Chemical Bonding PDF

# Chemical Bonding ## Syllabus - Electrovalent, covalent and co-ordinate bonding, structures of various compounds. Electron dot structure. - Electrovalent Bonding. - Electron dot Structure of Electrovalent compounds NaCl, MgCl2, CaO. - Characteristic properties of electrovalent compounds s...

# Chemical Bonding ## Syllabus - Electrovalent, covalent and co-ordinate bonding, structures of various compounds. Electron dot structure. - Electrovalent Bonding. - Electron dot Structure of Electrovalent compounds NaCl, MgCl2, CaO. - Characteristic properties of electrovalent compounds state of existence, melting and boiling points, conductivity (heat and electricity), dissociation in solution and in molten state to be linked with electrolysis. - Covalent Bonding. - Electron dot structure of covalent molecules on the basis of duplet and octet of electrons (example: hydrogen, oxygen, chlorine, nitrogen, ammonia, carbon tetrachloride, methane). - Polar covalent compounds – based on difference is electronegativity – Examples – HCl, NH3 and H2O including structures. - Characteristic properties of Covalent compounds state of existence, melting and boiling points, conductivity (heat and electricity), ionisation in solution. - Comparison of Electrovalent and Covalent compounds. - Coordinate Bonding. - Definition - The lone pair effect of the oxygen atom of the water molecule and the nitrogen atom of the ammonia molecule to explain the formation of H3O+ and OH- ions in water and NH4+ ion. - The meaning of lone pair; the formation of hydronium ion and ammonium ion must be explained with the help of electron dot diagrams. ## 2.1 Introduction Everything in this world wants stability, so is the case with atoms. For atoms, stability means having the electron arrangement of an inert gas, i.e., octet in its outermost shell. Helium has two electrons (DUPLET) while all other inert gases, i.e., Neon, Argon, Krypton, Xenon and Radon have eight electrons (OCTET) in their outermost shell, as given in Table 2.1 below. | Inert gas | Atomic No. | Electronic configuration | No. of valence electrons | | ----- | ----- | ----- | ----- | | He | 2 | 2 | 2 | | Ne | 10 | 2 8 | 8 | | Ar | 18 | 2 8 8 | 8 | | Kr | 36 | 2 8 18 8 | 8 | | Xe | 54 | 2 8 18 18 8 | 8 | | Rn | 86 | 2 8 18 32 18 8 | 8 | It is found that the elements with their complete outermost shell do not react or are least reactive. We, therefore, conclude that the atoms having 8 electrons (or 2 electrons, Helium configuration) in their outermost shells are very stable and unreactive. Therefore, to attain stability, atoms tend to combine chemically by redistribution of electrons in the outermost shell or valence electrons so that each is left with a stable electronic configuration (duplet or octet). Cause of chemical combination is the tendency of elements to acquire the nearest noble gas configuration in their outermost orbit and become stable. During redistribution of electrons, a force of attraction develops between atoms, which binds them together to form molecules. This force of attraction is known as the chemical bond. A chemical bond may be defined as the force of attraction between any two atoms, in a molecule, to maintain stability. There are three methods in which atoms can achieve a stable configuration: 1. The transfer of one or more electrons from one atom to the other to form an electrovalent (or an ionic) bond. 2. Sharing of one, two or three pairs of electrons between two atoms to form a covalent (or a molecular) bond. 3. When the shared electron pairs are contributed by only one of the combining atoms, the bond formed is known as coordinate (or dative) bond. ## 2.2 Electrovalent (Or Ionic) Bond Atoms of metallic elements that have 1, 2 or 3 valence electrons can lose electron(s) to atoms of non-metallic elements, which have 7, 6 or 5 electrons respectively in their outermost shell and thereby forming an electrovalent compound. After the transfer of electron(s), both the combining atoms acquire the electronic configuration of the nearest inert gas. A metallic atom, which loses electron(s), becomes a