UNIT 2 : Bonding, Reactions and Stoichiometry

Nuclear reactions involve the nucleus Chemical reactions involve 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

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

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

Dot Structures for Ionic Compounds: Ionic formulas depend on electron gain/loss (look at oxidation state) Examples: calcium ___ oxide ___ iodide ___ aluminum ___ Ca+2 O-2 I- Al+3 The formula for an ionic compound is neutral (+ = - ) Examples: Ca+2I- Ca+2I- CaI2 calcium iodide 2 Al+3O-2 Al+3O-2 Al2O3 aluminum oxide 2 3 Dot Structures for Ionic Compounds: x Example: Ca x I x x I [Ca]+2 [ ]- I [ ]- I

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

Covalent Bonds -- when 2 atoms share electrons Fluorine molecule -- F2 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

The bond has 2 opposite (+ and -) ends, or poles The result is a polar covalent bond 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

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 = polar covalent difference of 0 = nonpolar covalent Atoms share pairs of e- They can share 2 or 3 pairs to form double or triple bonds Ex., CO2, N2

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

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

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

Intermolecular Forces – Between molecules are important in molecular substances not ionic are weaker than ionic, covalent, or metallic bonds 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 dipole not a dipole

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

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 H2Te ...... -2 H2Se ..... -41 H2S ..... -61 H2O ..... +100 There are also forces that attract non-polar molecules to each other – but these are weaker than dipole or hydrogen bonds

Strength of intermolecular forces: 1. H bonds 2. regular dipole The intermolecular forces determine: 3. non-polar forces 1. melting and boiling points (strong forces = high) 2. hardness (strong forces = hard) 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 like salt

NAMING COMPOUNDS

REMEMBER THESE LAWS… In the 1700’s, scientists making accurate measurements discovered several new laws -- 1. Law of Conservation of Matter --matter is not created or destroyed during chemical or physical changes 2. Law of Definite Composition -- compounds have an unvarying composition 3. Law of Multiple Proportions -- elements combine in simple ratios

Then Amedeo Avogadro, Italy, 1800’s at the same temp and pressure, equal volumes of different gases have the same number of molecules Example: 1.0 L of H2 has the same number of molecules as 1.0 L of CO2. The CO2 weighs more because each CO2 molecule is heavier (formula mass H2 = 2, CO2 = 44) Since it’s difficult to work with single molecules, chemists usually work with groups of them: Avogadro’s number = 6.02 x 1023 6.02 x 1023 molecules = 1 mole Why 6.02 x 1023? --because 6.02 x 1023 molecules of any substance weigh the formula mass of that substance, in grams

Examples: 1 6.02 x 1023 Na atoms = ______moles = _______ grams 23 6.02 x 1023 H2O molecules = ______ moles 1 = _______ grams 18 3.01 x 1023 B atoms = _____ moles 0.5 = _______ grams 5.5 80 g calcium = _____ moles 2 12.04 x 1023 = _______________atoms 24.08 x 1023 4 48 g carbon = _____ moles = ______________ atoms 3.55 x 1023 16 g aluminum = ______ moles 0.59 = _____________ atoms grams . formula mass moles =

By Avogadro’s law, 1 mole of any gas must have the same volume The molar volume of a gas at STP = 22.4 L Examples: 89.6 L 1. What volume would 4 moles NO2 occupy, at STP? __________ 12.04 x 1023 2 2. 44.8 L of hydrogen at STP = ____ moles = ___________molecules 3. What is the mass of 22.4 L of O2 at STP? _________ 32 g 4. a. What volume does 1 mole NO occupy at STP? _________ 22.4 L 30 g b. What is the mass of 1 mole of NO? _________ 1.3 g/L c. What is the density of NO? __________ mass volume density = 5. A gas has a density of 1.78 g/L at STP a. what is the mass of 1 L of this gas? _________ 1.78 g b. how much does 22.4 L of this gas weigh? ________ 39.9 g c. what is the formula mass of this gas? ________ 39.9 g

Types of Chemical Reactions 1. Synthesis -- A + B  AB Ex.) Fe + O2  Fe2O3 4 3 2 Ex.) SO3 + H2O  H2SO4 2. Decomposition -- AB  A + B Ex.) H2O  H2 + O2 2 2 Ex.) CaCO3  CaO + CO2 Single Replacement -- A + BC  AC + B or BA + C Ex.) Zn + NaCl  ZnCl2 + Na 2 2 Ex.) Cl2 + KBr  KCl + Br2 2 2

4. Double Replacement -- AB + CD  AD + CB Ex.) MgI2 + K2S  MgS + KI 2 Ex.) Ba(NO3)2 + (NH4)2SO4  BaSO4 + NH4NO3 2