Unit 5 Acid-Base Equilibrium Flashcards Preview

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Flashcards in Unit 5 Acid-Base Equilibrium Deck (41):

Top 6 Strongest Acids (Bronsted-Lowry)

Perchloric acid HClO4
Hydriodic acid HI
Hydrobromic acid HBr
Hydrochloric acid HCl
Nitric acid HNO3
Sulphuric acid H2SO4


Naming acids

Binary Acids (hydrogen bonded to a simple ion)
hydro- + root of anion + -ic acid
**No oxygen

ex: HCl hydrochloric acid
Root of anions:
bromic (Br)
Iodic (I)
Sulfuric (S)
Phosphoric (P0
Fluoric (F)

Ternary Acids (hydrogen bonded to a complex/polyatomic ion)
Root of anion + -ic acid
**With oxygen (no hydro)

ex: HClO3 Chloric acid
Root of anions:
ClO3 Chloric
BrO3 Bromic
IO3 Iodic
SO4 Sulfuric
PO4 Phosphoric
NO3 Nitric
CO3 Carbonic

Varied Number of Oxygen Atoms:
Per- (one more)
-ous (one less)
hypo- -ous (two less)

Acids Containing COO- (organic acids) may have the hydrogen placed at the end i.e. Acetic acid, Benzoic acid


Naming bases

Typically bonded with hydroxide

Exceptions: Ammonia (NH3)


Definition of Acids and Bases - Arrhenius

Acids: substances that ionize in water to produce hydrogen ions (H+) or free protons

Bases: substances that ionize in water to produce hydroxide ions (OH-)

Although the Arrhenius Theory was correct, NOT ALL acids contain hydrogen ions and not all bases contain hydroxide ions.
Ex: NH3 and NaHCO3 (sodium bicarbonate)
Ex: CO2


Definition of Acids and Bases - Bronsted-Lowry

Acid: substance that DONATES (loses) a proton (H+)

Base: substance that ACCEPTS (gains) a proton (H+)

*all acids and bases included in the Arrhenius Theory are also bases and acids according to Bronsted-Lowry


How to identify acids and bases in a chemical reaction (Bronsted-Lowry)


Identify which substance contains a hydrogen ion. The acid donates the hydrogen ion to the base.

If both substances contain hydrogen, look at products to determine which substance loses the H+

If required to predict, then must use the Ka chart (substance can act as either an acid or base - amphoteric)
-look at either acid column or base column and identify two substances
-The stronger substance (for acid or base) is more prevalent

*the H+ or a free proton is never independent -->forms H3O+ (hydronium ion)
-->for convenience, usually use simplified version H+


Conjugate Acid-Base Pairs


ALWAYS have a conjugate acid-base pair

Conjugate Base of an acid is the particle that REMAINS after a proton has been released by the acid
-result of an acid donating proton

Conjugate Acid of a base is formed when a base ACQUIRES a proton that has been released by the acid
-result of a base accepting a proton

A --> Cb
B --> Ca

**Check by seeing the reverse reaction (Cb accepts while Ca donates)

***Conjugate A-B are related by the loss or gain of a single H+ (transfer)
ex: NH3 + H3O+ NH4+ + H2O

NH3 (b) and NH4+ (Ca)
H3O+ (a) and H20 (b)


Transferring of H+/Proton

Acids lose proton (become less positive): Neutral becomes negative, negative becomes more negative

Bases accept proton (become more positive): Neutral becomes positive, negative becomes less negative/neutral

**NOT an electron; PROTON (+)


Characteristics of amphoteric substances

Substances that act as either an acid or a base
-neutral or charged
(both donate and accept H+)

Must have a Hydrogen ion (neutral) or hydrogen ion with negative charge (charged)

Ex: H2O, NH3 -->amphoteric molecules
HCO3-, HSO4, HCO3-, HS- -->amphoteric ions


Bronsted-Lowry Acid and Base Reactions (predicting reactions)

