Objectives Define the properties of Acids and Bases Compare and
contrast acids and bases as defined by the theories of Arrhenius,
Bronsted-Lowry, and Lewis
Slide 3
Acids They taste sour Conduct electricity Strong or weak
electrolytes React with metals to form H 2 gas Change color of
indicators (shown in 2 slides) React with bases to form water and a
salt (not necessarily NaCl)
Slide 4
Acids Continued pH less than 7 (more later) React with
carbonates and bicarbonates to produce a salt, water and CO 2 How
can we tell something is an acid? Usually starts with H HCl, HNO 3,
H 2 SO 4 (not water!)
Slide 5
Acids and Indicators Blue litmus paper turns red
Slide 6
Acids have a pH less than 7
Slide 7
Acids and Metals React with active metals to form salts and
hydrogen gas HCl (aq) + Mg (s) MgCl 2(aq) + H 2(g) What type of
reaction is this?
Slide 8
Acids React with Carbonates and Bicarbonates HCl + NaHCO 3 NaCl
+ H 2 O + CO 2 Hydrochloric acid + sodium bicarbonate salt + water
+ carbon dioxide An old-time home remedy for relieving an upset
stomach
Slide 9
Effects of Acid Rain
Slide 10
Slide 11
Acids React with Bases HCl + NaOH NaCl + H 2 O Neutralization
Acid + Base = salt + water What type of reaction is this?
Slide 12
Bases React with acids to form salt and water Taste bitter Feel
slippery Strong or weak electrolytes Change color of
indicators
Bases and Indicators Red litmus turns blue Phenolphthalein
turns purple
Slide 15
Bases have a pH greater than 7
Slide 16
Theories 3 Acid-Base theories: Arrhenius Bronsted-Lowry
Lewis
Slide 17
Svante Arrhenius Swedish chemist Won Nobel Prize in Chemistry
in 1903 One of the first chemists to explain chemical behavior of
acids and bases
Slide 18
Svante Arrhenius (1859-1927)
Slide 19
Arrhenius Theory Acids produce hydrogen ions (H + ) in aqueous
solutions HCl H + + Cl - Bases produce hydroxide ions (OH - ) in
aqueous solutions NaOH Na + + OH - Limited to water solutions Only
one kind of base hydroxide
Slide 20
Polyprotic Acid Some compounds have more than one ionizable
hydrogen to release HNO 3 monoprotic H 2 SO 4 diprotic H 3 PO 4
triprotic More hydrogens does not mean stronger!
Slide 21
Acids Not all compounds that have hydrogen are acids Example:
water Not all the hydrogen in an acid will be released as ions
Slide 22
Arrhenius Examples HCl CH 3 COOH H 2 SO 4 H 3 PO 4 NaOH Ca(OH)
2 Al(OH) 3
Slide 23
Johannes Brnsted Thomas Lowry (1879-1947) (1874-1936) Denmark
England
Slide 24
Bronsted-Lowry Acid = H + donor Base = H + acceptor Acids and
bases always come in pairs
Slide 25
Why is Ammonia a Base? NH 3 is the H + acceptor (base) and
water is the H + donor (acid) This causes the OH - concentration to
increase and the ammonium solution becomes basic
Slide 26
Acid-Base Pairs Conjugate base the remainder of the original
acid after it donates its hydrogen ion Conjugate acid formed when
the original base gains a hydrogen ion Conjugate acid-base pair is
related by loss or gain of single hydrogen ion
Slide 27
Water Water is amphoteric Can act as both an acid and a
base
Slide 28
Gilbert Lewis (1875-1946)
Slide 29
Lewis Theory Lewis acid electron pair acceptor Lewis base
electron pair donor Most general definition Acids dont even need
hydrogen!