positively charged ion and is known as a cation and a non-metallic atom, which gains electron(s), becomes a negatively charged ion and is known as an anion. An ion is a charged particle which is formed due to the gain or the loss of one or more electrons by an atom. A metallic element, whose one atom readily loses electron(s) to form a positively charged ion, is an electropositive element. Na – e- → Na+ (cation) A non-metallic element, whose atom readily accepts electron(s) to form a negatively charged ion, is an electronegative element. Cl + e- → Cl- (anion) The cation and the anion being oppositely charged attract each other and form a chemical bond. Since this chemical bond formation is due to the electrostatic force of attraction between a cation and an anion, it is called an electrovalent (or an ionic) bond. ## 2.2.1 Conditions for the formation of an electrovalent (or ionic) bond When an ionic compound is formed, the neutral atom is changed to a cation or an anion. The formation of cations and anions depends on the following factors: 1. Low ionisation potential: If the ionisation potential of a particular atom is low, it will lose electron(s) easily, i.e., a cation is formed easily. 2. High electron affinity: If the electron affinity value is high, anion will be formed easily, i.e., a higher electron affinity value favours ionic bonding. 3. Large electronegativity difference: If the difference in the electronegativities of two elements is higher, then the transfer of electrons will be easier. Therefore, more the difference in electronegativity, more will be the ionic nature of the resulting compound. The metals of groups 1, 2 and 13 have a tendency to lose their valence electrons. So they combine with the non-metals of Groups 15, 16 and 17, which have a tendency to gain electron(s) and form ionic bonds. _Note:_ Group 1 elements are most electro-positive, i.e., they are metallic in nature. Their metallic nature increases down the group. Group 17 elements are most electronegative. Fluorine is the most electronegative element. Thus, caesium fluoride CsF is the most ionic compound. Bonds formed between metals and non-metals are ionic or electrovalent. ## Why are Ionic Compounds Stable? Ionic compounds are formed by ions but there also exists a repulsive force between ions for like charges. Since the electrostatic force of attraction between opposite charges is much higher, it makes the ionic compounds stable. ## 2.2.2 Structures of some electrovalent compounds ### 1. Sodium chloride (NaCl) The electronic configuration of a sodium atom is 2, 8, 1. It has one electron in excess of the stable electronic configuration of the nearest noble gas, neon, (2, 8). Therefore, an atom of sodium shows a tendency to give up the electron from its outermost shell, so as to acquire a stable electronic configuration of neon. | Sodium atom | Sodium cation | | ----- | ----- | | Na (2, 8, 1) | Na+ (2, 8) | However, after giving up one electron, the sodium atom is no more electrically neutral. It has eleven protons in its nucleus but only ten electrons revolving around it. Therefore, it has a net positive charge of +1. This positively charged atom is called sodium ion and is written as Na+ and its electronic configuration resembles that of the noble gas neon. The properties of Na are different from Na+. | Property | Sodium atom (Na) | Sodium cation (Na+) | | ----- | ----- | ----- | | Colour | Silvery white | Colourless | | Toxicity | Poisonous | Non-poisonous | | Chemical action | Very active | Inactive | | Valence shell | Incomplete outermost shell | Complete outermost shell | | Electrical state | Neutral | Positively charged | | Existence | Combined state | Independent existence | The electronic configuration of chlorine is 2, 8, 7. It has an electronic configuration with one electron less than that of the nearest noble gas, argon (2, 8, 8). Therefore, the chlorine atom shows a tendency to acquire an electron to attain octet in its outermost shell. | Chlorine atom | Chloride anion | | ----- | ----- | | Cl (2, 8, 7) | Cl- (2, 8, 8) | An atom of chlorine is electrically neutral, as