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AP Chemistry The Ultimate Chemical Equations Handbook HOMEWORK Do all exercises in this book on a separate sheet of paper. DO NOT WRITE IN THE BOOK. Chapter 1 Symbols elements There and Nomenclature of the is interesting info where the elements got their name, but nothing we will cover. Chapter 2 and 3 Naming Binary Compounds First, determine if you have an ionic compound or a covalent compound. A metal and a nonmetal will form an ionic bond. Compounds with Polyatomic ions form ionic bonds. Nonmetals bonding together or Nonmetals and a metalloid form covalent bonds. Covalent bonding is very similar to ionic naming You always name the one that is least electronegative first (furthest from fluorine) Most electronegative last, and gets the suffix “-ide”. Covalent bonding is very different from ionic naming Ionic names ignored the subscript because there was only one possible ratio of elements. Covalent gives several possibilities so we have to indicate how many of each atom is present in the name Prefixes you have to know prefix meaning prefix meaning *mono- 1 hex- 6 di- 2 hept- 7 tri- 3 oct- 8 tetr- 4 non- 9 pent- 5 dec- 10 * the first atom named does not get the prefix “mono-”, it just keeps its original name! Examples CO carbon monoxide CO2 carbon dioxide NI3 nitrogen triiodide P4O6 tetraphosphorus hexoxide Continuing I4O9 tetriodine nonoxide S2F10 disulfur decafluoride IF7 Iodine heptafluoride Si2Cl6 disilicon hexachloride Naming ionic compounds For monoatomic anions only drop the ending and add “-ide” so F fluoride Cl-, O2-, C4 chloride, oxide and carbide Continuing… cations keep the name of the element. When naming compounds always name the positive (cation) first and the negative (anion) last. so mixing ions of chlorine and sodium give you sodium chloride (positive) (negative) Determining the formula of ions Ionic compounds are neutral You need to find the lowest number of each ion to make it neutral for example: Na+ and O2 2 sodium for every one oxygen Na2O More examples Al3+ and O2- Al2O3 K+ and Cl- KCl the subscripts don’t effect the name if there is only one possibility still (cation)(anion) Aluminum oxide Potassium chloride Several atoms can form a couple of different ions. All of these are metals that are not in group 1, 2 or aluminum. for example iron can form Fe2+ or Fe3+ These are said as iron (II) and iron (III) Cu+ and Cu2+ is Copper (I) and Copper (II) Figuring out charge on these elements If the ion is named, the charge is in the name. If you have the formula, use the charges of the other ions present to determine the charge. Remember Alkali will always be +1 Alkaline Earth +2, Halogens -1, oxygen group -2 Aluminum will always be +3 Examples Copper (II) chloride Cobalt (III) sulfide NiF2 TiS2 Polyatomic ions Polyatomic ion You Ions- many atoms in one can NOT break these apart in this section. the “ide” suffix only applies to monoatomic anions Common polyatomic ions Ammonium NH4+ Perchlorate ClO4- Sulfate SO42- Acetate CH3CO2- Chlorate ClO3- Sulfite SO32- Nitrate NO3- Chlorite Nitrite NO2- Hypochlorite ClO- Dichromate Cr2O72- Cyanide CN- Iodate IO3- Chromate CrO42- Thiocyanate SCN- Bromate BrO3- Oxalate C2O42- Hydroxide OH- Silicate SiO32- HSO4- Permanganate MnO4- Phosphate PO43- HSO3- Thiosulfate S2O32- Arsenate AsO43- Hydrogen carbonate Hydrogen sulfate Hydrogen sulfite HCO3- ClO2- Carbonate CO32- YOU WILL HAVE TO MEMORIZE THESE! This is one of the big differences from last year. We will have a quiz just like the elements quiz last year over these! Determining the formula of ions Ionic compounds are neutral Remember– don’t break a polyatomic ion apart for example: Ammonium carbonate NH4+ and CO32 (NH4)2CO3 Chapter 4 acids and salts Oxyanions- negative ions containing oxygen. These have the suffix “-ate” or “-ite” “-ate” means it has more oxygen atoms bonded, “-ite” has less For example SO42- sulfate SO32- sulfite Oxyanions Oxyanions