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Cool Chemistry Show! Unit 2: Chapter 6 Chapter Vocabulary 1. Physical Change 2. Solution 3. Homogeneous 4. Solute 5. Solvent 6. Chemical Change 7. Product 8. Reactant 9. Chemical Reaction 10.Precipitate 11.Concentration 12.Molarity 13.Mole 14.Saturated Solution 15.Supersaturated Solution 16.Polymer 17.Chemical Test 18.Acid-Base Indicator 19.Rate of Reaction 20.Surface Area 21.Catalyst 22.Endothermic Change/Rxn 23.Exothermic Change/Rxn 24.Activation Energy 25.Kinetic Energy 26.Law of Conservation of Energy 27.Gravitational Potential Energy 28.Heat 29.Temperature Chapter Vocabulary 30.Compound 31.Ion 32.Anion 33.Cation 34.Ionic Compound 35.Polyatomic Ion 36.Molecular Compound 37.Covalent Bond 38.Oxidation Number 39.Synthesis Reaction 40.Decomposition Reaction 41.Single-Replacement Reaction 42.Double Replacement Reaction 43.Salts 44.Acid 45.Base 46.Neutralization 47.Titration 48.Endpoint 49.Buffer 50.pH 51.Oxidation 52.Reduction 53.Redox Reaction Chemical and Physical Changes Act 1 Properties of Matter Properties • Words that describe matter (adjectives) • Physical Properties- a property that can be observed and measured without changing the composition. • Examples: color, hardness, m.p., b.p. • Chemical Properties- a property that can only be observed by changing the composition of the material. • Ex: Iron rusting Classification of Matter Matter Mixtures Pure Substance Mixtures • Combination of 2 or more pure substances • Each substance retains its individual chemical properties • Composition variable • Most everyday matter Mixture examples • Sand and water • Salt and water Classification of Matter Matter Mixtures Homogeneous Pure Substance Heterogeneous Mixtures • Heterogeneous: • Mixture is not uniform in composition • EX: Chocolate chip cookie, gravel, soil. • Homogeneous: • Same composition throughout; called “solutions” • Ex: Kool-aid, air, salt water Solutions • Homogeneous mixture • Mixed molecule by molecule • Can occur between any state of matter: • gas in gas • liquid in gas • gas in liquid • solid in liquid • solid in solid (alloys), etc. Solutions: Parts • Solute: • Substance being dissolved • Salt, sugar • Solvent: • Substance that does the dissolving • Water, alcohol Saturated & Supersaturated Solutions • Concentration: ratio of solute to solvent • 1.5M (Molarity) : number of moles of solute dissolved in 1 liter of solution • Dilute: less solute • Concentrated: more solute • Saturated: no more solute will dissolve • Supersaturated: contains more solute than normal Pure Substance • Same fixed composition and properties • Ex: sand, water, oxygen, sodium 1. Elements 2. Compounds Classification of Matter Matter Mixtures Homogenous Pure Substance Heterogeneous Elements Compounds Elements • • • • • Simplest kind of matter All one kind of atom. On Earth 91 occur naturally Unique name and symbol Symbol 1, 2 or 3 letters • 1st: always capitalized • 2nd and 3rd: lower case Compounds • Made of two or more atoms, chemically combined • Can be broken down only by chemical methods • When broken down, the pieces have completely different properties than the original compound. • NaCl: Table salt • Na: Flammable metal Cl: poisonous gas • Formula: combination of chem symbols to show what makes up the compound Matter Mixtures Heterogeneous Dirt, blood, milk Physical Changes Pure Substances Homogenous Elements Compounds Lemonade, gas, steel Oxygen, gold, iron Salt baking soda, sugar Chemical Changes Compound or Mixture Compound Mixture Made of one kind of material Made of more than one kind of material Made by a chemical change Made by a physical change Definite composition Variable composition Which is it? Mixture Element Compound Physical Changes • A change that changes appearances, without changing the composition. • Ex. Boil, melt, cut, bend, split, crack Chemical Changes • A change where a new form