{"id":2302,"date":"2023-10-02T12:04:46","date_gmt":"2023-10-02T12:04:46","guid":{"rendered":"http:\/\/localhost\/ecole9ja\/?p=2302"},"modified":"2023-10-02T12:15:14","modified_gmt":"2023-10-02T12:15:14","slug":"week-6-ss1-second-term-chemistry-notes","status":"publish","type":"post","link":"https:\/\/ecolebooks.com\/nigeria\/posts\/week-6-ss1-second-term-chemistry-notes\/","title":{"rendered":"Week 6 – SS1 Second Term Chemistry Notes"},"content":{"rendered":"

\u00a0WEEK 6
\n<\/strong>CHEMICAL BONDING
\n<\/strong>INTRODUCTION
\n<\/strong>Though the periodic table has only 118 or so elements, there are obviously more substances in nature than 118 pure elements. This is because atoms can react with one another to form new substances called compounds . Formed when two or more atoms chemically bond together, the resulting compound is unique both chemically and physically from its parent atoms.
\nLet’s look at an example.\u00a0 The element sodium is a silver-colored metal that reacts so violently with water that flames are produced when sodium gets wet.\u00a0 The element chlorine is a greenish-colored gas that is so poisonous that it was used as a weapon in World War I.\u00a0 When chemically bonded together, these two dangerous substances form the compound sodium chloride, a compound so safe that we eat it every day – common table salt!
\n\"\"
\n\t\tIn 1916, the American chemist Gilbert Newton Lewis proposed that chemical bonds are formed between atoms because electrons from the atoms interact with each other. Lewis had observed that many elements are most stable when they contain eight electrons in their valence shell. He suggested that atoms with fewer than eight valence electrons bond together to share electrons and complete their valence shells.
\nWhile some of Lewis’ predictions have since been proven incorrect (he suggested that electrons occupy cube-shaped orbitals), his work established the basis of what is known today about chemical bonding. <\/p>\n

\u00a0
\n\u00a0
\n\u00a0
\n\u00a0MAIN TYPES OF CHEMICAL COMBINATION
\n<\/strong><\/p>\n

    \n
  1. Electrovalent\/ionic combination\n<\/li>\n
  2. \n
    Covalent classified into:\n<\/div>\n
      \n
    1. Ordinary covalent\n<\/li>\n
    2. \n
      Coordinate or dative covalent.\n<\/div>\n

      \u00a0<\/li>\n<\/ol>\n<\/li>\n<\/ol>\n

      EXAMPLES OF ELECTROVALENT
      \n<\/strong><\/p>\n

        \n
      1. \n
        Formation of NaCl<\/p>\n
        \n\n\n\n\n\n
        \u00a0<\/td>\nNa atom<\/td>\nCl atom<\/td>\n<\/tr>\n
        Before <\/td>\n11
        \n2,8,1<\/td>\n        		17
        \n2, 8, 7<\/td>\n<\/tr>\n
        After <\/td>\n11
        \n2,8<\/td>\n        		18
        \n2,8,8<\/td>\n<\/tr>\n<\/tbody>\n<\/table>\n<\/div>\n

        Equation: Na \u2013 e \u2192 Na+<\/sup>Cl + e \u2192Cl\n<\/div>\n

        Na + Cl Na+ <\/sup>+ Cl–
        \n<\/sup>
        \n\u00a0Diagram
        \n\"\"\/<\/p>\n

        \u00a0Na = 2, 8, 1 Cl = 2,8,7 Na+ <\/sup>= 2,8 Cl = 2,8,8\n<\/li>\n

      2. \n
        FORMATION OF MgO<\/strong><\/p>\n
        \n\n\n\n\n\n
        \u00a0<\/td>\nMg<\/td>\nO<\/td>\n<\/tr>\n
        Before<\/td>\n12 = 2,8,2<\/td>\n8= 2,6<\/td>\n<\/tr>\n
        After<\/td>\n12 = 2,8<\/td>\n8 = 2,8<\/td>\n<\/tr>\n<\/tbody>\n<\/table>\n<\/div>\n