may contain the prefix “hypo-”, less than, or “per-”, more than. For example ClO4Perchlorate ClO3Chlorate ClO2Chlorite ClOHypochlorite Acids Certain compounds produce H+ ions in water, these are called acids. You can recognize them because the neutral compound starts with “H”. For example HCl, H2SO4, and HNO3. Don’t confuse a polyatomic ion with a neutral compound. HCO3- is hydrogen carbonate (or bicarbonate), not an acid. Naming acids Does it contain oxygen? If it does not, it gets the prefix “hydro-” and the suffix “-ic acid” HCl Hydrochloric acid HF Hydroflouric acid HCN Hydrocyanic acid Naming Acids If it does contain an oxyanion, then replace the ending. If the ending was “–ate”, add “-ic acid” If the ending was “–ite”, add “-ous acid” H2SO4 Sulfuric Acid H2SO3 Sulfurous Acid Examples HNO3 HI H3AsO3 Chapter 5 Complex ions Complex ion- transition metal ion with attached ligands Fe(CN)63 Glance over this chapter. Skip the problems, this material is out of the test. Ch 6 Organic Alkanes- straight chain hydrocarbons with all single bonds Alkenes- hydrocarbons with a double bond Alkynes- hydrocarbons with a triple bond Cyclic hydrocarbons- rings Root words Meth # of C atoms 1 Hex # of C atoms 6 Eth 2 Hept 7 Prop 3 Oct 8 But 4 Non 9 Pent 5 Dec 10 Name this molecule And give its molecular formula 4 ethyl octane C10H22 4 propyl decane C13H28 Name and give the formula Methyl cyclohexane C7H14 Functional groups halogenated *R means any carbon chain Alcohols R-OH -al -oic acid Ketones NOT at the edge R-C-R =O Aldehydes at the edge R=O R-C=O -OH -ol Carboxylic Acids -one Predicting organic reactions Addition reactions occur by adding halogens or hydrogen to alkene or alkynes. In the reaction, the new molecule takes the place of the double or triple bond. Cl2 + CH3-CH=CH2 CH3-CClHCClH2 example 1- butene is reacted with fluorine C4 H8 + F2 C4H8F2 Predicting organic reactions Substitution reactions occur by adding halogens to an alkane. In the reaction, the new molecule takes the place of a hydrogen. Cl2 + CH3-CH3 CH3-CClH2 + HCl Cl2 + C2H6 C2ClH5 + HCl Predicting organic reactions Combustion reactions occur when an organic compound is burned in oxygen. The products of a complete combustion are water vapor and carbon dioxide. C6H12O6 + 6 O2 6 H2O+ 6 CO2 Predicting organic reactions Esterification reactions Made by reacting carboxylic acids with alcohols. Carboxylic acid + H-O-R R-C-O-R alcohol O= O= R-C-O-H Ester + H-O-H Examples Fluorine Ethanol is added to 2 propene is burned in oxygen Chlorine is added to propane Ethanoic acid is reacted with 1- butanol Ch 7 Predicting molecular equations Synthesis Reactions pg 42 Metals and nonmetals combine to form salts. CHECK CHARGES OF IONS!!!!! Metal Oxides and water form bases, bases have OH attached Nonmetal oxides and water form acids Metal oxides and nonmetal oxides form salts (you will make an Examples Magnesium Calcium burns in oxygen reacts with chlorine gas Magnesiumvoxide reacts with water Carbon dioxide is bubbled through Lithium oxide is added to carbon water Ch 7 Predicting molecular equations Decomposition Reactions pg 42 CHECK CHARGES OF IONS!!!!! Carbonates decompose into oxides and CO2 Chlorates decompose into chlorides and O2 Ammonium and a base makes ammonia and water Hydrogen peroxide decomposes into water and oxygen. Three that always appear on the AP test Hydrogen peroxide decomposes into water and oxygen Ammonium hydroxide decomposes into ammonia and water Carbonic acid decomposes into carbon dioxide and water Examples Titanium Copper (IV) chlorate decomposes (III) carbonate is heated Magnesium Carbonic chloride is electrolyzed acid is heated Hydrogen peroxide decomposes Ch 8 Single replacement reactions A + BX AX + B You will have a chart of activity series More active metals will replace less active metals from their compound in a solution A less active element will have no reaction when added to a more active element! Active metals replace