of matter is formed. • Ex. Rust, burn, decompose, ferment • Reactants: starting materials • Products: ending materials Evidence of a Chemical Reaction • Temp change • Release: light/heat • Absorb • • • • Color change Odor change Production of gas bubbles Formation of a precipitate (solid) Polymers • Giant molecules of repeating groups of atoms (monomers) • Linked through intermolecular forces • Can be molded into useful objects Polymers • Cross-linking • Short bridges b/w long polymer chains • Gives polymer new properties monomer polymer Cross-linking Polymers • Polyester: clothing • Polyethylene: blow-molded beverage bottles, auto gas tanks, and pipe • Epoxy adhesives Biological Polymers • Carbohydrates: sugars • Lipids: fatty acids • Proteins: amino acids • DNA/RNA: nucleotides Check your own Portfolio • Chapter Vocabulary • Act 1 • • • • • What do you think? Summary Checking up Q’s What do you think now? Chemical Essential Q’s Act 1 Quiz! 1. Is the property below chemical or physical ? Mercury is a silvery liquid. 2. Which of the following is NOT evidence of a chemical change? a) formation of a precipitate b) formation of a gas c) melting of a solid d) color change 3. The materials forming a solution are called the ____ and _____. 4. When no more solute will dissolve in the solvent the solution is __________. a) dilute b) concentrated c) saturated d) supersaturated 5. For each determine if the material is an element, compound, heterogeneous or homogeneous. a) pepperoni pizza b) gold c) steel d) soda pop Ion Name Ion NH4+ ammonium PO43- phosphate NO2- nitrite CO32- Carbonate NO3- nitrate SO32- Sulfite OH- hydroxide SO42- Sulfate CN- cyanide BrO3- bromate MnO4- permanganate IO3- Name iodate More Chemical Changes Act 2 Chemical Test for Gases • Chemical test: procedure or chemical rxn used to ID a substance • Oxygen • Glowing splint; burst into flames http://www.youtube.com/watch?v=NGEc6wMedDE • Hydrogen • Burning splint; loud pop http://www.youtube.com/watch?v=7s2dfXcoOyI&feature=related • Carbon Dioxide • Limewater: Ca(OH)2 Forms a ppt http://www.youtube.com/watch?v=21aBA0jUJXU Indicators for Acids and Bases • Indicators used to determine acids an bases • Substances that change color when react with acid or base • Phenol red • Turns yellow in presence of an acid • Litmus Paper Check your own Portfolio • Chapter Vocabulary • Act 1 • • • • • What do you think? Summary Checking up Q’s What do you think now? Chemical Essential Q’s • Act 2 • • • • • What do you think? Summary Checking up Q’s What do you think now? Chemical Essential Q’s Act 2 Quiz! Act 2 Quiz 1. Which of the following is a test for carbon dioxide? a) Burning splint causes a pop into flame c) Fmn of white ppt in limewater b) glowing splint burst d) global warming 2. What is the acid-base indicator used in lab? a) Thymol Blue c) Methyl yellow b) Methyl red d) Phenol red 3. The indicator in lab turned (color) of (acid or base) . 4. Write the formula. Sodium sulfate 5. Write the name. Ra(NO3)2 in the presence 1. Which of the following is NOT a physical process? a) Sharpening a pencil c) An explosion of dynamite b) melting ice d) popping a balloon 2. Which of the following is evidence of a chemical reaction? a) Formation of gas b) disappearance of liquid c) Formation of smaller pieces d) formation of crystals 3. Which of the following is a test for carbon dioxide? a) Burning splint causes a pop into flame c) Fmn of white ppt in limewater b) glowing splint burst d) global warming 4. What is the acid-base indicator used in lab? a) Thymol Blue c) Methyl yellow b) Methyl red d) Phenol red 5. The indicator in lab turned (color) of (acid or base) . in the presence Investigate Water Baking soda Baking Powder Antiacid Tablet Heated Vinegar Ammonia No More Intense No bubbles bubbles bubbles bubbles CO2 Slow Intense Intense No bubbles bubbles bubbles bubbles CO2 CO2 Fast Fast Intense No bubbles