Honor Chemistry Chapter 15 Mr. Pendolino
Acids and Bases Chapter 15 Notes
Properties of Acids
• Sour taste
• Change color of vegetable dyes
• React with “active” metals
– Like Al, Zn, Fe, but not Cu, Ag or Au
Zn + 2 HCl ®ZnCl2 + H2
– Corrosive
• React with carbonates, producing CO2
– Marble, baking soda, chalk
CaCO3 + 2 HCl ®CaCl2 + CO2 + H2O
• React with bases to form ionic salts
– And often water
Properties of Bases
• Also Known As Alkalis
• Taste bitter
• Feel slippery
• Change color of vegetable dyes
– Different color than acid
– Litmus = blue
• React with acids to form ionic salts
– And often water
– Neutralization
Arrhenius Theory
• Acids ionize in water to H+1 ions and anions
• Bases ionize in water to OH-1 ions and cations
• Neutralization reaction involves H+1 combining with OH-1 to make water
• H+ ions are protons
• Definition only good in water solution
• Definition does not explain why ammonia solutions turn litmus blue
– Basic without OH- ions
Brønsted-Lowery Theory
• H+ transfer reaction
– Since H+1 is a proton, also known as proton transfer reactions
• Acid is H+ donor; Base is H+ acceptor
– Base must contain an unshared pair of electrons
• In the reaction, a proton from the acid molecule is transferred to the base molecule
– H forms a bond to lone pair electrons on the base molecule
– We consider only 1 H transferred in each reaction
• Products are called the Conjugate Acid and Conjugate Base
– After reaction, the original acid is the conjugate base and the original base is changed to what is now called the conjugate acid
H-A + :B ® A-1 + H-B+1
A-1 is the conjugate base, H-B+1 is the conjugate acid
• Conjugate Acid-Base Pair is either the original acid and its conjugate base or the original base and its conjugate acid
– H-A and A-1 are a conjugate acid-base pair
– :B and H-B+1 are a conjugate acid-base pair
• The conjugate base is always more negative than the original acid; and the conjugate acid is always more positive than the original base
Example #1
Write the conjugate base for the acid H3PO4?
• Determine what species you will get if you remove 1 H+1 from the acid
– The Conjugate Base will have one more negative charge than the original acid
H3PO4 ® H+1 + H2PO4-1
Brønsted-Lowery Theory
• In this theory, instead of the acid, HA, dissociating into H+1(aq) and A-1(aq); The acid donates its H to a water molecule
HA + H2O ® A-1 + H3O+1
A1- is the conjugate base, H3O1+ is the conjugate acid
• H3O+1 is called hydronium ion
• In this theory, substances that do not have OH1- ions can act as a base if they can accept a H1+ from water
H2O + :B ® OH-1 + H-B+1
Strength of Acids & Bases
• The stronger the acid, the more willing it is to donate H
• Strong acids donate practically all their H’s
HCl + H2O ® H3O+1 + Cl-1
• Strong bases will react completely with water to form hydroxides
CO3-2 + H2O ®HCO3-1 + OH-1
• Weak acids donate a small fraction of their H’s
– The process is reversible, the conjugate acid and conjugate base can react to form the original acid and base
HC2H3O2 + H2O Û H3O+1 + C2H3O2-1
• Only small fraction of weak base molecules pull H off water
HCO3-1 + H2O ÛH2CO3 + OH-1
Multiprotic Acids
• Monoprotic acids have 1 acid H, diprotic 2, etc.
