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CHEMICAL BONDS
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CHEMICAL BONDS

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  1. CHEMICAL BONDS • Atoms or ions are held together in molecules or compounds by chemical bonds. • The type and number of electrons in the outer electronic shells of atoms or ions are instrumental in how atoms react with each other to form stable chemical bonds. • Over the last 150 years scientists developed several theories to explain why and how elements combine with each other.

  2. Bonding in Chemistry Central theme in chemistry: Why and How atoms attach together This will help us understand how to: Predict the shapes of molecules. Predict properties of substances. Design and build molecules with particular sets of chemical and physical properties.

  3. CHEMICAL BONDS Two of the most common substance on our dining table are salt and granulated sugar C12H22O11 NaCl The properties of substances are determined in large part by the chemical bonds that hold their atoms together

  4. Chemical Bonds All chemical reactions involve breaking of some bonds and formation of new ones which yield new products with different properties.

  5. Bonding Theories • Lewis bond Theory • Valence Bond Theory • Molecular Orbital Theory Gilbert Newton Lewis

  6. Lewis Bonding Theory Atoms ONLY come together to produce a more stable electron configuration. Atoms bond together by either transferring or sharing electrons. Many of atoms like to have 8 electrons in their outer shell. Octet rule. There are some exceptions to this rule—the key to remember is to try to get an electron configuration like a noble gas. Li and Be try to achieve the He electron arrangement.

  7. Lewis Symbols of Atoms Uses symbol of element to represent nucleus and inner electrons. Uses dots around the symbol to represent valence electrons. Puts one electron on each side first, then pair. Remember that elements in the same group have the same number of valence electrons; therefore, their Lewis dot symbols will look alike. • •• •• •• •• Li• Be• •B• •C• •N• •O::F: :Ne: • • • • • • ••

  8. Valence electrons

  9. Practice to write the Lewis symbol for Arsenic As is in group 15 (5), therefore it has 5 valence electrons.

  10. Using Lewis Theory to Predict Chemical Formulas Compounds ∙ ∙ ∙ ∙ ∙ Ca Ca Cl Cl Cl ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ Predict the formula of the compound that forms between calcium and chlorine. Draw the Lewis dot symbols of the elements. Transfer all the valance electrons from the metal to the nonmetal, adding more of each atom as you go, until all electrons are lost from the metal atoms and all nonmetal atoms have 8 electrons. Ca2+ CaCl2

  11. Examples for Lewis representation of some chemical bonds •• •• •• •• •• • • •• F F • • • • O O •• •• •• •• •• •• •• •• O •• •• •• •• O •• •• •• F F O O F F •• •• •• • • • • H H O •• •• •• H O •• H ••

  12. Total number of valence electrons = 6 + 4 + 6 = 16 Actually 24 electrons needed for completing the octet of each atom Thus 24 - 16 = 8 electrons are shared. Since two electrons make a bond, the molecule should have 4 bonds. The remaining 8 electrons are lone pair electrons. Information: Given: CO2 Find: Lewis structure Solution Map: formula → skeletal → electron distribution → Lewis Example:Write the Lewis structure of CO2. .. .. O C O .. ..

  13. Practice—Draw Lewis Resonance Structures for CNO−(C Is Central with N and O Attached) C = 4 N = 5 O = 6 (-) = 1 Total = 16 e-

  14. Example NO3─ 1. Write skeletal structure. N is central because it is the most metallic. - 2. Count valence electrons. N = 5 O3 = 3 x 6 = 18 (-) = 1 Total = 24 e-

  15. TYPES OF CHEMICAL BONDS • Ionic bonds • Covalent bonds • Metallic bonds

  16. The three possible types of bonds.

  17. Ionic compounds consist of a cation and an anion • the formula is always the same as the empirical formula • the sum of the charges on the cation and anion in each formula unit must equal zero. Lewis bonding theory is able to explain ionic bonds very well. The ionic compound NaCl

  18. Ionic bonding • Ionic substances are formed when an atom that loses electrons relatively easily react with an atom that has a high affinity for electrons. ex. metal-nonmetal compound

  19. Chemical Bonds Ionic bonds are formed by the attraction of oppositely charged ions.

  20. Ionic Bonds Metal to nonmetal. Metal loses electrons to form cation. Nonmetal gains electrons to form anion. The electronegativity between the metal and the nonmetal must be > than 2. Ionic bond results from + to − attraction. Larger charge = stronger attraction. Smaller ion = stronger attraction. Lewis theory allows us to predict the correct formulas of ionic compounds.

