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Chemical Reactions Types of Reactions • There are five types of chemical reactions we will talk about: 1. 2. 3. 4. 5. • Synthesis reactions _____________ reactions Single displacement reactions ________________ reactions Combustion reactions You need to be able to identify the type of reaction and predict the product(s) Steps to Writing Reactions • Some steps for doing reactions 1. 2. 3. Identify the type of reaction Predict the product(s) using the type of reaction as a model Balance it Don’t forget about the diatomic elements! (HOFBrINCl) For example, Oxygen is O2 as an element. In a compound, it can’t be a diatomic element because it’s not an element anymore, it’s a compound! Synthesis reactions • • Synthesis reactions occur when two substances (generally elements) combine and form a compound. (Sometimes these are called combination or addition reactions.) reactant + reactant 1 product Basically: A + B AB • • • • Example: 2H2 + O2 2H2O Example: C + O2 CO2 Sodium and Chlorine Reaction.mov Formation of AlBr3.MOV Synthesis Reactions • Here is another example of a synthesis reaction Practice • • • • Predict the products. Write and balance the following synthesis reaction equations. Sodium metal reacts with chlorine gas Na(s) + Cl2(g) Solid Magnesium reacts with fluorine gas Mg(s) + F2(g) Aluminum metal reacts with fluorine gas Al(s) + F2(g) Decomposition Reactions • Decomposition reactions occur when a compound breaks up into the elements or in a few to simpler compounds 1 Reactant Product + Product In general: AB A + B Example: 2 H2O 2H2 + O2 Example: 2 HgO 2Hg + O2 • Decomposition of HgO.mov • • • • Decomposition Reactions • Another view of a decomposition reaction: Decomposition Exceptions • Carbonates and chlorates are special case decomposition reactions that do not go to the elements. • Carbonates (CO32-) decompose to carbon dioxide and a metal oxide • • Example: CaCO3 CO2 + CaO Chlorates (ClO3-) decompose to oxygen gas and a metal chloride • Example: 2 Al(ClO3)3 2 AlCl3 + 9 O2 Decomposition of Water Water cannot be decomposed by heating (it boils, but does not change into a new substance). Instead, an electrical current must be passed through an acidified sample of the water to break it down into hydrogen and oxygen gases. Electrolysis of Water Electrolysis of Water.MOV Practice • • • Predict the products. Then, write and balance the following decomposition reaction equations: Solid Lead (IV) oxide decomposes PbO2(s) Aluminum nitride decomposes AlN(s) Practice Identify the type of reaction for each of the following synthesis or decomposition reactions, and write the balanced equation: N2(g) + O2(g) Nitrogen monoxide BaCO3(s) Co(s)+ S(s) (make Co be +3) NH3(g) + H2CO3(aq) NI3(s) Single Replacement Reactions • • • Single Replacement Reactions occur when one element replaces another in a compound. A metal can replace a metal (+) OR a nonmetal can replace a nonmetal (-). element + compound product + product A + BC AC + B (if A is a metal) OR A + BC BA + C (if A is a nonmetal) (remember the cation always goes first!) When H2O splits into ions, it splits into H+ and OH- (not H+ and O-2 !!) Sodium and Potassium Plus Water.mov Single Replacement Reactions • Another view: Single Replacement Reactions Write and balance the following single replacement reaction equation: • Zinc metal reacts with aqueous hydrochloric acid Zn(s) + 2 HCl(aq) ZnCl2 + H2(g) Note: Zinc replaces the hydrogen ion in the reaction • Single Replacement Reactions • Sodium chloride solid reacts with fluorine gas 2 NaCl(s) + F2(g) 2 NaF(s) + Cl2(g) Note that fluorine replaces chlorine in the compound • Aluminum metal reacts with aqueous copper (II) nitrate Al(s)+ Cu(NO3)2(aq) Double Replacement Reactions • • • Double Replacement Reactions occur when a metal replaces a metal in a compound and a nonmetal replaces a nonmetal in a compound Compound + compound product + product AB + CD AD + CB Double Replacement Reactions • • • Think about it like “foil”ing in algebra, first and last ions go together + inside ions go together Example: AgNO3(aq) + NaCl(s) AgCl(s) + NaNO3(aq) Another example: K2SO4(aq) + Ba(NO3)2(aq) 2 KNO3(aq) + BaSO4(s) Practice • Predict the products. Balance the equation 5. HCl(aq) + AgNO3(aq) CaCl2(aq) + Na3PO4(aq) Pb(NO3)2(aq) + BaCl2(aq) FeCl3(aq) + NaOH(aq) H2SO4(aq) + NaOH(aq) 6. KOH(aq) + CuSO4(aq) 1. 