Bonding

1. Bonding

2. What is a bond? • A bond is an attraction between two atoms that holds them together.

3. Chemical Bonds and Energy • The forces holding atoms together involve energy.

4. Bond Formation • The process of forming chemical bonds releases energy (exothermic).

5. A Spring as an Analogy for Bond Formation • Unbonded atoms that are apart are like a stretched spring. They have potential energy. As they are attracted and come closer to each other, they lose or release potential energy until they are bonded (like a spring at rest position).

6. It takes energy to break a bond holding atoms together. Bond breaking is an endothermic process. + energy  Bond Breaking

7. The Covalent Bond • When atoms share two or more electrons between them, this is called a covalent bond.

8. The Formation of a Covalent Bond between Hydrogen Atoms • The negative electrons of a hydrogen atom attract to its own nucleus and also to the nucleus of a neighbor hydrogen atom, bringing the atoms closer.

9. Covalent Bonding between Hydrogen Atoms • As hydrogen atoms approach each other, they overlap their orbitals, forming a molecular orbital. • In the molecular orbital there are two electrons which are shared between the atoms, giving each atom a full level, n=1.

10. Attracting and Repelling Forces between Hydrogen Atoms • Attracting forces are found between electrons and their own nucleus as well as the other atom’s nucleus. • Repelling forces are found between the two nuclei as well as the electrons with other electrons.

11. Potential Energy is Lost With Bond Formation

12. At lowest potential energy, atoms have a given bond distance between them. If they are forced closer, their identical positive nuclear charges force them apart.

13. Ways of Representing a Covalent Bond • A covalent bond can be shown as a probability distribution, as a Lewis Dot Diagram or as a line drawn between two atoms (structural formula). • A single line (covalent bond) always represents two shared electrons (an electron pair) between the atoms.

14. Valence Electrons The outermost electrons in an atom are called valence electrons.

15. Lewis Dot Diagrams A Lewis Dot Diagram is an element’s symbol with a number of dots around it equal to the number of its valence electrons.

16. Noble Gas Dot Diagrams

17. The Noble Gases The chemical family called the Noble Gases (Inert Gases) have full shells or shells with 8 outer electrons. All noble gases have 8 valence electrons except helium which has 2 valence electrons.

18. What is a Covalent Bond? Covalent bonds are the bonds formed when atoms share electrons between them. By sharing each other’s electron(s) they get octets or 2 electrons like helium

19. Sharing Electrons To Get Octets or Two (Like He) Covalent bonds are the bonds formed when atoms share electrons between them. By sharing each other’s electron(s) they get octets or 2 electrons (for hydrogen) like helium.

20. Representing Covalent Bonds Covalent bondscan be represented with a Lewis Diagram as two dots between atoms, or as a single line drawn between two atoms. A single line represents an electron pair (two electrons).

21. Sharing Electrons Creates Molecules When atoms share electrons, they form a unit or group called a molecule.

22. The Kinds of Atoms Forming Covalent Bonds Covalent bonds form when a nonmetal combines with another nonmetal. Ex: H with H to form H2

23. Covalent Compounds : Nonmetal + Nonmetal

24. Lewis Dot Diagrams Show Covalent Bonds Often Lewis Dot Diagrams are used to show covalent bonds.

25. Covalent Compounds and Lewis Diagrams Carbon’s 4 valence electrons are arranged around the symbol. Each hydrogen shares one valence electron, having 2 like He and C gains 4 by sharing to get an octet.

26. Sharing More Than One Pair Of Electrons Two oxygen atoms can each get an outer octet by sharing four electrons (2 pairs). This results in a double bond.

27. Sharing More Than One Pair Of Electrons Two nitrogen atoms can each get an outer octet by sharing six electrons (3 pairs). This results in a triple bond.

28. Examples of Covalent Compounds

29. Steps in Completing Lewis Dot Diagrams 1. Decide which atoms are bonded.For example, lets draw the Lewis structure for a SO3 molecule. • Count all valence electrons.There are a total of 24 valence electrons in SO3 (6 from the sulfur, and 6 each from the 3 oxygens). • Place two electrons in each bond. O S O O

30. Steps in Completing Lewis Dot Diagrams • Complete the octets of the atoms attached to the central atom by adding electrons in pairs. • Place any remaining electrons on the central atom in pairs.(already has 24, no remaining electrons in this example) • If the central atom does not have an octet, form double bonds. If necessary form triple bonds. O S O O

31. Electronegativity : An Atom’s “Electron-Pull” Power Electronegativity is the degree to which an atom attracts electrons. In the periodic chart, atoms have increasingelectronegativity moving to the right and from bottom to top.