bubbles bubbles bubbles CO2 CO2 Chemical Names and formulas Act 3 Ions • Form by losing or gaining electrons • Anions • Typically nonmetals • Gain electrons • Negatively charged • Cations • Typically metals • Lose electrons • Positively charged Ionic compound • Bond b/w cation and anion • Electrostatic force • Conducts electricity when liquid or dissolved in water Naming ions and ionic compounds • Name cation first and the anion second • Monatomic cations use the element name • Monatomic anions take their name from the root of the element name plus the suffix –ide • Oxidation numbers of transition metals are written as roman numerals in parentheses • If the compound contains a polyatomic ion, simply name to ion • Examples Examples • Write the formula for the pairs of elements • • • • Barium and oxygen Strontium and iodine Radium and chlorine Lithium and chlorine • Write the formula for each : • • • • Lithium oxide Calcium bromide Sodium oxide Aluminum sulfide Examples • Write the formula for each : • • • • Ammonium and sulfite ions (SO32-) Barium and nitrate ions (NO3-) Magnesium hydroxide (OH-) Sodium phosphate (PO43-) Examples • Write the formula for each : • • • • Copper (I) sulfide Tin (IV) fluoride Gold (III) cyanide Lead (II) sulfide • Write the names: • Mn2O3 • HgF2 Molecular Compounds • • • • Covalent bonds Sharing electrons Nonmetals only Typically not conductive Molecular Nomenclature Prefixes PREFIX monoditritetrapentahexaheptaoctanonadeca- NUMBER 1 2 3 4 5 6 7 8 9 10 Molecular Nomenclature: Examples • CCl4 • carbon tetrachloride • N2O • dinitrogen monoxide • SF6 • sulfur hexafluoride More Molecular Examples • arsenic trichloride • AsCl3 • dinitrogen pentoxide • N2O5 • tetraphosphorus decoxide • P4O10 Diatomic Molecules • • • • • • • • • H I Br O Cl F N Hi Bronclif! aka... HI BrONClF Check your own Portfolio • Chapter Vocabulary • Act 1 • • • • • What do you think? Summary Checking up Q’s What do you think now? Chemical Essential Q’s • Act 2 • • • • • What do you think? Summary Checking up Q’s What do you think now? Chemical Essential Q’s • Act 3 • • • • • What do you think? Summary Checking up Q’s What do you think now? Chemical Essential Q’s Act 3 Quiz! (1pt each) Name Formula Cation Anion Sodium Nitrate Br- MgBr2 Li + O2- K2S Calcium hydroxide Sr2+ Strontium bromate Fe2(CO3)3 Potassium sulfate (NH4)3PO4 Fe3+ CO32- Act 3 Quiz! (1.5pt each) Name Formula Cation Br- MgBr2 Li + K2S Calcium nitride Iron III carbonate Fe2(CO3)3 Ammonium Phosphate (NH4)3PO4 Anion NH4+ O2- Chemical Equations Act 4 Representing Chemical Reactions • Have two parts: • Reactants - the substances you start with • Products- the substances you end up with • Reactants Products In a chemical reaction • The way atoms are joined is changed • Atoms aren’t created or destroyed. • Can be described several ways: 1. Word equation Copper + chlorine copper (II) chloride 2. Skeleton equation Cu(s) + Cl2(g) CuCl2(l) Skeleton Equation • Uses formulas and symbols to describe a reaction • All chemical equations are sentences that describe reactions. Symbols in equations • Arrow separates the reactants from the products • Read “reacts to form” • The plus sign = “and” • (s) after the formula = solid • (g) after the formula = gas • (l) after the formula = liquid Symbols used in equations • (aq) after the formula dissolved in water, an aqueous solution Convert these to equations • Solid iron (III) sulfide reacts with gaseous hydrogen chloride to form iron (III) chloride and hydrogen sulfide gas. • Nitric acid (HNO3) dissolved in water reacts with solid sodium carbonate to form liquid water and carbon dioxide gas and sodium nitrate dissolved in Now, read these: • Fe(s) + O2(g) Fe2O3(s) • Cu(s) + AgNO3(aq) Ag(s) + Cu(NO3)2(aq) Pt • NO2 (g) N2(g) + O2(g) Balanced Equation • Atoms can’t be created or destroyed • All the atoms we start with we must end up with • A balanced equation has the