it contains 17 protons in its nucleus and 17 electrons revolving around it. But, after acquiring an electron from the sodium atom, the chlorine atom does not remain electrically neutral. It has one electron more than the number of protons in its nucleus and therefore has charge of -1 represented as Cl- i.e. chloride ion. The properties of chlorine atom are different from its ion. | Property | Chlorine atom (Cl) | Chloride anion (Cl-) | | ----- | ----- | ----- | | Colour | Yellowish green (as Cl2 gas) | Colourless | | Toxicity | Poisonous | Non-poisonous | | Odour | Suffocating | Odourless | | Chemical action | Very active | Inactive | | Valence shell | Incomplete outermost shell | Complete outermost shell | | Electrical state | Neutral | Negatively charged | | Existence | Not independent | Independent | Chloride ion has an octet of electrons in its outermost shell, and its electronic configuration resembles that of the noble gas argon (Fig. 2.1). Thus, when an atom of sodium combines with an atom of chlorine, one electron is transferred from the sodium atom to the chlorine atom, resulting in the formation of a sodium chloride molecule. *Sodium atom* Na (2, 8, 1) Na-e- Na+ *Chlorine atom* Cl (2, 8, 7) Cl + e- Cl- The cation Na+ and anion Cl- are attracted towards each other, due to opposite electrical charge or coulomb force existing between them and form an ionic compound. ### Electron dot symbol (Lewis symbol) The electron dot symbol for an atom consists of the symbol of the element surrounded by dots representing only the outermost shell electrons. The paired electrons are represented by a pair of dots, whereas the unpaired electron in the outermost orbit is represented by a single dot. Example: Electron dot symbol of Hydrogen is H and of Oxygen is:O: Symbols other than dots, such as circles and crosses can be used to distinguish between the electrons of different atoms in a molecule, for example: Ammonia (NH3) can be represented as ``` H N H X X H ``` ### Electron dot structure of NaCl ``` .. Na +Cl: → Na+ + Cl- or NaCl ``` ### 2. Magnesium chloride (MgCl2) The number of valence electrons of magnesium (atomic number 12) is 2 and that of chlorine (atomic number 17) is 7. Magnesium atom acquires a stable configuration of 8 electrons by losing two electrons from its outermost shell (one each to each atom of chlorine) and thus becomes a positive magnesium ion, Mg2+. | Magnesium atom | Magnesium cation | | ----- | ----- | | Mg (2, 8, 2) | Mg2+ (2, 8) | However, each chlorine atom, which contains 7 electrons in its outermost shell, can accept only 1 of the 2 electrons donated by a magnesium atom. Therefore, for each magnesium atom forming a magnesium ion, there must be two chlorine atoms to form two chloride ions. ``` 2C1 (2, 8, 7) 2 Cl atoms + 2e- 2Cl- (2, 8, 8) 2 chloride anions ``` Thus, the ratio of magnesium to chloride ions in magnesium chloride must be 1: 2, so the molecular formula of the compound magnesium chloride is MgCl2 (Fig. 2.2). ### Electron dot structure of magnesium chloride ``` 00 Mg + → Mg2+ 2[:C1:] 00 CI 00 ``` ### 3. Calcium oxide (CaO) The number of valence electrons of a calcium atom (atomic number 20) is 2, and that of an oxygen atom is 6, i.e., oxygen requires 2 electrons to attain octet. In the presence of oxygen, each calcium atom loses its 2 valence electrons to one oxygen atom (Fig. 2.3). As a result, the calcium atom forms a calcium ion with charge +2 (Ca2+), and the oxygen atom forms an oxide ion with charge -2 (02-). Since only one oxygen atom is needed to accept the 2 valence electrons donated by a calcium atom, the formula of calcium oxide is CaO and not Ca202. | Calcium atom | Calcium cation | | ----- | ----- | | Ca (2, 8, 8, 2) | Ca2+ (2, 8, 8) | | Oxygen atom | Oxide anion | | ----- | ----- | | O (2, 6) | O2- (2, 8) | In the formation of an electrovalent bond, the transfer of electron(s) is involved. The electropositive atom undergoes oxidation, while the electronegative atom undergoes reduction. This is known as REDOX PROCESS. ### For example: #### Formation of sodium chloride: Sodium chloride is formed by the combination of sodium and chlorine. 