32. Uneven Sharing Due To Different Electronegativities When a fluorine atom (e. neg.=4) shares electrons with a chlorine atom (e. neg.=3.0), the fluorine pulls the electrons more to itself, making the molecule slightly negative on the fluorine end and slightly positive on the chlorine end.

33. Uneven Sharing Causes A Polar Covalent Bond The positive and negative ends of the molecule are called poles. The bond is called a polar covalent bond, one which has uneven sharing of electrons.

34. Valence for Polar Covalent Compounds I In polar covalent compounds (uneven sharing of electrons), the element with the greater electronegativity is written second and the element with the lesser electronegativity is written first. Example: HF

35. Valence for Polar Covalent Compounds II Example: The most electronegative atom becomes negative (has the shared electrons most of the time) while the least electronegative atom becomes positive (having the shared pair less of the time).

36. Valence for Polar Covalent Compounds III The most electronegative atom is assigned a negative valence while the least electronegative atom is assigned a positive valence. Examples: ClF Cl+F - HF H+F - OF2 O+2F2-1

37. Valence for Polar Covalent Compounds IV The valence number assigned to a polar covalent element is the number of electrons that element is sharing while the valence sign indicates if it is stronger (negative) or weaker(positive) in attracting electrons. Example: compound NF3 (called nitrogen trifluoride) Valences Assigned: N+3F3-1 Note that this pattern is similar to the ionic compounds like Ca+2Cl2-1 where the + valence element is written first.

38. Ionic Bond : Extreme Uneven Electron Sharing When atoms have an extreme uneven sharing of electrons because their difference in electronegativity is so large, the bond is referred to as an ionic bond and the pair of electrons is virtually transferred from one atom to the other atom. .9 3.0 Ex: Na + Cl  NaCl

39. Determining Bond Character The difference in electronegativity of the two elements forming a bond will determine if the bond is covalent, polar covalent or ionic.

40. Table of Electronegativities (Heath p333)

41. Sample Problem What is the bond type in CH4 ? Step 1 : Find the electronegativities of C (2.5) and H (2.1). Step 2 : Find the difference in these electronegativities (.4) Step 3 : Assign the bond type based on EN diff. table Step 3 : This bond would be classified as a nonpolar covalent bond

42. Sample Problem 2 What is the bond type in Li3N ? Step 1 : Find the electronegativities of Li (1.0) and N (3.0). Step 2 : Find the difference in these electronegativities (2) Step 3 : The bond type would be ionic based on EN diff. table

43. Ionic Compounds and Ionic Bonds • Ionic compounds are produced when their atoms form ionic bonds. An ionic bond is an extreme polar covalent bond in which the electron is so unevenly shared that it is effectively transferred from one atom to another. If atoms have an electronegativity difference of 1.7 or larger, they will form ionic bonds.

44. Ionic Solids and Bond Strength • Ionic bonds are very strong bonds so compounds with these bonds have high melting and boiling points. • Ionic bonds produce solids in which oppositely-charged ions alternate in a crystal lattice. • Covalent bonds (polar and nonpolar) produce neutral molecules and molecular substances.

45. Comparison of Melting Points of Ionic Solids In the unit on periodicity, it was noted that within a chemical family, the nucleus of larger atoms (lower in the vertical column) has less ability to hold its outer valence electrons due to their greater distance from the nucleus and the screening effect of inner electrons. Thus larger atoms in a chemical family bond with weaker bonds than smaller atoms in the same family. Thus a compound like LiF has stronger bonds than a compound like NaCl since the nuclei of Li and F have stronger attractions than the nuclei of Na and Cl. This causes the melting point of LiF (845 C) to be higher than the melting point of NaCl (801 C) or KBr (734 C). Likewise, mp BeO: (2570 C) and MgS: (2000)

46. Elements forming Ionic or Covalent Compounds Metals and nonmetals form ionic bonds (substances). Nonmetals with nonmetals form polar covalent bonds.

47. Comparison of Bond Lengths • The distance between bonded atoms varies for single double and triple bonds. The interatomic distance becomes less as the bond goes from single to double to triple bonds.

48. Resonance (Heath p 347) of SO3 • When drawing Lewis structures for chemicals, often more than one structure can be composed. The question must be asked, “Which one is correct?”.

49. Resonance • Since the distance between double bonds is less, bond lengths can be measured experimentally to identify any differences. • Experimental measurements show that all bonds are identical and intermediate in length between single and double bonds.

50. Resonance • Resonance is a concept that tries to explain the intermediate bond length and different resonance structures.