Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Unit 11:
Acids and Bases
Student Name: _______________________________________
Class Period: ________
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Unit 11 Vocabulary:
1. Acidity: The property of exhibiting the qualities of an acid.
2. Alkalinity: The property of exhibiting the qualities of a base.
3. Amphiprotic: any substance (ionic or covalent) that can both accept
and donate at least one proton (H+).
4. Amphoteric: A species that can act either as a Brönsted-Lowry acid
or a Brönsted-Lowry base, depending on the other species it is
reacting with.
5. Arrhenius Acid: An electrolyte that ionizes in aqueous solution to
yield H+ as the only cation in the solution.
6. Arrhenius Base: An electrolyte that ionizes in aqueous solution to
yield OH- as the only anion in the solution.
7. Basicity: The property of exhibiting the qualities of a base.
8. Brönsted-Lowry (B-L) Acid: A species that donates H+ to a B/L base
in a chemical reaction.
9. Brönsted-Lowry (B-L) Base: A species that accepts H+ from a B/L
acid in a chemical reaction.
10. Buret: A calibrated precision (& expensive!) glass tube that
precisely measures the volume of a liquid dispensed.
11. Caustic: A substance that will destroy or irreversibly damage a
substance or surface that it contacts; usually used to describe bases.
12. Corrosive: A substance that will destroy or irreversibly damage a
substance or surface that it contacts; usually used to describe acids.
13. Electrolyte: A compound that ionizes (dissociates) in water, allowing
the solution to flow electrons freely and conduct electricity.
14. Hydrolysis: The process whereby a base reacts with a glycerol ester
(a fat) to produce soap.
15. Indicator: A substance whose color is sensitive to the pH of a
solution to which it is added.
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
16. Neutralization: A double-replacement reaction where an acid and a
base react to form water and a salt.
17. Nonelectrolytes: A molecular compound that does not ionize (not
dissociate) in water, preventing the solution from conducting
electricity.
18. pH: The negative logarithm of the hydrogen ion concentration. A pH
value less than 7 indicates an acidic solution, a pH value 7 is a
neutral solution, and a pH value greater than 7 indicates a basic
solution. The scale ranges from a hypothetical 0 (most acidic) to 14
(most basic).
19. Protonation: The addition of an acid’s H+ (proton) to a water
molecule to form a hydronium (H3O+) molecule.
20. Salt: An ionic compound formed when an acid and a base neutralize
each other. The salt compound consists of both the anion of the acid
and the cation of the base.
21. Titration: A process of controlled acid-base neutralization carried out
using burets dispensing either an acid or a base.
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Unit 11 Homework Assignments:
Assignment:
Date:
Due:
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Notes page:
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Unit 11: Acids & Bases-Lecture
Topic:
Regents Chemistry ’14-‘15
Mr. Murdoch
Arrhenius Acids and Bases
Objective: What is the definition of an Arrhenius Acid or Base?
Arrhenius Acids:
 A substance that contains H+ ions that ionize when dissolved in water
is known as an Arrhenius Acid.
 Acids (and bases) are the only molecules (different from ionic
crystals) that ionize (dissociate) when dissolved in water. Acids (and
bases) are electrolytes, unlike other molecular substances like water
(H2O) and sugar (C6H12O6). The H+ leaves the acid and bonds to the
water molecule to form a hydronium ion (H3O+).
Arrhenius Acid examples:
HCl(g) + H2O(l)  H3O+(aq) + Cl-(aq)
i.
The acid HCl contains one H+ ion which combines with one H2O
molecule to form one H3O+ ion.
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Unit 11: Acids & Bases-Lecture
ii.
iii.
Regents Chemistry ’14-‘15
Mr. Murdoch
The acid (HCl) ionizes to form a hydrogen cation (H+) and an
anion, in this example a chloride ion (Cl-).
The hydrogen ion (H+) bonds to the water molecule using a
“sneaky” bonding method. The H+ doesn’t have an electron to
share, and there are no unpaired valence electrons in the water
molecule. The hydrogen ion tends to “bogart” two electrons from
the oxygen in the water molecule. The extra hydrogen ion now has
a “claim” on the electrons on the oxygen, forming a coordinate
covalent bond.
H2SO4(l) + 2 H2O(l)  2 H3O+(aq) + SO4-2(aq)
iv.
The acid H2SO4 contains two H+ ions which combine with two
H2O molecules to form two H3O+ hydronium ions.
H3PO4(l) + 3 H2O(l)  3 H3O+(aq) + PO4-3(aq)
v.
The acid contains three H+ ions which combine with three H2O
molecules to form three H3O+ hydronium ions.
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Unit 11: Acids & Bases-Lecture
Topic:
Regents Chemistry ’14-‘15
Mr. Murdoch
Properties of Acids
Objective: How are ions situated within a solution?
Properties of Acids:
 Acids eat away (oxidize) active metals (metals above H2 on Table J).
 Metals such as Li, Mg, and Zn may be oxidized by an acid to produce
hydrogen gas (H2(g)). These are examples of metals above H2 on
Table J. We won’t discuss metals below H2 on Table J now.
 For these single replacement reactions:
2 Li(s) + 2 HCl(aq)  2 LiCl(aq) + H2(g)
Ca(s) + 2 HCl(aq)  CaCl2(aq) + H2(g)
i. Both of these are examples of how a more active metal (above H2)
will “kick” the hydrogen out of the acid (HCl), leaving a new
aqueous ionic compound and hydrogen gas.
1. Acids have a pH value less than 7.
i. The pH scale measures the relative acidity or alkalinity of a
solution. A pH value 7 is neutral, and acids have a pH value less
than 7. Each decrease of one whole number of pH is a tenfold
increase in acid strength. An acid with a pH value 3 is ten times
more acidic than a solution with a pH value 4, and an acid with a
pH value 3 is 100 times more acidic than a solution with a pH
value 5.