same number of each element on both sides of the equation. C + O O O C • C + O2 CO2 • This equation is already balanced • What if it isn’t? O C + O O C O • C + O2 CO • We need one more oxygen in the products. • Can’t change the formula, because it describes what it is (carbon monoxide in this example) C + O O C O C O • Must be used to make another CO • But where did the other C come from? C + C O O C O C O • Must have started with two C • 2 C + O2 2 CO Rules for balancing: Assemble, write the correct formulas for all the reactants and products Count the number of atoms of each type appearing on both sides Balance the elements one at a time by adding coefficients (the numbers in front) - save H and O until LAST! Check to make sure it is balanced. • Never change a subscript to balance an equation. • If you change the formula you are describing a different reaction. • H2O is a different compound than H2O2 • Never put a coefficient in the middle of a formula • 2 NaCl is okay, Na2Cl is not. Example H2 + O2 H2O Make a table to keep track of where you are at Example H2 + O2 H2O R P 2 H 2 2 O 1 Need twice as much O in the product Example H2 + O2 R P 2 H 2 2 O 1 Changes the O 2 H2O Example H2 + O2 2 H2O R P 2 H 2 2 O 1 2 Also changes the H Example H2 + O2 2 H2O R P 2 H 2 4 2 O 1 2 Need twice as much H in the reactant Example 2 H2 + O2 2 H2O R P 2 H 2 4 2 O 1 2 Recount Example 2 H2 + O2 2 H2O R P 4 2 H 2 4 2 O 1 2 The equation is balanced, has the same number of each kind of atom on both sides Example 2 H2 + O2 2 H2O R P 4 2 H 2 4 2 O 1 2 This is the answer Not this Balancing Examples • _AgNO3 + _Cu _Cu(NO3)2 + _Ag • _Mg + _N2 _Mg3N2 • _P + _O2 _P4O10 • _Na + _H2O _H2 + _NaOH • _CH4 + _O2 _CO2 + _H2O Balancing Examples • 2AgNO3 + Cu Cu(NO3)2 + 2Ag • 3Mg + N2 Mg3N2 • 4P + 5O2 P4O10 • 2Na + 2H2O H2 + 2NaOH • 2Na + 2H2O H2 + 2NaOH • CH4 + 2O2 CO2 + 2H2O Act 4 Quiz Part I • Write the skeleton equation for: 1. Solid potassium reacts with iron (II) sulfate dissolved in water to form potassium sulfate dissolved in water and solid iron. 2. Solid phosphorous and oxygen gas reacts to form solid diphosphorus trioxide. • Balance the following equations: 3. Li + AlCl3 LiCl + Al 4. Ca + Au(NO3)3 Ca(NO3)2 + Au • Write a sentence for the equation (remember roman numerals with transition metals) 5. AgNO3(aq) + BaCl2(aq) Ba(NO3)2 (aq) + AgCl Types of Chemical Reactions • OBJECTIVES: • Identify a reaction as combination, decomposition, single-replacement, doublereplacement, or combustion • Predict the products of combination, decomposition, single-replacement, doublereplacement, and combustion reactions. Types of Reactions • • • • • There are millions of reactions. Can’t remember them all Fall into several categories. We will learn 5 major types. Will be able to predict the products. • For some, we will be able to predict whether they will happen at all. #1 - Combination Reactions (Synthesis) • Combine - put together • 2 substances combine to make one compound. • Ca +O2 CaO • SO3 + H2O H2SO4 • We can predict the products if they are two elements. • Mg + N2 #1 - Combination Reactions (Synthesis) • Combine - put together • 2 substances combine to make one compound. • Ca +O2 CaO • SO3 + H2O H2SO4 • We can predict the products if they are two elements. • 3Mg + N2 Mg3N2 Write and balance • Ca + Cl2 • Fe + O2 iron (II) oxide • Al + O2 • Remember that the first step is to write the correct formulas • Then balance by using coefficients only #2 - Combustion • Means “add oxygen” • A compound composed of only C, H, and maybe O is reacted with oxygen • If the combustion is complete, the products will be CO2 and H2O. • If the combustion is incomplete, the products will be CO (possibly just C) and H2O. Examples • C4H10 + O2 • C6H12O6 + O2 Examples • 2C4H10 + 13O2 8CO2 + 10H2O • C6H12O6 + 6O2 6CO2 + 6H2O #3 - Decomposition Reactions • decompose = fall apart • one reactant falls apart into