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AP Chemistry 8.1 Introduction to Acids and Bases Study Notes

AP Chemistry 8.1 Introduction to Acids and Bases Study Notes - New Syllabus Effective fall 2024

AP Chemistry 8.1 Introduction to Acids and Bases Study Notes- New syllabus

AP Chemistry 8.1 Introduction to Acids and Bases Study Notes – AP Chemistry –  per latest AP Chemistry Syllabus.

LEARNING OBJECTIVE

Calculate the values of pH and pOH, based on Kw and the concentration of all species present in a neutral solution of water.

Key Concepts: 

  • pH & pOH
  • The Ionic Product of Water
  • pH & pOH of Strong Acids & Bases

AP Chemistry-Concise Summary Notes- All Topics

AP Chemistry Exam Questions MCQs and FRQs

8.1.A.1 pH, pOH, and the Hydronium Ion (H₃O⁺):

1. pH and pOH Definitions:

pH is a measure of the acidity of a solution, which is defined by:

pH=−log[H3​O+]

A lower pH means higher acidity (more hydronium ions).

pOH is a measure of the basicity of a solution, which is defined by:

pOH=−log[OH−]

A lower pOH means higher basicity (more hydroxide ions).

Both are logarithmic scales, thus a unit change represents a tenfold change in ion concentration.

2. Water Ionization and Kw:

i. Water Ionizes Slightly:
In pure water (or any aqueous solution), a small number of water molecules dissociate:

H2​O (l)⇌H3​O+(aq)+OH−(aq)

ii. Ion-Product Constant of Water (Kw):
Hydronium concentration times hydroxide ion concentration equals 1.0 x 10^(-14) at 25°C:

[H3O+][OH−]=1.0×10−14

This value is referred to as Kw, the ion-product constant.

iii. pH and pOH Relationship:
Because:

pH=−log[H3​O+]                         pOH=−log[OH−]             

Taking the negative log of both sides of Kw results in:

pH+pOH=14

This correspondence is especially useful when you have either pH or pOH and need to solve for the other.

3. Hydronium vs. Hydrogen Ion:

H⁺ (Hydrogen Ion) does not occur by itself in water because it is merely a proton—really small and extremely reactive.

H⁺ immediately bonds to a water molecule to create the more stable hydronium ion (H₃O⁺):

H++H2​O→H3​O+

H₃O⁺ is used by preference in chemistry to refer to the actual species in aqueous solutions.

However, for practical purposes, H⁺(aq) is commonly used and accepted—particularly on examinations like the AP Chemistry exam.

8.1.A.2  Water Autoionization and Kw (1.0 × 10⁻¹⁴ at 25°C):

1. Autoionization of Water:

Autoionization is when two water molecules react with each other to produce ions, even in pure water:

2H2​O⇌H3​O++OH−

One water molecule donates a proton (H⁺), and the other accepts it, producing:

Hydronium ion (H₃O⁺) – a hydrated proton
Hydroxide ion (OH⁻) – what’s left after a water molecule loses H⁺

This reaction is in equilibrium, that is, these ions are present in minute amounts even in neutral water.

2. Ion-Product Constant () and pH:

i. Ion-Product Constant for Water Kw:

Definition: Kw = [H^+][OH^-]
Value at 25°C: Kw = 1.0 times 10^{-14}

That is:
In pure water, the concentration of H^+ and the concentration of OH^- are both:

[H^+] = [OH^-] = √Kw = 1.0 × 10^{-7}, mol/L

ii. Relation to pH and pOH:

pH is expressed as: pH = -log[H^+]

pOH is:  pOH = -log[OH^-]

Their relationship:     pH + pOH = 14 quad {at 25°C}

iii. Neutral, Acidic, and Basic Solutions:

Neutral: [H^+] = [OH^-] ⇒ pH = 7
Acidic: [H^+] > [OH^-] Rightarrow   {pH} < 7
Basic: [H^+] < [OH^-] Rightarrow     {pH} > 7

 

8.1.A.3 Neutral Water at 25°C: pH = pOH = 7.0, pKw = 14:

In pure water, some water molecules spontaneously ionize into hydrogen ions H^+ and hydroxide ions OH^-:

H2​O⇌H++OH−

At 25°C, water ionizes enough so that:

[H+]=[OH−]=1.0×10^−7mol/L

i. pH and pOH:

The pH is a measure of hydrogen ion concentration:

pH = -log[H+] = -log(1.0 × 10^−7) = 7.0

The pOH is a measure of the concentration of hydroxide ions:

pOH = -log[OH^-] = -log(1.0 × 10^−7) = 7.0

Because pH = pOH, the solution is known as neutral.

ii. pKw Relationship:

The product of [H⁺] and [OH⁻] is equal at a given temperature (25°C):

Kw = [H^+][OH^-] = 1.0 × 10^−14

Taking the negative logarithm:

pKw = -log(Kw) = 14.0

Then,
pKw = pH + pOH = 7.0 + 7.0 = 14.0

8.1.A.4 Temperature Dependence of pH in Pure Water: 

1. Autoionization of Water and

Autoionization: Hydrogen molecules from water can react with themselves to form hydrogen ions (H⁺) and hydroxide ions (OH⁻).