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
2. Acidic solutions may conduct electricity.
i. Acids are electrolytes, because they dissociate to form ions in
solution. The stronger the acid, the more it ionizes, and therefore
could carry electron flow better and have better electrical
conductivity.
ii. Stronger electrolytes (stronger acids) include: HI, HCl, HNO3, and
H2SO4. H2SO4 is commonly used in lead-acid automotive batteries.
iii. Weaker electrolytes (weaker acids) include: H2CO3 (carbonic acid in
soda) and HC2H3O2 (acetic acid, or vinegar).
3. Dilute solutions of acids taste sour.
*NOTE: this statement does NOT give you permission
to taste ANY acid in lab to test this statement!
i. Citric acid is found in citrus fruits like lemons and grapefruits.
Citric acid is also used to give an extra sour “kick” to food and
candies.
ii. Acetic acid (5% in aqueous solution) is also known as “white
vinegar” and is yummy on salads or French Fries!
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
4. Acids react with carbonates to form carbon dioxide gas, a salt, and
water.
i. Baking soda and vinegar:
Spoiler Alert! - If you remember the mysterious “Enzyme M” from the required NYS Living
Environment Relationships and Biodiversity Lab, the powder was baking soda, and the 4 plant
“extract” solutions were food coloring and water, with three of the extracts containing vinegar.
NaHCO3(s) + HC2H3O2(aq)  CO2(g) + H2O(l) + NaC2H3O2(aq)
 This is the middle school (& VERY messy!) “volcano” reaction;
the CO2(g) forces the mixture of dry powder, liquid, and foam up
and out.
5. Acids may be formed by reaction of gaseous oxides with water.
i. Burning fossil fuels releases gaseous nonmetallic oxides (CO2,
*
NOx, SO2, and similar molecules) into the atmosphere. When
these gaseous nonmetallic oxides react with atmospheric water
vapor, they form weak aqueous acids that may form acidic
precipitation. Acid precipitation can lower the pH of bodies of
water, which can seriously affect aquatic ecosystems.
*
NOx is a generic symbol for the nitrogen oxides of NO, NO2, & N2O
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Naming of Acids
Topic:
Objective: How do we name acidic compounds?
Naming of Binary Acids:
 A binary acid has a name composed of the prefix (hydro-), the
nonmetallic ion name, and the last syllable (-ide) replaced with the
suffix (-ic) followed by the separate word “acid”.
 Binary Acid naming examples:
i. HCl(aq): (hydro-) + (chloride - ide) + (-ic) + “acid” = hydrochloric acid
ii. HBr(aq): (hydro-) + (bromide - ide) + (-ic) + “acid” = hydrobromic acid
iii. H2S(aq): (hydro-) + (sulfide - ide) + (-ic) + “acid” = hydrosulfic acid
iv. H3N(aq): (hydro-) + (nitride - ide) + (-ic) + “acid” = hydronitric acid
 Writing of Binary Acid formulas:
i. Write a binary acid formula just like any ionic compound, placing
the H+1 first, and the anion second.
1. Hydrofluoric acid:
One hydrogen (H+1) + one fluoride (F-1)  HF(aq)
2. Hydrophosphoric acid:
Three hydrogen (3 x H+1) + one phosphide (P-3)  H3P(aq)
ii. Ensure that you have the correct ratio of ions to equal the charges!
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Naming of Ternary Acids:
 A ternary acid has a name with NO (hydro-) prefix.
a) If the polyatomic ion ends in (-ide) or (-ate), replace the last
syllable of the polyatomic ion with the suffix (-ic) followed by the
word “acid”.
b) If the polyatomic ion ends in (-ite), replace the last syllable of the
polyatomic ion with the suffix (-ous) followed by the word “acid”.
 Ternary Acid naming examples:
i. HNO3(aq): (nitrate - ate) + (-ic) + “acid”  nitric acid
ii. HNO2(aq): (nitrite -ite) + (-ous) + “acid”  nitrous acid
iii. HClO3(aq): (chlorate - ate) + (-ic) + “acid”  chloric acid
iv. HClO2(aq): (chlorite - ite) + (-ous) + “acid”  chlorous acid
v. HCN(aq): (cyanide - ide) + (-ic) + “acid”  cyanic acid
 Writing of Ternary Acid formulas:
a) Remember that while there is no (hydro-) prefix, ternary acid
formulas still begin with hydrogen!
i. Sulfuric acid: hydrogen (2 x H+1) + sulfate (SO4-2)  H2SO4(aq)
ii. Sulfurous acid: hydrogen (2 x H+1) + sulfite (SO3-2)  H2SO3(aq)
a) Again, ensure that you have the correct ratio of ions to equal the
charges!
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Regents Practice Questions-Acids: (Ungraded)
1. The only positive ion found in an aqueous found in a aqueous
solution of sulfuric acid is the
a) Sulfite ion
c) Hydroxide ion
b) Sulfate ion
d) Hydronium ion
2. An Arrhenius acid has
a) Only hydrogen ions in solution
b) Only hydroxide ions in solution
c) Hydrogen ions as the only positive ions in solution
d) Hydrogen ions as the only negative ions in solution
3. Which substance is an Arrhenius acid?
a) LiF(aq)
c) CH3CHO
b) HBr(aq)
d) Mg(OH)2(aq)
4. What produces hydrogen ions as the only positive ions in aqueous
solution?
a) NH3
c) KOH
b) HBr
d) NaCl
5. When HCl is dissolved in water, the only positive ion present in the
solution is the
a) Hydride ion
c) Hydrogen ion
b) Chloride ion
d) Hydroxide ion
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Unit 11: Acids & Bases-Lecture
Topic:
Regents Chemistry ’14-‘15
Mr. Murdoch
Properties of Bases
Objective: How does temperature affect the solubility of a solution?