two or more elements or compounds. electricity • NaCl Na + Cl2 • CaCO3 CaO + CO2 • Note that energy is usually required to decompose #3 - Decomposition Reactions • Can predict the products if it is a binary compound • Made up of only two elements • Falls apart into its elements electricity • H2O • HgO #3 - Decomposition Reactions • Can predict the products if it is a binary compound • Made up of only two elements • Falls apart into its elements • 2H2O electricity 2H2 + O2 • 2HgO 2Hg + O2 #4 - Single Replacement • One element replaces another • Reactants must be an element and a compound. • Products will be a different element and a different compound • Na + KCl K + NaCl • F2 + LiCl LiF + Cl2 #4 Single Replacement • Metals replace other metals (and hydrogen) • K + AlN • Zn + HCl • Think of water as HOH • Metals replace one of the H, combine with hydroxide. • Na + HOH #4 Single Replacement • Metals replace other metals (and hydrogen) • 3K + AlN K3N + Al • Zn + 2HCl ZnCl2 + H2 • Think of water as HOH • Metals replace one of the H, combine with hydroxide. • Na + HOH NaOH + H2 #4 Single Replacement • We can tell whether a reaction will happen • Some chemicals are more “active” than others • More active replaces less active • There is a list on page 462 called the Activity Series of Metals • Higher on the list replaces lower. #4 Single Replacement • Note the * concerning Hydrogen • H can be replaced in acids by everything higher • Li, K, Ba, Ca, & Na replace H from acids and water Challenge! • • • • Fe + CuSO4 MgCl2 (aq) + Zn (s) Ag + Cu(NO3)2 F2 + NaBr • Br2 + NaF #4 Single Replacement • • • • Fe + CuSO4 FeSO4 + Cu MgCl2 + Zn no Ag + Cu(NO3)2 no F2 + 2NaBr 2NaF + Br2 • Br2 + 2NaF no #5 - Double Replacement • Two things replace each other. • Reactants must be two ionic compounds or acids. • Usually in aqueous solution • NaOH + FeCl3 • The positive ions change place. • NaOH + FeCl3 Fe+3 OH- + Na+1 Cl-1 • NaOH + FeCl3 Fe(OH)3 + NaCl #5 - Double Replacement • Has certain “driving forces” • Will only happen if one of the products: • doesn’t dissolve in water and forms a solid (a “precipitate”), or • is a gas that bubbles out, or • is a covalent compound (usually water). Solubility rules 1. All salts of alkali metals and ammonium are soluble 2. All chlorides, bromides and iodides are soluble with silver, lead, and mercury 3. All nitrates are soluble 4. All sulfates are soluble except with calcium, barium, strontium, and lead 5. All carbonates, phosphates, hydroxides, and sulfides are insoluble except with alkali metal or ammonium. Complete and balance • Assume all of the following reactions take place: • CaCl2 + NaOH • CuCl2 + K2S • KOH + Fe(NO3)3 • (NH4)2SO4 + BaF2 Complete and balance • Assume all of the following reactions take place: • CaCl2 + 2NaOH 2NaCl + Ca(OH)2 • CuCl2 + K2S 2KCl + CuS • 3KOH + Fe(NO3)3 Fe(OH)3 + 3K(NO3) • (NH4)2SO4 + BaF2 BaSO4 + 2NH4 F Act 4 Quiz Part 2 • For each equation a. b. c. d. 1. 2. 3. 4. 5. ID (1pt each equation) Predict the products (2pt each equation) Balance (1pt each equation) Write 2 of the 5 as sentences (2pt extra credit) K + B2O3 Al + S8 C10H22 + O2 H3PO4 + Ca(OH)2 Mg3P2 Act 4 ReQuiz Part 2 • For each equation a. b. c. d. 1. 2. 3. 4. 5. ID (1pt each equation) Predict the products (2pt each equation) Balance (1pt each equation) Write 2 of the 5 as sentences (2pt extra credit) Na + O2 Ca + Al2S3 C10H22 + O2 KCl LiNO3 + Mg(OH)2 Chemical Energy Act 5 Endothermic and Exothermic • Endothermic • Process that absorbs heat energy • Internal energy of system increases • Example from lab? • Cold pack made with NH4NO3 • 2H2O + energy 2H2 + O2 Endothermic and Exothermic • Exothermic: • Process that releases heat energy • Internal energy of system decreases • Example from lab? • Hot pack made with Na2CO3 • 2H2 + O2 2H2O + energy Importance of energy • In order for chem rxn to take place, particles from the reactants must collide • Particles must have enough kinetic energy to break existing bonds • Activation energy: minimum amount of energy required for a chem rxn to start • Bond breaking • Endothermic; add energy • Bond making • Exothermic; release energy Exothermic • Bond forming > bond breaking • Products have less energy than the reactants • ∆H: heat of reaction • Energy difference b/w reactants and products Endothermic • Bond forming < bond breaking • Reactants have less energy than the products • Ea: activation energy Conservation of Energy • Law of conservation of Energy • Energy absorbed from or released to surroundings is equal to the change in the energy system • Total energy in a closed system remains the same. Conservation of Energy • Law of conservation of Energy • Energy is not created or destroyed, only transferred • Some types of energy: • Light • Heat • Sound • Nuclear • Kinetic • Chemical Heat and Temperature • Related but not the same • Heat is a type of energy; can be transferred • Temperature related to kinetic energy of the molecules; what is measured 1. Act 5 Quiz The minimum energy required for a chemical reaction to begin is _________. 2. __________ is the measurement of the kinetic energy of particles. 3. What is the Law of conservation of Energy? 4. The process of breaking bonds is ________ and the process of making bonds is ________. 5. Give an example of an endothermic reaction from lab. EC. Fill in the missing slots AND ID the graph as endothermic or exothermic A D C B Reaction Rates Act 6 Factors Affecting Rate of Reaction • Rate of reaction • How fast a rxn takes place • Decrease of concentration of reactants over time • Increase of concentration of products over time • Factors (Increase particles colliding) • • • • Surface area Concentration of reactants Temperature Catalyst Factors Affecting Rate of Reaction • Surface Area • When the surface area of reactants increases, the reaction rate increases. • • • • • • • Chemical reactions occur when reactants collide at the surface of other reactants. Smaller particle size Greater surface area Greater chance for collisions to occur Chemical reaction will proceed faster Lab example? Antacid tablets: crushed vs. whole Factors Affecting Rate of Reaction • Concentration • When reactants are more concentrated, the rate of a chemical reaction will increase. • There is a greater chance that reactant particles will collide when they are more concentrated. • More collisions mean a faster reaction rate. • Lab example? • Vinegar: 3 different concentrations • HCl: 2 different concentrations Factors Affecting Rate of Reaction Temperature • • Temperature increases, the rate of a chemical reaction increases. Average kinetic energy of molecules of reactants increases as temperatures increase. More reactant particles with enough energy produces more successful collisions The reaction will proceed faster Lab example? Tea and antacid tablets • Hot vs. Cold water Factors Affecting Rate of Reaction • Catalyst • The presence of a catalyst will speed up a chemical reaction. • • • • • • A catalyst lowers the amount of energy needed to start a reaction (activation energy). Energy needed for successful collisions is less There will be more successful collisions The chemical reaction will proceed faster Lab example? MnO2 and H2O2 Act 6 Quiz 1. List the four factors that influence the rate of a reaction. 2. For each reaction, predict it will proceed faster (F) or slower (S) as a result. a. The reactants are made into smaller pieces. b. The reactants are diluted. c. The reactants are heated. d. The reactants are kept in large pieces. 3. Pick one condition from above and explain why it effects the reaction like it does. 4. A catalyst speeds up a reaction by ______. 