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Acids/Bases/Salts

Acids/Bases/Salts. Properties . Common Acids. Lactic sour milk Acetic vinegar Phosphoric tart taste in soda Citric citrus fruits Malic apples Tartaric grapes Formic ant bites. Common Base. Ammonia window cleaner Sodium hydroxide lye (drain and oven cleaners)

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Acids/Bases/Salts

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  1. Acids/Bases/Salts Properties

  2. Common Acids Lactic sour milk Acetic vinegar Phosphoric tart taste in soda Citric citrus fruits Malic apples Tartaric grapes Formic ant bites

  3. Common Base Ammonia window cleaner Sodium hydroxide lye (drain and oven cleaners) Sodium Bicarbonate Baking soda antacid Milk of magnesia antacid

  4. electrolytes Properties BASES ACIDS • electrolytes • bitter taste • sour taste • turn litmus red • turn litmus blue • react with metals to form H2 gas • slippery feel • ammonia, lye, antacid, baking soda • vinegar, soda, apples, citrus fruits • pH more than 7 • pH less than 7

  5. Acid nomenclature –naming acids • Binary acids – is an acid that contains only two different elements. Acids are composed of hydrogen (H+) followed by an anion (negative ion). • Oxyacids- is an acid that is a compound of hydrogen, oxygen and third element (usually a non-metal) • If the acid formula contains oxygen in the anion, such as in H2SO4, it is known as an oxyacid.

  6. 3 Rules To Naming Acids • If H + anion ending in –ide: Acid name is “hydro_____ic acid” • Take the root from the anion name and fill in the blank.

  7. Acid Naming Example • Example: HCl • Cl is the anion, its name is chloride • Name of acid is: hydrochloric acid • Example: HF • F is the anion, its name is fluoride • Name of acid is: hydrofluoric acid

  8. 3 Rules To Naming Acids • H + anion ending in –ate: Acid name is “_____ic acid” • Take the root from the anion name and fill in the blank. • “What I ATE was ICky”

  9. Acid Naming Example • Example: HNO3 • NO3 1- is the anion, its name is nitrate • Name of acid is: nitric acid • Example: H2CO3 • CO32- is the anion, its name is carbonate • Name of acid is: carbonic acid

  10. Exceptions • Sulfate (SO42-) • Root is not sulf, but sulfur • Sulfuric acid • Phosphate (PO43-) • Root is not phosph, but phosphor • Phosphoric acid

  11. 3 Rules To Naming Acids • H + anion ending in –ite: Acid name is “_____ous acid” • Take the root from the anion name and fill in the blank. • “Don’t bITE; it’s infectiOUS”

  12. Acid Naming Example • Example: HNO2 • NO21- is the anion, its name is nitrite • Name of acid is: nitrous acid • Example: HClO2 • ClO21- is the anion, its name is chlorite • Name of acid is: chlorous acid

  13. Strength of Acids • Strong Acid - is one that ionizes completely in aqueous solution. - completely dissociates. - a strong acid is a strong electrolyte. - increases with increasing polarity and decreasing bond energy.

  14. Strength of Acids • Weak acids - are weak electrolytes - it contains hydronium ions, anions and dissolved acid molecules - Organic acids (acidic carboxyl group) -COOH

  15. Aqueous solutions for base • Most bases are ionic compounds containing metal cations and the hydroxide anion. • When a base completely dissociates in water to yield aqueous OH- ions, the solution is called alkaline. NaOH (s)water Na+(aq) + OH-(aq)

  16. Aqueous solutions for base • Not all bases are ionic compounds. • Ammonia is one example because it produces hydroxide ions when it reacts with water molecules. NH3 (g) + H20 (l) NH4+(aq) + OH- (aq)

  17. Strength of bases • The strength of the base depends on the extent to which it dissociates to its ions. • Strong bases are strong electrolytes • Bases that are not very soluble do not produce a large number of hydroxide ions when added to water

  18. Base Naming Example • NaOH • name of base: sodium hydroxide • Mg(OH)2 • name of base: magnesium hydroxide • Fe(OH)2 • name of base: iron (II) hydroxide

  19. pH and pH scale 0 14 7 INCREASING ACIDITY INCREASING BASICITY NEUTRAL Whether or not a solution is acidic, basic, or neutral depends on the balance of H+ and OH- ions: Neutral: [H+] = [OH-] Acid: [H+] > [OH-] Base: [H+] < [OH-]

  20. pH and pOH Calculations • pH is the negative base 10 logarithm of the hydrogen ion concentration: pH = - log10 [H+] • pH is the expression of the acidity or alkalinity of a solution in terms of its hydronium ion concentration. pH = - log [H3O+] • The sum of the pH and the pOH always equals 14. • pH + pOH = 14

  21. Example • Calculate the pH, if the [H3O] = 2.4 X 10-6 M pH = - log [H3O+] = - log(2.4 X 10-6) = -(-5.6) = 5.6 • Find the pH, the pOH = 5.3 pH + pOH = 14 pH = 14 - 5.3 pH = 8.7

  22. Use the reverse of the equation to calculate the [H3O+] when pH is known. [H3O+] = antilog (-pH) = 10-pH pH calculations • Use an identical equation to calculate pOH. pOH = - log [OH-] • Calculate the [H3O+], if the pH is 4.71. • [H3O+] = antilog (-pH) • = antilog (-4.71) • = 1.95 X 10-5 M

  23. Acid- Base Theories

  24. Acid-Bases Theories Arrhenius Acids and Bases Arrhenius acid – is a chemical compound that increases the concentration of hydrogen ions H+, in an aqueous solution. Arrhenius base – is a substance that increases the concentration of hydroxide ions OH-, in aqueous solution.