Properties of Bases:
 Bases are substances that contain aqueous hydroxide (OH-1(aq)) ions
in solution.
1. Bases have a pH value greater than 7.
i. The pH scale measures the relative acidity or alkalinity of a
solution. A pH value 7 is neutral, and bases have a pH value
greater than 7. Each increase of one whole number of pH is a
tenfold increase in base strength. A base with a pH value 9 is ten
times more basic than a solution with a pH value 8, and a base with
a pH value 9 is 100 times more basic than a solution with a pH
value 7.
2. Basic solutions may conduct electricity.
i. Bases are electrolytes, because they dissociate to form ions in
solution. The stronger the base, the more it ionizes, and therefore
could carry electron flow better and have better conductivity.
ii. Stronger electrolytes (stronger bases) include: Group 1 metal
hydroxides (LiOH, NaOH, RbOH, and CsOH).
iii. Weaker electrolytes (weaker bases) include: Ca(OH)2, Mg(OH)2,
and Al(OH)3, all of which may be found in common antacids. We
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
will learn SOON how basic antacids neutralize excess stomach
acid.
3. Bases taste bitter.
AGAIN, we will NOT test this statement in lab!
i. However, many alkaloids are found in medicines, and are also in
coffee.
ii. Theobromine is an alkaloid found in chocolate, and the compound
that makes chocolate dangerous to many pets.
4. Bases may be formed when Group 1 or Group 2 metals react with
water, releasing hydrogen gas in the process.
i. 2 Na(s) + 2 H2O(l)  2 NaOH(aq) + H2(g)
ii. Mg(s) + 2 H2O(l)  Mg(OH)2 + H2(g)
5. Bases hydrolyze fats during saponification to form soap.
i. Drain cleaners usually contain sodium hydroxide which reacts
with the nonpolar molecules in grease clogging household drains.
The NaOH converts the nonpolar grease molecules into ionic soap
compounds, which then become soluble in water and may be
washed down the drain.
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Unit 11: Acids & Bases-Lecture
Topic:
Regents Chemistry ’14-‘15
Mr. Murdoch
Naming of Bases
Objective: How do we name acidic compounds?
Naming of Bases: (Reference Table L)
1. Most common bases are formed when a metal (especially Groups 1
and 2) bonds with a hydroxide. Naming bases is easy once you've
identified them!
2. Name the metal CATION first; it keeps its name as listed in the
Periodic Table.
3. The polyatomic ion "hydroxide" (OH-) also keeps its name.
 Naming of Bases examples:
i. LiOH is "lithium hydroxide"
ii. NH4OH is "ammonium hydroxide"
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Regents Practice Questions-Bases: (Ungraded)
1. According to the Arrhenius theory, when a base dissolves in water it produces
a) H+ as the only positive ion in solution
b) NH4+ as the only positive ion in solution
c) OH- as the only negative ion in solution
d) CO3-2 as the only negative ion in solution
2. A sample of Ca(OH)2 is considered to be an Arrhenius base because it dissolves
in water to yield
a) H- as the only negative ions in solution
b) Ca+2 ions as the only positive ions in solution
c) OH- ions as the only negative ions in solution
d) H3O+ ions as the only positive ions in solution
3. Which ion is produced when an Arrhenius base is dissolved in water?
a) H+ as the only positive ion in solution
b) OH- ions as the only negative ion in solution
c) H3O+ ions as the only positive ion in solution
d) H- as the only negative ion in solution
4. According to the Arrhenius theory, which list of compounds includes only
bases?
a) KOH, NaOH, and LiOH
b) KOH, Ca(OH)2, and CH3OH
c) LiOH, Ca(OH)2, and C2H4(OH)2 (C2 H4(OH)2 is an ethylene glycol molecule)
d) NaOH, Ca(OH)2, and CH3COOH
5. According to the Arrhenius theory, when a base is dissolved in water it
produces a solution containing only one kind of anion. What is the name of this
anion?
a) Hydride ion
c) Hydrogen sulfate ion
b) Hydroxide ion
d) Hydrogen carbonate ion
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Al(OH)
hydroxide”for pH
Topic: 3 is “aluminum
Testing
Objective: How do we know if a solution is Acidic or Basic?
How do we know if a solution is Acidic or Basic?
 Electronic pH testers:
i. Electronic pH testers have probes containing electrodes that detect
electrical conductivity. Once calibrated (a process comparing
known standards to the probe output), the electronic signal may be
interpreted to determine an ionization level, and therefore the pH
value.
 Acid-Base Indicators: (Reference Table M)
i. Indicators are chemicals that have certain color characteristics for a
narrow range of pH values. Indicators may be used to determine
acidity or alkalinity, or even to find a specific range of pH values.
An indicator won’t give an exact pH value, but a range of pH
values.
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
 Litmus and pH paper:
i. Red litmus paper is used to test a base. When red litmus paper is
immersed in a base, the red litmus paper turns blue indicating the
given solution as alkaline.
ii. Blue litmus paper is used to test an acid. When blue litmus paper is
immersed in an acid, the blue litmus paper turns red indicating the
given solution as acidic.
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
 Full-range pH paper contains a mixture of indicators that will react
with a solution to give an approximation of pH value for that
solution. The color that the pH paper strip turns is compared to a
color chart for the paper that gives a quantitative pH value based on
the qualitative color.
