on a watch-glass and try to cut it with a knife.

• What do you observe?

• Place any one element on a block of iron and strike it four or five times with a hammer. What do you observe?

• Repeat above steps with other elements. 

• Record the change in the shape of these elements.

• Which of the above elements are available in the form of wires?

1.12.1 Properties of Metals

ACTIVITY 1.8: Illustrating Conductivity of Heat and Electricity of Metals

• Take an aluminum or copper wire. Clamp this wire on a stand, as shown in Figure (a).

• Fix a pin to the free end of the wire using wax.

• Heat the wire with a spirit lamp, candle or a burner near the place where it is clamped. 

• What do you observe after some time?

• Repeat the same with carbon or Sulphur.

• Note your observations.  Does the element melt?

                                       

• Consider elements aluminum, copper, Sulphur and carbon.

• Set up an electric circuit as shown in Figure (b).

                                       

• Place the element to be tested in the circuit between terminals A and B as shown. Does the bulb glow? What does this indicate?

• Compile your observations regarding properties of elements in your exercise book.

 • Metals are good conductors of heat and electricity. This means that metals allow heat and electricity to pass through them easily. Silver metal is the best conductor of heat. Copper metal is a better conductor of heat than aluminum metal.

• Metals are lustrous (or shiny). This means that metals have a shiny appearance.
Chemical Bonding 17
 • Metals are usually strong. For example, iron metal (in the form of steel) is very strong  when freshly cut and is used in the construction of bridges, buildings and vehicles. Some metals are not strong. For example, sodium and potassium.

• Metals are ductile. This means that metals can be drawn (or stretched) into thin wires.

• Gold and silver are among the best ductile metals.

• Metals are malleable. This means that metals can be hammered into thin sheets.
The

The cooking utensils are made of metals because metals are good conductors of heat.

EXERCISE 1.11

1. Name the metal which is the best conductor of electricity.

2. Aluminum is a better conductor of heat than copper.  (True or False)

3. Metals are _______. This means that they can be hammered into thin sheets.

4. Why are cooking utensils made of metals?

5. Which of the following statement(s) is/are correct for metals? 

(a) Metals such as sodium and potassium are not strong. 

(b) Iron is used in the construction of buildings. 

(c) Gold and Silver are among the best  ductile metals. 

(d) All of these.

1.13 SUMMARY

• An atom achieves a stable electronic configuration by losing, gaining or sharing electrons.

♦ Metal atoms with one, two or three electrons in the outermost shell lose  electron(s) to form positively charged ions (cations).

♦ Non-metal atoms with five, six or seven electrons in the outermost shell gain three, two and one electron(s) to form negatively charged ions (anions).

♦ Non-metal atoms with four to seven outermost electrons may gain electrons by sharing them with each other.

• A chemical bond is a force that holds ions, molecules or atoms together. A bond is formed when each atom acquires a stable electronic configuration like noble gas.

• The  electrostatic binding force is called an ionic bond or electrovalent bond.

• Ionic compounds are formed by attraction of positive and negative ions. These compounds are crystalline solid. They conduct electricity. Ionic compounds have high melting and boiling points. 

• A covalent bond forms between two or more atoms of non-metals that are unable to form ions.
 • Covalent compound is formed when atoms achieve a stable electronic configuration by sharing of electrons. Covalent compounds are solids, liquid or gases. Covalent compounds have low melting and boiling points.

• The two forms of carbon that join covalently to form giant structure are diamond and graphite.

• The force which binds various metal atoms together is called metallic bond.

• Metals are generally hard, lustrous, strong, malleable and ductile. They conduct heat and electricity in both molten and solid state.

1.14 GLOSSARY

• Anion: a negatively charged ion.

• Cation: a positively charged ion.

• Crystal: a solid where the atoms form a periodic arrangement.

• Diamond: one of the known all otropes of carbon.

• Ductile: able to be drawn out into a thin wire.

• Electronic configuration: the distribution of electrons of an atom.

• Graphite: a grey crystalline allotropic form of carbon which occurs as a mineral in some rocks.

• Malleable: able to be hammered or pressed into shape without breaking or cracking.

• Noble gas: the gaseous elements helium, neon, argon, krypton, xenon, and radon.

1.15 UNIT ASSESSMENT

I. Multiple Choice Questions

1. The number of electrons gained by non-metals to achieve noble gas electronic configuration is

(a) one (b) two (c) three (d) all of these

2. The electronic configuration of sodium ion is

(a) 2,8,1 (b) 2,8,8 (c) 2,8 (d) 2,8,2

3. The electronic configuration of chloride ion is

(a) 2,8 (b) 2,8,8 (c) 2,8,7 (d) 2,8,3 

4. Choose the ionic compound.

(a) Calcium chloride   (b) Copper sulphate (c) Sodium hydroxide   (d) All of these

5. Most ionic compounds are soluble in

(a) water (b) petrol (c) kerosene (d) all of these

 6. Which of these is not a covalent compound?

(a) Carbon dioxide (b) Methane (c) Ammonia (d) None of these

7. Choose the correct statement.

(a) Covalent compounds have low melting points (b) Ionic compounds have high melting points (c) Urea and glucose are solid covalent compounds (d) All of these

8. Graphite is used for making ____________.

(a) lubricant oils (b) pencil leads (c) both (a) and (b) (d) jewelry

9. If we burn diamond, the product formed is ____________.

(a) carbon dioxide   (b) hydrogen gas (c) hydrogen chloride gas (d) oxygen gas

10. The force which binds various metal atoms together is called ____________.

(a) metallic bond (b) covalent bond (c) ionic bond (d) none of these

II. Open Ended Questions

 1. How can an atom achieve stability?

2. Distinguish between covalent and ionic bond.

3. Compare between the properties of ionic and covalent compounds.

4. Explain the formation of sodium ion.

5. Give five examples of each

       (a) Ionic compounds            (b) Covalent compounds

6. Compare the conductivity of distilled water with sodium chloride solution.

7. Write two uses of diamond.

8. Draw the structure of graphite.

9. Illustrate the physical properties of metals

. III. Practical-based Questions

1. Look at the figures and choose the correct statement.

 (a) Figure A is an example of  ionic compound

(b) Figure B is not an example of covalent compound

(c) Both Figure A and Figure B are covalent compounds

(d) None of these

2. The following figure illustrates the electronic configuration of

 (a) Lithium                 (b) Sodium                      (c) Chlorine                            (d) Helium

3. shows the structure of ...........................

                                               
 (a) the hardest substance known

(b) an allotrope of carbon

(c)  both (a) and (b)  

(d) none of these

4. Which of the following materials makes the circuit complete when inserted in between the crocodile clips?

                         

                          (a) Aluminum foil                                (b) Copper wire

                          (c) Both (a) and (b)                              (d) Sulphur

 5. In the given figure, arrow shows the

                                           

 (a) carbon rod                (b) iron rod               (c) brass rod                         (d) copper rod

6. Which of the following depicts the molecule of water?