2. 3. 4. Combustion Reactions • • Combustion reactions occur when a hydrocarbon reacts with oxygen gas. This is also called burning!!! In order to burn something you need the 3 things in the “fire triangle”: 1) A Fuel (hydrocarbon) 2) Oxygen to burn it with 3) Something to ignite the reaction (spark) Combustion Reactions • • • In general: CxHy + O2 CO2 + H2O Products in combustion are ALWAYS carbon dioxide and water. (although incomplete burning does cause some byproducts like carbon monoxide) Combustion is used to heat homes and run automobiles (octane, as in gasoline, is C8H18) Combustion Reactions Edgar Allen Poe’s drooping eyes and mouth are potential signs of CO poisoning. Combustion • Example • • C5H12 + 8 O2 5 CO2 + 6 H2O Write the products and balance the following combustion reaction: • C10H22 + O2 Combustion Substances other than hydrocarbons can also combust. However, you may not be able to tell whether it’s combustion from the chemical equation alone. Remember that combustion must have O2 as a reactant and must release (exothermic) heat and light energy. Reactions with O2.mov Mixed Practice • 1. 2. 3. 4. 5. State the type, predict the products, and balance the following reactions: BaCl2 + H2SO4 C6H12 + O2 Zn + CuSO4 Cs + Br2 FeCO3 Total Ionic Equations Once you write the molecular equation (synthesis, decomposition, etc.), you should check for reactants and products that are soluble or insoluble. We usually assume the reaction is in water We can use a solubility table to tell us what compounds dissolve in water. If the compound is soluble (does dissolve in water), then splits the compound into its component ions If the compound is insoluble (does NOT dissolve in water), then it remains as a compound Solubility Rules Soluble Salts 1. The Na+, K+, and NH4+ ions form soluble salts. Thus, NaCl, KNO3, (NH4)2SO4, Na2S, and (NH4)2CO3 are soluble. 2. The nitrate (NO3-) ion forms soluble salts. Thus, Cu(NO3)2 and Fe(NO3)3 are soluble. 3. The chloride (Cl-), bromide (Br-), and iodide (I-) ions generally form soluble salts. Exceptions to this rule include salts of the Pb2+, Hg22+, Ag+, and Cu+ ions. ZnCl2 is soluble, but CuBr is not. 4. The sulfate (SO42-) ion generally forms soluble salts. Exceptions include BaSO4, SrSO4, and PbSO4, which are insoluble, and Ag2SO4, CaSO4, and Hg2SO4, which are slightly soluble. Insoluble Salts 1. Sulfides (S2-) are usually insoluble. Exceptions include Na2S, K2S, (NH4)2S, MgS, CaS, SrS, and BaS. 2. Oxides (O2-) are usually insoluble. Exceptions include Na2O, K2O, SrO, and BaO, which are soluble, and CaO, which is slightly soluble. 3. Hydroxides (OH-) are usually insoluble. Exceptions include NaOH, KOH, Sr(OH)2, and Ba(OH)2, which are soluble, and Ca(OH)2, which is slightly soluble. 4. Chromates (CrO42-) are usually insoluble. Exceptions include Na2CrO4, K2CrO4, (NH4)2CrO4, and MgCrO4. 5. Phosphates (PO43-) and carbonates (CO32-) are usually insoluble. Exceptions include salts of the Na+, K+, and NH4+ ions. Solubilities Not on the Table! Gases only slightly dissolve in water Strong acids and bases dissolve in water Hydrochloric, Hydrobromic, Hydroiodic, Nitric, Sulfuric, Perchloric Acids Group I hydroxides (should be on your chart anyway) Water slightly dissolves (ionizes) in water! (H+ and OH-) There are other tables and rules that cover more compounds than your table, but for the most part you can use the table in reference packet. Total Ionic Equations Molecular Equation: K2CrO4 + Pb(NO3)2 PbCrO4 + 2 KNO3 Soluble Insoluble Soluble Soluble Total Ionic Equation: 2 K+ + CrO4 -2 + Pb+2 + 2 NO3- PbCrO4 (s) + 2 K+ + 2 NO3- Net Ionic Equations These are the same as total ionic equations, but you should cancel out ions that appear on BOTH sides of the equation Total Ionic Equation: 2 K+ + CrO4 -2 + Pb+2 + 2 NO3- PbCrO4 (s) + 2 K+ + 2 NO3Net Ionic Equation: CrO4 -2 + Pb+2 PbCrO4 (s) Net Ionic Equations Try this one! Write the molecular, total ionic, and net ionic equations for this reaction: Silver nitrate reacts with Lead (II) Chloride in hot water. Molecular: Total Ionic: Net Ionic: “Oxidation-Reduction Reactions” LEO SAYS GER Oxidation and Reduction (Redox) Early chemists saw “oxidation” reactions only as the combination of a material with oxygen to