hydrogen in water Active metals replace hydrogen in acids Active nonmetals replace less active nonmetals from their compounds in solutions Activity series chart Old chart from test Zinc Examples is added to a solution of cobalt (II) chloride Cadmium is added to a solution of barium iodide Lithium is added to a solution of copper (II) chlorate Chromium is left in water Examples Potassium Silver is added to sulfuric acid is added to hydrochloric acid Chlorine gas is bubbled through a solution of sodium bromide Chapter 9 Double replacement reactions AY + BX AX + BY These reactions occur in solution Remember in solution the ions are free floating. For a reaction to occur, the ions have to come together and leave their dissolved state. Formation of a precipitate Formation of a gas Formation of a molecular species Solubility Rules Acids are soluble. Compounds of: alkali metals, ammonium, and nitrate are soluble. This needs to be memorized. Other soluble compounds All acetates are soluble except Fe3+ All chlorates are soluble. All binary compounds of the halogens (other than F) with metals are soluble, except those of Silver, Mercury(I), and Lead. Fluorides are insoluble except for rule 1 and 2. All sulfates are soluble, except those of barium, strontium, calcium, lead, silver, and mercury (I). Insoluble compounds Carbonates, oxalates, sulfites, chromates, oxides, silicates, and phosphates are insoluble. Hydroxides are insoluble except Ba, Sr, and Ca Sulfides are insoluble except for calcium, barium, strontium, magnesium. The exception is with alkali metals or ammonium. Net Ionic Equation Determine what dissolved and precipitated Mg(NO3)2(aq)+Na2CO3(aq) MgCO3(s) + 2NaNO3(aq) Dissociate everything that is aqueous, not solid Mg++ + 2NO3- + 2Na+ + CO3-– MgCO3(s) + 2 Na+ + 2NO3- Now cancel out everything that is the same on both sides of the equation These are called spectator ions The remaining part is the net Mg++ + CO3-- MgCO3(s) ionic equation Examples Hydrochloric nitrate acid reacts with silver Potassium carbonate reacts with calcium chlorate Sodium chloride reacts with ammonium oxalate Scandium chromate acetate reacts with lithium H2S Gases (hydrogen sulfide) is formed from any sulfide reacting with an acid CO2 (Carbon dioxide) is formed from any carbonate reacting with an acid, water is also produced SO2 (sulfur dioxide) is formed from any sulfite reacting with an acid, water is also produced NH3 (ammonia) is formed from ammonium reacting with a soluble hydroxide Examples Ammonium chloride reacts with calcium hydroxide Sodium sulfide is combined with nitric acid Ammonium carbonate is combined with barium chlorate Lithium acid sulfite reacts with phosphoric Formation of a molecular species It is the same as precipitates or gases except a liquid is formed. Acid base neutralization reactions will produce water. NaOH + HNO3 H2O (l) + NaNO3 (aq) Strong acids Acid formula Acid Hydrochloric HCl acid Hydrobromic HBr acid Hydriodic HI acid Formula Sulfuric Acid Nitric Acid H2SO4 Perchloric Acid HClO4 Chloric Acid HClO3 HNO3 Strong Bases these make a lightning bolt on the periodic table! Name Formula Name Formula Sodium NaOH Hydroxide Calcium Ca(OH)2 Hydroxide Potassium KOH Hydroxide Strontium Sr(OH)2 Hydroxide Barium Ba(OH)2 Hydroxide Strong acids and bases Strong acids and bases are not at equilibrium, there is no reverse reaction. Strong acids and bases will never be formed in a net ionic equation. All other acids/bases can be formed, and will be formed by reacting the appropriate ion with a strong acid/base. *Most other bases are insoluble Examples Calcium hydroxide reacts with chloric acid Hydrochloric acid reacts with calcium nitrite Nitric acid reacts with sodium chlorite Sodium chloride is mixed with sulfuric acid Chapter 10 Aqueous Solutions and Ionic Equations This chapter has already been covered *Only dissociate soluble ionic compounds Molecular equation Na2S + CrCl2 CrS + NaCl Full Ionic Equation 