Chemical Bonding

1. Chemical Bonding Atoms interact with other atoms to form molecules, this is chemical bonding Bonding theories – are models that predict how atoms bond together to form molecules Bonding Theories are applied to design molecules that will interfere with the active site of HIV-protease. This delays or inhibits the onset of AIDS.

2. OUTLINE

3. CHEMICAL BOND • The nature and type of the chemical bond is directly responsible for many physical and chemical properties of a substance: (e.g. melting point, conductivity). • Most matter in nature is found in form of compounds: 2 or more elements held together through a chemical bond. • Elements combine together (bond) to fill their outer energy levels and achieve a stable structure (low energy). • Noble gases are un-reactive since their energy levels are complete.

4. CHEMICAL BOND • When the conductivity apparatus is placed in salt solution, the bulb will light. • But when it is placed in sugar solution, the bulb does not light. • This difference in conductivity between salt and sugar is due to the different types of bonds between their atoms. • Two common types of bonding are present: ionic & covalent.

5. Gilbert Newton Lewis (1875 - 1946) was a famous American physical chemist known for the discovery of the covalent bond (see his Lewis dot structures and his 1916 paper "The Atom and the Molecule") Other major contributions were his theory of Lewis acids and bases and Lewis coined the term "photon" for the smallest unit of radiant energy.

6. The Origin of Lewis Symbols of Atoms Drawings of cubical atoms, the corners of the cube represented possible electron positions Lewis later cited these notes in his classic 1916 paper on chemical bonding, as being the first expression of his ideas.

7. LEWIS SYMBOLS OF ATOMS • Lewis structures use Lewis symbols to show valence electrons in molecules and ions of compounds. • In Lewis symbols, valence electrons for each element are shown as a dot. • Lewis symbols for the first 3 periods of representative elements are shown below:

8. Lewis Bonding Theory • atoms bond because it results in a more stable electron configuration • atoms bond together by either transferring or sharing electrons so that all atoms obtain an outer shell with 8 electrons • 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

9. •• •• Li• Li+1:F: [:F:]-1 • •• Lewis Symbols of Ions • Cations have Lewis symbols without valence electrons • Lost in the cation formation • Anions have Lewis symbols with 8 valence electrons • Electrons gained in the formation of the anion

10. Ionic Bonds • metal to nonmetal • metal loses electrons to form cation • nonmetal gains electrons to form anion • ionic bond results from + to - attraction • larger charge = stronger attraction • smaller ion = stronger attraction • Lewis Theory allow us to predict the correct formulas of ionic compounds

11. Metal Nonmetal IONIC BOND • Ionic bonds occur between metals and non-metals. • Ionic bonds occur when electrons are transferred between two atoms. • After bonding, each atom achieves a complete shell (noble gas configuration).

12. IONIC BOND • The smallest particles of ionic compounds are ions (not atoms). • Atoms that lose electrons (metals) form positive ions (cations). • Atoms that gain electrons (non-metals) form negative ions (anions). Anion Cation

13. ∙ ∙ ∙ ∙ ∙ Ca Ca Cl Cl Cl ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ Using Lewis Theory to Predict Chemical Formulas of Ionic Compounds 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

14. Ionic Bonding • The transfer of electrons in an ionic bond is from an atom of low ionization energy (usually a metal) to an atom of high electron affinity (usually a nonmetal). • The electrostatic attraction between two oppositely charge ions constitutes an ionic bond.

15. Ionic Bonding & Lattice Structures • A lattice is a stable, ordered, solid three dimensional array of ions associated with ionic compounds. • Lattice energy, ∆Hlattice is the energy required to completely separate a mole of solid ionic compound into its gaseous ions.

16. Lattice Energy Example • For example, the lattice energy of potassium fluoride is 808 kJ/mol. KF(s) → K+(g)+ F- (g) ∆Hlattice = +808 kJ/mol • Lattice energies can be found in a table as known values. • Lattice energy increases with smaller, more highly charged ions

17. Potential Energy in Bonds • The potential energy of two interacting charged particles is given by the equation E = KQ1Q2/d Where Q1 and Q2 are the charges on the ions and d is the distance between them.