        \u00a0<\/p><\/div>\n

        Equation: Mg \u2013 2e \u2192 Mg 2+<\/sup> O + 2e– <\/sup>\u2192 O2-
        \n<\/sup> Mg + O \u2192 Mg2+<\/sup>O2-
        \n<\/sup>
        \n\u00a0<\/li>\n<\/ol>\n

        FORMATION OF MgCl2<\/sub><\/strong><\/p>\n

        \n\n\n\n\n\n
        \u00a0<\/td>\nMg<\/td>\nCl<\/td>\n<\/tr>\n
        Before <\/td>\n12 = 2,8,2<\/td>\n17 = 2,8,7<\/td>\n<\/tr>\n
        After <\/td>\n12 = 2,8<\/td>\n17 = 2,8,8<\/td>\n<\/tr>\n<\/tbody>\n<\/table>\n<\/div>\n

        \u00a0Equation <\/strong>
        \n\t\t\tMg \u2013 2e \u2192 Mg2+<\/sup>Cl + 2e \u2192 2Cl<\/p>\n

        \u00a0
        \n\u00a0
        \n\u00a0
        \n\u00a0
        \n\u00a0
        \n\u00a0
        \n\u00a0
        \n\u00a0
        \n\u00a0
        \n\u00a0Diagram
        \n\"\"\/
        \n\t\t\tPROPERTIES
        \n<\/strong><\/p>\n

          \n
        1. Electrostatic forces of attraction are strong\n<\/li>\n<\/ol>\n
            \n
          1. Consist of ions\n<\/li>\n
          2. In terms of structure, they exist as solids at room temperature, arranged in orderly manner to form crystals ie. =ve ion\n<\/li>\n
          3. \n
            surrounded be \u2013ve ion and vice verse e.g.\n<\/div>\n

            NaCl crystals
            \n=Cl– <\/sup>ion
            \n=Na+ <\/sup>ion
            \nThey are hard. If force is applied, it gets scattered but does not change shape\n<\/li>\n

          4. High mpt and bpt because of strong bond between ions. E.g. NaCl melts at 8010 <\/sup>C and its bpt 14670<\/sup>C.\n<\/li>\n
          5. Soluble in polar solvents e.g. water, ethanol but insoluble in non-polar solvent e.g. benzene, CCl4<\/sub>\n\t\t\t<\/li>\n<\/ol>\n

            \u00a0
            \n\u00a0HOW?
            \n<\/strong>When NaCl for example is placed in water, the <\/p>\n

            \u00a0
            \n\u00a0water surrounds individual ions (Na+<\/sup>Cl–<\/sup>) in the surface and exposes the inner layers of NaCl ions.<\/p>\n

              \n
            1. \n
              Good conductors or electrolytes of electricity.\n<\/div>\n

              Reason:<\/strong> the ions are free to move about when in a liquid state or in solution.<\/p>\n

              \u00a0<\/li>\n<\/ol>\n

              \n\n\n\n
              \n

              Covalent Bonds
              \n<\/h2>\n

              Covalent chemical bonds involve the sharing of a pair of valence electrons by two atoms, in contrast to the transfer of electrons in ionic bonds. Such bonds lead to stable molecules if they share electrons in such a way as to create a noble gas configuration for each atom.
              \nHydrogen gas forms the simplest covalent bond in the diatomichydrogen molecule. The halogens such as chlorine also exist as diatomic gases by forming covalent bonds. The nitrogen and oxygen which makes up the bulk of the atmosphere also exhibits covalent bonding in forming diatomic molecules. <\/p>\n

              \n\n\n\n
              \"\"\/<\/td>\nCovalent bonding can be visualized with the aid of Lewis diagrams. <\/td>\n<\/tr>\n<\/tbody>\n<\/table>\n<\/div>\n\n\n\n\n
              \u00a0<\/td>\n<\/tr>\n<\/tbody>\n<\/table>\n<\/td>\n<\/tr>\n
              \u00a0 <\/p>\n\n\n\n
              \u00a0<\/td>\n<\/tr>\n<\/tbody>\n<\/table>\n<\/td>\n<\/tr>\n<\/tbody>\n<\/table>\n<\/div>\n