Testing for an Acid or Base example:
An unknown solution gives the following results when tested with the
following four indicators:
 Thymol Blue = yellow
 Methyl Orange = yellow
 Bromcresol Green = blue
 Phenolphthalein = colorless
1. Which of the following pH values could the unknown solution have?
a) 2.8
b) 4.8
c) 6.5
d) 8.5
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
 Using Reference Table M, Common Acid-Base Indicators, we
should be able to narrow down our possibilities.
i. Thymol Blue = turned yellow
Thymol Blue works between pH values of 8.0 (yellow) - 9.6
(blue). As it was yellow in this test, the unknown solution could
be at the lower end of Thymol Blue’s range, so a pH of 8.5 is a
possibility, but it also could be below a pH value 8.
ii. Methyl Orange = turned yellow
Methyl Orange works between pH values of 3.1 (red) - 4.4
(yellow). As it was yellow in this test, the unknown solution has to
be ABOVE a pH value of 3.1, so we may eliminate choice ‘A’.
iii. Bromcresol Green = turned blue
Bromcresol Green works between pH values of 3.8 (yellow) - 5.4
(blue). As it was not yellow in this test, the unknown solution has
to be ABOVE a pH value of 3.1, so we may eliminate choice ‘A’.
However, Bromcresol Green turns blue with a pH value ABOVE
5.4, so a pH of 4.8 is eliminated (along with choice ‘B’), or it
would have been green.
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
iv. Phenolphthalein = showed colorless (no change)
Phenolphthalein works between pH values of 8 (colorless) - 9
(pink). As it was colorless in test, the unknown solution has to be
BELOW a pH value of 8, so we may eliminate choice ‘D’.
(Geesh…I sound like a witty Sicilian…)
v. Methyl Orange eliminated a pH value 2.8, Bromcresol Green
eliminated pH value 4.8, and Phenolphthalein eliminated pH value
8.5. In this case Thymol Blue could only say that the unknown pH
was below 8, so that leaves choice ‘C’ with a pH value 6.5.
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Regents Practice Questions-Indicators: (Ungraded)
1. Which solution below when mixed with a drop of bromthymol blue will
cause the indicator to change from blue to yellow?
a) 0.1 M HCl(aq)
c) 0.1 M NaOH(aq)
b) 0.1 M NH3(aq)
d) 0.1 M CH3OH(aq)
2. A solution with pH value 11 is first tested with phenolphthalein and then
with red litmus. What will be the color of each indicator in this solution?
a) Phenolphthalein is pink and red litmus is red
b) Phenolphthalein is pink and red litmus is blue
c) Phenolphthalein is colorless and red litmus is red
d) Phenolphthalein is colorless and red litmus is blue
3. The ability of H2SO4(aq) to change blue litmus paper red is mainly due to
the presence of
a) H3O+(aq) ions
c) SO4-2(aq) ions
b) SO2(aq) molecules
d) H2O(l) molecules
4. Pure water containing phenolphthalein will change from colorless to pink
with the addition of
a) KCl(aq)
c) HOH(aq)
b) HCl(aq)
d) KOH(aq)
5. A student records the following observations about an unknown solution:
 Conducts electricity
 Turns blue litmus red
The student should conclude that the unknown solution is most likely
a) A base
b) An acid
c) Normal
d) Nonpolar
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Student name: _________________________
Class Period: _______
Please carefully remove this page from your packet to hand in.
Acids & Bases Homework (1 pt. ea.)
Identify each below as an acid or base based on their formulas and properties.
Property
Turns litmus paper red
Acid or Base Property
Turns bromthymol blue
to yellow
Tastes sour
Tastes bitter
Hydrolyzes fats into soap
Reacts with active
metals forming H2(g)
HCl(aq)
KOH(aq)
pH value 12
Forms H3O+ in water
Acid or Base
Write the correct name for the following acids and bases.
1. HCl(aq):
__________________________________________________
2. HNO3(aq):
__________________________________________________
3. H2SO4(aq):
__________________________________________________
4. HC2H3O2(aq):
__________________________________________________
5. KOH(aq):
__________________________________________________
6. Ca(OH)2(aq):
__________________________________________________
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Write the correct formula for the following acids and bases. (As (aq))
7. Perchloric acid:________________________________________________
8. Hypochlorous acid: ___________________________________________
9. Chromic acid: ________________________________________________
10. Thiosulfuric acid: ____________________________________________
11. Aluminum hydroxide: ________________________________________
12. Barium hydroxide :___________________________________________
13. The results for of testing a colorless solution with three different
indicators are shown in the table below.
Indicator
Red litmus paper
Blue litmus paper
phenolphthalein
Result
Blue
Blue
Pink
Which formula could represent the solution that was tested?
a) HCl(aq)
b) NaOH(aq)
c) C6H12O6(aq)
d) C12H22O11(aq)
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Topic: Acid and Base Neutralization
Objective: What occurs during acid-base neutralization reactions?
Neutralization:
i. The combining of an acidic (H+ donor) and basic (OH- supplier)
forms a common molecule, HOH, or water (H2O). The anion from
the acid and the cation from the base join to form a salt. A salt is
simply an ionic compound that may be formed during acid-base
neutralization.
ii. Acid-base neutralization is a simple double-replacement reaction,
but in this case the water molecule is the precipitate. It may be
difficult to think of water as a “precipitate” WITHIN water, but in
this case the water formed is in excess.
iii. One mole of H+ ions exactly neutralizes one mole of OH- ions.
 Completing neutralization reactions:
i. Determining the products of acid-base neutralization is the SAME
process as determining the products of a double-replacement
reaction.