  25. Acid-Bases Theories Bronsted-Lowry acid -Is a molecule or an ion that is a proton donor (H+) - Donates proton to water HCl + H2O H3O+ + Cl- HCl + NH3 NH4+ + Cl – In both reaction HCl is a Bronsted-Lowry acid

  26. Bronsted-Lowry Acid Water could be a Bronsted-Lowry acid also. H2O (l) + NH3 (g) NH4+ (aq)+ OH-(aq) The waterdonates a hydrogen ion (proton) to the ammonia molecule.

  27. Bronsted-Lowry Base A molecule or ion that is a proton acceptor (H+) H2O (l) + NH3 (g) NH4+ (aq)+ OH-(aq) Ammonium (NH4) is a Bronsted-Lowry Base because it accepts the proton from the water.

  28. Bronsted-Lowry Acid-Base Reaction • Protons are transferred from one reactant (the acid) to another (base). H2O (l) + NH3 (g) NH4+ (aq)+ OH-(aq

  29. Bronsted-Lowry Acid-Base • Monoprotic Acids – an acid that can donate only one proton (HCl, HNO3) • Polyprotic Acids – an acid that can donate more than one proton per molecule. - Diprotic – can donate two protons per molecule (Sulfuric acid, H2SO4) - Triprotic – can donate three protons per molecule(Phosphoric acid, H3PO4)

  30. Acid –Base Reaction The Bronsted-Lowry definition s of acids and bases provide the basis for studying proton (H+) transfer reaction. Suppose a Bronsted-Lowry acid gives up a proton, the remaining ion or molecule can re-accept that proton and can act as a base- a conjugate base.

  31. Acid –Base Reaction • Conjugate base- the ion or molecule that remains after a Bronsted-Lowry acid has given up a proton is the conjugate base of that acid. • Conjugate acid- the ion or molecule that is formed when a Bronsted-Lowry base gains a proton is the conjugate acid of that base.

  32. Acid –Base Reaction HF + H2O F- + H3O+ AcidBaseConjugateBase Conjugate Acid The species remaining after a Brønsted-Lowry acid gives up its proton is the conjugate base of that acid: Take off one H from the acid. The species remaining after a Brønsted-Lowry base accepts its proton is the conjugate acid of that base: Take off one H. Add an H to the base.

  33. Acid –Base Reaction HCO3- (aq) + H2O (l)H2CO3(aq) + OH-(aq) baseacid conjugate acidconjugate base (proton acceptor) HF (aq) + H2O (l)F-(aq) + H3O+ (aq) acidbaseconjugatebase conjugate acid (proton donor)

  34. Strength of conjugate acids and bases • The stronger an acid is, the weaker its conjugate base; the stronger a base is, the weaker its conjugate acid. • The weaker an acid is, the stronger its conjugate base; the weaker a base is, the stronger its conjugate acid.

  35. Lewis Acids and Base • A Lewis acid is an atom, ion, or a molecule that accepts an electron pair to form a covalent bond. • A Lewis base is an atom, ion or a molecule that donates an electron pair to form a covalent bond.

  36. Lewis Acids and Base H+ (aq) + : NH3(aq) [H-NH3+ ] (aq) H+ (aq) + : NH3(aq) [NH4+ ] (aq) A bare proton (H+) is a Lewis acid in reactions in which it forms a covalent bond.

  37. Lewis Acids and Base . . BF3(aq) + : F : -(aq) BF4– (aq) . . An anion is a Lewis Base in a reaction in which it forms a covalent bond by donating an electron pair.

  38. Indicators • Chemical dyes that change color as pH changes. • Different indicators change colors at different pH levels • choose an indicator that will show a color change at the pH that you are interested in. • Indicators can be on a strip of paper • called pH or litmus paper • Other indicators can be added to the solution directly. • Some indicators change color more than once and can be added to solutions so that we can see what is happening over time.

  39. Acid/Bases/SaltsNeutralization/Titrations

  40. Neutralization • Chemical reaction between an acid and a base. • Products are a salt (ionic compound) and water.

  41. Neutralization ACID + BASE  SALT + WATER HCl + NaOH  NaCl + H2O strong strong neutral HC2H3O2 + NaOH  NaC2H3O2 + H2O weak strong basic • Salts can be neutral, acidic, or basic. • Neutralization does not mean pH = 7.

  42. Neutralization +1 -1 • HCl + NaOH  • H2SO4 + KOH  • HNO3 + Ca(OH)2 NaCl + H2O Hydrocholoric acid Sodium chloride Water Sodium hydroxide +1 -2 2 K2SO4 + H2O 2 Potassium hydroxide Potassium sulfate Sulfuric acid Water +2 2 -1 Ca(NO3)2 + H2O 2 Calcium nitrate water Calcium hydroxide Nitric acid

  43. standard solution unknown solution Titration • Titration • Analytical method in which a standard solution is used to determine the concentration of an unknown solution.

  44. Titration • End Point – • point at which an indicator changes color during a titration • Equivalence point • Point at which equal amounts of H3O+ and OH- have been added. • when mole ratio exactly equals mole ratio required by reaction • Determined by… • indicator color change • dramatic change in pH

  45. Titration moles H3O+ = moles OH- MVn = MVn M: Molarity V: volume n: # of H+ ions in the acid or OH- ions in the base

  46. Titration • 42.5 mL of 1.3M KOH are required to neutralize 50.0 mL of H2SO4. Find the molarity of H2SO4. H3O+ M = ? V = 50.0 mL n = 2 OH- M = 1.3M V = 42.5 mL n = 1 MV# = MV# M(50.0mL)(2) =(1.3M)(42.5mL)(1) M = 0.55M H2SO4

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