2Na++S2-+Cr2++2Cl- CrS +Na++Cl Net Ionic Equation Chapter 11 Redox Reactions Redox or oxidation-reduction reactions are reactions that involve a transfer of electrons. Oxidation is the loss of electrons. Reduction is the gain of electrons. (think of the charge, OIL RIG) 4 K + O2 → 4 K+ + 2 O2 Potassium get oxidized, oxygen get reduced Using oxidation states In the reaction… 2 Na +2 H2O →2 NaOH + H2 0 +1 -2 +1 -2 +1 0 Note the changes Sodium went from 0 to 1 2 of the hydrogen atoms went from +1 to 0 (the other two were unchanged) Breaking into two half reactions Sodium must have lost 2 electrons 2 Na → 2Na+ + 2 e And Hydrogen gained two electrons 2 H2O +2 e-→ 2 OH- + H2 Sodium is oxidized, hydrogen is reduced in this reaction Oxidation is an increase in oxidation state Reduction is a decrease in oxidation state Balancing Redox Equations by Half Reactions Method or oxidation state method The book does not separate these into half reactions, although it adds another step I think it makes it easier Half reactions Ce4+ + Sn2+ → Ce3+ + Sn4+ Half reactions Ce4+ + e- → Ce3+ Sn2+ → 2e- + Sn4+ Electrons lost must equal electrons gained! 2 Ce4+ +2 e- →2 Ce3+ Merge the two half reactions 2 Ce4+ + Sn2+ → 2 Ce3+ + Sn4+ Redox reactions in acidic solutions It will be noted in the problem Balance all elements except hydrogen and oxygen. Balance oxygen by adding H2O (which is always prevalent in an acidic solution) Balance hydrogen by adding H+ Then balance the charge adding electrons and proceed as normal. Example In an acidic solution Cr2O7 2- + Cl- → Cr3+ + Cl2 Half reactions Cr2O7 2- → Cr3+ Cl- → Cl2 Reduction side Cr2O7 2 Cr2O7 2 Cr2O7 2- Cr2O7 2 Cr2O7 2- H2O → Cr3+ → 2 Cr3+ → 2 Cr3+ + 7 H2O + 14 H+→ 2 Cr3+ + 7 H2O + 14 H++ 6 e- →2Cr3++7 Oxidation side Cl- → Cl2 2 Cl- → Cl2 2 Cl- → Cl2 + 2 e I have to equal 6 e- so multiply by 3 6 Cl- → 3 Cl2 + 6 e- Combine my half reactions Cr2O7 2- + 14 H++ 6 e- → 2 Cr3+ + 7 H2O 6 Cl- → 3 Cl2 + 6 e And you get Cr2O7 2-+14 H++6Cl-→2Cr3++3 Cl2+7H2O The electrons cancel out . Example In an acidic solution MnO4- + H2O2 → Mn2+ + O2 Balancing Redox Equations in a basic solution Look for the words basic or alkaline Follow all rules for an acidic solution. After you have completed the acidic reaction add OH- to each side to neutralize any H+. Combine OH- and H+ to make H2O. Cancel out any extra waters from both sides of the equation. Example We will use the same equation as before In a basic solution MnO4- + H2O2 → Mn2+ + O2 2 MnO4- + 6 H++ 5 H2O2 → 2 Mn2+ + 5 O2 + 8 H2O Another example In a basic solution MnO4 − + SO32-→MnO4 2− + SO42 Half reactions MnO4 − → MnO4 2− SO32-→ SO42- Predicting redox reactions Single replacement reactions we went over a few chapters ago are all redox reactions. Here are the same examples, this time write out the net ionic equation The change will be the charge!! Examples Zinc is added to a solution of cobalt (II) chloride Cadmium is added to a solution of barium iodide Lithium is added to a solution of copper (II) chlorate Chromium is left in water Examples Potassium Silver is added to sulfuric acid is added to hydrochloric acid Chlorine gas is bubbled through a solution of sodium bromide Common Oxidizing/reducing agents Oxidizing agent Product Reducing Agent Product MnO4- in acid Mn2+ H2O2 O2 MnO2 in acid Mn2+ MnO4- in base MnO2 Halogens (dilute basic) Hypohalite ion (hypochlorite) CrO42- in acid Cr3+ Cr2O72- in acid Cr3+ halogens (conc basic) Halate ion (chlorate) HClO4 Cl- Free metals Metal ions Na2O2 NaOH H2O2 H2O C2O42- CO2 H2SO4 conc. SO2 Sulfite or SO2 SO4 2- Free halogens Halide ion Halide ion Free halogen HNO3 conc. NO2 NO2- NO3- HNO3 dilute NO Metal ions Lower oxidation number Metal ion Higher oxidation number IMPORTANT For every redox equation, you have to have both an oxidizing agent and a reducing agent!! Otherwise it doesn’t work. Examples Bromate reacts with bromide in an acidic solution Permanganate reacts with oxalate in an acidic solution Calcium