Reactions represent a competition between two bases for a proton (H+)

Goal: to determine which direction of reaction will be favoured (forward or reverse)

1. Identify reactants as acid or base
-if both can be either (amphoteric), then use Ka chart (shows relative strengths of acids) and focus on acid column (left) or base column (right)
-the substance that is stronger is more prevalent

2. Predict products knowing the acid and base.
-acid will donate a proton and the base will accept a proton
-label conjugate acid and conjugate base

3. Determine direction the reaction will proceed
-compare either both bases (b and Cb) or both acids (a and Ca)
-the strong acid and strong base will always be on the same side of the equation and the weak acid and weak base will always be on the same side of the equation
-the reaction will proceed in the direction of the weaker acid/base pair
*competing for proton (stronger acid/base will win, therefore shift towards weaker)


Strengths of Acids and Bases

Strength (strong or weak) depends on the degree to which they IONIZE in water

*Electrolyte: a substance which conducts electricity in a solution due to the presence of ions

Strong Electrolytes:
-completely ionize/dissociate as they dissolve in water (0% s, 100% ions)
-excellent conductors of electricity
Ex: HCl, NaOH, LiCl

Weak Electrolytes: -->used for Ka and Kb calculations
-partially ionize/dissociate as they dissolve in water (99% s, 1% ions)
-poor conductors of electricity
ex: acetic acid, NH3 + H20

Non-Electrolytes: substances that do not dissociate into ions and therefore do not conduct electricity
ex: sugar


Classifying Acids and Bases
(Strong vs. Weak and Dilute vs. Concentrated)

By solution concentration - dilute or concentrated (amount per L)
Concentrated: high # of moles
Dilute: low # of moles

By extent of ionization - strong or weak
Strong: 100% dissociation
Weak: low % dissociation

ex: 0.001 M of HCl is dilute and strong


Strong bases

all group 1 hydroxides (NaOH, LiOH, KOH etc.)
Last group 2 hydroxides (Ca(OH)2, Sr(OH)2, Ba(OH)1)


Weak acids

anything that is not in the top 6

Ex: vinegar/acetic acid, carbonic acid


Weak bases

ammonia derivatives

ex: NH3, CH3NH2


Distinguishing between strong and weak acids and bases

1) pH Measurement - *same concentration (equimolar)
Strong acids have LOWER pH than weak acids
Strong bases have HIGHER pH than weak bases
Why? Because strong acids produce a higher concentration of H+ ions in the solution than a weak acid.

2) Conductivity Measurement
Strong acids and bases: higher readings on a conductivity meter than weak acids and bases of equal concentration (due to high ionization; contain more ions)


Self-ionization of Water

Water molecules are polar and occasionally collide with enough energy to react - self-ionization
-One water molecule acts as the acid and one water molecule acts as the base.
-->hydrogen ion is transferred from on water molecule to another, forming hydronium ion and OH-

Simplified: H2O H+ + OH-
Neutral solution contains equal concentration of both ions (1x10^-7)
-->ions are interdependent (as one increases the other decreases) -->Le Chatelier's principle

Addition of [H+] or [OH-] shifts the equilibrium (reverse), thus decreasing the concentration of the other ion
-pH will change as well
More [H+] than [OH-] = lower pH (acidic)
More [OH-] than [H+] = higher pH (basic)
if equal concentration = neutral


Ion Product Constant for water, Kw

Kw = [H+][OH-] in pure water
Kw always equal to 1x10^-14 at 25 degrees for every aq solution
(1x10^-7)(1x10^-7) -->Related to pH scale (7 for each)

-no units



Every solution is neutral, acidic, or basic

-->Compare the concentrations of [H+] and [OH-] to determine if solution is neutral, acidic, or basic