              \u00a0
              \n\u00a0
              \n\u00a0
              \n\u00a0
              \n\u00a0
              \n\u00a0
              \n\u00a0
              \n\u00a0
              \n\u00a0
              \n\u00a0EXAMPLES OF COVALENT COMPOUNDS
              \n<\/strong><\/p>\n

                \n
              1. Formation of a hydrogen molecule H2<\/sub>\n\t\t\t<\/li>\n
              2. Formation of HCl\n<\/li>\n
              3. \n
                Formation of H2<\/sub>O\n<\/div>\n

                2H2<\/sub>O + O2 <\/sub>\u2192 2H2<\/sub>O\n<\/li>\n

              4. \n
                C + 2H2<\/sub> \u2192 CH4<\/sub>\n\t\t\t\t<\/div>\n

                \"\"\/\n\t\t\t\t<\/li>\n

              5. \n
                Formation with double bond \u2013 CO2 <\/sub>\n\t\t\t\t<\/div>\n

                C + O2<\/sub>\u2192 CO2 <\/sub>O = C = O
                \nCovalent bonds i.e. where each atom contributes a pair of electrons
                \nPROPERTIES<\/strong>\n\t\t\t\t<\/li>\n<\/ol>\n

                  \n
                1. Consists of molecules\n<\/li>\n
                2. Electrostatic forces are not strong\n<\/li>\n
                3. Low m.pt and b.pt\n<\/li>\n
                4. Usually dissolve in non-polar solvent and insoluble in polar solvent\n<\/li>\n
                5. \n
                  Poor conductor of electricity and heat\n<\/div>\n

                  Reason:<\/strong> Molecules do not contain charged ions\n<\/li>\n

                6. They are often gases or volatile liquids e.g. iodine molecule.\n<\/li>\n
                7. \n
                  \"\"\n\t\t\t\t<\/div>\n

                  \u00a0
                  \n\u00a0<\/li>\n<\/ol>\n

                  \u00a0
                  \n\u00a0HYDROGEN BOND<\/strong>
                  \n\t\tBond between hydrogen and any strongly electro negative elements e.g. N, O, and F. They are covalently bonded.
                  \nThey form dipole i.e. coming together of a positive pole (H) and negative pole (N) or (F). The bond is weak.
                  \n\"\"\/
                  \n\t\tExercise: <\/strong>HF, H2<\/sub>O, NH3
                  \n<\/sub>The hydrogen bond is responsible for:
                  \n\t\t\t<\/sub><\/p>\n

                    \n
                  1. B.pt of water (1000<\/sup>C) is high than that of H2S (-650<\/sup>C.). The oxygen in water has more affinity for e–<\/sup> than sulphur hence H2<\/sub>O has a stronger bond.\n<\/li>\n
                  2. \n
                    Easy logue faction e.g NH3<\/sub>\n\t\t\t\t<\/div>\n<\/li>\n
                  3. \n
                    Solubility of some organic compounds in water e.g. ethanol, sugar.\n<\/div>\n<\/li>\n<\/ol>\n

                    \u00a0
                    \n\u00a0OTHER BINDING FORCES
                    \n<\/strong><\/p>\n

                      \n
                    1. \n
                      Metallic Bond: <\/strong> holds metal atoms together in crystal lattices.\n<\/div>\n

                      \"\"\/
                      \n\t\t\t\tNote:<\/strong> the outer most e–<\/sup> in the metal form the electron cloud which determines how strong the metallic bond is going to be i.e. the larger the number of e–<\/sup> in the electron cloud the stronger the metallic bonding.
                      \nMetallic bonding is responsible for metals:<\/strong><\/p>\n