metal reacts with permanganate in a sodium hydroxide solution Example Chromate reacts with chloride in an acidic solution Chlorine reacts with permanganate in a concentrated solution of sodium hydroxide Ch 12 Electrolysis in water Electrolysis is a fairly simple process. There are two plates in a solution, and an electric current is sent through. The plates are the cathode, where reduction takes place, and the anode, where oxidation takes place. cathode-reduction anodeoxidation Rules for cathode reaction A cation may be reduced to a metal Cu+ + 1 e- Cu Or water way be reduce to hydrogen 2 H2O + 2 e- H2 + 2 OH Transition metals tend to reduce before water, main group metals tend to reduce after Rules for anode reactions An anion nonmetal may be oxidized to a nonmetal 2 Cl- Cl2+ 2 e Water may be oxidized to oxygen 2 H2O O2 + 4 H+ + 4 e Chlorine, bromine and iodine will oxidize before oxygen. That is it. Rules for molten binary salts Molten means melted, with no water. These are straightforward and easy! Molten magnesium chloride is electrolyzed MgCl2 Mg + Cl2 examples Aqueous calcium bromide is electrolyzed Aqueous chromium (III) nitrate is electrolyzed Aqueous cobalt (II) bromide electrolyzed Molten sodium chloride is electrolyzed Ch 13 Complex ion reactions Formation of complex ions Common complex ions metals Fe Co Ni Cr Cu Zn Ag Al Common ligands NH3 CN- OHSCN General rule: the number of ligands will be twice the charge of the metal ion Example Iron (III) chloride reacts with potassium cyanide Fe3+ + CN- Fe(CN)63 How did I get the charge? Iron is 3+ , 6 cyanides at 1- Examples Zinc (I) fluoride reacts with sodium thiocyanate Concentrated ammonia is reacted with cobalt (III) iodide Barium hydroxide reacts with nickel (II) nitrate Solution Stoichiometry Ch 4 Homework Due with test Pg 171 Chapter 4 review 1-87 odd Water and the Nature of Aqueous Solutions Water is the foundation of all life on Earth. Each O-H covalent bond is highly polar because the electronegativity of oxygen is so much greater than hydrogen. The dipole moment creates a slightly negative O and a slightly positive H as the electrons are pulled strongly toward O. Water is a bent molecule with o a bond angle of 109.5 109.5o Hydrogen bond Polar molecules are attracted to one another by dipole attractions. In water, it is hydrogen bonding. The molecules can’t easily slide past one another, they are held in place. The strength of the hydrogen bonds accounts for water’s high surface tension, its low vapor pressure, its high specific heat, its high heat of vaporization, and its high boiling point. Hydrogen bonding in water Solvents and Solutes Chemically pure water never exists in nature because water dissolves so many substances * An aqueous solution is one in which water samples contain dissolved substances. Components of a solution. A solvent is the dissolving medium. A solute is the dissolved particles. Solutions Solutions are homogeneous mixtures in which solute particles are usually less than 1 nm in diameter. The solute and solvent are not capable of being separated by filtration. Substances that dissolve most readily in water include ionic compounds and polar covalent molecules. Nonpolar molecules do not dissolve in water (a polar molecule). Rule: Like dissolves like The Solution Process Water molecules are in constant motion as a result of their kinetic energy. The molecules collide with solute particles. The solvent molecules attract the solute ions. The ionic crystal breaks apart by the action of the solvent. Solvation is the process that occurs when a solute dissolves. http://mw2.concord.org/public/student/solution/dissolve.cml Insoluble In some ionic compounds, the attractions between the ions in the crystal are stronger than the attractions exerted by water. * These compounds are insoluble. Nonpolar substances form a solution because there are no repulsive forces between them, not because the solute and solvent are attracted. Electrolytes