Slide 30
Lewis Acid-Base
Slide 31
Slide 32
Hydrogen Ions and Acidity
Slide 33
Objectives Describe how [H + ] and [OH - ] are related in an
aqueous solution Classify a solution as neutral, acidic, or basic
given the hydrogen-ion and hydroxide ion concentration Convert
hydrogen ion concentrations into pH and hydroxide ion
concentrations to pOH Describe the purpose of an acid-base
indicator
Slide 34
Ions from Water Water ionizes into H + and OH - H 2 O H + + OH
- Self-ionization of water Occurs at a very small extent [H + ] =
[OH - ] = 1x10 -7 M Since the concentrations are equal the solution
is neutral
Slide 35
Ion Product Constant K w = ion product constant of water K w =
[H + ][OH - ] K w = 1x10 -14 M 2 If [H + ] > 10 -7 acidic
solution More H + than OH - If [OH - ] > 10 -7 basic solution
More OH - than H +
Slide 36
Problem If the [H + ] of a cola solution is 2.0 x 10 -6 M, is
the solution acidic, basic or neutral? What is the [OH - ] of this
solution?
Slide 37
On Your Own Classify each solution as acidic, basic or neutral:
[H + ] = 6.0 x 10 -10 M [OH - ] = 3.0 x 10 -2 M [H + ] = 2.0 x 10
-7 M [OH - ] = 1.0 x 10 -7 M
Measuring pH Steps: Moisten the pH indicator paper strip with a
few drops of solution, by using a stirring rod. Compare the color
to the chart on the vial then read the pH value.
Slide 41
Acid-Base Indicator Limitations: Usually given for a certain
temperature may change at different temps The solution may already
have a color (i.e. paint) The ability to the human eye to
distinguish colors is limited
Slide 42
Acid-Base Indicator A pH meter may give more definitive results
Measures voltage between 2 electrodes Typically accurate to within
0.01 pH unit of true pH Instruments must be calibrated
Slide 43
pOH pOH = -log[OH - ] pH + pOH = 14 Thus, a solution with pOH 7
is acidic Not commonly used
Slide 44
pH and Significant Figures For pH calculations, the hydrogen
ion concentration is usually expressed in scientific notation [H 1+
] = 0.0010 M = 1.0 x 10 -3 M, and 0.0010 has 2 significant figures
the pH = 3.00, with the two numbers to the right of the decimal
corresponding to the two significant figures
Slide 45
Problem What is the pH of a solution with a hydrogen-ion
concentration of 4.2 x 10 -10 M? The pH of an unknown solution is
6.35. what is the hydrogen ion concentration?
Slide 46
Another What is the pH of a solution if [OH - ] = 4.0 x 10 -11
M?
Slide 47
On Your Own Calculate pH of each solution [H + ] = 5.0 x 10 -5
M [OH - ] = 4.5 x 10 -11 M What are the hydrogen ion concentrations
for solutions with following pH values: 4.00 11.55
Slide 48
Strengths of Acids and Bases
Slide 49
Objectives Define strong acids and weak acids Describe how an
acids strength is related to the value of its acid dissociation
constant Calculate an acid dissociation constant (K a ) from
concentration and pH measurements Order acids by strength according
to their acid dissociation constants (K a ) Order bases by strength
according to their base dissociation constants (K b )
Slide 50
Strength Acid-base strength is classified according to how much
they ionize in water Strong = completely ionize Weak = slightly
ionize Strength concentration
Slide 51
Strength Strong means it will make many ions when dissolved
Mg(OH) 2 is a strong base it falls completely apart nearly 100% But
not much of it dissolves in the first place not concentrated
Slide 52
Problem assuming 100% dissociation! A 1.45x10 -3 M aqueous
solution of sulfuric acid is made. Calculate the pH.
Slide 53
Try it 1 g of phosphoric acid is added to water to make a 2 L
solution. Calculate the pH of the solution.
Slide 54
Strong Acid Dissociation (makes 100 % ions)
Slide 55
Weak Acid Dissociation (only partially ionizes)
Slide 56
Measuring Strength
Slide 57
Same for Bases
Slide 58
Strength vs. Concentration The words concentrated and dilute
tell how much of an acid or base is dissolved in solution - refers
to the number of moles of acid or base in a given volume The words
strong and weak refer to the extent of ionization of an acid or
base
Slide 59
Practice Problem A 0.1000 M solution of acetic acid is only
partially ionized. From measurements of the pH of the solution, [H
+ ] is determined to be 1.34 x 10 -3 M. What is the acid
dissociation constant (K a ) of acetic acid?
Slide 60
Try it In an exactly 0.1 M solution of HCOOH, [H + ] = 4.2 x 10
-3 M. Calculate the K a of the HCOOH.