2Na + Cl2 → 2 NaCl The reaction can be written as two half reactions: 2Na → 2Na+ + 2e (Oxidation) Cl2 + 2e- → 2Cl- (Reduction) Oxidation 2Na + Cl2 → 2Na+ + 2Cl- (Redox Reaction) 2 Reduction _Note:_ Oxidation and reduction always occur simultaneously because the electron(s) lost by the reducing agent must be gained by the oxidising agent. #### For example: | Oxidising agent | Oxidised | Reducing agent | Reduced | | ----- | ----- | ----- | ----- | | Fe2O3 | 2Fe | 2A1 | Al¬2O3 | In this reaction, Aluminium acting as a reducing agent reduces Iron (III) oxide to Iron. This is a reduction reaction. Fe3+ + 3e- → Fe (Reduction) * Orbit structure not in syllabus At the same time, Aluminium is oxidised to Aluminium oxide by the oxidising agent Iron (III) oxide, and this is an oxidation reaction. Al-3e Al3+ (Oxidation) Thus, the net reaction is a redox reaction. It can be inferred from the above example that an oxidising agent is an acceptor of electron(s) and a reducing agent is a donor of electron(s). ## Intext Questions 1. How do atoms attain noble gas configuration? 2. Define: - a chemical bond - an electrovalent bond - a covalent bond 3. What are the conditions for the formation of an electrovalent bond? 4. An atom X has three electrons more than the noble gas configuration. What type of ion will it form? Write the formula of its (i) sulphate (ii) nitrate (iii) phosphate (iv) carbonate (v) hydroxide. 5. Mention the basic tendency of an atom which makes it to combine with other atoms. 6. The element X has the electronic configuration 2, 8, 18, 8, 1. Without identifying X, - predict the sign and charge on a simple ion of X. - write if X will be an oxidising agent or a reducing agent and why. 7. In the formation of the compound XY2, an atom X gives one electron to each Y atom. What is the nature of bond in XY2? Draw the electron dot structure of this compound. 8. An atom X has 2,8,7 electrons in its shell. It combines with Y having 1 electron in its outermost shell. - What type of bond will be formed between X and Y? - Write the formula of the compound formed. 9. Explain with the help of ionic equation and electron dot structural diagram the formation of the following electrovalent compounds. - NaCl - MgCl¬2 - CaO 10. Compare: - sodium atom and sodium ion - chlorine atom and chloride ion, with respect to - atomic structure - electrical state - chemical action - toxicity. 11. The electronic configuration of Fluoride ion is the same as that of a neon atom. What is the difference between the two? 12. State which of the following are reduction reaction and which are oxidation - Pb → Pb2+ + 2e- - Fe2+ → e-Fe3+ - A3+ + e → A2+ - Cu→ Cu2+ 13. What do you understand by redox reactions? 14. Explain: - oxidation and reduction in terms of loss or gain of electrons. - Formation of electrovalent compound is a redox reaction. 15. Divide the following redox reactions into oxidation and reduction half reactions. - Zn + Pb2+ → Zn2+ + Pb - Zn + Cu2+ → Zn2+ + Cu - Cl2 + 2Br→ Br2 + 2Cl- - Sn2+ + 2Hg2+ → Sn4+ + Hg22+ - 2Cu+ Cu + Cu2+ 16. Potassium (at No. 19) and chlorine (at No. 17) react to form a compound. Explain on the basis of electronic concept: - oxidation - reduction - oxidising agent - reducing agent. ## 2.3 Covalent (Molecular) Bond The chemical bond that is formed between two combining atoms by mutual sharing of one or more pairs of electrons is called a covalent (or a molecular) bond and the compound formed due to this bond is called a covalent compound. The molecule formed due to the sharing of electrons (covalent bond) is called a covalent molecule. The atoms of non-metals usually have 5, 6 or 7 electrons in their outermost shell (except carbon which has 4 and hydrogen which has just 1 electron in the outermost shell). The atoms of such elements do not favour the loss of its electrons due to energy considerations and thus the transfer of electrons is not possible. Therefore, this atom can complete its octet only by mutually sharing one or more pairs of electrons. Each atom contributes equal number of electron(s). So, whenever a non-metal combines