1. 1 Unit 3 – Chemical Bonds

2. 2 Nuclear reactions  nucleus Chemical reactions  valence electrons Chemical bonds form when electrons are attracted to 2 different atoms When bonds form, energy is released ex.) Ca + S  CaS + energy When bonds break, energy is absorbed ex.) CaS + energy  Ca + S

3. 3 Like the noble gases He Ne Ar Kr Xe Rn Atoms are most stable when their last electron energy level is full— usually 8 valence electrons (2 for 1 st energy level) Atoms can reach a stable octet by forming chemical bonds: 1. Atoms gain or lose electrons ionic bond 2. Atoms share electrons covalent bond

4. 4 Ionic Bonds -- form when electrons are transferred from one atom to another Na 2 - 8 - 1 more stable if 1 e - is lost ( 2 – 8 matches Ne) +11 p -10 e +1 Na + sodium ion Cl 2 - 8 -7 more stable if 1 e - is gained ( 2 – 8 – 8 matches Ar) +17 p -18 e Cl - - chloride ion Since + and – attract, the Na + and Cl - bond to form NaCl, sodium chloride An electron has transferred from Na + to Cl - Example:

5. 5 Ionic formulas depend on electron gain/loss Examples: calcium ___ oxide ___ iodide ___ aluminum ___ I-I- Ca +2 O -2 Al +3 The formula for an ionic compound is neutral(+ = - ) Examples: Ca +2 I - Ca +2 I - 2 CaI 2 calcium iodide Al +3 O -2 Al 2 O 3 aluminum oxide 2 3 Dot Structures for Ionic Compounds: Example: Ca I I [Ca] +2 I[ ] - I

6. 6 Ionic Substances --high melting and boiling points -- hard and brittle --poor electrical conductors when solid --good electrical conductors when melted or dissolved – they ionize: separate into + and - ions

7. 7 Covalent Bonds -- when 2 atoms share electrons When 2 identical atoms share electrons, they share equally --the result is a non-polar covalent bond When 2 different atoms share electrons, they share unequally --one of the atoms attracts the electrons more strongly -- the “stronger” atom becomes slightly negative, the “weaker” atom becomes slightly positive Fluorine molecule -- F 2

8. 8 The result is a polar covalent bond The bond has 2 opposite (+ and -) ends, or poles The tendency of an atom to attract electrons is measured by its Electronegativity (see ref table S) High electronegativity = strong attraction for e - - F = 4.0 Low electronegativity = weak attraction for e - - Fr = 0.7

9. 9 Bonds can be classified by the difference in electronegativities of the atoms difference of 1.7 or more = ionic difference of less than 1.7 = covalent Atoms share pairs of e- -- they can share 2 or 3 pairs to form double or triple bonds Ex., CO 2, N 2

10. 10 Covalent bonds produce molecular substances -- groups of atoms, or molecules, are separate from each other -- soft -- poor conductors of heat and electricity -- low melting and boiling points

11. 11 Metals and Metallic Bonding -- positive metal ions form a framework, surrounded by mobile electrons -- good conductors of heat and electricity -- malleable (can be shaped by hammering; not brittle) -- most have high melting points

12. 12 Summary of Bonds and Substances BOND SUBSTANCE Ionic Ionic Polar Covalent Non-Polar Covalent } Molecular Metallic Metal

13. 13 Intermolecular Forces are weaker than ionic, covalent, or metallic bonds are important in molecular substances depend on the attraction between + and - charge A molecule with a separation of + and – is called a polar molecule, or a dipole To be a dipole, a molecule must have: 1. polar covalent bonds 2. an asymmetric distribution of charge not a dipole dipole

14. 14 The + end of one dipole attracts the – end of another H—N—H | H | H—C—H | H H—ClCl—Cl Examples:

15. 15 Hydrogen Bonds -- a dipole force formed when polar molecules contain H bonded to a small, electronegative atom: N, O, or F -- are stronger than other dipole forces Example: Boiling Point, ˚C H 2 Te...... -2 H 2 Se..... -41 H 2 S..... -61 H 2 O..... +100 There are also forces that attract non-polar molecules to each other – but these are weaker than dipole or hydrogen bonds

16. 16 Strength of intermolecular forces:1. H bonds 2. regular dipole 3. non-polar forces The intermolecular forces determine: 1. melting and boiling points 2. hardness 3. solubility -- “like dissolves like” polar substances mix with polar substances non-polar substances mix with non-polar substances Example: water, a polar liquid, dissolves ionic solids (strong = high) (strong = hard)

17. 17 Types of Chemical Reactions 1. Synthesis -- A + B  Ex.) Fe + O 2  Fe 2 O 3 432 Ex.) SO 3 + H 2 O  H 2 SO 4 2. Decomposition -- AB  Ex.) H 2 O  H 2 + O 2 22 Ex.) CaCO 3  CaO + CO 2 3.Single Replacement -- A + BC  Ex.) Zn + NaCl  ZnCl 2 + Na 22 Ex.) Cl 2 + KBr  KCl + Br 2 22 AC + B or BA + C A + B AB

18. 18 4. Double Replacement -- AB + CD  Ex.) MgI 2 + K 2 S  MgS + KI 2 Ex.) Ba(NO 3 ) 2 + (NH 4 ) 2 SO 4  BaSO 4 + NH 4 NO 3 2 AD + CB