                        \n
                      1. High m.pt e.g. iron is 15350<\/sup>C, Na = 980<\/sup>C,\n<\/li>\n
                      2. because of few e–<\/sup>s in the outer most shell.\n<\/li>\n
                      3. Malleability and ductility. Because layers of metallic atoms can slide over each other\n<\/li>\n
                      4. High electrical and heat conductivity because close packing of particles\n<\/li>\n
                      5. Light densities e.g. Iron = 7.8gcm-3<\/sup>\n\t\t\t\t\t<\/li>\n<\/ol>\n<\/li>\n<\/ol>\n

                        v.High boiling point<\/p>\n

                        Metallic Bond
                        \n<\/h2>\n

                        Home \u2192 Metallic Bond
                        \nMetals constitute about three-fourth of all the known elements. They have characteristic properties such as bright lustre, high electrical and thermal conductivity, malleability and ductility and high tensile strength. The attractive force which binds various metal atoms together is called metallic bond. The metallic bond is neither a covalent bond nor an ionic bond because neither of these bonds are able to explain the known properties of metals. For example, neither ionic nor covalent compounds conduct electricity in the solid stale but metals are very good conductors of electricity. In order to explain bonding in metals different theories have been put forward. We shall be studying here electron gas model or electron sea model for metallic bonding.
                        \n\u00a0
                        \n\u00a0ELECTRON GAS MODEL OR ELECTRON SEA MODEL<\/strong>
                        \n\tThis is the simplest model that explains the properties of metals. This model was proposed by Lorentz. The main features of this modal are:
                        \n1. A metal atom is supposed to consist of two parts, valence electrons and the remaining part (the nucleus and inner shells) which is called kernel.
                        \n2. The metallic crystal consists of crystal packed metal atoms in three dimensions. The kernels of metal atoms occupy fixed positions called Lattice sites while space between the kernels is occupied by valence electrons. The arrangement of kernels and valence electrons is shown in Fig. 7.16.
                        \n\"\"
                        \n\tFig. 7.16. Arrangement of metallic kernels.<\/strong>
                        \n\t\u00a0<\/strong>3. Due to smaller ionisation energy, the valence electrons of metal atoms are not held by the nucleus very firmly. Therefore, they can leave the field of influence of one kernel and enter the field of influence of the other. This movement can take place through the vacant valence orbitals. Thus, the valence electrons are not localised but are mobile or delocalised. As the movement of electrons in metallic crystal is just like gas molecules, hence, the model is called electron gas model.
                        \n4. The simultaneous force of attraction between the mobile electrons and the positive kernels is responsible for holding the metal atoms together and is known as metallic bond.
                        \n\u00a0
                        \n\u00a0The metallic bond is non-directional and is weaker than the covalent bond.
                        \n\u00a0
                        \n\u00a0STRENGHT OF METALLI C BONO<\/strong>
                        \n\tStrength of metallic bond depends on the magnitude of attractive force between positive kernels and mobile valence electrons.
                        \nThe average attractive force and metal bond strength increases with the decrease in atomic radius and increase in number of valence electrons. It must be noted carefully that both these factors at the same time decrease the metal character because of the tendency to form metallic crystal decreases. For example when we move along the period from left to right metallic character decreases. Among the elements of 3rd period metal character decreases from left to right. The metallic elements are only Na, Mg, AI, but strength of metallic bond increases from Na -7 AI. It is reflected from their melting points.
                        \n\"\"
                        \n\tSimilarly, as we move down the group among alkali metals, the atomic radius increases. Consequently metal bond strength decreases and this causes decrease in the melting points among alkali metals from top to bottom, i.e., from Li\u00e0Cs
                        \nLi\u00a0\u00a0 \u00a0\u00a0\u00a0\u00a0 Na \u00a0\u00a0\u00a0\u00a0\u00a0 \u00a0K \u00a0\u00a0\u00a0\u00a0\u00a0\u00a0 Rb\u00a0 \u00a0\u00a0\u00a0\u00a0 Cs
                        \nMelting points (K)\u00a0\u00a0\u00a0\u00a0 454 \u00a0\u00a0\u00a0 371 \u00a0\u00a0\u00a0\u00a0 336\u00a0\u00a0\u00a0\u00a0\u00a0 312 \u00a0\u00a0\u00a0 302