More [H+] than [OH-] = lower pH (acidic)
More [OH-] than [H+] = higher pH (basic)
if equal concentration = neutral

pH scale (potency of hydrogen) - base 10 scale
1 = 10^-1
2 = 10^-2
(as pH value increases, the [H+] decreases)-->more basic


Kw calculations

Kw = 1.0 x 10^-14 at 25oC

Kw = [H+][OH-]

Rearrange Kw expression to find concentrations-->compare to determine if the solution is acidic, basic, or neutral.
**Need a statement and units for concentration

Question types:
-given [H+] or [OH-], asked to find other concentration (rearrange Kw)
-given the concentration of a solution with either [H+] or [OH-], asked to solve the other concentration (mole ratio to find concentration of one, then rearrange Kw)
-given solid with mass dissolved in a given volume (containing either H+ or OH-), asked to find the [H+] and [OH-] (convert to moles m/mm, find concentration of solution, dissociation/mole ratio, rearrange Kw)


Rainbow Reaction tube experiment

Universal indicator changed colour to show the pH of the substance.
-Universal indicator in water (neutral - green)
-added equal amounts of acid (red) and base (blue/purple)

Less than 3 (strong acid) = red
3-6 (weak acid) = orange/yellow
7 (neutral) = green
8-11 (weak base) = blue
Greater than 11 (strong base) = purple

**4 is orange


pH Calculations

pH = - log [H+]
[H+] = inverse log (-pH) OR = 10^-pH

pOH = - log [OH-]
[OH-] = inverse log (-pOH) OR = 10^-pOH

pH + pOH = 14

pH for [H+] = 1.0 x _____ = exponent
ex: [H+] = 1 x 10^-8 pH = 8

Two step problems:
no formula connecting [OH-] and pH
-use Kw first to solve for [H+]. Then solve for pH
-solve for pOH and then pH + pOH = 14


Relationship between pH and Concentration

**pH and pOH not concentration!
pH = 13
= 10^-13 (low concentration of H+)
Therefore basic

If the concentration of H+ is increased by a factor of ten, then the pH decreases by one unit
Ex; 10^-4 to 10^-3 -->4 to 3

If the concentration of OH- is increased by a factor of ten, the hydrogen ions decreased by a factor of ten and then the pH INCREASES by one unit
Ex: 10^-6 to 10^-5 OH- means 10^-8 to 10^-9 = 9

Hydrogen ions decrease -->pH increases
Hydrogen ions increase -->pH decreases

pH changes by one for every 10-fold change in [H+]
o Ex: solution with pH 3 has a [H+] that is 10x greater than a solution with pH 4 and 100x greater than a solution with pH 5

**pH changes by a factor of one when [H+] changes by a factor of ten**


Ionization Equations: Acids

*all equations should produce one H+ per step*
*state aq*
-Ionize to produce positive and negative ion

Monoprotic Acids: produce one H+ per molecule as they ionize
Ex: HCl

Diprotic Acids: produce two H+ per molecule as they ionize
Ex: H2SO4

Triprotic Acids: produce three H+ per molecule as they ionize
Ex: H3PO4

Polyprotic acids (di- and tri-) acids ionize in steps (produce one H+ per step) --> STEPWISE Ionization
--> similar to a reaction mechanism. The molecule that loses the H+ is similar to the intermediate (goes to next step and continues to ionize).

There is a Ka expression for every step (for weak acids). Use the net dissociation to see if correctly ionized in steps

Ex: H2SO4 should have 2H+ in net


Ionization Equations: Bases

-ionize to produce positive and negative ion
*state aq*

Strong Bases (have hydroxide, therefore do not need water):
CaOH(aq) --> Ca2+(aq) + 2OH-(aq)

Weak Bases (that do not have hydroxide need to show reaction with water, which acts as the acid):
->dissociates into OH- and other positive ion


When writing ionization equations...

