Acids and Bases

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Acids and Bases. Some Properties of Acids. Produce H + (H 3 O + ) ions in water. The Hydronium Ion (H 3 O + ) is an H + (proton) attached to a water molecule. Taste Sour. React with certain metals to produce H 2 gas and a salt.
Transcript
Acids and Bases Some Properties of Acids
  • Produce H+ (H3O+) ions in water.
  • The Hydronium Ion (H3O+) is an H+ (proton) attached to a water molecule.
  • Taste Sour.
  • React with certain metals to produce H2 gas and a salt.
  • Salt – ionic-metal or a positive polyatomic ion bonded with a negative ion other than OH-
  • Example: MgCl2, NH4Cl
  • Aqueous solutions of acids conduct electricity.
  • Electrolytes – the greater the concentration of ions in solution, the greater the electrical conductivity.
  • Strong Acids & weak Acids
  • Some Properties of Acids
  • React with carbonates and bicarbonates to produce carbon dioxide gas.
  • HCl(aq) + NaHCO3(aq) NaCl(aq) + H2O(l) + CO2(g)
  • React with bases to form a salt and water.
  • Neutralization Reaction (Double Replacement)
  • HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
  • pH is less than 7
  • pH scale expresses the amount of H+ as a number from 0 to 14
  • pH of 0 is strongly acidic and has the highest amount of H+ ions, pH of 7 is neutral, pH of 14 is strongly basic and has the fewest H+ ions.
  • “ABC easy as 123”
  • Cause acid-base indicators to change color.
  • Acids turn Blue litmus Red
  • Acid Nomenclature Acid Nomenclature
  • Use the flowchart to name the following acids.
  • HBrHydrobromic Acid
  • H2CO3 Carbonic Acid
  • H2SO3 Sulfurous Acid
  • Going Backwards…
  • Write H first.
  • Write the 2nd ion. (you may have to check table E)
  • Assign charges.
  • Criss Cross, if necessary.
  • Examples
  • Sulfuric Acid _____ H2SO4__________
  • Nitrous Acid _______HNO2 __________
  • Oxalic Acid ______ H2C2O4_________
  • Name ‘Em
  • HI(aq)
  • HCl(aq)
  • H2SO3
  • HNO3
  • HClO4
  • Some Properties of Bases
  • Produce OH- ions in water.
  • Taste Bitter, chalky.
  • Aqueous solutions of bases conduct electricity.
  • Electrolytes – the greater the concentration of ions in solution, the greater the electrical conductivity.
  • Strong bases & weak bases
  • Feel soapy, slippery.
  • This is because they break down the normal body fat in your hands or whatever part of your body they come into contact with.
  • NaOH
  • Before After Some Properties of Bases
  • React with acids to form a salt and water.
  • Neutralization Reaction (Double Replacement)
  • HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
  • pH is greater than 7
  • pH scale expresses the amount of H+ as a number from 0 to 14
  • pH of 0 is strongly acidic and has the fewest OH- ions, pH of 7 is neutral, pH of 14 is strongly basic and has the greatest amount of OH- ions.
  • “ABC is easy as 123”
  • Cause acid-base indicators to change color.
  • Bases turn Red litmus Blue
  • Naming Bases
  • Name the Metal.
  • If the metal has only one possible charge, just write it’s name.
  • If the metal has more than one possible charge, use Roman Numerals to indicate the charge.
  • Follow with Hydroxide
  • Examples:
  • LiOHLithium Hydroxide
  • Fe(OH)3Iron(III) Hydroxide
  • Going Backwards…
  • Write the symbol of the metal.
  • Write OH-
  • Assign Charges.
  • Criss Cross, if necessary
  • Examples
  • Cesium Hydroxide CsOH
  • Chromium(III) Hydroxide Cr(OH)3
  • Strontium Hydroxide SrOH
  • Practice Name each of the following…
  • HBr
  • H2SO3
  • H2C2O4
  • HClO
  • Ca(OH)2
  • AgOH
  • HgOH
  • HF
  • HI
  • HClO4
  • HCl
  • LiOH
  • Sn(OH)2
  • Ti(OH)3
  • Practice Now go backwards….
  • nitric acid
  • carbonic acid
  • dichromic acid
  • acetic acid
  • nitrous acid
  • potassium hydroxide
  • cesium hydroxide
  • barium(II) hydroxide
  • aluminum(III) hydroxide
  • strontium(III) hydroxide
  • Explaining Acids and Bases
  • There have been several attempts to explain the properties of acids and bases.
  • These explanations define how acids and bases behave.
  • There are three such definitions.
  • Arrhenius Theory
  • Brønsted – Lowry
  • Lewis Acids & Bases
  • Arrhenius Theory
  • Acids – produce H+ ions (or hydronium ions H3O+) as the onlypositive ion. HCl(l) Cl- + H+
  • A substance with a carboxyl group(COOH) looks like a base when you look at the chemical formula but it is an acid. (Acetic Acid = HC2H3O2 = CH3COOH) CH3COOH + H2O  CH3COO- + H+
  • Arrhenius Theory
  • Bases – produce OH- ions (or hydroxide ions).