Watch Crash Course Chemistry Acid & Base Reactions YouTube video - 11:17
 Neutralization examples:
1. HCl(aq) + NaOH(aq)  NaCl(aq) + HOH(l)
Acid
Base
Salt
Water
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
2. H2SO4(aq) + 2 KOH(aq)  K2SO4(aq) + 2 HOH(l)
Acid
Base
Salt
Water
3. 2 HNO3(aq) + Ca(OH)2(aq)  Ca(NO3)2(aq) + 2 HOH(l)
Acid
Base
Salt
Water
Antacid neutralization:
Antacids are bases used to neutralize the acid that causes heartburn.
 Almost all antacids act on excess stomach acid by neutralizing it
with weak bases. The most common of these bases are hydroxides,
carbonates, or bicarbonates.
 The following table contains a list of the active ingredients found in
several common commercial antacids, and the reactions by which
these antacids neutralize the HCl in stomach acid.
Compound
Formula
Chemical Reaction
Aluminum
hydroxide
Al(OH)3
Al(OH)3(s) + 3 HCl(aq)  AlCl3(aq) + 3 H2O(l)
Calcium
carbonate
CaCO3
CaCO3(s) + 2 HCl(aq)  CaCl2(aq) + H2O(l) + CO2(g)
Magnesium
carbonate
MgCO3
MgCO3(s) + 2 HCl(aq)  MgCl2(aq) + H2O(l) + CO2(g)
Magnesium
hydroxide
Mg(OH)2 Mg(OH)2(s) + 2 HCl(aq)  MgCl2(aq) + 2 H2O(l)
Sodium
bicarbonate
NaHCO3
NaHCO3(aq) + HCl(aq)  NaCl(aq) + H2O(l) + CO2(g)
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Buffers
Topic:
Objective: How may we experimentally determine concentration?
Acid-Base Buffers:
 Buffers are solutions consisting of a weak acid and its conjugate
base; these solutions resist pH changes when either acid or base is
added to it.
i. A buffered solution may contain acetic acid as its weak acid and
sodium acetate as its conjugate base. If we add some hydrochloric
acid to this solution, the sodium acetate would react with it by the
following reaction:
HCl + NaC2H3O2 ⇔ C2H3O2H + NaCl
ii. The strong HCl added to the solution has been converted to acetic
acid, a weak acid. Weak acids cause a much smaller disruption in
pH than strong acids, and the pH of the solution will decrease much
less than if it contained no sodium acetate.
iii. Likewise, if we were to add sodium hydroxide to this solution, the
acetic acid would react to it by the following process:
NaOH + C2H3O2H ⇔ NaC2H3O2 + H2O
iv. The strong base NaOH has been converted to the weak base sodium
acetate, and the pH of the solution won't rise nearly as much as if the
acetic acid weren't present in the first place.
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Unit 11: Acids & Bases-Lecture
Topic:
Regents Chemistry ’14-‘15
Mr. Murdoch
Titration
Objective: How may we experimentally determine concentration?
Titration:
 Titration is the controlled process of acid-base neutralization that
may be used to determine the unknown molarity of an acid or base
using a precisely measured volume of a base or acid of known
molarity.
 We’ll be doing a titration lab, but here are the basic steps:
1. Place a base of unknown molarity in the buret, and record the
starting volume.
2. Add phenolphthalein indicator to a measured volume of a known
molarity acid in a flask.
3. Add small quantities of the unknown molarity base to the acid until
you start to see some semi-persistent color change. The base will
cause “blooms” of bright pink in the solution when you start the
titration.
4. Once the color begins to last longer than a few moments while
gently swirling the flask, start adding the base drop-by-drop,
swirling the contents of the flask after each drop.
5. Once the flask has ANY sign of a permanent light-pink tinge,
STOP adding base, and record the ending volume in the buret.
6. For best results, do several trials. Usually the first trial is overtitrated, and additional trials will refine your technique to get more
precise results.
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Mr. Murdoch
 The volume of added base (or acid if titrating acid into base instead of
base into acid) can be used along with the original volume of the acid
in the flask as well as the acid’s molarity to determine the molarity of
the unknown.
 The point in the titration where the color of the phenolphthalein
indicator barely changes to a just noticeable pink is called the
ENDPOINT of the titration. Since we are adding a base to an acid,
one might think that the neutralization point would be near the
neutral pH value 7. Phenolphthalein actually changes color a little
higher on the pH scale, around pH value 8.2, a bit more basic than
neutral.
 If you want to find the equivalence point, or the amount of added base
that brings the solution in the flask to exactly pH value 7, the best
way is with a pH probe. Of course, multiple trials are usually the best
way to get closest to either an endpoint or an equivalence point in a
titration.
 Here is an online Titration Simulation where you can practice
working with titration calculations. Note that the “endpoint” in the
simulation occurs immediately, and is NOT the way it would work in
the lab.
http://group.chem.iastate.edu/Greenbowe/sections/projectfolder/flashfiles/stoichiometry/a_b_phtitr.html
Online Titration Simulation
Watch Crash Course Chemistry Buffers (w/ Titration) YouTube video - 11:40
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Mr. Murdoch
Titration Equations:
i. Remember, one mole of H+ neutralizes one mole of OH-.
ii. Therefore, # of moles of H+ = # of moles of OHiii. Since we are dealing with solutions and molarities, use the formula:
M=
𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒
𝐿 𝑜𝑓 𝑠𝑜𝑙𝑣𝑒𝑛𝑡
(remember this equation is “Concentration” on Reference Table T!)
iv. Rearrange the equation to:
moles =
𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒
𝐿
x liters.
v. Liters are our standard unit of volume, and
𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒
𝐿
is molarity
(M), so the equation becomes:
moles = M x V.
vi. Moles of acid = Macid x Vacid and moles of base = Mbase x Vbase
vii. These provide the formula:
MAcid x VAcid = MBase x VBase
viii. We can abbreviate this to:
MA x VA = MB x VB
ix. A pretty simple equation to remember, but if not, it is on Reference
Table T for your use anytime as: MAVA = MBVB
x. This titration equation ONLY works if the acid used has only one H+
per acid molecule, and ONLY if the base used has only one OH- per
base molecule. We can then modify the titration equation by adding
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Regents Chemistry ’14-‘15
Mr. Murdoch
coefficients for both the number of hydrogens in an acid (# of H) and
also the number of hydroxides in an acid (# of OH).