18. Characteristics of Ionic Compounds • Ionic Compounds tend to have high lattice energies • High lattice energies make them hard and brittle • High lattice energies make them have relatively high melting points • Ionization energies increase rapidly for each successive electron removed

19. Transition Metals • Lose their valence level “s” electrons first and then one or two “d” orbital elctrons • Having two outermost “s” electons is the major reason why transition metals commonly form 2+ ions • For example: Fe: (Ar) 3d64s2 Fe2+: (Ar) 3d64s0 or (Ar) 3d6 Fe3+: (Ar) 3d54s0 or (Ar) 3d5

20. Covalent Bonds • Formed when two atoms share one or more pairs of electrons • Sharing one: Single bond • Sharing two: Double bond • Sharing three: Triple bond • sharing pairs of electrons to attain octets • molecules generally weakly attracted to each other • observed physical properties of molecular substance due to these attractions

21. COVALENT BOND • Covalent bonds form when electrons are shared between two atoms. • Covalent bonds form between twonon-metals. Electrons shared

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

23. •• •• • • • • O O •• •• •• •• O •• •• •• •• O O O Double Covalent Bond • two atoms sharing two pairs of electrons • 4 electrons • shorter and stronger than single bond

24. •• •• • • • • N N • • N N •• •• •• •• •• N N Triple Covalent Bond • two atoms sharing 3 pairs of electrons • 6 electrons • shorter and stronger than single or double bond

25. POLAR & NON-POLARBONDS Electrons shared equally • Two types of covalent bonds exist: • Non-polar covalent bonds occur between similar atoms. Polar & Nonpolar • In these bonds the electron pair is shared equally between the two protons.

26. HF POLAR & NON-POLARBONDS • Polar covalent bonds occur between different atoms. • In these bonds the electron pair is shared unequally between the two atoms. • As a result there is a charge separation in the molecule, and partial charges on each atom. + 

27. Dipole Moments • A dipole is a material with positively and negatively charged ends • Polar bonds or molecules have one end slightly positive, d+; and the other slightly negative, d- • not “full” charges • come from nonsymmetrical electron distribution • Dipole Moment, m, is a measure of the size of the polarity • measured in Debyes, D

28. ELECTRONEGATIVITY • Linus Pauling derived a relative Electronegativity Scale based on Bond Energies. • Electronegativity (E.N.) is the ability of an atom involved in a covalent bond to attract the bonding electrons to itself. F 4.0 Cs 0.7 Most electronegative Least electronegative

29. ELECTRONEGATIVITY(THIS IS REVIEW) Electronegativity increases

30. BOND POLARITY &ELECTRONEGATIVITY Polarity is a measure of the inequality in the sharing of bonding electrons The more different the electronegativity of the elements forming the bond The larger the electronegativity difference(EN) The more polar the bond formed

31. As difference in electronegativity increases Bond polarity increases POLARITY &ELECTRONEGATIVITY Most polar Least polar

32. SERIOUSLY NEED TO MEMORIZE THIS!!! POLARITY &ELECTRONEGATIVITY EN = 0 Non-polar covalent 0 < EN <1.7 Polar covalent 1.7 < EN Ionic

33. H H Electronegativity 2.1 Electronegativity 2.1 Hydrogen Molecule POLARITY &ELECTRONEGATIVITY EXAMPLES The molecule is nonpolar covalent EN = 0

34. + - H Cl Electronegativity 2.1 Electronegativity 3.0 Hydrogen Chloride Molecule POLARITY &ELECTRONEGATIVITY EXAMPLES The molecule is polar covalent EN = 0.9

35. Na+ Cl- Electronegativity 0.9 Electronegativity 3.0 Sodium Chloride POLARITY &ELECTRONEGATIVITY EXAMPLES No molecule exists The bond is ionic EN = 2.1

