1. Chemical Equations and Reactions Chemical Reaction: one or more substances are changed into one or more different substance Original substances- reactants New stuff- products Indication of a Chemical Reaction: 1) Evolution of heat or light 2) Production of a gas 3) Formation of a precipitate (a solid that is produced during a chemical reaction 4) Color change 5) Odor production

2. Chemical Equations The law of conservation of mass must be satisfied Coefficients: small whole number that appears in front of a formula in a chemical equation Word equation: methane + oxygen carbon dioxide + water Formula equation: CH 4 (g) + O 2 (g) CO 2 (g) + H 2 O(g) (not balanced) CH 4 (g) + 2O 2 (g) CO 2 (g) + 2H 2 O(g) (balanced) Balancing: make sure that the number of moles of an element are equal on both sides of the equation Catalyst: substance that a changes the rate of a chemical reaction but can be recovered unchanged.

3. Symbols used in Chemical Equations Yeilds reversible reaction (s) a solid (l) a liquid (g) a gas (aq) aqueous solution (dissolved in water) a gas a precipitate heat reactants are heated  MnO 4 formula of a catalyst used to alter the rate 2 atm pressure at which the reaction is carried out 0 C temperature at which the reaction is carried out

4. H 2 + Cl 2 2HCl 1 mol H 2 1 mol Cl 2 2 mol HCl 2.02 g H 2 70.90 g Cl 2 2 x 36.46 = 72.92 g HCl Diatomic molecules: H 2, F 2, Cl 2, Br 2, I 2, O 2, N 2 Balancing: 1) write out word equation is the problem is a written 2) balance atoms one at a time 3) first balance atoms that are combined into a formula 4) then balance polyatomic ions 5) Balance H atoms and O atoms after all other elements are balanced 6) Check the number for all atoms on both sides!! They have to be equal

5. Types of Chemical Reactions Combustion Reactions: substance combines with oxygen, releasing a large amount of energy (light/ heat) - Carbon Compounds + O 2 : always get CO 2 + H 2 O C 3 H 8 (g) + 5O 2 (g) 3CO 2 (g) + 4H 2 O(g) Synthesis Reaction: (composition reaction) two or more substances combine to form a new compound. A + X AX Reactions of elements with Oxygen and Sulfur: - 2Mg(s) + O 2 (g) 2MgO(s) Metals and Halogens : - 2Na(s) + Cl 2 (g) 2NaCl(s)

6. Metal Oxides and Water: CaO(s) + H 2 O(l) Ca(OH) 2 (s) Decomposition Reaction: single compound breaks down into two or more simpler substances - Opposite of synthesis reaction - usually only takes place if energy (electricity/ heat) is added AX A + X Electrolysis: decomposition of a substance by an electric current If H in the compound you get H 2 O If C in the compound you will get CO 2 If O and no H or C, then you get O 2 H 2 SO 3 SO 2 + H 2 O

7. Single Replacement (displacement) Reaction: when one element replaces a similar element in a compound A + BX AX + B Y + BX BY + X Mg(s) + 2HCl(aq) H 2 (g) + MgCl 2 (aq) Group 1 & 2 metals + H 2 O metal hydroxides + H 2 (g) d-block metal + H 2 O metal oxides + H 2 (g) Double Replacement Reactions: ions of 2 compounds exchange places in an aqueous solution to form new AX + BY AY + BX FeS(s) + 2HCl(aq) H 2 S(g) + FeCl 2 (aq)

8. Solubility Rules 1. Alkali metals (gr. 1) and Ammonium compounds are soluble in water (aq). 2. Nitrates, acetates, and chlorates are soluble (aq). 3. Most chlorides, bromides, and iodides are soluble (aq), except those of silver, mercury (I), and lead (s) 4. Most sulfates are soluble (aq), except those of barium, strontium, calcium, lead, and mercury (s). 5. Most carbonates, phosphates, and sulfites are insoluble (s), except those of alkali metals (gr. 1) and ammonium (aq). 6. Most sulfides are insoluble (s), except those of calcium, strontium, alkali metals (gr.1) and ammonium (aq). 7. Most oxides are insoluble (s), except those of calcium, barium, and alkali metals (gr. 1) and ammonium (aq) 8. Most hydroxides are insoluble (s), except those of barium, and alkali metals (gr. 1) and ammonium (aq)

