1. CHEMICAL BONDING

2. ***Occurs when atoms of elements combine together to form compounds.*****

3. Formation of compounds  Involves valence electrons.  PE is lower in bonded atoms.  Attractive force that develops is called "chemical bond“  Occurs during chem. reactions

4. Two (or three) methods Ionic bonding - attraction of ions Covalent bonding - shared pairs of electrons Metallic bonding - alloys (metals) not compounds

5. An ionic bond -(electrovalent) a. Definition - An attraction that forms between oppositely charged ions b. Pos. ions + Neg. ions → neutral compound ∆EN > 1.7

6. Ionic Bonds Cations: Metals lose electron(s) and become positive ions(+). Examples: Na +1, K +1, Mg +2, Ca +2, Al +3 Anions: Nonmetals gain electron(s) & becomes negative ions(-). Name ends in “–ide” Examples: Cl 1-, Br 1-, O 2- Choride, bromide, oxide...

7. Ionic bonds - (crystals)

8. Characteristics of Ionic Solids Made from + and - ions Metal-nonmetal (or polyatomic ions) Compound is neutral. Tend to be solid, Brittle and crystalline High MP and BP Have strong attractions in all directions Non-conductors as solids, but will conduct when molten or dissolved Some dissolve readily

9. Ionic solids are brittle due to Crystal lattice +-+- + - +- +-+- + - +-

10. Ionic solids are brittle + - + - + - +- +-+- + - +- Strong Repulsion breaks crystal apart.

11. Lattice Energy (ionic) More correct than bond energy for crystals.

12. Metallic Bonding Sea of Electrons In metals, the valence electrons are not bonded to any specific atom. (delocalized) Able to move freely over the positive centers. Causes unique properties.

13. Characteristics of Metals A. Malleable: dent when hammered. B. Ductile: draw into a wire C. Conductivity: electricity and heat D. Alloy: a blend of metals

14. Sea of Electrons ++++ ++++ ++++ Electrons are free to move through the solid. Metals conduct electricity and heat well.

15. Malleable ++++ ++++ ++++

16. ++++ ++++ ++++ Electrons allow atoms to slide by.

17. Alloys Alloys: mixtures of two or more metals. Important because their properties are often better than the individual elements. Examples: –Bronze is made from copper & tin. It is harder than copper & more easily cast. –Sterling silver: Ag (92.5%) & Cu (7.5%) –Stainless steel: Fe (80.6%), Cr (18%), C (0.4%), & Ni (1%) -- Brass, Pewter, and others.

18. Covalent Bond a. an attractive force that develops between atoms that are sharing pairs of electrons ∆EN < 1.7 b. Hydrogen – H 2 H + H → H : H (dot diagram) Structural formulas use a dash H - H

19. Characteristics of Covalent Compounds Nonmetal-nonmetal combinations Can be gases, liquids, or solids Low to med. MP and BP Insulators/Nonconductors (except for acids) Molecular (a few are crystalline) Generally not soluble (some polar exceptions)

20. Bond energy Energy required to break a bond. Bonds form to lower PE, so breaking bonds will increase PE.

21. Breaking Bonds is always endothermic Energy is required. ALWAYS.

23. Why not ionic? The difference in electronegativity is less than 1.7 Electrons are not pulled away from either atom. They are shared.

24. Fluorine – F 2

25. Ammonia -NH 3

26. Special Types of Covalent Bonds Multiple bonds – Occasionally atoms share more than one pair of electrons –Double bond – two shared pairs  Ex. O 2 O=O or

27. - Triple bonds - two atoms share three pairs of electrons. Ex. N 2 N N or

28. How to write Lewis Structures 1. Set out the atoms – think symmetry. 2. Count all the valence e - 3. Insert single bonds first, then fill rest. A ll the e - are paired. each nonmetal atom requires an octet. H only requires 2 e -. Multiple bonds may be needed. Readily formed by C, N, O, S, and P.

29. Polarity An unequal sharing of electrons due to difference in electronegativity. Polar bond – Any bond with ∆EN 0.5 - 1.7 Polar molecule – Has a positive end and a negative end. Occurs in water and ammonia (**know these two) Causes intermolecular attraction to increase

30. Molecular Substances Covalently bonded substances – show more variety in phases and properties. Tend to be insulators(nonconductors). Nonmetal-nonmetal combinations 1. Nonpolar – tends to be gases at room temp. – have only dispersion (Vanderwaals) forces. Have low MP and BP and high VP (vapor pressure) 2. Polar –(dipoles) tend to be liquids or solids at room temp. Have ↑MP and BP and↓VP 3. H- bonding – very strong type of polar

31. VSEPR VALENCE SHELL ELECTRON PAIR REPULSION A theory which describes the shapes of molecules based on the idea that pairs of electrons will repel each other as much as possible.

32. Polar molecules are asymmetrical and have polar bonds. Nonpolar molecules are symmetrical and may or may not have polar bonds.

33. Intermolecular Attractions ( Vanderwaals forces) Attractive forces between molecules Happens to covalent molecules. Strongest - hydrogen bonds Medium - dipoles/ polar Weakest - dispersion or London forces ( All molecules have London forces)

34. Hydrogen bonding In some highly polar compounds, a H atom is attracted to, and forms a weak bond with, an adjacent molecule. Only occurs in compounds where there is: H-F (strongest) H-O H-N (weakest) (know these three)

35. Hydrogen Bonding H H O ++ -- ++ H H O ++ -- ++

36. Hydrogen bonding H H O H H O H H O H H O H H O H H O H H O

37. H-bonding Gives water its very unusual properties. High MP and BP Holds the DNA molecule together. Provides stability and shapes for proteins, enzymes, etc. Strongest type of intermolecular attraction.

38. London Dispersion Forces (Van der waals) Weakest type of intermolecular attraction. Develops between nonpolar molecules due to temporary shifts in the electron positions. Strength of attraction is directly proportional to the number of electrons (wax is a nonpolar molecule – but large so it is solid)

39. Network Solids Large arrays of covalently bonded crystals. Do not conduct, hardest solids Very high MP and BP Examples: diamond, graphite, SiO 2 (very few – easiest to just memorize) More on this in 2 nd semester

40. Bond length and energy Bond length depends on the size of the atoms/ions and the number of bonds between them C-C is longer than C=C is longer than C  C Shorter bonds are stronger. Measured in kJ/mol

41. Breaking Bonds is always endothermic Energy is required. ALWAYS.

42. Endothermic reactions

43. Forming bonds is exothermic Energy is released. ALWAYS

44. Exothermic reactions