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Chemical reactions
 

Chemical reactions

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    Chemical reactions Chemical reactions Presentation Transcript

    • Unit 9:Chemical Reactions
    • Bellringer 10/15• List 4 signs that let you know a chemical reaction has occurred.
    • Objectives• Follow directions• Identify a cause and effect relationship in balancing chemical reactions• Balance Chemical reactions• List the signs of a chemical reaction
    • Introduction• Chemical reactions occur when bonds between the outermost parts of atoms are formed or broken• Chemical reactions involve changes in matter, the making of new materials with new properties, and energy changes.• Symbols represent elements, formulas describe compounds, chemical equations describe a chemical reaction
    • Parts of a ReactionEquation• Chemical equations show the conversion of reactants (the molecules shown on the left of the arrow) into products (the molecules shown on the right of the arrow). • A + sign separates molecules on the same side • The arrow is read as “yields” • Example C + O2  CO2 • This reads “carbon plus oxygen react to yield carbon dioxide”
    • Chemical EquationsTheir Job: Depict the kind of reactants and products and their relative amounts in a reaction.4 Al (s) + 3 O2 (g) ---> 2 Al2O3 (s)The numbers in the front are called CoefficientsThe letters (s), (g), and (l) are the physical states of compounds.
    • Chemical EquationsContinued• You can indicate the physical state of a substance by putting a symbol after each formula.• (s) – solid• (l) – liquid• (g) – gas• (aq) – aqueous solution (in water) For example: K(s) + Cl(g) → KCl(s)
    • Conservation of Matter• Matter cannot be created or destroyed. • It can change forms• Total mass in reactants must equal total mass of products
    • Terms with Equations• Activation energy—minimum energy colliding particles must have in order to react• Endothermic reaction—process that absorbs heat from the surroundings• Exothermic reaction—process that releases heat to the surroundings
    • Preview/Predict• Looking at the two rxns, label one as endothermic & one as exothermic. Explain why. SP
    • Catalysts• A substance used to speed up the rate of a reaction.• Neither a product nor a reactant.• Written above the arrow.
    • Symbols Used inEquations• Solid ___• Liquid ___• Gas ___• Aqueous solution ___ H2SO4• Catalyst• Escaping gas ()• Change of temperature ()
    • • The charcoal used in a grill is basically carbon. The carbon reacts with oxygen to yield carbon dioxide. The chemical equation for this reaction, C + O2  CO2, contains the same information as the English sentence but has quantitative meaning as well.
    • Chemical EquationsBecause of the principle of the conservation of matter,an equation must be balanced.It must have the same number of atoms of the same kind on both sides. Lavoisier, 1788
    • Balancing Equations• When balancing a chemical reaction you may add coefficients in front of the compounds to balance the reaction, but you may not change the subscripts. • Changing the subscripts changes the compound. Subscripts are determined by the valence electrons (charges for ionic or sharing for covalent)
    • Subscripts vs. Coefficients • The subscripts tell you how many atoms of a particular element are in a compound. The coefficient tells you about the quantity, or number, of molecules of the compound.
    • Chemical Equations4 Al(s) + 3 O2(g) ---> 2 Al2O3(s)This equation means4 Al atoms + 3 O2 molecules ---produces---> 2 molecules of Al2O3 AND/OR4 moles of Al + 3 moles of O2 ---produces---> 2 moles of Al2O3
    • Steps to Balancing EquationsThere are four basic steps to balancing a chemical equation. 1. Write the correct formula for the reactants and the products. DO NOT TRY TO BALANCE IT YET! You must write the correct formulas first. And most importantly, once you write them correctly DO NOT CHANGE THE FORMULAS! 2. Find the number of atoms for each element on the left side. Compare those against the number of the atoms of the same element on the right side. 3. Determine where to place coefficients in front of formulas so that the left side has the same number of atoms as the right side for EACH element in order to balance the equation. 4. Check your answer to see if: − The numbers of atoms on both sides of the equation are now balanced. − The coefficients are in the lowest possible whole number ratios. (reduced)
    • Some Suggestions to Help YouSome Helpful Hints for balancing equations: • Take one element at a time, working left to right except for H and O. Save H for next to last, and O until last. • IF everything balances except for O, and there is no way to balance O with a whole number, double all the coefficients and try again. (Because O is diatomic as an element) • (Shortcut) Polyatomic ions that appear on both sides of the equation should be balanced as independent units
    • Balancing Equations 2 2___ H2(g) + ___ O2(g) ---> ___ H2O(l) What Happened to the Other Oxygen Atom????? This equation is not balanced! Two hydrogen atoms from a hydrogen molecule (H2) combines with one of the oxygen atoms from an oxygen molecule (O2) to form H2O. Then, the remaining oxygen atom combines with two more hydrogen atoms (from another H2 molecule) to make a second H2O molecule.