and Nonelectrolytes Compounds that conduct an electric current in aqueous solution or the molten state are called electrolytes. All ionic compounds are electrolytes. Soluble ionic compounds conduct electricity in both a solution and in the molten state. Insoluble ionic compounds only conduct electricity in the molten state. Nonelectrolytes Compounds that do not conduct an electric current in either aqueous solution or the molten state are called nonelectrolytes. Many molecular compounds are nonelectrolytes because they do not contain ions. Most compounds of carbon are nonelectrolytes. Some very polar molecular compounds are nonelectrolytes in the pure state, but are electrolytes when they dissolve in water. This occurs because such compounds ionize in solution. Weak Electrolytes Not all electrolytes conduct electricity to the same degree. A weak electrolyte conducts electricity poorly because only a fraction of the solute exists as ions. * Most of the compound is in the original form. * The most common weak electrolytes are weak acids and weak bases. Strong Electrolytes A strong electrolyte conducts electricity very well because almost all of the solute exists as separated ions. * Very little of the original compound remains intact. * Classes of electrolytes include soluble salts, strong acids and strong bases. The Composition of Solutions To perform stoichiometric calculations when two solutions are mixed, two things must be known. The nature of the reaction, which depends on the exact forms the chemical takes when dissolved. The amounts of chemical present in solution, usually expressed as concentrations. Molarity Molarity (M) is defined as the moles of solute per volume of solution in liters. M = molarity = moles of solute liters of solution Calculate the molarity of a solution prepared by dissolving 11.5 g of solid NaOH in enough water to make 1.50 L of solution. Calculation of Molarity II Calculate the molarity of a solution prepared by dissolving 1.56 g of gaseous HCl in enough water to make 26.8 mL of solution. Concentrations of Ions Give the concentration of each type of ion in the following solutions: a. 0.50 M Co(NO3)2 b. 1 M Fe(ClO4)3 Concentrations of Ions II Calculate the number of moles of Clions in 1.75 L of 1.0 x 10-3M ZnCl2. Concentration and Volume I. Typical blood serum is about 0.14 M NaCl. What volume of blood contains 1.0 mg NaCl? Solutions of Known Concentration .To analyze the alcohol content of a certain wine, a chemist needs 1.00 L of an aqueous 0.200 M K2Cr2O7 (potassium dichromate) solution. How much solid K2Cr2O7 must be weighed out to make this solution? Dilution .To save money and space in a laboratory, solutions are often in concentrated form. Water is added to achieve the molarity desired for a particular solution. This process is called dilution. Dilution with water does not alter the numbers of moles of solute present. Concentration and Volume What volume of 16M sulfuric acid must be used to prepare 1.5 L of a 0.10 M H2SO4 solution? Another way to express the dilution process is: M1V1 = M2V2 What volume of 14.8M ammonia must be used to prepare .75 L of a 1.5 M H2SO4 solution? Stoichiometry of Precipitation Reactions 1. Identify the species present in the combined solution, and determine what reaction occurs. 2. Write the balanced net ionic equation for the reaction. 3. Calculate the moles of reactants. 4. Determine which reactant is limiting. 5. Calculate the moles of reactants or products, as required. 6. Convert to grams or other units, as required. Determining the Mass of Products Formed Calculate the mass of solid NaCl that must be added to 1.50 L of a 0.100 M AgNO3 solution to precipitate all the Ag+ ions in the form of AgCl. Stoichiometry of Precipitation Reactions 1. Identify the species present in the combined solution, and determine what reaction occurs. 2. Write the balanced net ionic equation for the reaction. 3. Calculate the moles of reactants. 