                        \n\u00a0
                        \n\u00a0FACTORS THAT FAVOUR THE METALLIC BONDING<\/strong>
                        \n\tMetallic bonding is generally favoured by the following factors:
                        \n1. The atomic size of the element should be large.
                        \n2. Ionisation energy of the element should be low
                        \n3. Electron affinity of the element should be low
                        \n4. The Umber of valence electrons should be small (usually 1-2)
                        \n5. The number of vacant orbitals in the valence shell should be large.
                        \n\u00a0
                        \n\u00a0EXPLANATION ON OF PHYSICAL PROPERTIES OF METALS<\/strong>
                        \n\t1. Metallic Lustre. <\/strong>When light falls on the surface of the metal, the free electrons absorb the photons of light and are set into vibrations. These vibrating electrons immediately emit energy and become a source of light. Thus, incident light appears to be reflected from the surface of the metal. Consequently the metallic surface acquires a shining appearance which is referred to as metallic lustre.
                        \n\u00a02. Electrical Conductivity.<\/strong> Whenever a difference is applied across the metallic strip, the free mobile electrons in the metal start moving towm-ds positive terminal At the same time the electrons from the negative terminal enter into the metallic crystal. Thus, metallic crystal maintains flow of electron from negative to positive terminal.
                        \nAt high temperature, the metallic kernels start due to increase of the kinetic energy. This restricts the free movement of the electrons. Consequently, the resistance of\u00a0 metals increases with the increase in the temperature.
                        \n\"\"
                        \n\t\u00a0\u00a0\u00a0 Fig. 7 .17. Electrical conductivity of metals.<\/strong>
                        \n\t\u00a0
                        \n\u00a03. Thermal Conductivity.<\/strong> The conduction of through the metals can also be explained on the basic electron gas mode\\. On heating a part of the metal, the kinetic energy of the electrons in that region increases. These energetic electrons move rapidly to the cooler parts and transfer their kinetic energy by means of collisions with other electrons . In this way, the heat travels from hotter to cooler parts of the metals.
                        \n\u00a0
                        \n\u00a04. Malleability and Ductility. <\/strong>Malleability is the property of metals by virtue of which they can be beaten into sheets whereas ductility is the property by virtue of which they\u00a0 \u00a0can. be drawn into wires. These properties are exhibited by metals on account of the \u00a0of non-directional nature of metallic bond. Whenever any stress is applied on metal, the position of metallic kernels is altered without destroying the crystal The crystal lattice gets deformed by slippage of the layers of kernels moving past to another as shown in fig 7.18. whenone layer of kernels moves past another, the positive on metal ions are shielded from each other by the electrons.
                        \n\u00a0
                        \n\u00a0The electron sea model could explain the properties metals qualitatively. However, the properties of metals be explained more quantitatively by molecular orbital which is beyond the scope of this book.
                        \n\"\"
                        \n\tFig. 7.18. Malleability and ductility of metals.<\/strong>
                        \n\t\u00a0
                        \n\u00a0The general properties associated with three primary interatomic bonds are being summarized in tabular form as follows.
                        \n\u00a0
                        \n\u00a0\u00a0General Characteristics of Substances with different Interatomic Bonding<\/strong>
                        \n\t\"\"<\/p>\n

                        \u00a0
                        \n\u00a0
                        \n\u00a0\u00a0 Van der waal forces<\/strong><\/p>\n

                          \n
                        • \n
                          Very weak\n<\/div>\n<\/li>\n
                        • \n
                          Very important in the liquifaction of gases and in the formation of molecular lattices e.g. in iodine and Naphthalene crystals.\n<\/div>\n<\/li>\n<\/ul>\n

                          \u00a0
                          \n\u00a0
                          \n\u00a0COORDINATE COVALENT COMBINATION
                          \n<\/strong>Electrons to be shared are donated by only one of the participating atoms. Such pair of e–<\/sup> s are called Ione pair. This combination always leads to the formation of complex ions.<\/p>\n