# H MA VA = # OH MB VB
xi. If you are only solving for moles of solute instead of molarity, the
formula may be simplified to:
# H MolesA = # OH MolesB
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Topic: Titration Problem Examples
Objective: How may we calculate the molarity of an unknown?
Titration Problem examples:
1. If it takes 15.0 mL of 0.40 M NaOH(aq) to neutralize 5.0 mL of
HCl(aq), what is the molar concentration of the HCl(aq) solution?
i. Since we are given both molarity and volume of the base, and
volume of the acid, use the equation: # H MA VA = # OH MB VB
ii. # H MA VA = # OH MB VB rearranged to solve for the molarity of
the acid is:
MA = # OH MBVB / # H VA
[(1)𝑥 (0.40 𝑀)𝑥 (15.0 𝑚𝐿)
= 1.2 M HCl(aq)
[(1)𝑥 (5.0 𝑚𝐿)
2. If it takes 10.0 mL of 2.0 M H2SO4(aq) to neutralize 30.0 mL of
KOH(aq), what is the molar concentration of the KOH(aq) solution?
i. Since we are given both molarity and volume of the acid, and the
volume of the base, use the equation: # H MA VA = # OH MB VB
ii. # H MA VA = # OH MB VB rearranged to solve for molarity of the
base is:
MB = # H MAVA / # OH VB
[(2) x (2.0 M) x (10.0 mL)]
[(1) x (30.0 mL)]
= 1.3 M KOH(aq)
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Unit 11: Acids & Bases-Lecture
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Mr. Murdoch
3. How many mL of 2.0 M H2SO4(aq) are required to neutralize 30.0 mL
of 1.0 M NaOH?
i. Since we are given both molarity and volume of the base, and the
molarity of the acid, use the equation: # H MA VA = # OH MB VB
ii. # H MA VA = # OH MB VB rearranged to solve for volume of the
acid is:
VA = # OH MBVB / # H MA
[(1) x (1.0 M) x (30.0 mL)]
[(2) x (2.0 M)]
= 7.5 mL of 2.0 M H2SO4(aq)
4. How many mL of 0.10 M Ca(OH)2(aq) are required to neutralize 25.0
mL of 0.50 M HNO3(aq)?
i. Since we are given both molarity and volume of the acid, and the
molarity of the base, use the equation: # H MA VA = # OH MB VB
ii. # H MA VA = # OH MB VB rearranged to solve for volume of the
base is:
VB = # H MAVA / # OH MB
[(1) x (0.50 M) x (25.0 mL)]
[(2) x (0.10 M)]
= 63 mL of 0.10 M Ca(OH)2(aq)
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Regents Practice Questions-Titration: (Ungraded)
1. A student neutralized 16.4 mL of HCl(aq) by adding 12.7 mL of 0.620
M KOH(aq). What was the original molarity of the HCl(aq)?
a) 0.168 M
c) 0.620 M
b) 0.480 M
d) 0.801 M
2. How many mL of 0.600 M H2SO4(aq) are required to exactly neutralize
100. mL of 0.300 M Ba(OH)2(aq)?
a) 25.0 mL
c) 100. mL
b) 50.0 mL
d) 200. mL
3. The pH of a solution that is formed by the neutralization of 1.0M
H2SO4(aq) and 1.0 M KOH(aq) is closest to
a) 1
b) 4
c) 7
d) 10
4. If equal volumes of 0.1 M NaOH(aq) and 0.1 M HCl(aq) are
combined, the resulting solution will contain a salt and
a) H2O(l)
c) NaCl(aq)
b) HCl(aq)
d) NaOH(aq)
5. The following data were collected by a student performing an acidbase titration:
Volume of aqueous acid (HCl) used:
20.0 mL
Molarity of aqueous acid (HCl) used:
0.50 M
Volume of aqueous base (NaOH) used:
40.0 mL
From the collected data, the concentration of the base should be
calculated as
a) 1.0 M
b) 2.0 M
c) 0.25 M
d) 0.50 M
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Student name: _________________________
Class Period: _______
Please carefully remove this page from your packet to hand in.
Acid and Base Neutralization Homework
Write the formula of the salt formed from each reaction below. (1 pt. ea.)
1. H2SO4(aq) + Mg(OH)2(aq)  2 H2O(l) + ______________________________ (___)
2. H2CO3(aq) + 2 KOH(aq)  2 H2O (l) + ______________________________ (___)
Write the formula of the unknown used in each reaction and balance. (4 pts. ea.)
3. ____________________(aq) + ___ Al(OH)3(aq)  ___ H2O (l) + ___ Al2(SO4)3(aq)
4. ___ HCl(aq) + ___________________________(aq)  ___ CaCl2(aq) + ___ H2O(l)
Solve the following titration problems, showing ALL steps. (2 pts. ea.)
5. How many moles of KOH(aq) are needed to completely neutralize 1.5 moles of
H2SO4(aq)?
Cont’d on back:
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
6. What volume of 5.0 M NaOH(aq) is needed to neutralize 40. mL of 2.0 M
HCl(aq)?
7. What is the molarity of a NaOH(aq) solution if it takes 100. mL of a NaOH to
completely neutralize 50. mL of 0.10 M H2SO4(aq)?