36. SUMMARYOF BONDING Ionic Bond (large EN) EN > 1.7 Non-polar (similar electronegativities) EN = 0 Polar (moderate EN) Covalent Bond (small to moderate EN) 0 < EN < 1.7

37. Bonding & Lone Pair Electrons • Electrons that are shared by atoms are called bonding pairs • Electrons that are not shared by atoms but belong to a particular atom are called lone pairs • also known as nonbonding pairs

38. LEWIS STRUCTURES • In a Lewis structure, a shared electron pair is indicated by two dots between the atoms, or by a dash connecting them. • Unshared pairs of valence electrons (called lone pairs) are shown as belonging to individual atoms or ions.

39. LEWIS STRUCTURES • Covalent molecules are best represented with electron-dot or Lewis structures. • Structures must satisfy octet rule (8 electrons around each atom). • Hydrogen is one of the few exceptions and forms a doublet (2 electrons).

40. LEWISSTRUCTURES • Non-bonding electrons must be displayed as dots. • Bonding electrons can be displayed by a dashed line.

41. Polyatomic Ions • The polyatomic ions are attracted to opposite ions by ionic bonds • Form crystal lattices • Atoms in the polyatomic ion are held together by covalent bonds

42. Lewis Formulas of Molecules • shows pattern of valence electron distribution in the molecule • useful for understanding the bonding in many compounds • allows us to predict shapes of molecules • allows us to predict properties of molecules and how they will interact together

43. LEWISSTRUCTURES • More complex Lewis structures can be drawn by following a stepwise method: 1. Count the number of electrons in the structure. • Draw a skeleton structure. - most metallic element generally central - halogens and hydrogen are generally terminal - many molecules tend to be symmetrical - in oxyacids, the acid hydrogens are attached to an oxygen

44. LEWISSTRUCTURES • More complex Lewis structures can be drawn by following a stepwise method: 3. Connect atoms by bonds (dashes or dots). 4. Distribute electrons to achieve Octet rule. 5. Form multiple bonds if necessary.

45.     Example 1: Write Lewis structure for H2O Step 1: H2O = 8 electrons 2 (1) + 6 = 8 Step 2: H O H Step 3: Skeleton structure should be symmetrical Hydrogen has doublet 4 electrons used4 electrons remaining Octet rule is satisfied Step 4:

46. Example 2: Write Lewis structure for CO2 Step 1: CO2 = 16 electrons 4 + 2(6) = 16 Step 2:     O C O   Step 3:       Skeleton structure should be symmetrical Step 4: Octet rule is NOT satisfied 10 electrons used6 electrons remaining 4 electrons used12 electrons remaining Octet rule is satisfied Step 5:

47. Writing Lewis Structures forPolyatomic Ions • the procedure is the same, the only difference is in counting the valence electrons • for polyatomic cations, take away one electron from the total for each positive charge • for polyatomic anions, add one electron to the total for each negative charge

48.                 Example 3: Write Lewis structure for CO32- Step 1: CO32- = 24 electrons 4+3(6)+2 = 24 Step 2: 0 electrons remaining 18 electrons remaining 12 electrons remaining 6 electrons remaining O C O O Step 3:   Step 4: Step 5: Octet rule is NOT satisfied Octet rule is satisfied

49. Example 4: Determine if each of the following Lewis structures are correct or incorrect. If incorrect, rewrite the correct structure. Structure has 14 electrons Only 12 electrons shown 2(5) + 4(1) = 14 2 4 2 Structure is incorrect 2 2 Octets are complete

50. Exceptions to the Octet Rule • H & Li, lose one electron to form cation • Li now has electron configuration like He • H can also share or gain one electron to have configuration like He • Be shares 2 electrons to form two single bonds • B shares 3 electrons to form three single bonds • expanded octets for elements in Period 3 or below • using empty valence d orbitals • some molecules have odd numbers of electrons • NO