Balanced Chemical Equations: Represent reactants, products, and their amounts Make use of chemical formulas i.e. H 2 O can not be altered as they represent chemical compounds. Changing the formula changes the compound. Supports the Law of Conservation of Matter as no atoms are created or destroyed

When writing equations remember: Diatomic elements, when free and uncombined include: H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2 memorize these ↔, shows the reaction is reversible →, means "yields" and gives the reaction direction. , shows heat has been added. → Use the symbols in parenthesis for solids (s), liquids (l), gases (g) or aqueous (aq) reactants or products

Writing Reactions, Hints for Success Use the Activity Series before writing replacement reactions. Be sure the reaction actually occurs.Activity Series Recognize unstable products that decompose. Write the decomposition products instead. H 2 CO 3(aq) → H 2 O (l) + CO 2(g) Carbonic acid, such as found in soft drinks H 2 SO 3(aq) → H 2 O (l) + SO 2(g) Sulfurous acid decomposes as it is formed. NH 4 OH (aq) → NH 3(g) + H 2 O (l) ammonium hydroxide smells like ammonia gas because it decomposes into ammonia.

Classifying Reactions…Four Types exist. A. Synthesis (combination): two or more elements or compounds combine. A + X → AX Examples Metal + oxygen → metal oxide Mg (s) + O 2(g) → 2MgO (s) Nonmetal + oxygen → nonmetallic oxide C (s) + O 2(g) → CO 2(g) Metal oxide + water → metallic hydroxide MgO (s) + H 2 O (l) → Mg(OH) 2(s) Nonmetallic oxide + water → acid CO 2(g) + H 2 O (l) → H 2 CO 3(aq) Metal + nonmetal → salt 2 Na (s) + Cl 2(g) → 2NaCl (s) A few nonmetals combine with each other. 2P (s) + 3Cl 2(g) → 2PCl 3(g) These two synthesis reactions should be memorized: N 2(g) + 3H 2(g) → 2NH 3(g) NH 3(g) + H 2 O (l) → NH 4 OH (aq)

B. Decomposition: A single compound breaks down into its component parts or simpler compounds. AX → A + X Examples: Metallic carbonates, when heated, form metallic oxides and CO 2(g). CaCO 3(s) → CaO (s) + CO 2(g) Most metallic hydroxides, when heated, decompose into metallic oxides and water. Ca(OH) 2(s) → CaO (s) + H 2 O (g) Heating metallic chlorates produces metallic chlorides and oxygen. 2KClO 3(s) → 2KCl (s) + 3O 2(g) Some acids, when heated, decompose into nonmetallic oxides and water. H 2 SO 4 → H 2 O (l) + SO 3(g) Some oxides, when heated, decompose. 2HgO (s) → 2Hg (l) + O 2(g) Some decomposition reactions are produced by electricity. 2H 2 O (l) → 2H 2(g) + O 2(g) 2NaCl (l) → 2Na (s) + Cl 2(g)

Replacement of a metal in a compound by a more active metal. Fe (s) + CuSO 4(aq) → FeSO 4(aq) + Cu (s) Replacement of hydrogen in water by an active metal. 2Na (s) + 2H 2 O (l) → 2NaOH (aq) + H 2(g) Mg (s) + H 2 O (g) → MgO (s) + H 2(g) Replacement of hydrogen in acids by active metals. Zn (s) + 2HCl (aq) → ZnCl 2(aq) + H 2(g) Replacement of nonmetals by more active nonmetals. Cl 2(g) + 2NaBr (aq) → 2NaCl (aq) + Br 2(l)

Activity Series of Metals & Nonmetals: This is an aid to predicting products of replacement reactions and other reactions. Each element on the list replaces (from a compound) any elements below it. The larger the gap between elements, the more vigorous the reaction. lithium to sodium are very active metals; they react with cold water to produce the metal hydroxide and hydrogen gas. magnesium to chromium are active metals; they react with very hot water or steam to form the metal oxide and hydrogen gas. The oxides of all of these first metals resist reduction by H 2. Iron to lead replace hydrogen from HCl and dilute sulfuric and nitric acids. Their oxides undergo reduction by heating with H 2, carbon, and carbon monoxide. lithium to copper combine directly with oxygen to form the metal oxide. Mercury to gold are often found free in nature (chemically uncombined). Their oxides decompose with mild heating, and they form oxides only indirectly. lithium potassium strontium calcium sodium magnesium aluminum zinc chromium iron cadmium cobalt nickel tin lead HYDROGEN antimony arsenic bismuth copper Mercury silver palladium platinum gold

Ionic, or double replacement/replacement: Basic form: AX + BY → AY + BX Occurs in solutions of aqueous ions that: Form a precipitate. NaCl (aq) + AgNO 3(aq) → NaNO 3(aq) + AgCl (s) BaCl 2(aq) + Na 2 SO 4(aq) → 2NaCl (aq) + BaSO 4(s) Form a gas. HCl (aq) + FeS (s) → FeCl 2(aq) + H 2 S (g) Form water. (acid and base reactions are neutralization reactions.) HCl (aq) + NaOH (aq) → NaCl (aq) + H 2 O (l) Form a product which decomposes. CaCO 3(s) + HCl (aq) → CaCl 2(aq) + CO 2(g) + H 2 O (l) Use solubility rules to decide if a product precipitates (is insoluble in water).solubility rules Show soluble products as aqueous or as separate aqueous ions. Showing only ions taking part in the reaction shows a net ionic equation

Combustion of Hydrocarbons: hydrocarbons burned with enough O 2 produce CO 2 and water vapor. (complete combustion). Otherwise incomplete combustion occurs in which CO and water vapor are products. Hydrocarbon (C x H y ) + O 2(g) → CO 2(g) + H 2 O (g) (complete) CH 4(g) + 2O 2(g) → CO 2(g) + 2H 2 O (g) 2C 4 H 10(g) + 13O 2(g) → 8CO 2(g) + 10H 2 O (g)

Solubility Rules 1.All common compounds of Group I and ammonium ions are soluble. 2.All nitrates, acetates, and chlorates are soluble. 3.All binary compounds of the halogens (other than F) with metals are soluble, except those of Ag, Hg(I), and Pb. Pb halides are soluble in hot water.) 4.All sulfates are soluble, except those of barium, strontium, calcium, lead, silver, and mercury (I). The latter three are slightly soluble. 5.Except for rule 1, carbonates, hydroxides, oxides, silicates, and phosphates are insoluble. 6.Sulfides are insoluble except for calcium, barium, strontium, magnesium, sodium, potassium, and ammonium.