with another non-metal (to attain stable configuration), the sharing of electrons takes place between their atoms and a covalent bond is formed. For example, hydrogen is a non-metal and chlorine is also a non-metal. When hydrogen combines with chlorine to form hydrogen chloride (HCl), the sharing of electrons takes place between hydrogen and chlorine atoms and a covalent bond is formed. H+ + :Cl: → H:Cl: It should be noted that a covalent bond can also be formed between two atoms of the same non-metal. For example, two chlorine atoms combine together by the sharing of electrons to form a chlorine molecule (Cl¬2) and a covalent bond is formed between the two chlorine atoms. ## Covalent bonds are of following three types: 1. Single covalent bond 2. Double covalent bond 3. Triple covalent bond A single covalent bond is formed by the sharing of one pair of electrons between the atoms, each atom contributing one electron. It is denoted by putting a short line (-) between the two atoms. Molecules of hydrogen, chlorine, hydrogen chloride, water, ammonia, methane and carbon tetrachloride are examples of single covalent bonds. - H-H Hydrogen - CI-CI Chlorine - H-CI Hydrogen chloride - H-O-H Water - H H-N-H H Ammonia - H H-C-H H Methane - CI CI-C-CI | CI Carbon tetrachloride Similarly, a double bond is formed by the sharing of two pairs of electrons between two atoms. A double bond is actually a combination of two single bonds, so it is represented by putting two short lines (=) between the two atoms. For example, oxygen molecule, O2, contains a double bond between two atoms and it can be written as O=O. Carbon dioxide contains two double bonds. CO2 O=C=O A triple bond is a combination of three single bonds. Nitrogen molecule, N2, contains a triple bond, so it is written as N = N. Some molecules have a combination of single bond as well as a double or a triple bond. For example, ethene (C2H4) molecule has one double covalent bond and four single covalent bonds. - H H >C=C H H Ethene (C2H4) Similarly, ethyne (C2H2) molecule has one triple covalent bond and two single covalent bonds. Ethyne (C2H2) H − C = C - H The covalency of an atom is the number of its electrons taking part in the formation of shared pairs. Thus, the covalency of hydrogen is 1, oxygen 2, nitrogen 3 and carbon 4. ## 2.3.1. Non-polar covalent compounds Covalent compounds are non polar when shared pair of electron(s) are equally distributed between the two atoms. No charge separation takes place. The molecule is symmetrical and electrically neutral. If the two covalently bonded atoms are identical the shared electron pair(s) is at equal distance from the combining atoms i.e., the shared electron pair(s) is equally attracted by the nuclei of the two types of charge, such molecules are non-polar. For example hydrogen (H2), chlorine (Cl2), oxygen (O2), etc., are perfectly non-polar compounds. These compounds do not ionise in water due to lack of charge separation. The bond formed between dissimilar atoms can be non-polar if their electronegativity difference is little and their structure permits the shared pair of electrons to attract equally the linked atoms and thus the molecule becomes symmetrical. For example methane, carbon tetrachloride, etc. (see structure 4 and 5, article 2.3.4). The shared electron pairs in methane are at equal distance from the carbon and the hydrogen atoms, because the two have nearly equal electronegativities (carbon = 2.5, hydrogen = 2.1). ## 2.3.2. Polar covalent compounds The covalent compounds are said to be polar when the shared pair of electrons are not at equal distance between the two atoms. This results in the development of fractional positive and negative charges on them. They ionise in water. For example, hydrogen chloride, hydrogen fluoride. In hydrogen chloride, the strong nuclear charge of the chlorine atom (the electro-negativity of chlorine is 3) attracts the shared electron pair towards itself, i.e., negative charge shifts towards the chlorine atom thereby developing a slight negative charge (δ-) on it. The hydrogen atom (electronegativity 2.1) develops a