    • Balancing Equations 2 3 ___ Al(s) + ___ Br2(l) ---> ___ Al2Br6(s)
    • Balancing Equations____C3H8(g) + _____ O2(g) ----> _____CO2(g) + _____ H2O(g)
    • ____B4H10(g) + _____ O2(g) ----> ___ B2O3(g) + _____ H2O(g)
    • Balancing EquationsSodium phosphate + iron (III) oxide  sodium oxide + iron (III) phosphate Na3PO4 + Fe2O3 ----> Na2O + FePO4
    • Diatomic Elements• Fluorine • F2• Chlorine • Cl2• Bromine • Br2• Iodine • I2• Hydrogen • H2• Nitrogen • N2• Oxygen • O2
    • Bellringer• List and describe the five types of chemical reactions listed in your book.
    • Types of Reactions• There are five types of chemical reactions we will talk about: 1. Synthesis reactions 2. Decomposition reactions 3. Single displacement reactions 4. Double displacement reactions 5. 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. Identify the type of reaction 2. Predict the product(s) using the type of reaction as a model 3. Balance it Don’t forget about the diatomic elements! (BrINClHOF) 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!
    • 1. Combination or Synthesis • Two or more substances combine to form a single substance. • Forms only ONE PRODUCT! R + S → RS For Example: Al + N2 → AlN
    • 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) 
    • 2. Decomposition• A single compound is broken down into two or more products.• Has only ONE REACTANT! RS → R + S
    • 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) 
    • PracticeIdentify the type of reaction for each of the following synthesis or decomposition reactions, and write the balanced equation: Nitrogen monoxideN2(g) + O2(g) BaCO3(s) (make Co be +3)Co(s)+ S(s) NI3(s) 
    • Bellringer 10/19• Predict the products of the following single replacement reaction, then balance the reaction, and correctly name each product: NaCl(s) + F2(g)  When you are done pick up your lab book.
    • Free Response• IN Your Lab Book:• Write an outline for the following prompt: When exposed to the natural elements Iron metal reacts with oxygen gas. In this natural state Iron usually has an oxidation state of +3. Tell what type reaction this is, write a balanced chemical equation, give the oxidation state for oxygen, and identify what substance is reduced in this reaction, and tell why that substance is reduced.