  • Some bases DO NOT have hydroxide ions attached.
  • Amines – organic compounds containing C and N. Amines are bases even though they do not have an hydroxide ion. Instead they react with water to produce the OH-ion. NH3 + H2O  NH4+ + OH- ~Caution~
  • Alcohols – contain an –OH but ARE NOTbases.
  • Example: CH3OH (hydroxyl group on a carbon chain)
  • Types of Acids and Bases
  • Acids
  • Monoprotic: produce one H+ ion.
  • HCl
  • Diprotic: produce two H+ ions.
  • H2SO4
  • Triprotic: produce three H+ ions.
  • H3PO4
  • Bases
  • Monohydroxy: produce one OH- ion.
  • NaOH
  • Dihydroxy: produce two OH- ions.
  • Ba(OH)2
  • Trihydroxy: produce three OH- ions.
  • Al(OH)3
  • Strength of Acids and Bases Determined by the amount of Ionization.
  • Strong Acids
  • 100% dissociation in water.
  • Great conductors of electricity. HNO3(aq) + H2O(l) H3O+(aq) + NO3-(aq)
  • HI, HCl, HBr, H2SO4, and HClO4 are strong acids.
  • Weak Acids
  • Much less than 100% dissociation.
  • Poor conductors of electricity. CH3COOH(aq) + H2O(l) CH3COO-(aq)+ H3O+(aq)
  • Acetic Acid(CH3COOH)
  • Strength of Acids and Bases Determined by the amount of Ionization.
  • Strong Bases
  • 100% dissociation (ionization) in water.
  • Great conductors of electricity. NaOH(aq) Na+(aq) + OH-(aq)
  • KOH, Ca(OH)2, Group 1 or 2 metals with hydroxide!!
  • Weak Bases
  • Much less than 100% dissociation (ionization).
  • Poor conductors of electricity. NH3(aq) + H2O(l) NH4+(aq)+ OH-(aq)
  • Ammonia (NH3)
  • Brønsted-Lowry Acids and Bases
  • Acids – Proton Donors
  • According to the Brønsted-Lowy concept, an acid is the chemical species that donates the proton in a proton transfer reaction.
  • Bases – Proton Acceptors
  • According to the Brønsted-Lowy concept, a base is the chemical species that accepts the proton in a proton transfer reaction. A “proton” is really just a hydrogen that has lost its electron…H+
  • Conjugate Pairs
  • The Brønsted-Lowry concept defines a species as an acid or a base according to its function in the proton-transfer reaction.
  • Consider the Reaction of NH3 + H2O
  • In the forward reaction, NH3 accepts a proton donated by H2O. Thus, NH3 is a base and H2O is an acid.
  • Conjugate Pairs
  • In the reverse reaction, NH4+ donates a proton to OH- which accepts it. Thus, NH4+ is acid and OH- is the base.
  • Conjugate Pairs
  • The species NH4+ and NH3 are a conjugate acid-base pair.
  • A conjugate acid-base pair consists of two species in an acid-base reaction, one acid and one base, that differ by the loss or gain of a proton.
  • NH4+ is the conjugate acid of NH3
  • NH3 is the conjugate base of NH4+
  • The species OH- and H2O are a conjugate-acid base pair as well.
  • OH- is the conjugate base of H2O
  • H2Ois the conjugate acid of OH-
  • Conjugate Pairs Conjugate Pairs…Practice Problems
  • Label the Acid, Base, Conjugate Acid, Conjugate Base in each reaction.
  • Strength of Acid-Base Conjugate Pairs
  • Strong Acids (Proton Donors) have weak conjugate bases.
  • Strong Bases (Proton Acceptors) have weak conjugate acids.
  • Strong acids and strong bases are always on the same side of an equation.
  • An acid can donate it H+ to any base EXCEPT it’s conjugate base.
  • Example: H3PO4 can donate to F-, but not to PO43-
  • Practice Problems
  • Write the conjugate base for each.
  • HCl ______________
  • H2CrO4 ___________
  • NH4+ _____________
  • NH3 ______________
  • Write the conjugate acid for each.
  • F- _________________
  • H2PO4- ___________
  • NH3 ______________
  • HSO4- ____________
  • Practice Problems
  • CH3COO- + H30+ CH3COOH + H2O
  • HCl + H2O  H3O+ + Cl-
  • NH2- + H2O  NH3 + OH-
  • H3O + OH-  H2O +H2O
  • CN- + H2O  HCN + OH-
  • HClO4 + CH3COOH  ClO4- + CH3COOH2+
  • HCN + H2O  H3O+ + CN-
  • Practice Problems
  • HSO4- + HCl H2SO4 +Cl-
  • SO42- + HNO3  HSO4- + NO3-
  • NH4+ + HSO4-  NH3 + H2SO4
  • HCl + Al(H2O)5(OH)2+  Cl- + Al(H2O)63+
  • NH3 + NH3  NH4+ + NH2-
  • Amphoteric (Amphiprotic) Species
  • Substances that act as both an acid or a base.
  • Depends on chemical environment.