slight positive charge (δ+). Therefore, a polar covalent bond is formed. This is shown below. H+ + :Cl: → H*CI: → H+ δ+ → Cl δ- The arrow on the line indicates that the shared pair is shifted towards that atom. Other examples of polar covalent compounds that ionise are water (see structure 6 & 7, article 2.3.4), ammonia and hydrogen fluoride. ``` Oδ- Hδ+ Hδ+ Water molecule N Hδ+ Hδ+ Hδ+ Ammonia molecule Hδ+ -Fδ- Hydrogen fluoride molecule ``` In solution, the fractional charges of polar covalent compounds are converted to complete charges and ions are produced. The process by which covalent compounds are converted into ions is called ionisation. _Note:_ 1. The more the electronegativity difference between two atoms forming a bond, the more is the polar nature of the molecule. 2. The bond formed between two atoms - with same electronegativity is non-polar - with slightly different electronegativity is polar - with much electronegativity difference is ionic (more than 1-7). 3. Since a polar covalent molecule has both positive and negative poles, it is also known as a 'dipole molecule'. A molecule that has both, slight positive and slight negative charge is called a Dipole molecule. ## 2.3.3 Conditions for The Formation of a Covalent Bond 1. Both atoms should have four or more electrons in their outermost shells, i.e., non-metals (exceptions are H, Be, B, Al, etc.). 2. Both the atoms should have high electronegativity. 3. Both the atoms should have high electron affinity. 4. Both the atoms should have high ionization energy. 5. The electronegativity difference between the combining atoms should either be zero or negligible. ## 2.3.4 Some Covalent Molecules and Their Structures ### 1. Hydrogen molecule (Non-polar compound) A hydrogen atom has one electron in its only shell. It needs one more electron to attain duplet. To meet this need, hydrogen atom shares its electron with another hydrogen atom. Thus, one electron each is contributed by the two atoms of hydrogen, and the resulting pair of electrons is mutually shared by both the atoms to form a hydrogen molecule. #### Formation of a hydrogen molecule *Electron dot structure* *Before combination* H. + H. H-atom H-atom *After combination* HH or H-H (H2) One shared pair of electrons Single covalent bond [A hydrogen molecule contains two atoms of hydrogen] ### 2. Chlorine molecule (Non-polar compound) | Electronic configuration | Nearest noble gas | | ----- | ----- | | 17CI [2, 8, 7] | Argon (18Ar) [2, 8, 8] | To attain the stable electronic configuration of the nearest noble gas, chlorine needs one electron. When two chlorine atoms come closer, each contributes one electron and form one shared pair of electrons between them. Both the atoms of chlorine thus attain an octet. A single covalent bond [Cl - Cl] is formed between them. #### Formation of a chlorine molecule *Electron dot structure* *Before combination* :CI + CI: Cl-atom Cl-atom *After combination* ::: or [C1-C1] (Cl2) One shared pair of electrons Single covalent bond [A chlorine molecule contains two atoms of chlorine] ### 3. Nitrogen molecule (Non-polar compound) | Electronic configuration | Nearest noble gas | | ----- | ----- | | Nitrogen (7N) [2,5] | Neon (10Ne) [2, 8] | To attain the stable electronic configuration of the nearest noble gas, nitrogen needs three electrons. When two nitrogen atoms come closer, each contributes three electrons and so they have three shared pairs of electrons between them. Both atoms attain an octet, resulting in the formation of a triple covalent bond [N = N] between them. #### Formation of a nitrogen molecule *Electron dot structure* *Before combination* ·· N : + : N :: N-atom N-atom *After combination* ··· N≡N or [N=N] (N2) Mutual sharing of three pairs of electrons Triple covalent bond Three shared pairs of electrons [A nitrogen molecule contains two atoms of nitrogen] ### 4. Carbon tetrachloride molecule (Non-polar compound) | Atoms involved | Electronic configuration | Nearest noble gas | | ----- | ----- | ----- | | Carbon | C [2, 4] | Neon [2, 8] | | Chlorine | 17C1 [2, 8, 7] | Argon [2, 8, 8] | To attain the stable