    • Bellringer• Pick up a worksheet from the front desk.• Tear a sheet of paper in half, widthwise, and label it January PreTest• Read the passages and answer the questions in the order they occur (#’s are messed up)
    • 3. Single-ReplacementReactions• An element replaces another element in a compound.• Whether one metal will displace another metal is determined by the activity series of metals chart.• A reactive metal will replace any metal listed below it in the activity series. For example, Mg will replace Zn. T + RS → TS + R When water splits it splits into H & OH
    • Single ReplacementReactions• Write and balance the following single replacement reaction equation:• Sodium metal reacts with aqueous hydrochloric acid Na(s) + HCl(aq) Note: Sodium replaces the hydrogen ion in the reaction
    • Single Replacement Reactions• Sodium chloride solid reacts with fluorine gas NaCl(s) + F2(g) Note that fluorine replaces chlorine in the compound• Aluminum metal reacts with aqueous copper (II) nitrate Al(s)+ Cu(NO3)2(aq)
    • 4. Double-Replacement• Two ionic compounds react by exchanging cations to form two different compounds.• Again, whether one metal will replace another depends on the activity series of metals chart. RS + TU → RU + TS
    • 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) 
    • More on dbl replacement • Takes place in aqueous solutions • Usually produces a precipitate, gas, or molecular compound • Aqueous solutions often disassociate (break apart) in water
    • Practice• Predict the products. Balance the equation1. HCl(aq) + AgNO3(aq) 2. CaCl2(aq) + Na3PO4(aq) 3. Pb(NO3)2(aq) + BaCl2(aq) 4. FeCl3(aq) + NaOH(aq) 5. H2SO4(aq) + NaOH(aq) 6. KOH(aq) + CuSO4(aq) 
    • 5. Combustion• A hydrocarbon reacts with oxygen producing energy as light and heat.• Hydrocarbons are compounds composed of C, H, and sometimes O.• Always forms carbon dioxide and water!! CHO + O2 → CO2 + H2O
    • Combustion Reactions• Products in combustion are ALWAYS carbon dioxide and water. (although incomplete burning does cause some by-products like carbon monoxide)• Combustion is used to heat homes and run automobiles (octane, as in gasoline, is C8H18)
    • Combustion ReactionsEdgar Allen Poe’sdrooping eyes andmouth are potentialsigns of COpoisoning.
    • Bellringer 10/18• Identify the following reactions as decomposition, combustion, single replacement, double replacement, or combination.1. T + RS → TS + R2. CHO + O2 → CO2 + H2O3. RS → R + S4. Al + N2 → AlN
    • Bellringer• give the products1. T+ + RS →2. CHO + O2 →3. RS →4. Al + N2 →
    • 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 Table
    • Solubilities Not on the Table!• Gases only slightly dissolve in water• Strong acids and bases dissolve in water • Hydrochloric, Hydrobromic, Hydroiodic, Nitri c, Sulfuric, Perchloric Acids • Group I hydroxides (should be on your chart anyway)• Water slightly dissolves in water! (H+ and OH-)• For the homework… SrSO4 is insoluble; BeI2 and the products are soluble• There are other tables and rules that cover more compounds than your table!
    • Total Ionic EquationsMolecular Equation:K2CrO4 + Pb(NO3)2  PbCrO4 + 2 KNO3Soluble Soluble Insoluble SolubleTotal 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 equationTotal Ionic Equation:2 K+ + CrO4 -2 + Pb+2 + 2 NO3-  PbCrO4 (s) + 2 K+ + 2 NO3-Net Ionic Equation:CrO4 -2 + Pb+2  PbCrO4 (s)
    • Net Ionic Equations• Shows the molecules broken apartAgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq) vs.Ag+ + NO3- + Na+ + Cl-  AgCl(s) + Na+ + NO3-Which ions are the same on both sides of the reaction? These are spectator ions! Na+ NO3-So, the Net equation is: Ag+ + Cl-  AgCl
    • 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:
    • Compounds SolubilitySalts of alkali Solublemetals &ammonia SolubilityNitrate salts Solubleand chloratesalts • Precipitates formSulfate salts, Soluble when an insolubleexceptcompounds solid is formed.with Pb2+, Ag+,Hg22+, Ba2+, • This can beSr2+, and Ca2+Chloride salts, Soluble predicted by usingexcept the solubility tablecompoundswith Pb2+, Ag+ on pg 344!and, Hg 2+ • Not soluble = 2Carbonates, Most arephosphates,chromates, insoluble precipitatesulfides, andhydroxides
    • Will a Precipitate Form?•2AgNO3 + H2S  Ag2S + 2 HNO3 (s)• Yes salts with Ag are not soluble!• The solid Ag2S will precipitate out!
    • Formation of a precipitate
    • Bellringer• Using complete sentences tell the difference between a net ionic equation and a complete ionic equation, then give an example of each.
    • Recall!1. Write 1 characteristic about each of the 5 reactions discussed in class to help you remember that reaction.
    • Recall1. Comb/synth: two (single substances) reactants become one product2. Decomp: one reactant breaks into two products3. Combust: combines with O24. SR: Has 1 single ion plus a compound on reactants side5. DR: Has 2 compounds on reactants side
    • Exit• What type of practice has helped you best understand writing and balancing equations?