* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
Download Reactions
Stability constants of complexes wikipedia , lookup
Coupled cluster wikipedia , lookup
Equilibrium chemistry wikipedia , lookup
Ionic compound wikipedia , lookup
Surface properties of transition metal oxides wikipedia , lookup
Nanofluidic circuitry wikipedia , lookup
Woodward–Hoffmann rules wikipedia , lookup
Marcus theory wikipedia , lookup
Physical organic chemistry wikipedia , lookup
Detailed balance wikipedia , lookup
Chemical equilibrium wikipedia , lookup
Chemical thermodynamics wikipedia , lookup
Acid–base reaction wikipedia , lookup
Reaction progress kinetic analysis wikipedia , lookup
Electrolysis of water wikipedia , lookup
Ene reaction wikipedia , lookup
Hydrogen-bond catalysis wikipedia , lookup
Electrochemistry wikipedia , lookup
George S. Hammond wikipedia , lookup
Transition state theory wikipedia , lookup
Reactions Sentences in Chemistry I. Formulas show chemistry at a standstill. Equations show chemistry in action. A. Equations show: 1. the reactants which enter into a reaction. 2. the products which are formed by the reaction. 3. the amounts of each substance used and each substance produced. B. Two important principles to remember: 1. Every chemical compound has a formula which cannot be altered. 2. A chemical reaction must account for every atom that is used. This is an application of the Law of Conservation of Matter which states that in a chemical reaction atoms are neither created nor destroyed. Some things to remember about writing equations: 1. The diatomic elements when they stand alone are always written H2, N2, O2, F2, Cl2, Br2, I2 2. The sign, ----->, means "yields" and shows the direction of the action. 3. A small delta, (∆), above the arrow shows that heat has been added. 4. A double arrow, <----->, shows that the reaction is reversible and can go in both directions. 5. Before beginning to balance an equation, check each formula to see that it is correct. NEVER change a formula during the balancing of an equation. 6. Balancing is done by placing coefficients in front of the formulas to insure the same number of atoms of each element on both sides of the arrow. 7. Always consult the Activity Series of metals and non-metals before attempting to write equations for replacement reactions. 8. If a reactant or product is solid, place (s) after the formula 9. If the reactant or product is a liquid, place (l) after the formula 10. If the reactant or product is a gas, place (g) after the formula 11. If the reactant or product is in water, place (aq) after the formula 12. A category of reaction will produce an unstable product which decomposes: H2CO3 (aq) H2O (l) + CO2 (g) carbonic acid H2SO3 (aq) H2O (l) + SO2 (g) sulfurous acid NH4OH (aq) NH3 (g) + H2O (l) ammonium hydroxide Rules for Writing Chemical Equations: 1. Write down the formula(s) for the reactants and add a “+” between them then put a “yields” arrow () at the end 2. Examine the formula(s) to determine which of four basic types of reactions will occur. On the basis of your decision, write down correct formula(s) for the products on the right side of the yield arrow. 3. Place a coefficient in front of each species in the reactants and the products to ensure that the conservation of matter is observed: “Balance” the reaction II. Four Basic Types of Chemical reactions A. Synthesis two or more elements or compounds combine to give a more complex product Examples of Synthesis reactions: 1. Metal + oxygen -----> metal oxide 2Mg(s) + O2(g) ----> 2MgO(s) 2. Nonmetal + oxygen -----> nonmetallic oxide C(s) + O2(g) ----> CO2(g) 3. Metal oxide + water -----> metallic hydroxide MgO(s) + H2O(l) ----> Mg(OH)2(s) 4. Nonmetallic oxide + water -----> acid CO2(g) + H2O(l) ----> ; H2CO3(aq) 5. Metal + nonmetal -----> salt 2 Na(s) + Cl2(g) ----> 2NaCl(s) 6. A few nonmetals combine with each other. 2P(s) + 3Cl2(g) ----> 2PCl3(g) B. Decomposition: • A single compound breaks down into simpler compounds. • Basic form: AX -----> A + X Examples of decomposition reactions (when heated): 1. Metallic carbonates form metallic oxides and CO2(g). CaCO3(s) ----> CaO(s) + CO2(g) 2. Most metallic hydroxides decompose into metal oxides and water. Ca(OH)2(s) ----> CaO(s) + H2O(g) 3. Metallic chlorates decompose into metallic chlorides and oxygen. 2KClO3(s) ----> 2KCl(s) + 3O2(g) 4. Some acids decompose into nonmetallic oxides and water. H2SO4 ----> H2O(l) + SO3(g) 5. Some oxides decompose. 2HgO(s) ----> 2Hg(l) + O2(g) 6. Some decomposition reactions are produced by electricity. 2H2O(l) ----> 2H2(g) + O2(g) 2NaCl(l) ----> 2Na(s) + Cl2(g) C. Replacement: • a more active element takes the place of another element and frees the less active one. • Basic form: A + BX -----> AX + B or AX + Y -----> AY + X Examples of replacement reactions: 1. Replacement of a metal in a compound by a more active metal. Fe(s) + CuSO4(aq) ----> FeSO4(aq) + Cu(s) 2. Replacement of hydrogen in water by an active metal. 2Na(s) + 2H2O(l) ----> 2NaOH(aq) + H2(g) Mg(s) + H2O(g) ----> MgO(s) + H2(g) 3. Replacement of hydrogen in acids by active metals. Zn(s) + 2HCl(aq) ----> ZnCl2(aq) + H2(g) 4. Replacement of nonmetals by more active nonmetals. Cl2(g) + 2NaBr(aq) ----> 2NaCl(aq) + Br2(l) NOTE: Refer to the Activity Series for metals and nonmetals to predict products of replacement reactions. If the free element is above the element to be replaced in the compound, then the reaction will occur. If it is below, then no reaction occurs. If the free element is above the element to be replaced in the compound, Then the reaction will occur. If the free element is below it, no reaction occurs. D. Ionic or Double Displacement: • occurs between ions in aqueous solution. A reaction will occur when a pair of ions come together to produce at least one of the following: 1. a precipitate 2. a gas 3. water or some other non-ionized substance. • Basic form: AX + BY -----> AY + BX Examples of ionic reactions: 1. Formation of precipitate. NaCl (aq) + AgNO3(aq) ----> NaNO3(aq) + AgCl(s) BaCl2(aq) + Na2 SO4(aq) ----> 2NaCl(aq) + BaSO4(s) 2. Formation of a gas. HCl(aq) + FeS(s) ----> FeCl2(aq) + H2S(g) 3. Formation of water. (If the reaction is between an acid and a base it is called a neutralization reaction.) HCl(aq) + NaOH(aq) ----> NaCl(aq) + H2O(l) 4. Formation of a product which decomposes. CaCO3(s) + HCl(aq) ----> CaCl2(aq) + CO2(g) + H2O(l) Use the Solubility Rules to decide whether a product of an ionic reaction is insoluble in water and will thus form a precipitate. If a compound is soluble in water then it should be shown as being in aqueous solution, or left as separate ions. It is, in fact, often more desirable to show only those ions that are actually taking part in the actual reaction. Equations of this type are called Net Ionic equations.