4. Determine which reactant is limiting. 5. Calculate the moles of reactants or products, as required. 6. Convert to grams or other units, as required. Determining the Mass of Products Formed Calculate the mass of solid NaCl that must be added to 1.50 L of a 0.100 M AgNO3 solution to precipitate all the Ag+ ions in the form of AgCl. Acid-Base Reactions The Bronsted-Lowry Definitions: An acid is a proton donor. A base is a proton acceptor. Arrhenius definitions: Acid generate protons (hydronium) Bases generate hydroxide Lewis definitions Acids are electron pair acceptors Bases are electron pair donors Performing Calculations for AcidBase Reactions 1.List the species present in the combined solution before any reaction occurs, and decide what reaction will occur. 2. Write the balance net ionic equation for this reaction. 3. Calculate the moles of reactants. For reactions in solution, use the volumes of the original solutions and their molarities. 4. Determine the limiting reactant. 5. Calculate the moles of the required reactant and product. 6. Convert to grams or volume (of solution), as required. Example Problem What volume of a 0.100 M HCl solution is needed to neutralize 25.0 mL of 0.350 M NaOH? In a certain experiment, 28.0 mL of 0.250 M HNO3 and 53.0 mL of 0.320 M KOH are mixed. Calculate the amount of water formed in the resulting reaction. What is the concentration of H+ or OH- ions in excess after the reaction goes to completion? Acid-Base Titrations Volumetric analysis is a technique for determining the amount of a certain substance by doing a titration. A titration is a process in which a solution of known concentration is used to determine the concentration of another solution through a monitored reaction. The equivalence point is reached when all of the moles of H+ ions present in the original volume of acid solution have reacted with an equivalent number of moles of OH- ions added. This is also known as the stoichiometric point. Titration An acid-base indicator is a chemical that changes color once the equivalence point is reached. The endpoint of a titration is the point in a titration where the indicator actually changes color. Three requirements for a titration 1. The exact reaction between titrant (known solution) and analyte (unknown solution or one you are analyzing) must be known (and rapid). Three requirement (cont.) 2. The stoichiometric (equivalence) point must be marked accurately. 3. The volume of titrant required to reach the stoichiometric point must be known accurately. Problem solving The first step in the analysis of a complex solution is to write down the components and focus on the chemistry of each one. Take a big problem and look at the small problems within. Neutralization Titration A student carries out an experiment to standardize (determine the exact concentration of) a sodium hydroxide solution. To do this, the student weighs out a 1.3009-g sample of potassium hydrogen phthalate (KHC8H4O4, often abbreviated KHP. KHP (molar mass 204.22 g/mol) has one acidic hydrogen. The student dissolves the KHP in distilled water, adds phenolphthalein as an indicator, and titrates the resulting solution with the sodium hydroxide solution to the phenolphthalein endpoint. The difference between the final and initial buret readings indicates that 41.20 mL of the sodium hydroxide solution is required to react exactly with the 1.3009 g KHP. Calculate the concentration of the sodium hydroxide solution. Neutralization Analysis An environmental chemist analyzed the effluent (the released waste material) from an industrial process known to produce the compounds carbon tetrachloride (CCl4) and benzoic acid (HC7H5O2), a weak acid that has one acidic hydrogen atom per molecule. A sample of this effluent weighing 0.3518 g was shaken with water, and the resulting aqueous solution required 10.59 mL of 0.1546 M NaOH for neutralization. Calculate the mass percent of HC7H5O2 in the original sample.