  21. Ions that pack as spheres in a very regularpattern form crystalline substances .

  22. Formation of an Ionic Solid • 1. Sublimation of the solid metal M(s) → M(g) [endothermic] • 2. Ionization of the metal atoms M(g) →M+(g) + e- [endothermic] • 3. Dissociation of the nonmetal 1/2X2(g) → X(g) [endothermic]

  23. Electron affinity of F Dissociation of F2 Ionization of Li Formation of solid Sublimation of Li

  24. Lattice Energy Calculations k: a proportionality constant that depends on the structure of the solid and the electron configuration of the ions Q1 and Q2: charges on the ions r: the shortest distance between the centers of cations and anions

  25. More Gains and Losses • Can elements lose or gain more than one electron? • The element magnesium, Mg, in Group 2 can lose two electron and element oxygen in Group 6 can gain two electrons to form stable Nobel gas configurations. The ions can come together to form a crystal structure.

  26. Relative sizes of some ions and their parent atoms.

  27. Structure of ionic crystals Different types of crystals are formed depending on the ionic radii and the charge of the ions involved.

  28. How about the bonds between atoms that have the same electronegativity (as in H-H molecule) or when the electonegativuty difference is < 1.0 (as in C-H)?

  29. Convalent Bonds—Sharing • Some atoms are unlikely to lose or gain electrons because the number of electrons in their outer levels makes this difficult. • Consider the Lewis dot structure of carbon . C C+4 + 4e- . . . • The alternative is sharing electrons.

  30. Covalent Bonds Often found between two nonmetals. Typical of molecular species. Atoms bonded together to form molecules. Strong attraction. Atoms share pairs of electronsto attain octets. Molecules generally weakly attracted to each other. Observed physical properties of molecular substance due to these attractions.

  31. Covalent Bonding • Electron are shared by nuclei

  32. The Convalent Bond • Shared electrons are attracted to the nuclei of both atoms. • They move back and forth between the outer energy levels of each atom in the covalent bond. • So, each atom has a stable outer energy level some of the time.

  33. The formation of a bond between two atoms.

  34. Electron Density for the H2 molecule An electron density plot for the H2 molecule shows that the shared electrons occupy a volume equally distributed over BOTH Hydrogen atoms.

  35. Chemical Bonds Covalent bonds form when atoms share 2 or more valence electrons. Covalent bond strength depends on the number of electron pairs shared by the atoms. single bond double bond triple bond < <

  36. Examples of Convalent Bond • The neutral particle is formed when atoms share electrons is called a molecule

  37. Single Covalent Bonds Two atoms share one pair of electrons. 2 electrons. One atom may have more than one single bond. •• •• • • •• F F •• •• •• •• •• •• •• •• F F F F •• •• •• • • • • H H O •• •• •• H H •• O ••

  38. Double Covalent Bond Two atoms sharing two pairs of electrons. 4 electrons. Shorter and stronger than single bond. •• •• • • • • O O •• •• •• •• O •• •• •• •• O O O

  39. Chemical Bonds

  40. Bond Polarity Bonding between unlike atoms results in unequal sharing of the electrons. One atom pulls the electrons in the bond closer to its side. One end of the bond has larger electron density than the other. The result is bond polarity. The end with the larger electron density gets a partial negative charge and the end that is electron deficient gets a partial positive charge. • • d+ H Cl d-

  41. Nonpolar and polar covalent bonds

  42. Probability representations of the electron sharing in HF.

  43. Trends in electronegativity across a period and down a group

  44. Nature of bonds and electronegativity Electronegativity Bond difference (∆) ∆ > 2 Ionic 0.4 < ∆ < 2 Polar covalent ∆ < 0.4 Covalent