5. Pick one of the 4 factors. Give the lab example and explain why the rate increased. Acids and Bases Act 7 Red and Blue Litmus Paper Phenolthalien Magnesium • Reacts with acid Bromothymol Blue Acidic Neutral Basic Some Definitions • Arrhenius • Acid: when dissolved in water, increases the concentration of hydrogen ions. • Base: when dissolved in water, increases the concentration of hydroxide ions. Some Definitions • Brønsted–Lowry • Acid: Proton donor • Base: Proton acceptor What Happens When an Acid Dissolves in Water? • Water acts as a Brønsted–Lowry base and pulls a proton (H+) from the acid. • As a result, the conjugate base of the acid and a hydronium ion are formed. Conjugate Acids and Bases: • From the Latin word conjugare, meaning “to join together.” • Reactions between acids and bases always yield their conjugate bases and acids. Properties of Acids • Reacts with metals to produce hydrogen gas • The effect of acids on indicators (from lab) • Neutralization of bases • Sour taste • Have a pH less than 7 Acid and Base Strength • Strong acids are completely dissociated in water. • Their conjugate bases are quite weak. • Weak acids only dissociate partially in water. • Their conjugate bases are weak bases. Strong Acids • Recognize and write ionization equations for the following five strong acids • HI HI (aq) H+ + I• HBr • HCl • HNO3 • H2SO4 Properties of Bases • • • • Bases are electrolytes The effect of bases on indicators Neutralization of acids Water solutions of bases taste bitter and feel slippery • Have a pH greater than 7 Acid and Base Strength • Substances with negligible acidity do not dissociate in water. • Their conjugate bases are exceedingly strong. Strong Bases • Recognize and write ionization equations for the following four strong bases • NaOH Na+ + OH• Ca(OH)2 • Ba(OH)2 • KOH Neutralization • When acids and bases react hydrogen ion and hydroxide ion form water • Remaining ions form a salt • HCl(aq) + NaOH(aq) HOH(l)+ NaCl(aq) • Resulting solution is neutral • Buffer: • A solution that resist change in pH when acid or base is added. pH • pH is defined as the negative base-10 logarithm (ten fold) of the hydrogen ion concentration. pH = −log [H+] pOH • pOH is defined as the negative base-10 logarithm (ten fold) of the hydroxide ion concentration. pOH = −log [OH-] Act 7 Quiz 1. Label as an acid or base and write the ionization reactions and balance for each a) Ca(OH)2 b)H2S c)HCl d)LiOH 2. Write the equation and ID the acid, base, conjugate acid and conjugate base. H2SO4(aq) + H2O(aq) H3O+ + HSO4-(aq) 3. What does pH stand for? 4. When acids and bases react ______ and a ______ are formed. 5. Phenolthalein is color in the presence of an acid. 6. Red litmus paper is color in the presence of a base. Color Reactions that Involve the Transfer of Electrons Act 8 Redox reactions • Oxidation • When substance loses electrons • Lab example: Iron lost electrons to form Fe2+ that dissolve in solution • Reduction • When substances gains electrons • Lab example: Cu2+ gained electrons from the iron to form copper atoms that plate out as solid • Redox rxn • Process of oxidation and reduction together • Remember! • Lose Electrons Oxidation; Gain Electrons Reduction Redox Reactions • Rusting • Water and oxygen corrode iron metal to rust • Iron loses electrons to b/c Fe3+ • O2 gain electrons to b/c O2• Prevention by protective covering, i.e. painting Lab 1. Did a rxn take place? How do you know? • • • Zn + CuSO4 What type of rxn? What is oxidized? Why? What is reduced? Why? 2. Did a rxn take place? How do you know? • • • Al + CuCl2 What type of rxn? What is oxidized? Reduced? Lab 3. Did a rxn take place? How do you know? • • • Mg + CuCl2 What type of rxn? What is oxidized? Reduced? 4. Did a rxn take place? How do you know? • • • Cu + Al(NO3)2 What type of rxn? What is oxidized? Reduced? Act 8 Quiz 1. Na + Cu(NO3)2 a. What type of rxn (of the 5 types)? b. Predict the products. c. What is oxidized? Why? What is reduced? Why? 2. What is needed for iron to rust? What can be done to protect iron from rusting?