produce an oxide. • For example, when methane burns in air, it oxidizes and forms oxides of carbon and hydrogen. Oxidation and Reduction (Redox) But, not all oxidation processes that use oxygen involve burning: •Elemental iron slowly oxidizes to compounds such as iron (III) oxide, commonly called “rust” •Bleaching stains in fabrics •Hydrogen peroxide also releases oxygen when it decomposes Oxidation and Reduction (Redox) A process called “reduction” is the opposite of oxidation, and originally meant the loss of oxygen from a compound Oxidation and reduction always occur simultaneously The substance gaining oxygen (or losing electrons) is oxidized, while the substance losing oxygen (or gaining electrons) is reduced. Transfer of Electrons Today, many of these reactions may not even involve oxygen Redox currently says that electrons are transferred between reactants Mg + S→ Mg2+ + S2- (MgS) •The magnesium atom (which has zero charge) changes to a magnesium ion by losing 2 electrons, and is oxidized to Mg2+ •The sulfur atom (which has no charge) is changed to a sulfide ion by gaining 2 electrons, and is reduced to S2- Assigning Oxidation Numbers 0 1 0 1 2 Na Cl 2 2 Na Cl Each sodium atom loses one electron: 1 0 Na Na e Each chlorine atom gains one electron: 0 1 Cl e Cl LEO says GER : Lose Electrons = Oxidation 1 0 Na Na e Sodium is oxidized Gain Electrons = Reduction 0 1 Cl e Cl Chlorine is reduced LEO says GER : - Losing electrons is oxidation, and the substance that loses the electrons is called the reducing agent. - Gaining electrons is reduction, and the substance that gains the electrons is called the oxidizing agent. Mg is the reducing agent Mg is oxidized: loses e-, becomes a Mg2+ ion Mg(s) + S(s) → MgS(s) S is the oxidizing agent S is reduced: gains e- = S2- ion It is easy to see the loss and gain of electrons in ionic compounds, but what about covalent compounds? In water, oxygen is highly electronegative, so: the oxygen gains electrons (is reduced and is the oxidizing agent), and the hydrogen loses electrons (is oxidized and is the reducing agent) Not All Reactions are Redox Reactions - Reactions in which there has been no change in oxidation number are NOT redox reactions. Examples: 1 5 2 1 1 1 1 1 5 2 Ag N O3 (aq) Na Cl (aq) Ag Cl ( s) Na N O3 (aq) 1 2 1 1 6 2 1 6 2 1 2 2 Na O H (aq) H 2 S O 4 (aq) Na 2 S O 4 (aq) H 2 O(l ) Assigning Oxidation Numbers • An “oxidation number” is a positive or negative number assigned to an atom to indicate its degree of oxidation or reduction. • Generally, a bonded atom’s oxidation number is the charge it would have if the electrons in the bond were assigned to the atom of the more electronegative element Rules for Assigning Oxidation Numbers 1) The oxidation number of any uncombined element is zero. 2) The oxidation number of a monatomic ion equals its charge. 0 0 1 1 2 Na Cl 2 2 Na Cl Rules for Assigning Oxidation Numbers 3) The oxidation number of oxygen in compounds is -2, except in peroxides, such as H2O2 where it is -1. 4) The oxidation number of hydrogen in compounds is +1, except in metal hydrides, like NaH, where it is -1. 1 2 H2O Rules for Assigning Oxidation Numbers 5) The sum of the oxidation numbers of the atoms in the compound must equal 0. 1 2 H2O 2(+1) + (-2) = 0 H O 2 2 1 Ca(O H ) 2 (+2) + 2(-2) + 2(+1) = 0 Ca O H Rules for Assigning Oxidation Numbers 6) The sum of the oxidation numbers in the formula of a polyatomic ion is equal to its ionic charge. ? 2 N O3 X + 3(-2) = -1 N O thus X = +5 ? 2 S O4 2 X + 4(-2) = -2 S O thus X = +6 Reducing Agents and Oxidizing Agents •An increase in oxidation number = oxidation • A decrease in oxidation number = reduction 1 0 Na Na e Sodium is oxidized – it is the reducing agent 0 1 Cl e Cl Chlorine is reduced – it is the oxidizing agent Trends in Oxidation and Reduction Active metals: Lose electrons easily Are easily oxidized Are strong reducing agents Active nonmetals: Gain electrons easily Are easily reduced Are strong oxidizing agents Other Types of Reactions as Redox? Synthesis Reactions from elements are Redox. Ex: 2Al (s) + 3Cl2 (g) 2AlCl3 (s) Decomposition Reactions to elements are Redox. Ex: 2H2O (l) 2H2 (g) + O2 (g) All Single Replacement Reactions are Redox. No Double Replacement Reactions are Redox. All Combusiton Reactions are Redox. Dehydration of Sugar.MOV