-Label substance as strong or weak acid/base
-react ammonia derivative bases with water
-polyprotic acids ionize in steps

Write Ka expressions for weak acids and Kb expressions for weak bases

**strong acids and bases do not expressions because their values are very large (due to high percentage of ionization)

*Weak A/B have double arrow
*Strong A/B have single arrow (complete ionization)


Percentage Ionization for acids and bases

= ( [H+]/[acid]i ) x 100

= ( [OH-]/[base]i ) x 100

**remember that acids and bases both have hydrogen and hydroxide ions (but very small value due to water's self ionization).


Acid dissociation constant, Ka
(Acid Ionization Constant)

Reflects the fraction of acid that is in the ionized form
*based on ionization equation (No Liquids)

Ka = products/reactants
*no exponents (acids ionize one H+ at a time

*for polyprotic acids there is a Ka expression and value for each step


Base dissociation constant, Kb
(Base Ionization Constant)

Reflects the fraction of the base that is in the ionized form
*based on ionization equation (no liquids)

Kb = products/reactants


Ka and Kb values

A small Ka or Kb value indicates the acid or base ionizes only slightly

Small Ka/Kb = small % ionization
Large Ka/Kb = large % ionization

*since strong acids and bases ionize completely and thus have a 100% ionization, the Ka/Kb is not usually calculated.


Ka and Kb calculations (including % ionization, Kw, pH, pOH)

*ALWAYS identify if strong or weak acid/base

Problem type 1: Strong Acids and Bases
-Dissociate 100%. Therefore ion concentrations are derived from the concentration of the acid/base dissociated (using stoichiometry)

Problem type 2: Weak Acids and Bases (finding Ka/Kb)
Write ICE table to identify what you have and what you need
-To calculate Ka/Kb, you need the [H+] for acids and [OH-] for bases => sub into Kb expression knowing initial concentration (ion concentration always have 1:1:1 ratio)

Problem type 3: Weak Acids and Bases (finding concentration)
Write ICE table with X for change
-sub X into Ka/Kb expression and solve for X knowing initial concentration (1:1:1: ratio)

**you can ignore the change for the solution concentration because the change is very small compared to the concentration.

**You can also calculate...
-Initial concentration by rearranging the Ka/Kb expression once calculated ion concentration
-% Ionization
-pH and pOH

^can use info to solve for each problem type as well

**if question gives you % ionization = weak acid/base
**ICE table always is 1:1:1 and 0 for initial ion concentration


Summary for Ka, Kb, Strong and Weak acids/bases

Stronger acid => higher % ionization => higher [H+] in solution => larger Ka

Weaker acid => lower % ionization => lower [H+] in solution => smaller Ka

Stronger base => higher % ionization => higher [OH-] in solution => larger Kb

Weaker base => lower % ionization = lower [OH-] in solution => smaller Kb


Titration: what is it?

Titration is a process of neutralization commonly used to determine the amount of acid or base in a solution

The process involves delivering a KNOWN concentration from a burette into an UNKNOWN solution containing an INDICATOR until the substance being analyzed is just consumed (equivalence point)

Standard solution/Titrant: known solution

Equivalence point: the point at which the acid has completely reacted with or been neutralized by the base
[mol H+] = [mol OH-]
-->signaled by colour change
-->midpoint of steep part

Half-equivalence point: half of volume required to neutralize

End point: the pH at which the indicator changes colour (JUST PAST the equivalence point)


Titration pt 2 (info gathered from titration)

Volume of titrant used is recorded (Vf - Vi), which can be used to calculate MOLES of TITRANT

Use mole ratio to calculate other info on unknown (i.e. concentration)

Can plot the pH of the solution analyzed to create a Titration Curve (y = pH x = volume of titrant added)



Reactions between an acid and base

Molecular (double displacement and balanced), ionic (show dissociation of ions) and net ionic equations (without spectator ions) -->include physical states

Calculations: The word neutralize = chemical reaction with stoichiometry
*diluting and mixing are physical changes