electronic configuration of the nearest noble gas, carbon needs four electrons and chlorine needs one electron. When a molecule of carbon tetrachloride is to be formed, one atom of carbon shares four electron pairs, one with each of the four atoms of chlorine. Though the bond between carbon and chlorine is polar but molecule of CCl4 as a whole is non-polar, as the molecule is symmetrical in shape. #### Formation of a carbon tetrachloride molecule *Electron dot structure* *Before combination* :CI: + :CI: + C + :CI: + :CI: *After combination* :CI: XCX Ci: X :CI: or Cl CI-C-CI (CCI4) CI One shared pair of electrons with each chlorine atom Four single covalent bonds [One molecule of carbon tetrachloride contains five atoms in all, i.e. one atom of carbon and four atoms of chlorine] ### 5. Methane molecule (Non-polar compound) | Atoms involved | Electronic configuration | Nearest noble gas | | ----- | ----- | ----- | | Carbon | C [2, 4] | Neon [2, 8] | | Hydrogen | H[1] | Helium [2] | To attain the stable electronic configuration of the nearest noble gas, carbon needs four electrons and hydrogen needs one electron. When a molecule of methane is to be formed, one atom of carbon shares four electron pairs, one with each of the four atoms of hydrogen. #### Formation of a methane molecule *Electron dot diagram* *Before combination* H X H x+C+ X H x+C+ X H *After combination* - H H-C-H H or H H-C-H | H Four single covalent bonds [A methane molecule contains a total of five atoms, ie. one atom of carbon and four atoms of hydrogen] ### 6. Water molecule (Polar compound) | Atoms involved | Electronic configuration | Nearest noble gas | | ----- | ----- | ----- | | Hydrogen | Η [1] | Helium [2] | | Oxygen | Ο [2, 6] | Neon [2, 8] | To attain the stable electronic configuration of the nearest noble gas, hydrogen needs one electron and oxygen needs two electrons. In the case of a water molecule, each of the two hydrogen atoms shares an electron pair with the oxygen atom such that hydrogen acquires a duplet configuration and oxygen an octet, resulting in the formation of two single covalent bonds. #### Formation of a water molecule *Electron dot structure* *Before combination* H x H + O + x H x H H-atom O-atom H-atom *After combination* H X O X H or H-O-H (H2O) One shared pair of electrons on each side Two single covalent bonds [One molecule of water contains a total of three atoms, ie. one atom of oxygen and two atoms of hydrogen] ### 7. Ammonia molecule (Polar compound) | Atoms involved | Electronic configuration | Nearest noble gas | | ----- | ----- | ----- | | Nitrogen | N [2, 5] | Neon [2, 8] | | Hydrogen | Η [1] | Helium [2] | To attain the electronic configuration of the nearest noble gas, nitrogen needs three electrons and hydrogen needs one electron. When a molecule of ammonia is to be formed, one atom of nitrogen shares three electron pairs, one with each of the three atoms of hydrogen. #### Formation of an ammonia molecule *Electron dot structure* *Before combination* H + H * + *N * + *H *After combination* - H HNH H OR H H-N-H H One shared pair of electrons with each hydrogen atom Three single covalent bonds [The ammonia molecule contains a total of four atoms, ie, one atom of nitrogen and three atoms of hydrogen] ## 2.4 Properties and Comparison of Electrovalent and Covalent Compounds | Property | Electrovalent compounds | Reason | Property | Covalent compounds | Reason | | ----- | ----- | ----- | ----- | ----- | ----- | | 1. Nature | (i) Their constituent particles are ions. (ii) They are hard solids consisting of ions. | These have strong electrostatic forces of attraction between their ions, which cannot be separated easily. | (i) Their constituent particles are molecules. (ii) These are gases or liquids or soft solids. | They have weak forces of attraction between their molecules. | | 2. Boiling point and melting point | These are non-volatile, with high boiling and high melting points. | There exists a strong force of attraction between the oppositely charged ions, so a large amount of energy is required to break the strong bonding force between ions. | These