9. Solubility Examples Ex: Predict compounds, solubility, and then balance 1. (NH 4 ) 2 S (aq) + Cd(NO 3 ) 2 (aq)  2. Mercury (II) chloride (aq) + potassium sulfide (aq)  3. Sodium carbonate (aq) + calcium chloride (aq) 

10. Activity Series Activity: ability of an element to react the more an element reacts with other substances, the greater the activity is. Metals: the greater the activity, the greater it loses electrons (to form cations) Non-metals: the greater the activity, the greater it gains electrons (to form anions) Activity series: a list of which elements a particular element can replace in a single replacement reaction - predicts whether or not a chemical reaction will occur 2Al(s) + 3ZnCl 2 (aq) 3Zn(s) + 2AlCl 3 (aq) Co(s) + 2NaCl(aq) no reaction

11. React vigorously with Acids to produce H 2 gas React with Acids to produce H 2 gas Do not react with Acids Li MOST REACTIVE React vigorously with liquid water to form H 2 gas React slowly with liquid water but readily with steam to form H 2 gas LEAST REACTIVE K Ba Ca Na Mg Al Zn Cr Fe Cd Co Ni Sn Pb H2H2 Cu Hg Ag Au

12. Oxidation Numbers - used for covalent bonds and instead of charges (ions) 1) Atoms of a pure element have an oxidation number “0”. Pure Na or O 2 have an oxidation # equal to “0” 2) More electronegative element has oxid # = to its charge (if it were an anion). Least electronegative element has oxid # = to its charge (if it were a cation). 3) Oxygen is -2, except when it is O 2, then it is -1 4) Hydrogen is +1, except when with a metal, it is -1 5) the sum of the oxid. #s is equal to zero (neutral), except for a polyatomic ion, when the sum is equal to the ion’s overall charge. (NO 3 -1 )

13. Redox Reactions Reduction and Oxidation Reactions Reduction Reaction: gain electrons –Gaining electrons results in a more negative charge and therefore reduced! GER –Cl 2 + 2e - 2Cl -, the chlorine diatomic molecule goes from oxidation state 0 to a chlorine ion with a –1 charge Oxidation Reaction: lose electrons –Losing electrons results in less negative charge LEO –Increase in oxidation state (charge) because more positive (ie. Less negative) –Na Na + + e -, sodium goes from oxidation state 0 to a sodium ion with a +1 charge

14. Reduction: when an atom/molecule turns into a negative ion it is reduced, made more negative. Oxidation: when an atom of/molecule turns into a positive ion it is oxidized, made more positive. H 2 + Cl 2 2HCl 00+1 -1

15. 2H 2 O + 2Cl 2  4HCl + O 2 Academic: look in book for examples and go over on board.

16. Stoichiometry Grams X (molar mass) mol X mol X (mol ratio) mol Y mol Y (molar mass) grams Y Mole ratio: conversion factor that relates the amounts in moles of any two substances involved in a chemical reaction 2Al 2 O 3 (l) 4Al(s) + 3O 2 (g) 2 mol Al 2 O 3 or 4 mol Al 4 mol Al 2 mol Al 2 O 3 3 mol O 2 or 2 mol Al 2 O 3 2 mol Al 2 O 3 3 mol O 2

17. Limiting Reactants The reactant that limits the amounts of the other reactants that can combine in a chemical reaction (it also limits the amount of product that can form) look at moles, not grams Excess reactant: The substance that is not used up To determine the limiting reactant- must convert to moles –moles for all substances must be for the same reactant –convert moles of Y to moles of X, then compare to original moles of X

18. Percent Yeild The ratio of the actual yield to the theoretical yield percent yield = actual yield x 100 theoretical yield Given: the actual yield of a product the mass (grams) of some reactant 1) first change the mass of reactant to grams of product using molar mass, then mole ratio, then molar mass 2) This will be theoretical yield