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Mr. Murdoch
37
g of NaCl x 3.55 = 131.55, or about
130 grams of NaCl are soluble in
Topic:
pH
355 grams of water at 10°C
Objective: What do we mean by the expression of pH?
Power of Hydronium Ion in a Solution:
 The unit pH is a measure of the hydrogen ion concentration in an
aqueous solution. (paired [ ] means concentration)
i. Pure water is in equilibrium (remember that?) where a small
amount of water molecules dissociate from HOH to form H+ and
OH- ions.
ii. H2O  H+ + OH-, and then the free H+ form a coordinate bond
with another H2O molecule to form a H3O+ hydronium ion.
iii. This is a SMALL number; in neutral water, the concentration of H+
= 1.0 x 10-7M, so the pH of pure water is given as 7.
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Unit 11: Acids & Bases-Lecture
Topic:
Regents Chemistry ’14-‘15
Mr. Murdoch
Adding Acids to Water
Objective: What happens to pH when we add an acid to water?
Adding Acids to Water:
 When adding an acid to pure water, the added acid increases the
H3O+ concentration, so that the pH increases by one with each tenfold
increase in acid strength. (paired [ ] means concentration)
i. If the concentration of H3O+ = 10-1 M, the pH value is 1 (absolute
value of the exponent equals the pH value for an acid)
ii. If the concentration of H3O+ = 10-3 M, the pH value is 3
iii. If the concentration of H3O+ = 10-9 M, the pH value is 9
iv. If the concentration of H3O+ = 10-11 M, the pH value is 11
v. If the concentration of H3O+ = 5 x 10-6 M, the pH value is 6.5
vi. If the concentration of H3O+ = 2 x 10-2 M, the pH value is 2.2
vii. A solution with a pH value 3 is tenfold more acidic than a solution
with a pH value 4.
viii. A solution with a pH value 3 is a thousand fold more acidic than a
solution with a pH value 6.
ix. A solution with a pH value 4.4 is four times more acidic than a
solution with a pH value 4.
x. A solution with a pH value 4.7 is seven times more acidic than a
solution with a pH value 4.
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Unit 11: Acids & Bases-Lecture
Topic:
Regents Chemistry ’14-‘15
Mr. Murdoch
Adding Bases to Water
Objective: What happens to pH when we add a base to water?
Adding Bases to Water:
 When adding a base to pure water, the added base increases the OHconcentration, so that the pH increases by one with each tenfold
increase in base strength. (paired [ ] means concentration)
i. If the concentration of OH- = 10-1 M, the pH value is 13 (value of the
exponent subtracted from 14 equals the pH value for an base)
ii. If the concentration of OH- = 10-3 M, the pH value is 11
iii. If the concentration of OH- = 10-5 M, the pH value is 9
iv. If the concentration of OH- = 5 x 10-6 M, the pH value is 7.5
v. If the concentration of OH- = 7 x 10-5 M, the pH value is 8.3
vi. A solution with a pH value 9 is tenfold more basic than a solution
with a pH value 8.
vii. A solution with a pH value 11 is a thousand fold more basic than a
solution with a pH value 8.
viii. A solution with a pH value 8.8 is eight times more basic than a
solution with a pH value 8.
ix. A solution with a pH value 11.2 is two times more basic than a
solution with a pH value 11.
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Comparing Common examples of pH values:
Watch Bozeman Chemistry Acids, Bases, & pH YouTube video - 8:53
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Regents Practice Questions-pH: (Ungraded)
1. Which statement below correctly describes a solution with a pH of 9?
a) It has a higher concentration of H3O+ than OH- and causes red litmus paper
to turn blue
b) It has a higher concentration of OH- than H3O+ and causes red litmus paper
to turn blue
c) It has a higher concentration of OH- than H3O+ and causes methyl orange to
turn red
d) It has a higher concentration of H3O+ than OH- and causes methyl orange to
turn yellow
2. Which pH change represents a hundredfold (100x) increase in the concentration of
H3O+?
a) pH 3 to pH 1
c) pH 5 to pH 7
b) pH 4 to pH 3
d) pH 13 to pH 14
3. Given the following solutions:
i. Solution A: pH of 5
ii. Solution B: pH of 7
iii. Solution C: pH of 10
Which list below has the solutions placed in order of increasing H+
concentration?
a) A  B  C
b) C  A  B
c) B  A  C
d) C  B  A
4. Which relationship is present in a solution with a pH of 7?
a) [H+] = [OH-]
c) [H+] < [OH-]
b) [H+] > [OH-]
d) [H+] + [OH-] = 7
5. What is the H3O+ ion concentration of a solution that has an OH- ion concentration
of 1.0 x 10-3 M?
a) 1.0 x 10-3M
c) 1.0 x 10-11M
b) 1.0 x 10-7M
d) 1.0 x 10-14M
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Mr. Murdoch
Notes page:
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Student name: _________________________
Class Period: _______
Please carefully remove this page from your packet to hand in.
pH Homework
Circle the correct answer for each multiple choice question. (1 pt. ea.)
1. Which of the following pH values is the most acidic?
a) 5
b) 7
c) 9
d) 11
2. Which of the following pH values is the most basic?
a) 5
b) 7
c) 9
d) 11
3. Which of the following 1 x 10-5 M in aqueous solution could have a pH of 5?
a) CH4
b) HCl
c) NaCl
d) NaOH
4. What would be the pH of a solution of made from equal volumes of 0.1 M
HCl(aq) and 0.1 M NaOH(aq)?
a) 0.2
c) Less than 7
b) Exactly 7
d) Greater than 7
5. Which pH value would indicate a solution as the better acidic electrolyte?
a) 3
b) 5
c) 9
d) 12
6. Which pH value would indicate a solution as the poorer basic electrolyte?
a) 3
b) 5
c) 9
d) 12
Short answer questions. (1 pt. ea.)
