Acids and Bases Chapter 19. Acids pH less than 7 Sour taste Conduct electricity Reacts with metals...
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Transcript of Acids and Bases Chapter 19. Acids pH less than 7 Sour taste Conduct electricity Reacts with metals...
![Page 1: Acids and Bases Chapter 19. Acids pH less than 7 Sour taste Conduct electricity Reacts with metals to produce hydrogen gas Higher [H + ] concentration.](https://reader036.fdocuments.net/reader036/viewer/2022062407/56649d305503460f94a0811a/html5/thumbnails/1.jpg)
Acids and Bases
Chapter 19
![Page 2: Acids and Bases Chapter 19. Acids pH less than 7 Sour taste Conduct electricity Reacts with metals to produce hydrogen gas Higher [H + ] concentration.](https://reader036.fdocuments.net/reader036/viewer/2022062407/56649d305503460f94a0811a/html5/thumbnails/2.jpg)
Acids
• pH less than 7• Sour taste• Conduct electricity• Reacts with metals to
produce hydrogen gas• Higher [H+]
concentration
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Brønsted-Lowry AcidSome Common Acids
Name Formula
Hydrochloric acid HCl
Nitric acid HNO3
Sulfuric acid H2SO4
Phosphoric acid H3PO4
Ethanoic acid CH3COOH
Carbonic acid H2CO3
• Compounds that donate protons (H+)
• Polyprotic compounds
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Acids
•a HYDRONIUM ION is the ion formed when a water molecule gains a hydrogen ion
•H3O+
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To summarize...
•An acid is a chemical that dissolves in water to create more H+ ions than there are in neutral water
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Bases
• pH greater than 7
• Bitter taste• Slippery feel• Higher [OH-]
concentration
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Brønsted-Lowry Bases
•compounds that accept protons (H+)
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Some common basesSome Common Bases
Name Formula Solubility in Water
Sodium hydroxide NaOH High
Potassium hydroxide KOH High
Calcium hydroxide Ca(OH)2 Very low
Magnesium hydroxide Mg(OH)2 Very low
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to summarize...
•A base is a chemical that dissolves in water to create fewer H+ ions than there are in neutral water, or more OH- ions
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Bronsted-Lowry Theory
•Ammonia gains a proton, so it is a base
•Water donates a proton, so it is an acid
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Conjugate Acid•the ion or molecule formed
when a base gains a hydrogen ion.
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Conjugate Base
•the ion or molecule that remains after an acid loses a hydrogen ion.
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Conjugate acids and bases
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Some Conjugate Acid-Base Pairs
Acid Base
HCl Cl–
H2SO4 HSO4–
H3O+ H2O
HSO4– SO4
2–
CH3COOH CH3COO–
H2CO3 HCO3−
HCO3– CO3
2–
NH4+ NH3
H2O OH–
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Amphoteric Substances
•A substance is AMPHOTERIC if it can act as either an acid or a base, such as water
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Identify each reactant as a Bronsted-Lowry acid (hydrogen-ion donor) or base (hydrogen-ion acceptor).
1.HNO3 + H2O H3O+ + NO3-
2.CH3COOH + H2O ↔ H3O+ + CH3COO-
3.NH3 + H2O ↔ NH4+ + OH-
4.H2O + CH3COO- ↔ CH3COOH + OH-
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Identify the conjugate acid-base pairs of each reaction.
1.HNO3 + H2O H3O+ + NO3-
2.CH3COOH + H2O ↔ H3O+ + CH3COO-
3.NH3 + H2O ↔ NH4+ + OH-
4.H2O + CH3COO- ↔ CH3COOH + OH-
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Acid/Base Reactions• Acids and bases react to produce
salt and water in a neutralization reaction.
H2SO4 + 2NaOH Na2SO4 + 2H2O acid base salt water
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Self-ionization of water
•the reaction in which water molecules produce ions
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Self-ionization of water
• In pure water at 25°C, the concentration of hydrogen ions is only 1 × 10−7M.
• The concentration of OH− is also 1 × 10−7M because the numbers of H+ and OH− ions are equal in pure water.
• Any aqueous solution in which [H+] and [OH−] are equal is a neutral solution.
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Ion-product constant for water•The product of the concentrations
of the hydrogen ions and the hydroxide ions in water is called the ion-product constant for water (Kw).
Kw = [H+] × [OH−] = 1.0 × 10−14
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in acidic solutions...
•When acids dissolve in water hydrogen ions are released:
•HCl(aq) → H+(aq) + Cl−(aq)
•The H+ concentration is greater than the OH- concentration.
•A solution in which [H+] is greater than [OH−] is an acidic solution.
–The [H+] is greater than 1 × 10−7M.
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in basic solutions...
