5.3 pH and Buffers for A-Level Chemistry (OCR)

In Brønsted–Lowry theory, what particle does an acid donate in a reaction?

Proton

A proton is a hydrogen ion (H⁺).

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In Brønsted–Lowry theory, what type of particle does a base accept?

Proton

Bases accept H⁺ ions from acids during reactions.

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Fill in the blank:

A Brønsted–Lowry acid–base reaction involves the transfer of a _______ between species.

proton

Acid–base equilibria are based on proton transfer.

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True or False:

A Brønsted–Lowry base is defined as a substance that donates a proton.

False

A base accepts a proton, while an acid donates one.

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What type of equilibrium is formed when an acid transfers a proton to a base in solution?

Acid–base equilibrium

The reaction can proceed in both forward and reverse directions.

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True or False:

In Brønsted–Lowry theory, acids and bases always react in pairs.

True

A proton donor must transfer H⁺ to a proton acceptor.

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Fill in the blank:

When an acid loses a proton, it forms its _______ base.

conjugate

The conjugate base is the species remaining after proton donation.

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What is formed when a base accepts a proton?

Conjugate acid

The base becomes its conjugate acid after gaining H⁺.

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True or False:

Acid–base equilibria can shift depending on the relative strengths of acids and bases.

True

Stronger acids donate protons more readily than weaker ones.

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What type of particle transfer distinguishes Brønsted–Lowry reactions from other acid–base theories?

Proton transfer

This definition focuses specifically on H⁺ transfer between species.

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What scale is used to measure the acidity of aqueous solutions based on hydrogen ion concentration?

pH scale

The pH scale is logarithmic because [H⁺] values span many orders of magnitude.

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What does the symbol [H⁺] represent in pH calculations?

Hydrogen ion concentration

Concentration is measured in mol dm^-3.

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Fill in the blank:

The mathematical expression for pH is pH = −log₁₀[_______].

H⁺

The brackets indicate concentration of hydrogen ions in solution.

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True or False:

The pH scale is logarithmic rather than linear.

True

Each unit change in pH corresponds to a tenfold change in [H⁺].

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If the hydrogen ion concentration increases, what happens to the pH value?

Decreases

Higher acidity corresponds to lower pH values.

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Fill in the blank:

A tenfold increase in hydrogen ion concentration causes the pH to decrease by _______ unit.

one

Because the scale is base-10 logarithmic.

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True or False:

A solution with pH 3 has ten times more H⁺ ions than a solution with pH 4.

True

Each pH unit represents a factor of 10 change in [H⁺].

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How can the hydrogen ion concentration be calculated from pH?

[H⁺] = 10^-pH

This rearranges the definition pH = −log₁₀[H⁺].

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True or False:

The pH of a strong acid can be calculated directly from its concentration.

True

Strong acids fully dissociate, so [H⁺] ≈ acid concentration.

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Fill in the blank:

When calculating pH using logarithms, the concentration of H⁺ must be expressed in _______ dm^-3.

mol

This ensures consistency in pH calculations.

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Why is a logarithmic scale used for pH?

Wide range

Hydrogen ion concentrations in aqueous solutions vary over many powers of ten.

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What type of acid is typically assumed to fully dissociate when calculating pH directly from concentration?

Strong acid

For example, HCl or HNO₃ in dilute aqueous solution.

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What equilibrium constant describes the self-ionisation of water?

It represents the equilibrium between H₂O, H⁺ and OH^- ions.

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What term is used for the equilibrium constant of water dissociation?

Ionic product of water

It is derived from the equilibrium expression for water splitting into ions.

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# Fill in the blank: The **ionic product of water** is expressed as Kw = [H⁺][_\_\_\_\_\_\_].

OH^- ## Footnote The concentrations are measured in mol dm^-3.

# True or False: Pure **water** contains both **hydrogen ions and hydroxide ions**.

True ## Footnote Water slightly dissociates into H⁺ and OH^- ions.

What type of **equilibrium** produces **H⁺ and OH^- ions** in pure water?

Self-ionisation ## Footnote Also called auto-ionisation of water.

# True or False: The value of **Kw** remains constant regardless of **temperature**.

False ## Footnote Kw changes with temperature because the equilibrium position shifts.

# Fill in the blank: In a **neutral solution**, the concentrations of **H⁺ and OH^-** are _\_\_\_\_\_\_.

equal ## Footnote This results in pH 7 at 298 K.

How can the **hydrogen ion concentration** be calculated if the **hydroxide ion concentration** is known?

Kw divided by [OH^-] ## Footnote Rearranged from Kw = [H⁺][OH^-].

# True or False: The **pH of a strong base** can be calculated using **Kw**.

True ## Footnote First find [H⁺] from Kw, then calculate pH using −log₁₀[H⁺].

# Fill in the blank: To find **pH** from Kw calculations, the **hydrogen ion concentration** must be converted using a _\_\_\_\_\_\_ scale.

logarithmic ## Footnote pH is calculated using −log₁₀[H⁺].

If the **concentration of OH^- increases**, what happens to the **concentration of H⁺**?

Decreases ## Footnote Because the product [H⁺][OH^-] must remain equal to Kw at a given temperature.

What **unit** is used for **concentrations** in the Kw expression?

mol dm^-3 ## Footnote Both [H⁺] and [OH^-] are expressed in mol dm^-3.

How do **weak acids** behave in aqueous solution compared with **strong acids**?

Partially dissociate ## Footnote Only a small fraction of molecules ionise to produce H⁺ ions.

What symbol represents the **acid dissociation constant** for a weak acid?