H2O⇌H++OH−

Ion-product constant: Measures the extent of this ionization.

Kw=[H+][OH−]

At 25°C:

Kw=1.0×10−14

So, in pure water

[H+]=[OH−]=1.0×10−7 mol/L

Importance: is the pH definition and central theme of acid-base chemistry.

2. Effect of Temperature on K_w

Why Kw Varies with Temperature: The autoionization of water is an endothermic reaction:

H2O (l)⇌H+(aq)+OH−(aq)

Because the process absorbs heat, increasing the temperature promotes the forward reaction (produces more H⁺ and OH⁻), by Le Châtelier’s Principle.

 Effect on 𝐾𝑤:

As temperature increases, 𝐾𝑤 increases.
This leads to [H⁺] and [OH⁻] both rising, so pure water is ionized more at high temperatures.

For instance:

At 25°C:

Kw=1.0×10−14

At 50°C:Kw≈5.5×10−14

Even though [H⁺] rises, water may still be neutral if [H⁺] = [OH⁻], but its pH will be below 7.

3. Impact on pH and Neutrality:

As temperature increases, 𝐾𝑤 increases, so both [H⁺] and [OH⁻] increase.

Neutrality still requires:

[H+]=[OH−][\text{H}^+] = [\text{OH}^-]But pH varies because it’s derived from:

pH=−log⁡[H+]

Neutral pH Shifts with Temperature:

At 25°C, neutral pH = 7.00
At 50°C,  rises → neutral pH ≈ 6.63
At 0°C,  falls → neutral pH ≈ 7.47

Although [H⁺] = [OH⁻], the numerical value of neutral pH varies with temperature.

OLD Content

Introduction to Acids and Bases

Acids

  • All acids start with H
  • Covalent bonds bcuz hydrogen bonds with nonmetals

Naming acids

  • Binary: hydro___ic acid
  • Polyatomic: NO hydro
    • -ate = __ic acid
    • -ite = __ous acid

Bases

  • Ionic bonds because has cation with anion

Naming Bases

  • Name first element and end with hydroxide

Models of acids and bases

  • At equilibrium:
    • There is a competition for H+ between H2O and A-
    • The stronger base controls direction → The direction of equilibrium depends on if the acid is weak or strong
      • If H2O is a stronger base than A-, (H2O attracts H+ more) → forward reaction favored → most of the acid dissolved will be in the ionized form
      • If A- is a much stronger base than H2O → reverse reaction favored → at equilibrium most of the acid will exist as HA
  • Monoprotic acids: one acidic hydrogen
  • Polyprotic acids: more than one acidic hydrogen which can donate to the solution
  • Oxyacids: acidic hydrogen is attached to the oxygen of an ion
  • Organic Acids: acids that contain carbon and usually a carboxyl group, generally very weak  

Water as an Acid and a Base

  • Amphoteric: it can behave either as an acid or a base (ex: water, HSO4-)
    • As an acid:
    • As a base:
  • Water is amphoteric and autoionizes:
    • Ex:
  • (on RFS)
    • Remember that solids and liquids are not included; temperature dependent
  • Neutral solution: [H+] = [OH-]
  • Acidic solution: [H+] > [OH-]
  • Basic solution: [H+] < [OH-]

The pH Scale

  • As pH decreases, [H+] increases exponentially
    • The pH changes by 1 for every power of 10 change in [H+]
  • Sig Figs & pH: number of sig figs of molarity = number of decimal places in pH
    • Ex:

Relationships (All on RFS)


    • P = (-) log of …

    • Ionization increases with increasing temperature (kW will be greater value)

    • Subtract 14 from pOH to find pH
  • → Given any one of these → can find the other three

Estimating Values When Not Given a Calculator 

  • Finding the -log of smthn places answer close to exponent
  • If first term is exactly 1.0 → -log of it will be the same as exponent
  • The greater the first term, the greater the answer is going to fall below the exponent
  • If first term is pi, -log of it will fall halfway between the exponent
    • Ex:
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