•When bases dissolves in water, it forms hydroxide ions in solution.
•NaOH(aq) → Na+(aq) + OH−(aq)
•The H+ concentration is less than the OH- concentration.
•A basic solution is one in which [H+] is less than [OH−].
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let’s practice...
•If the [H+] in a solution is 1.0 × 10−5M, is the solution acidic, basic, or neutral? What is the [OH−] of this solution?
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The pH concept•The negative logarithm of the
hydrogen-ion concentration.
pH = -log [H+]
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Calculating pH from [H+]
IF given in scientific notation in this form:
1.0 x 10 –x (x being any negative number)
then the exponent on the 10 is the pH value.
Example:1.0 x 10-3 M pH= 3
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ExampleWhat are the pH values of the following three solutions, based on their hydrogen ion concentrations?
1.[H+]= 1.0 x 10-5 M2.[H+]= 1.0 x 10-9 M3.[H+]= .001 M
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ExampleWhat are the pH values of the following three solutions, based on their hydrogen ion concentrations?
1.[H+]= 1.0 x 10-5 M pH =52.[H+]= 1.0 x 10-9 M pH = 93.[H+]= .001 M pH= 3
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Calculating pH from [H+]
If the coefficient is not 1.0 then do the math.
Example:What is the pH of a solution with a hydrogen-ion concentration of 4.2 x10-
10 M?
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Calculating [H+] from pH
IF the pH value is an integer (i.e. 1, 2, 3, 4, 5…) then the [H+]= 1.0 x 10 -1,-2,-3,-4,-5…
ExamplepH = 4 [H+] = 1.0 x 10-4
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If NOT, then math[H+] = antilog(-pH)
In calculator, the antilog function: Second, log
Example: The pH of an unknown solution is 6.35. What is hydrogen-ion [H+] concentration?
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Calculating pH from [OH-]
Use the ion-product constant for water to solve for [H+] and then find pH.
Kw = [H+] × [OH−] = 1.0 × 10−14
Example: What is the pH of a solution if [OH-]= 4.0 x 10 -11 ?
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Strong/Weak Acids
Strong acids will completely ionize in aqueous solutions.
Weak acids will not ionize completely in aqueous solutions CH3COOH (aq) + H20 (l) ↔ H3O+ (aq) + CH3OO- (aq)
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Strong/Weak Acids/Bases
Strong bases dissociates completely into metal ions and hydroxide ions in aqueous solutions
Weak bases reacts with water to form the conjugate acid of the base and hydroxide ions
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Substance Formula
Hydrochloric acidNitric acidSulfuric acid
HClHNO3
H2SO4
Phosphoric acid H3PO4
Ethanoic acid CH3COOH
Carbonic acid H2CO3
Hypochlorous acid HClO
Ammonia NH3
Sodium silicate Na2SiO3
Calcium hydroxideSodium hydroxidePotassium hydroxide
Ca(OH)2
NaOHKOH
Strong acids
Strongbases
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Identify each compound as a strong or weak acid or base.
1.NaOH2.NH3
3.H2SO4
4.HCl
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Nomenclature for Acids
•The naming system depends on the suffix of the anion: -ide, -ite, and –ate.
•HnX
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Nomenclature for Acids
1. When the name of the anion ends in –ide, the acid name begins with the prefix hydro-. The stem of the anion has the suffix –ic and is followed by the word acid.
Ex.HCl anion= chlorideHydrochloric acid
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Nomenclature for Acids
2. When the anion name ends in –ite, the acid name is the stem of the anion with the suffix –ous, followed by the word acid.
Ex.
H2SO3 anion= sulfite
Sulfurous acid
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Nomenclature for Acids
3. When the anion name ends in –ate, the acid name is the stem of the anion with the suffix –ic, followed by the word acid.
Ex.
HNO3 anion= nitrate
Nitric acid
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Write the name of the following acids
1.HF2.HNO3
3.H2SO3
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Nomenclature for acids•Look at the ions being used and write the
formula like an ionic compound
Ex.
Hydrobromic acid anion= bromide (rule 1)
H+ Br- = HBr
Phosphoric acid anion= phosphite (rule 2)
H+ PO3-3 = H3PO3
Sulfuric acid anion= sulfate (rule 3)
H+ SO4-2 = H2SO4
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Write the formula of the following
acids
1.Perchloric acid2.Hydroiodic acid3.Chlorous acid
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Nomenclature for Bases
•The formula and names are the same as ionic compounds
NaOH = sodium hydroxide
Aluminum hydroxide = Al+3 OH-
Al(OH)3
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Write the name of the following bases
1.Ba(OH)2
2.Ca(OH)2
3.RbOH
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Write the formulas of the following
bases1.Cesium hydroxide2.Beryllium hydroxide3.Manganese
hydroxide