– In oxyacids only the H on the O is acidic
• In strong multiprotic acids, like H2SO4, only the first H is strong; transferring the second H is usually weak
H2SO4 + H2O ® H3O+1 + HSO4-1
HSO4-1 + H2O Û H3O+1 + SO4-2
Water as an Acid and a Base
• Amphoteric substances can act as either an acid or a base
– Water as an acid, NH3 + H2O Û NH4+1 + OH-1
– Water as a base, HCl + H2O ® H3O+1 + Cl-1
• Water can even react with itself
H2O + H2O Û H3O +1 + OH-1
Autoionization of Water
• Water is an extremely weak electrolyte
– therefore there must be a few ions present
H2O + H2O Û H3O+1 + OH-1
• all water solutions contain both H3O+1 and OH-1
– the concentration of H3O+1 and OH-1 are equal
– [H3O+1] = [OH-1] = 10-7M @ 25°C
• Kw = [H3O+1] x [OH-1] = 1 x 10-14 @ 25°C
– Kw is called the ion product constant for water
– as [H3O+1] increases, [OH-] decreases
Acidic and Basic Solutions
• acidic solutions have a larger [H+1] than [OH-1]
• basic solutions have a larger [OH-1] than [H+1]
• neutral solutions have [H+1]=[OH-1]= 1 x 10-7 M
[H+1] = / 1 x 10-14 / [OH-1] = / 1 x 10-14[OH-1] / [H+1]
Example #2
Determine the [H+1] and [OH-1] in a 10.0 M H+1 solution
¬ Determine the given information and the information you need to find
Given [H+1] = 10.0 M Find [OH-1]
Solve the Equation for the Unknown Amount
® Convert all the information to Scientific Notation and Plug the given information into the equation.
Given [H+1] = 10.0 M = 1.00 x 101 M
Kw = 1.0 x 10-14
pH & pOH
• The acidity/basicity of a solution is often expressed as pH or pOH
• pH = -log[H3O+1] pOH = -log[OH-1]
– pHwater = -log[10-7] = 7 = pOHwater
• [H+1] = 10-pH [OH-1] = 10-pOH
• pH < 7 is acidic; pH > 7 is basic, pH = 7 is neutral
• The lower the pH, the more acidic the solution; The higher the pH, the more basic the solution
• 1 pH unit corresponds to a factor of 10 difference in acidity
• pOH = 14 - pH
Example #3
Calculate the pH of a solution with a [OH-1] = 1.0 x 10-6 M
¬ Find the concentration of [H+1]
Enter the [H+1] concentration into your calculator and press the log key
log(1.0 x 10-6) = -6.0
® Change the sign to get the pH
pH = -(-6.0) = 6.0
Example #4
Calculate the pH and pOH of a solution with a [OH-1] = 1.0 x 10-3 M
¬ Enter the [H+1] or [OH-1]concentration into your calculator and press the log key
log(1.0 x 10-3) = -3.0
Change the sign to get the pH or pH
pOH = -(-3) = 3.0
® Subtract the calculated pH or pOH from 14.00 to get the other value
pH = 14.00 – 3.0 = 11.0
Example #5
Calculate the [OH-1] of a solution with a pH of 7.41
¬ If you want to calculate [OH-1] use pOH, if you want [H+1] use pH. It may be necessary to convert one to the other using 14 = [H+1] + [OH-1]
pOH = 14.00 – 7.41 = 6.59
Enter the pH or pOH concentration into your calculator
® Change the sign of the pH or pOH
-pOH = - (6.59)
¯ Press the button(s) on your calculator to take the inverse log or 10x
[OH-1] = 10-6.59 = 2.6 x 10-7
Calculating the pH of a Strong, Monoprotic Acid
• A strong acid will dissociate 100%
HA ® H+1 + A-1
• Therefore the molarity of H+1 ions will be the same as the molarity of the acid
• Once the H+1 molarity is determined, the pH can be determined
pH = -log[H+1]
Example #6
Calculate the pH of a 0.10 M HNO3 solution
¬ Determine the [H+1] from the acid concentration
HNO3 ® H+1 + NO3-1
0.10 M HNO3 = 0.10 M H+1
Enter the [H1+] concentration into your calculator and press the log key
log (0.10) = -1.00
® Change the sign to get the pH
pH = -(-1.00) = 1.00
Buffered Solutions
• Buffered Solutions resist change in pH when an acid or base is added to it.
• Used when need to maintain a certain pH in the system
– Blood
• A buffer solution contains a weak acid and its conjugate base
• Buffers work by reacting with added H+1 or OH-1 ions so they do not accumulate and change the pH
• Buffers will only work as long as there is sufficient weak acid and conjugate base molecules present
Page 4 of 5 Acids and Bases 5/14/2012