  • Examples: H2O, HSO4-, HS-
  • In the reaction between NH3 and H2O, water is an acid.
  • In the reaction between HNO2 and H2O, water is a base.
  • Water (H2O) is an amphoteric substance.
  • Lewis Acids and Bases
  • Lewis Acid – a substance that ACCEPTS an electron(e-) pair.
  • Lewis Base – a substance that DONATES an electron(e-) pair.
  • Formation of the Hydronium Ion is an excellent example. Lewis Acid/Base Reaction Reactions Involving Acids
  • Steps…
  • Check the metal on Table J. If it is above H2 proceed.
  • Write H2 as a product.
  • Combine the metal with the negative (-) ion to form an ionic salt. (write the metal first, followed by the negative ion.)
  • Assign charges
  • Criss Cross, if necessary
  • Balance the equation.
  • Reactions Involving Acids, Examples…
  • HCl + Sr __________ + __________
  • H3PO4 + Zn  __________ + __________
  • HNO3 + Au  __________ + __________
  • HC2H3O2 + K  __________ + __________
  • HF + Cu  __________ + __________
  • Neutralization Acid + Base  Salt + H2O
  • Steps…
  • Form water, H2O.
  • Get rid of all H+ on the acid, and all OH- on the base.
  • Write the metal 1st and the negative(-) ion 2nd.
  • Assign Charges.
  • CrissCross, if necessary.
  • Balance the equation.
  • Neutralization…Examples
  • CH3COOH + NaOH __________ + __________
  • KOH + HCl  __________ + __________
  • HCl + NaOH  __________ + __________
  • H2SO4 + NaOH  __________ + __________
  • HCl + Ba(OH)2  __________ + __________
  • HNO3 + LiOH  __________ + __________
  • Spectator Ions
  • Ions found on both sides of the equation that are not involved in making water.
  • Not part of the Net Arrhenius Equation.
  • Spectator Ions for the previous example…
  • CH3COO-, Na+
  • Net Arrhenius Equation
  • Does NOT include spectator ions!!
  • Net Arrhenius Equation is always… H+ + OH- H2O
  • The pH Scale
  • The pH Scale expresses the strength of acids and bases.
  • Logarithmic Scale – one jump on the scale represents a tenfold change in [H+]
  • [ ] = concentration (usually Molarity)
  • pH > 7 is a Base
  • pH = 7 is Neutral
  • pH < 7 is an Acid
  • As pH ↑ [H+] ↓
  • The stronger the acid the more H+ ions it produces.
  • The stronger the base the more OH- ions it produces.
  • pH Examples
  • A solution with a pH of 1 has how many times the amount of H+ compared to a solution with a ph of 6?
  • A solution with a pH of 2 has how many times the amount of H+ compared to a solution with a pH of 5?
  • If the [H+] increases, the [OH-] decreases by the same amount.
  • If the pH changes from 8 to 13, the [H+] decreases _______ times and the [OH-] increases _______ times.
  • If the pH changes from 6 to 2, the [H+] increases _______ times and the [OH-] decreases _______ times.
  • Using Table M Using Table M
  • Used to approximate the pH of a substance.
  • Indicators tend to have a distinct color at the ends of their useful pH range.
  • So a solution with methyl orange indicator is red at a pH of 0.
  • As a base is added pH begins to rise.
  • At a pH of about 3.2 the red color begins to change yellow but you first enter a intermediate color region that is a mixture of these two colors.
  • At a pH of about 4.4 the intermediate color region ends and the solution is now completely yellow all the way through to pH of 14.
  • Using Table M…Questions
  • What color will bromcresol green be in an acid with a pH of 2.2?
  • What color will bromcresol green be in a base with a pH of 10.0?
  • Can bromcresol green determine whether an unknown substance is an acid or base? Why or why not?
  • What is the pH range for a substance that turns methyl orange yellow and bromothymol blue yellow?
  • Titration Titration
  • Process of adding measured volumes of an acid or a base of known concentration to an acid or base of unknown concentration until neutralization occurs.
  • Data gathered from a titration experiment is used to calculate the concentration (Molarity) of a solution.
  • Knowing the volumes of the acid and the base used in the titration, together with the known concentration of the standard solution, it is possible to calculate the concentration of the unknown solution.
  • Titration Formula #MAVA = #MBVB # = the number of H+ or OH- produced per mole of acid or base. M = Molarity V = Volume SubscriptsA or B = Acid or Base respectively. Example: For H2CO3 you would plug in 2 for the # on the acid side. Titration…Practice Problem 1
  • 35.6 mL of NaOH is neutralized with 25.2 mL of 0.0998 M HCl by titration to an equivalence point. What is the concentration of the NaOH? #MAVA = #MBVB (1)(0.0998 M)(25.2 mL) = (1)MB(35.6 mL) MB = 0.071 M
  • Titration…Practice Problem 1
  • 60.3 mL of Ca(OH)2 is neutralized with 33.2 mL of 1.0 M HF by titration to an equivalence point. What is the concentration of the Ca(OH)2? #MAVA = #MBVB (1)(1.0 M)(33.2 mL) = (2)MB(60.3 mL) MB = 0.275 M
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