7. A pH of 4 is how many times more acidic than a pH of 6? __________
8. A pH of 11 is how many times more acidic than a pH of 8? __________
9. Neutral pH is how many times more acidic than a pH of 2? __________
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Mr. Murdoch
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Topic: Brönsted-Lowry Acids & Bases
Objective: Describe acids and bases using only H+ ions.
Brönsted-Lowry Acids and Bases:
The discussion we have had so far is concerning the Svante Arrhenius
theory on what makes an acid and acid, and a base a base. Two other
scientists, Johannes Brönsted and Thomas Lowry published separate
(but almost simultaneously) alternative theories about a more general
theory of acids and bases. For the Brönsted-Lowry theory of acids and
bases, a specific subset of what Brönsted-Lowry call acids and bases
include the Arrhenius acids and bases, so the Arrhenius classification is
more restrictive as what makes an acid or a base than the BrönstedLowry classification.
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Unit 11: Acids & Bases-Lecture
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Mr. Murdoch
Topic: Comparing Acid-Base Theory
Objective: How do Arrhenius and Brönsted-Lowry theories compare?
1. According to Arrhenius:
i. Arrhenius Acid: A compound that dissociates in water to produce H+
that is the ONLY positive cation in the solution.
ii. Arrhenius Base: A compound that dissociates in water to produce
OH- as the ONLY negative anion in the solution.
2. According to Brönsted-Lowry:
i. Brönsted-Lowry Acid: Any substance that donates a proton (H+).
a) A neutral hydrogen atom (11H0) has only one proton and one
electron to start with. Remove the e-, and all you have is the
proton, which we write as H+. The Brönsted-Lowry theory states
any substance that gives a H+ cation (a proton) is therefore a
Brönsted-Lowry acid.
ii. Brönsted-Lowry Base: Any substance that accepts a proton (H+).
+
b) The Brönsted-Lowry theory states any substance that accepts a H
cation (a proton) is therefore a Brönsted-Lowry base.
Watch The Fuse School Bronsted-Lowry Theory video - 3:55
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Unit 11: Acids & Bases-Lecture
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Mr. Murdoch
Brönsted-Lowry theory examples:
1. HCl(g) + H2O(l)  H3O+1(aq) + Cl-1(aq)
Acid
Base
a) The HCl donated its H+ to the H2O forming H3O+
b) As HCl lost (donated) an H+, it is the Brönsted-Lowry acid. As
H2O gained (accepted) the H+, it is the Brönsted-Lowry base.
2. NH3(g) + H2O(l)  NH4+1(aq) + OH-1
Base
Acid
a) The H2O donated an H+ to the NH3 forming NH4+
b) As H2O lost (donated) an H+, it is the Brönsted-Lowry acid. As
NH3 gained (accepted) the H+, it is the Brönsted-Lowry base.
Yup…water may act as both acid and base. Fun, eh?
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
3. HC2H3O2(aq) + H2O(l)  H3O+1(aq) + C2H3O2-1(aq)
Acid
Base
a) The HC2H3O2(aq) donated an H+ to the H2O
+
b) As HC2H3O2 lost (donated) an H , it is the Brönsted-Lowry acid.
As H2O gained (accepted) the H+, it is the Brönsted-Lowry base.
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Regents Practice - Brönsted-Lowry Acids and Bases: (Ungraded)
1. Given the reaction of NH3(g) + H2O(l)  NH4+(aq) + OH-(aq):
Water acts in this reaction as the
a) Acid
b) Base
c) Electron donor
d) Proton acceptor
2. According to the “alternative theory” of acids and bases, an acid is
any species that
a) Releases oxide ions into solution
b) Donates protons to another species
c) Releases hydroxide ions into solution
d) Accepts protons from another species
3. In the reaction NH3 + HCl  NH4+1 + Cl-1:
The NH3 acts in this reaction as
a) A Brönsted-Lowry acid, only
b) A Brönsted-Lowry base, only
c) Both a Brönsted-Lowry acid and a Brönsted-Lowry base
d) Neither a Brönsted-Lowry acid nor a Brönsted-Lowry base
4. In which forward reaction below is water acting only as a proton
acceptor?
a) H2O(l) + H2O(l)  H3O+(aq) + OH-(aq)
b) NH3(g) + H2O(l)  NH4+(aq) + OH-(aq)
c) H2SO4(aq) + H2O(l)  HSO4-1(aq) + H3O+1(aq)
d) CH3COO-1(aq) + H2O(l)  CH3COOH(aq) + OH-(aq)
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Mr. Murdoch
Notes page:
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Unit 11: Acids & Bases-Lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Student name: _________________________
Class Period: _______
Please carefully remove this page from your packet to hand in.
Brönsted-Lowry Acids and Bases Homework
For each of the following equilibrium systems, identify the Brönsted-Lowry acid
and the Brönsted-Lowry base on the reactant side. Place a capital “A” over the
Brönsted-Lowry acid, and a capital “B” over the base.
1.
HBr + H2O  H3O+ + Br-
2.
H2O + H2O  H3O+ + OH-
3.
NH3 + OH-  NH2- + H2O
4.
H2O + HPO4-2  PO4-3 + H3O+
5.
H3PO4 + H2O  H2PO4- + H3O+
6.
CH3COO- + H3O+  HCH3COO + H2O
7.
H2PO4- + CH3COO-  HCH3COO + HPO4-2
8.
H2O + S-2  HS- + OH-
9.
CN- + HCH3COO  HCN + CH3COO-
10.
OH- + NH4+  H2O + NH3
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Mr. Murdoch
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Mr. Murdoch
Notes page:
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