-Write BALANCED chemical equation
-Ensure that volume is in L
-Use C=n/v, n=m/mm other formulas to calculate unknown
-Use stoichiometry and mole ratio

**Differentiate between a single solution of acid or base (would be pH, Ka, Kb % Dissociation problems)


How to differentiate between Bronsted-Lowry and neutralization

SA + SB => WA + WB (Bronsted-Lowry - will say in question)

Other reacting acid and base together = neutralization (will say in question)
*titration is a form of neutralization

If just single acid or base, then use ionization equation


Interpreting Titration Curves

X axis: volume of titrant in mL
Y axis: pH

1) Identify the type of titration first
*Strong acid with strong base (pH low, equivalence point has a pH of 7)

*Weak acid with strong base (pH low, equivalence point has a basic pH, meaning base is STRONGER)

*Weak base with strong Acid (pH high, equivalence point has an acidic pH, meaning acid is STRONGER)

**weak acid and weak base titrations are too complicated (i.e. curve) and are almost never carried out

2) Write the balanced equation

3) From the graph you can determine volume of titrant required to neutralize (equivalence point is the midpoint of the steep part), the half equivalence point (half the volume required to neutralize) and pH

2) Typically given either mass or volume of unknown, concentration of titrant/standard solution
*Water does not change the number of moles (therefore ignore)

Typical Questions:
-Find the concentration of the unknown solution => use stoichiometry to calculate moles of unknown using moles of titrant, then solve for concentration)
*stoichiometry works because chemically related

-Suitable indicator (using chart) =>locate one drop past equivalence point (end point) and identify the range. Choose the pH that will change DURING this range or AFTER Equivalence point (not before equivalence point)

-Ka and Kb =>
Ka = [H+] at half equivalence point (convert pH at 1/2eqv point to [H+])

Kb = [OH-] at half equivalence point (pH to pOH to OH-)

*Kw = KaKb

-Molar mass (mm = m/n)



Indicators are weak organic acids in which the ACID and its CONJUGATE BASE are different colours. Indicators are used to identify the nature of a substance as acidic, basic or neutral

Common indicators:
Phenolphthalein - colourless in acid; pinky-purple in base
Litmus - turns red in acid; turns blue in base
Methyl Orange - orange in acid; yellow in base
Bromothymol Blue - yellow in acid; blue in base
Universal indicator (general) =>
0 - red
4 - orange
7 - green
10 - blue
14 - purple

*Choose indicator whose end point range lies on the steep part of the titration curve

Indicators are also related to Le Chatelier's Principle:
Presence of an acid increases [H+ also H3O+]

Presence of a base increases [OH-] and decreases [H+ also H3O+]


What are buffers?

Buffers consist of two main solutes that limit large changes in pH through neutralization

Buffer = to protect

Buffer solution is made from EQUAL moles of weak acid and its conjugate base (supplied by a soluble salt) or a weak base and its conjugate acid (supplied by a soluble salt)

*Strong acids and bases ionize 100%, therefore have a lot of conjugate ions but there is no way to have the other form due to dissolving (CANNOT BE USED AS BUFFERS)

One solute neutralizes acids and one solute neutralizes bases

Strong acids are sources of H+
Strong bases are sources of OH-

The product of the net ionic reaction is the opposite and water (neutral substance which maintain the pH)

The buffers are regenerated when strong acid and base is added to the buffer!

However, there is still a buffer capacity
(only protects a solution from small additions of strong acid or base. Increasing the concentration of the buffer team increases the capacity of the buffer)


Buffers: why do you need a soluble salt?

o Since weak acids and bases do not ionize a lot, a soluble salt containing the ion is required to have equal concentration
• Na+
• Cl-
• *Ignore spectator ions


Acidic and basic buffers

Acidic buffers: Acid and its conjugate base (supplied by a soluble salt)

Basic buffers: Base and its conjugate acid
*always ammonia derivatives (because OH- are strong bases)