                          Co-ordinate-Covalent Bond or Dative Bond
                          \n<\/h2>\n

                          Home \u2192 Co-ordinate-Covalent Bond or Dative Bond
                          \nIt is a special case of covalent bond the formation of which was postulated by Perkins (1921). It is formed by mutual sharing of electrons between the two atoms but the shared pair of electrons is contributed only by one of the two atoms, the other atom simply participates in sharing. The atom which donates an electron pair for sharing is called donor and it must have already completed its octet. On the other hand, the atom which accepts the electron pair in order to complete its octet is called acceptor. The bond is represented by an arrow\u00a0\u00a0\u00a0\u00a0\u00a0 pointing from the donor towards the acceptor. Let us consider the formation of ozone molecule. A molecule of, -oxygen contains two oxygen atoms which share four electrons and complete their octets. Now, if an atom of oxygen having six valence electrons comes close to oxygen molecule, it\u00a0 shares a lone pair of electrons with one of the oxygen of the molecule. It can be represented as follows:
                          \n\"\"
                          \n\tIt is important to note that co-ordinate bond one formed, cannot be distinguished from covalent bond.
                          \nSome more examples of molecules\/polyatomic ions having co-ordinate bond are as follows:
                          \n(i) SO2<\/sub> molecule
                          \n\"\"
                          \n\t(ii) Ammonium ion
                          \n\"\"
                          \n\t(iii) Hydronium ion
                          \n\"\"
                          \n\t(iv) Carbon (II) oxide. In this molecule carbon and oxygen atoms contribute two electrons forming pure covalent bonds. At the same time oxygen also act as donor atom to form co-ordinate covalent bond.
                          \n\"\"<\/p>\n

                          \u00a0
                          \n\u00a0
                          \n\u00a0
                          \n\u00a0Exercise
                          \n<\/strong>1. Formation of ammonium ion
                          \nReaction between NH3<\/sub> and HCl acid. The H+<\/sup> of
                          \nthe acid reacts with NH3<\/sub>
                          \n\t\t\t<\/sup>and NH4<\/sub>+<\/sup> is formed.
                          \n2 .Formation of oxonium ion or Hydroxonium ion
                          \nQUESTIONS
                          \n<\/strong>What are the types of bonding that exist in the following compounds.
                          \n1.HCl
                          \n2. NH4<\/sub>Cl
                          \nWith diagram and equations only , illustrate the
                          \nformation of:
                          \n\t\t\t<\/em><\/p>\n

                            \n
                          1. Oxygen molecule
                            \n\t\t\t\t<\/em><\/li>\n
                          2. Ethane molecule
                            \n\t\t\t\t<\/em><\/li>\n
                          3. Ammonia molecule
                            \n\t\t\t\t<\/em><\/li>\n
                          4. Nitrogen molecule
                            \n\t\t\t\t<\/em><\/li>\n
                          5. Al2<\/sub>O3<\/sub>
                            \n\t\t\t\t<\/em><\/li>\n<\/ol>\n