Ka ## Footnote It measures the extent of dissociation of a weak acid in solution.

# Fill in the blank: The relationship between **Ka and pKa** is pKa = −log₁₀ _\_\_\_\_\_\_.

Ka ## Footnote pKa is simply the negative logarithm of the acid dissociation constant.

# True or False: A **larger value of Ka** indicates a **stronger weak acid**.

True ## Footnote Greater dissociation means the acid donates protons more readily.

What does a **small Ka value** indicate about a **weak acid**?

Weak dissociation ## Footnote The equilibrium lies mainly towards the undissociated acid.

# True or False: A **larger pKa value** indicates a **stronger acid**.

False ## Footnote Larger pKa corresponds to weaker acids because Ka is smaller.

# Fill in the blank: For a **weak acid HA**, the equilibrium expression is Ka = [H⁺][_\_\_\_\_\_\_] / [HA].

A^- ## Footnote A^- represents the conjugate base formed when HA dissociates.

What happens to **pKa** when the value of **Ka increases**?

Decreases ## Footnote Because pKa = −log₁₀Ka.

# True or False: **Weak acids** fully ionise in **aqueous solution**.

False ## Footnote Only a small proportion of molecules dissociate.

What **experimental point in a titration** can be used to determine **Ka of a weak acid**?

Half neutralisation ## Footnote At this point pH = pKa.

# Fill in the blank: At **half neutralisation** in a weak acid titration, **pH equals** _\_\_\_\_\_\_.

pKa ## Footnote This relationship is commonly used to determine Ka experimentally.

What **species** is produced when a **weak acid donates a proton**?

Conjugate base ## Footnote The remaining species after proton donation is the conjugate base.

What **experimental technique** is used to determine the **concentration of an acid or base** using a solution of known concentration?

Titration ## Footnote A measured volume of one solution reacts exactly with another at the equivalence point.

What type of **graph** shows how **pH changes** as a titration progresses?

pH curve ## Footnote It plots pH against volume of titrant added.

# Fill in the blank: A **pH curve** plots pH against the _\_\_\_\_\_\_ of solution added during a titration.

volume ## Footnote Usually the volume of the titrant added from a burette.

# True or False: The **equivalence point** is where exactly the required **stoichiometric amounts of acid and base** have reacted.

True ## Footnote At this point the number of moles of acid equals the number of moles of base required by the equation.

What **region of a titration curve** shows the most **rapid change in pH**?

Vertical section ## Footnote This steep region occurs near the equivalence point.

# True or False: In a **strong acid–strong base titration**, the equivalence point occurs at **pH 7**.

True ## Footnote At 298 K, the solution is neutral at the equivalence point.

# Fill in the blank: A substance used to show when a **titration** has reached its end point is called an _\_\_\_\_\_\_.

indicator ## Footnote Indicators change colour over a narrow pH range.

What **property of an indicator** determines whether it is suitable for a particular titration?

pH range of colour change ## Footnote The indicator must change colour within the steep region of the pH curve.

# True or False: In a **weak acid–strong base titration**, the equivalence point is **above pH 7**.

True ## Footnote The conjugate base formed makes the solution slightly alkaline.

In a **strong acid–weak base titration**, is the equivalence point **above or below pH 7**?

Below ## Footnote The conjugate acid formed makes the solution slightly acidic.

# Fill in the blank: The point where the **indicator changes colour** during a titration is called the _\_\_\_\_\_\_ point.

end ## Footnote It should closely match the equivalence point.

Why must the chosen **indicator** change colour within the **steep part of the pH curve**?

Accuracy ## Footnote This ensures the colour change occurs close to the true equivalence point.

What type of **solution** resists **changes in pH** when small amounts of acid or base are added?

Buffer solution ## Footnote Buffers maintain a nearly constant pH even when diluted or when small amounts of acid/base are added.

What two **components** make up an **acidic buffer solution**?

Weak acid and its salt ## Footnote The salt provides the conjugate base of the weak acid.

# Fill in the blank: An **acidic buffer** contains a **weak acid** and the _\_\_\_\_\_\_ base of that acid.

conjugate ## Footnote The conjugate base usually comes from the salt of the weak acid.

# True or False: A **buffer solution** completely prevents any **change in pH**.

False ## Footnote Buffers only minimise pH changes; they do not eliminate them entirely.

What type of **buffer** contains a **weak base** and the salt of that weak base?

Basic buffer ## Footnote The salt supplies the conjugate acid of the weak base.

# True or False: A **buffer** can resist **changes in pH** when small amounts of acid are added.

True ## Footnote The conjugate base reacts with added H⁺ ions.

# Fill in the blank: In an **acidic buffer**, added **H⁺ ions** are removed by reaction with the _\_\_\_\_\_\_ base present.

conjugate ## Footnote The conjugate base reacts with H⁺ to form the weak acid.

What happens when small amounts of **OH^-** are added to an **acidic buffer**?

Neutralised by weak acid ## Footnote The weak acid donates H⁺ to react with OH^- to form water.

# True or False: **Buffer solutions** are important in **biological systems** where stable pH is required.

True ## Footnote For example, blood pH must remain within a narrow range.

# Fill in the blank: A **buffer** resists **pH change** because it contains a weak acid and its _\_\_\_\_\_\_ base.

conjugate ## Footnote This pair can neutralise added acid or base.

Why do **buffers** only work effectively when **small amounts of acid or base** are added?

Limited capacity ## Footnote Once the weak acid or conjugate base is used up, the buffer can no longer resist pH change.

5.3 pH and Buffers Flashcards

Calculate pH and analyse buffer systems in maintaining equilibrium. (69 cards)