Chemical Bonding Chapter 11

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1 Chemical Bonding Chapter 11

2 Chemical Bond Lewis (1916): Not all the electrons in an atom participate in chemical bonding. Electrons occupy shells surrounding the nucleus: The inner shells Core electrons, not significantly involved in chemical reactions. The valence shell, outermost shell Valence electrons, significantly involved in chemical reactions and chemical bonding. 2

3 Filled Shells and Becoming Chemically Stable Filled shell has high chemical stability. Filled shells can be found whenever a noble gas element is reached in the periodic table. Stability: Atom forms an ion whose outermost shell has the same number of electrons as the outermost shell of a noble gas atom. Isoelectronic: Having equal numbers of electrons or the same electronic configuration. 3

4 Cont d Chemical Stability Hydrogen and helium have valence shells completed with two electrons. Atoms of the first few periods of the periodic table have a maximum of 8 electrons in their valence shells. 4

5 Cont d Chemical Stability Tendency of atoms to achieve valence octets describes chemical stability. Atoms of elements in Groups I and II achieve octet by losing electrons to form cations. Atoms of elements in Groups VI and VII achieve octet by gaining electrons to form anions. 5

6 Octet Rule Atoms bond in such a way that each atom acquires eight electrons in its outer shell. An atom may achieve an octet by two ways:. (a) By transfer of electrons from one atom to another. (b) By sharing one or more pairs of electrons. 6

7 Bond Types Covalent Bonds: Formed when two atoms share electrons to form molecular compounds. Ionic Bonds: Formed from the complete transfer of electrons between atoms. Attraction of positively and negatively charged ions. As a result ionic compounds forms. Electrostatic forces are 7

8 Formation of Cations Cations form when an atom loses valence electrons to become positively charged. Most main group metals achieve a noble gas electron configuration by losing their valence electrons and becoming isoelectronic with a noble gas. 8

9 Cont d Formation of Cations Example: Magnesium (Group IIA/2) loses its two valence electrons to become Mg 2+. A magnesium ion has 10 electrons (12 2 = 10 e - ) and is isoelectronic with neon. 9

10 Cont d Formation of Cations Each of the metals in Period 3 form cations by losing 1, 2, or 3 electrons, respectively. Each metal atom becomes isoelectronic with the preceding noble gas, neon. More Examples: Use electron dot formulas to check the formation of cations. 10

11 Cont d Form Cations 11

12 Formation of Anions Anions form when an atom gains electrons and becomes negatively charged. Most nonmetals achieve a noble gas electron configuration by gaining electrons to become isoelectronic with a noble gas. 12

13 Cont d Formation of Anions Example: Chlorine (Group VIIA/17) gains one valence electron and becomes Cl. A chloride ion has 18 electrons ( = 18 e - ) and becomes isoelectronic with argon. 13

14 Cont d Formation of Anions The nonmetals in Period 3 gain 1, 2, or 3 electrons, respectively to form anions. Each nonmetal ion is isoelectronic with the following noble gas, argon. More Examples: Use electron dot formulas to check the formation of anions. 14

15 Cont d Form Anions 15

16 Ionic Radii A cation radius is smaller than the radius of its starting atom. An anion radius is larger than the radius of its starting atom. 16

17 Covalent Bonds Covalent bonds form: when two nonmetal atoms share electrons and the shared electrons in the covalent bond belong to both atoms. 17

18 Covalent Bonds Single Covalent Bond the atoms share 2 electrons, (1 pair) Double Covalent Bond the atoms share 4 electrons, (2 pairs) Triple Covalent Bond the atoms share 6 electrons, (3 pairs) Bond Strength = Triple > Double > Single For bonds between same atoms, C=N > C=N > C N Though Double not 2x the strength of Single and Triple not 3x the strength of Single Bond Length = Single > Double > Triple For bonds between same atoms, C N > C=N > CN 18

19 Cont d Covalent Bonds Example: Form HCl (hydrogen chloride): Hydrogen shares its one valence electron with chlorine making the chlorine atom eight electrons in its valence shell making it isoelectronic with argon. The chlorine atom shares one of its valence electrons with the hydrogen, making it two electrons in its valence shell making it isoelectronic with helium. 19

20 Bond Length In the covalent bond, the valence shells of the two atoms overlap with each other and atoms get closer to each other. This distance, when atoms are closer to each other is the bond length. Bond length is smaller than the sum of the atomic radii. 20

21 Cont d Bond Length Example: In HCl, hydrogen 1s energy sublevel overlaps with chlorine 3p energy sublevel. 21

22 Bond Energy Energy to form a covalent bond (energy is released): H(g) + Cl(g) HCl(g) + heat Energy to break a covalent bond (energy is needed): HCl(g) + heat H(g) + Cl(g) The amount of energy required to break a covalent bond is the same as the amount of energy released when the bond is formed. 22

23 Electron Dot Formulas of Molecules The number of dots around each atom is = the number of valence electrons. Lewis structures are electron dot formulas for molecules. Drawing Lewis structure, bonds between atoms and the arrangement of the atoms in a molecule can be visualized. 23

24 Drawing Lewis Structures 1. Calculate the total number of valence electrons: Add all of the valence electrons for each atom in the molecule. 2. Find the number of electron pairs in the molecule: Divide the total valence electrons by 2. 24

25 Cont d Drawing Lewis Structures 3.Octets around the atoms should be completed: Surround the central atom with 4 electron pairs. Use remaining electron pairs to complete the octets around the other atoms. The only exception is hydrogen which only needs two electrons. 25

26 Cont d Drawing Lewis Structures 4. Bonding electrons: Electron pairs that are shared by atoms. Nonbonding electrons, or lone pairs: The other electrons that complete octets. 5. If there are not enough electron pairs to provide each atom with an octet, move a nonbonding electron pair between two atoms that already share an electron pair. 26

27 Electron Dot Formula for H 2 O 1. Count the total number of valence electrons (e - ): oxygen: 6 e - and each hydrogen: 1 e -, 6 + 2(1) = 8 e - (total of 8 e - ). The number of e - pairs is 8/2 = Place 4 pair, 8 e - around the central oxygen atom. 3. Place the two hydrogen atoms in any of the four e - pair positions. Notice there - 27

28 Cont d Electron Dot Formula for H 2 O Single bond is used to represent bonding electron pairs with a single dash line. Structural formula is the resulting O structure the of the molecule. H H 28

29 Practice: Draw the Lewis structure (electron formula) for SO 3 29

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