                            Which of the compounds in 2 above is\/are Triple covalently bonded?
                            \n1. Which of the following species the constituent atoms are held by non-directional bonds?
                            \n(a) NH3\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0 (b) CsCl
                            \n(c) NF3\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0 (d) BeF2.
                            \n\u00a0
                            \n\u00a02. X and Y atoms have 2 and 6 valance electrons in their outermost shells respectively, the compound which X and Y are likely to form is:
                            \n(a) XY 2\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0 (b) XY
                            \n(c) YX2 \u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0 (d) YX3.
                            \n\u00a0
                            \n\u00a03. Which of the following substance is a polar covalent molecule?
                            \n(a) Hydrogen sulphide \u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0 (b) Nitrogen
                            \n(c) Potassium chloride\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0 (d) Oxygen.
                            \n\u00a0
                            \n\u00a04. Which of the following statement is\/are correct?
                            \nI. Energy is absorbed when a chemical bond is formed
                            \nII. Energy is released when a chemical bond is formed
                            \nIII. SF 6 is a super octet molecule
                            \n(a) I and III\u00a0\u00a0\u00a0\u00a0\u00a0 (b) III and II
                            \n(c) II only \u00a0\u00a0\u00a0\u00a0\u00a0\u00a0 \u00a0\u00a0(d) III only
                            \n5.In electrovalency,valence electrons are transferred and the atomic number is
                            \n(A) .also reduced (B).stabilized (C) .unaffected (D).destabilized
                            \n6.Arrangement of ions in a regular pattern in a solid crystal is called
                            \n(A).configuration(B).atomic structure(C).lattice(D).Buffer
                            \n7.The bond type in a diatomic nitrogen gas is
                            \n(A).double covalent bond (B).triple covalent bond
                            \n(C).single covalent bond (D).double covalent bond
                            \n8.The bond type between copper( ii )ion and water molecules is
                            \n(A).electrovalent bond (B).covalent bond (C).Dative covalent bond
                            \n(D).Hydrogen bond
                            \n9.The bond between two iodine molecules is
                            \n(A).co-ordinate bond (B) electrovalent bond (C).ionic bond
                            \n(D)Van der waal’s forces
                            \n10.Bonds between a highly electronegative atom and a hydrogen from another molecule is called
                            \n(A).hydrogen bond (B).covalent bond (C).intermolecular forces (D).Ligand<\/p>\n

                            \u00a0
                            \n\u00a0
                            \n\u00a0
                            \n\u00a0
                            \n\u00a0
                            \n\u00a0
                            \n\u00a0
                            \n\u00a0
                            \n\u00a0
                            \n\t\t\t<\/strong>\u00a0<\/p>\n","protected":false},"excerpt":{"rendered":"

                            \u00a0WEEK 6 CHEMICAL BONDING INTRODUCTION Though the periodic table has only 118 or so elements,…<\/p>\n","protected":false},"author":1,"featured_media":0,"comment_status":"open","ping_status":"open","sticky":false,"template":"","format":"standard","meta":{"footnotes":""},"categories":[1,196],"tags":[],"class_list":["post-2302","post","type-post","status-publish","format-standard","hentry","category-posts","category-second-term-ss1-chemistry"],"aioseo_notices":[],"_links":{"self":[{"href":"https:\/\/ecolebooks.com\/nigeria\/wp-json\/wp\/v2\/posts\/2302","targetHints":{"allow":["GET"]}}],"collection":[{"href":"https:\/\/ecolebooks.com\/nigeria\/wp-json\/wp\/v2\/posts"}],"about":[{"href":"https:\/\/ecolebooks.com\/nigeria\/wp-json\/wp\/v2\/types\/post"}],"author":[{"embeddable":true,"href":"https:\/\/ecolebooks.com\/nigeria\/wp-json\/wp\/v2\/users\/1"}],"replies":[{"embeddable":true,"href":"https:\/\/ecolebooks.com\/nigeria\/wp-json\/wp\/v2\/comments?post=2302"}],"version-history":[{"count":1,"href":"https:\/\/ecolebooks.com\/nigeria\/wp-json\/wp\/v2\/posts\/2302\/revisions"}],"predecessor-version":[{"id":2303,"href":"https:\/\/ecolebooks.com\/nigeria\/wp-json\/wp\/v2\/posts\/2302\/revisions\/2303"}],"wp:attachment":[{"href":"https:\/\/ecolebooks.com\/nigeria\/wp-json\/wp\/v2\/media?parent=2302"}],"wp:term":[{"taxonomy":"category","embeddable":true,"href":"https:\/\/ecolebooks.com\/nigeria\/wp-json\/wp\/v2\/categories?post=2302"},{"taxonomy":"post_tag","embeddable":true,"href":"https:\/\/ecolebooks.com\/nigeria\/wp-json\/wp\/v2\/tags?post=2302"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}