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Acids and Bases

Acids and Bases. pH Scale. pH scale – numbered from 0 to 14, measures acidity and alkalinity (how acidic or basic a solution is). Neutral. Strengths of Acids and Bases. The lower the pH the STRONGER the acid More hydrogen ions (H + )in solution The higher the pH the STRONGER the base

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Acids and Bases

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  1. Acids and Bases

  2. pH Scale • pH scale – numbered from 0 to 14, measures acidity and alkalinity (how acidic or basic a solution is) Neutral

  3. Strengths of Acids and Bases • The lower the pH the STRONGER the acid • More hydrogen ions (H+)in solution • The higher the pH the STRONGER the base • More hydroxide ions (OH-) in solution

  4. I. Acid-Base Models A) Arrhenius Acids and Bases • Acids – substances contain hydrogen and form H+ ions in aqueous solution Example - HCl H+ (aq) + Cl- (aq)

  5. 2. Bases – substances contain hydroxide (OH–) ions and form OH– ions in aqueous solution NaOH Na+ (aq) + OH– (aq)

  6. Problem with this model – some acids and bases don’t fit Arrhenius definitions Example: Is NH3 an acid or a base? • It is not likely that NH3 will lose a H+, but it will produce OH– ions in water, like this: NH3 + H2O NH4+1 (aq) + OH–1 (aq)

  7. B) Brønsted-Lowry Acids and Bases 1. Acids – hydrogen ion (proton) donors 2. Bases – hydrogen ion (proton) acceptors; do not necessarily contain OH– NH3 + H2O NH4+1 (aq) + OH–1 (aq)

  8. II. Auto-Ionization of Water In pure water, 2 out of every billion molecules transfer a proton and break apart into ions according to this reaction: HOH  H+ + OH-

  9. B) Ion-product constant for water (Kw) Kw = [H+] ∙ [OH-] = 1.0 x 10-14 M2 • this value is a constant that can be used to determine acidity or alkalinity (basic) of a solution • In neutral solutions, [H+] = [OH-] • in acidic solutions, [H+] > [OH-] • in basic solutions, [OH-] > [H+]

  10. Ion Concentration Examples [H+] [OH-] = 1 x 10-14 M2 1. What are the ion concentrations of a 0.000453M solution of HCl? is the solution an acid or base? 2. What are the ion concentrations of a 0.00250M solution of KOH? Acid or base?

  11. pH CalculationspH stands for power of Hydrogen • pH = – log[H+] Example: What is the pH of a solution with a hydrogen ion concentration of 3.5 x 10-9 M? pH = – log [H+] pH = -log[3.5 x 10-9 ] pH = 8.46

  12. How do I get back from pH to ion concentrations? [H+] = antilog (-pH) What’s antilog? 2nd function of log key (10x)

  13. [H+] = antilog (–pH) Example: Determine the concentration of hydrogen ions in a solution with a pH of 3.5 [H+] = antilog (–3.5) [H+] = 10(–3.5) [H+] = 3.16 x 10-4 M

  14. What is redox? • Any chemical process in which elements experience a change in charge is an oxidation-reduction reaction (redox reaction) • Redox reactions involve a transfer of electrons. • Oxidation – involves losing electrons • Reduction – involves gaining electrons • LEO the lion says GER

  15. Rules for Assigning Charges • All unbonded elements have a charge of 0 • Oxygen is always a -2 (for this class) • Fluorine is always a -1 • Hydrogen is always +1 (for this class) • Alkali metals are always +1 • Alkaline earth metals are always +2 • Aluminum is always +3 • Neutral compounds have a charge of 0 • The sum of the charges of the elements in a polyatomic ion equals the charge of the polyatomic ion.

  16. Example • Identify if the following equations are redox reactions or not. For each redox reaction, determine which element is oxidized and which element is reduced. • Ca(OH)2 + 2 HCl→ CaCl2+ 2 H2O • CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g) • 2 Al(s) + 3 CuCl2(aq) → 2 AlCl3(aq) + 3 Cu(s)

  17. Neutralization • When a strong acid reacts with a strong base to form a salt and water. • Example: 2 NaOH + H2SO4 Na2SO4 + 2 H2O

  18. Precipitation Reactions • Precipitate – solid that is produced when two solutions are mixed together • These reactions are double-replacement reactions (switch dance partners) that form solids • Example: • 2 KI (aq) + Pb(NO3)2 (aq)  PbI2 (s) + 2 KNO3 (aq)

  19. Titration Problems • Titration – using a solution of known concentration to determine the concentration of an unknown solution • Equivalence Point – the point when there are equal amounts of moles of the acid and base • Indicator – substance that changes color when the reaction is complete, end point.

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