are volatile, with low boiling and low melting points. | They have weak forces of attraction between the binding molecules, thus less energy is required to break the force of bonding. | | 3. Electricity conducting nature | (i) They do not conduct electricity in the solid state. (ii) They are good conductors of electricity in the fused or in aqueous state. | Electrostatic forces of attraction between ions in the solid state are very strong. These forces weaken in fused state or in solution state. Hence, ions become mobile. | They are non-conductors of electricity in solid, molten or aqueous state. | Due to the absence of free ions. | | 4. Dissociation | Electrovalent compounds are composed of ions. In solution, these ions become mobile. Or in molten state these ions dissociate. Their ions dissociate and migrate when an electric current passes through them in their molten or aqueous solution state. e.g. NaCl Na+ + Cl- • Electrovalent compounds are good conductors of heat. | Water being a polar covalent compound decreases the electrostatic forces of attraction, resulting in free ions in aqueous solution. NaCl Na+(aq) + Cl-(aq) Ions dissociate in water or in molten state. | 4. Ionisation in solution | On passing electric current, non-polar covalent compounds do not ionise. Some covalent compounds are polar in nature. They ionize in their solutions and can act as an electrolyte. e.g. HCI + H2O HO+ + Cl- The dissociation of molecules into ions does not take place in covalent molecules. • Covalent compounds are poor conductors of heat. | Covalent compounds do not have ions. Polar covalent molecules form covalent ions in their solutions. | | 5. Solubility | These are soluble in water but insoluble in organic solvents. | As water is a polar compound, it decreases the electrostatic forces of attraction, resulting in free ions in aqueous solution. Hence they dissolve. | 5. Solubility | These are insoluble in water but dissolve in organic solvents. | Covalent compounds do not have ions, so they do not dissociate. As organic solvents are non-polar, hence, these dissolve in non-polar covalent compounds. | | 6. Rate of reaction | They show rapid speed of chemical reactions in aqueous solutions. | Since free ions are easily formed in different solutions, they unite very fast forming compounds. | 6. Rate of reaction | They show slow speed of chemical reactions in aqueous solutions. | In covalent molecules, old bonds are broken and new bonds are formed, thus the reaction is slow between covalent compounds. | ## 2.5 Effect of Electricity on Electro-Valent and Covalent Compounds _Experiment:_ Arrange an electrolytic cell as shown in Fig. 2.4. The electric circuit contains a 6-volt battery, an ammeter bulb and platinum electrodes connected in series. **6 VOLT BATTERY** (-)* **AMMETER** (PLATINUM WIRE) ANODE ELECTROLYTIC CELL (+) (-) **BULB** CATHODE WATER OR SOME LIQUID _(As a convention, on a diagram, anode is shown on left and cathode on right)._ *Fig. 2.4 Effects of electric current in different solutions.* Take alcohol, chloroform, benzene, petrol and the solutions of sugar, magnesium chloride, sodium chloride, sodium hydroxide, copper sulphate in separate beakers and dip the platinum electrodes in them for a moment, one by one, and note the change in ammeter reading. **Observations:** When the current is passed in the solutions of magnesium chloride, sodium chloride, sodium hydroxide and copper sulphate the bulb glows. This shows that electrovalent compounds allow electric current to pass through them. _Thus, ionic compounds are good conductors of electricity in molten or aqueous states due to free mobile ions._ When the current is passed through the solutions of covalent compounds, i.e. distilled water, sugar solution, alcohol, chloroform, benzene and petrol the bulb does not glow. This shows that they do not conduct electricity. This happens because solutions of covalent compounds contain only molecules and no ions. ## 1. What are the conditions necessary for the formation of covalent molecules? ## 2. Elements A, B and C have atomic numbers 17,19 and 10 respectively.