Slide 61
A Bit Harder Calculate [H + ], pH, and % dissociation in 0.1 M
solution of weak acid HNO 2. Ka = 5.0 x 10 -5
Slide 62
On Your Own Calculate the pH and % dissociation of a 0.98 M
solution of acetic acid. Ka = 1.77 x 10 -5
Slide 63
HARDER! (You will see it in AP) The K a of HF is 6.46 x 10 -4.
Calculate the pH of 0.0100 M solution of HF.
Slide 64
Neutralization Reactions
Slide 65
Objectives Define the products of an acid-base reaction Explain
how acid base titration is used to calculate the concentration of
an acid or a base Explain the concept of equivalence in
neutralization reactions Describe the relationship between
equivalence point and the end point of a titration
Slide 66
Neutralization Reaction A reaction in which an acid and a base
react in aqueous solution to produce a salt and water HCl (aq) +
NaOH (aq) NaCl (aq) + H 2 O (l) H 2 SO 4(aq) + 2KOH (aq) K 2 SO
4(aq) + 2 H 2 O (l)
Slide 67
Titration Process of adding a known amount of solution of known
concentration to determine the concentration of another solution
Balanced equation gives mole ratio Equivalence point when the moles
of hydrogen ions equals moles of hydroxide ions (neutralized)
Slide 68
Sample Problem How many moles of sulfuric acid are required to
neutralize 0.50 M sodium hydroxide solution? The solution is 2 L.
What is the concentration of sulfuric acid?
Slide 69
Another What concentration of 15 mL of potassium hydroxide are
needed to completely neutralize 25 mL of 1.56 M phosphoric
acid?
Slide 70
Titration The concentration of an acid (or base) in solution
can be determined by performing a neutralization reaction Indicator
is used shows neutralization occurs Phenolphthalein often used
because colorless in acid and pink in base
Slide 71
Neutralization Steps 1. Measured volume of unknown acid
concentration is added to flask 2. Several drops of indicator added
3. A base of known concentration is added SLOWLY until the
indicator changes color 4. Measure volume of base added
Slide 72
Neutralization The solution of known concentration is called
the standard solution added by using a buret Continue adding until
the indicator changes color called the end point of the
titration
Slide 73
Salts in Solution
Slide 74
Objectives Describe when a solution of salt is acidic or basic
Demonstrate with equations how buffers resist change in pH
Slide 75
Salt Salt ionic compound Comes from anion of an acid Comes from
cation of base Formed during neutralization reaction Some are
neutral while others are acidic or basic
Slide 76
Salt Hydrolysis Salt hydrolysis a salt that reacts with water
to produce an acid or base Hydrolyzing salts come from: Strong acid
+ weak base A weak acid + strong base Strong degree of
ionization
Slide 77
Slide 78
Salt Hydrolysis To see if the resulting salt is acidic or basic
you need to look at the parent acid and base Lets look HCl + NaOH
NaCl + H 2 O H 2 SO 4 + NH 4 OH (NH 4 ) 2 SO 4 + H 2 O CH 3 COOH +
KOH CH 3 COOK + H 2 O
Slide 79
Predict whether the following salts will from a solution that
is acidic, basic or neutral. a. FeCl 3 e. NaF b. NH 4 CN*f. LiClO 4
c. K 2 CO 3 d. CaBr 2 h. NH 4 Br
Slide 80
Slide 81
Buffers Buffer solution in which the pH remains constant even
when acid and bases are added to them Made from a pair of
chemicals: Weak acid and one of its salts Weak base and one of its
salts
Slide 82
Buffers Better able to resist change in pH than pure water
Since it is a pair of chemicals: One chemical neutralizes any acid
added while the other chemical would neutralize any base added
Slide 83
Buffer Capacity Buffers resist change in pH HOWEVER if you add
enough acid or base the pH WILL change Buffer capacity the amount
of acid or base that can be added before a significant change in
pH
Slide 84
Buffers The 2 buffers crucial to maintaining the pH of human
blood: Carbonic acid (H 2 CO 3 ) & Hydrogen carbonate (HCO 3 -1
) Dihydrogen phosphate (H 2 PO 4 -1 ) & monohydrogen phosphate
(HPO 4 -2 )