What is the difference between the Arrhenius and Bronsted-Lowry definitions of an acid?

Arrhenius defines an acid as a substance that releases H+ in water. Bronsted-Lowry defines an acid as a proton (H+) donor, which is a broader definition that also covers reactions outside water.

Why is ammonia, NH3, classified as a base under the Bronsted-Lowry definition even though it has no OH- in its formula?

Under the Bronsted-Lowry definition, a base is a proton acceptor. NH3 accepts an H+ from another substance, so it acts as a base even though it does not release OH- on its own.

What does it mean for water to be amphoteric?

Water can act as an acid by donating H+ or as a base by accepting H+, depending on the substance it reacts with.

In the reaction HCl + H2O → H3O+ + Cl-, which two conjugate pairs are present?

HCl and Cl- are a pair, because HCl gives up one H+ to become Cl-. H3O+ and H2O are a pair, because H2O accepts an H+ to become H3O+.

How does the strength of an acid relate to the strength of its conjugate base?

The stronger the acid, the weaker its conjugate base. A strong acid like HCl gives a very weak conjugate base like Cl-.

Why does a change of one pH unit represent a tenfold change in hydrogen ion concentration?

pH is a logarithmic scale based on base 10. Each step of 1 in pH means the [H+] changes by a factor of 10, so pH 3 has 100 times more H+ than pH 5.

How do you find the pH of a 0.0010 M NaOH solution at 25 °C?

NaOH is a strong base that gives one OH- per formula unit, so [OH-] = 1.0 × 10^-3 M. pOH = 3.00, so pH = 14.00 - 3.00 = 11.00.

What is the difference between a strong acid and a weak acid in terms of ionization?

A strong acid ionizes essentially 100 percent in water, while a weak acid ionizes only partly and reaches an equilibrium. A larger Ka means a stronger acid.

What is the net ionic equation for the neutralization of a strong acid with a strong base?

H+ + OH- → H2O. The spectator ions, such as Na+ and Cl- in HCl + NaOH, are left out of the net equation.

What pH does the equivalence point have when a weak acid is titrated with a strong base, and why?

The pH is above 7, because the salt formed from a weak acid and a strong base produces a basic solution.

What is a buffer, and how does it resist changes in pH?

A buffer is a weak acid plus its conjugate base, such as acetic acid with sodium acetate. The base part absorbs added H+, and the acid part neutralizes added OH-, so the pH changes little.

Why does a good indicator matter in a titration?

A good indicator changes color at a pH close to the equivalence point, so the endpoint you observe matches the point where moles of H+ equal moles of OH-.

Chemistry

Acids and bases: flashcards, quiz and summary

A free study pack Chomp made from a page of notes on this topic: 12 flashcards, a 6-question quiz, the test traps and a summary. No account needed.

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The short version

Acids and bases are defined two ways: Arrhenius (H+ or OH- released in water) and Bronsted-Lowry (proton donors and acceptors). The main skills are calculating pH and pOH, identifying conjugate pairs, separating strong from weak acids and bases, and doing titration math.

Flashcards

1What is the difference between the Arrhenius and Bronsted-Lowry definitions of an acid?
Arrhenius defines an acid as a substance that releases H+ in water. Bronsted-Lowry defines an acid as a proton (H+) donor, which is a broader definition that also covers reactions outside water.
2Why is ammonia, NH3, classified as a base under the Bronsted-Lowry definition even though it has no OH- in its formula?
Under the Bronsted-Lowry definition, a base is a proton acceptor. NH3 accepts an H+ from another substance, so it acts as a base even though it does not release OH- on its own.
3What does it mean for water to be amphoteric?
Water can act as an acid by donating H+ or as a base by accepting H+, depending on the substance it reacts with.
4In the reaction HCl + H2O → H3O+ + Cl-, which two conjugate pairs are present?
HCl and Cl- are a pair, because HCl gives up one H+ to become Cl-. H3O+ and H2O are a pair, because H2O accepts an H+ to become H3O+.
5How does the strength of an acid relate to the strength of its conjugate base?
The stronger the acid, the weaker its conjugate base. A strong acid like HCl gives a very weak conjugate base like Cl-.
6Why does a change of one pH unit represent a tenfold change in hydrogen ion concentration?
pH is a logarithmic scale based on base 10. Each step of 1 in pH means the [H+] changes by a factor of 10, so pH 3 has 100 times more H+ than pH 5.
7How do you find the pH of a 0.0010 M NaOH solution at 25 °C?
NaOH is a strong base that gives one OH- per formula unit, so [OH-] = 1.0 × 10^-3 M. pOH = 3.00, so pH = 14.00 - 3.00 = 11.00.
8What is the difference between a strong acid and a weak acid in terms of ionization?
A strong acid ionizes essentially 100 percent in water, while a weak acid ionizes only partly and reaches an equilibrium. A larger Ka means a stronger acid.
9What is the net ionic equation for the neutralization of a strong acid with a strong base?
H+ + OH- → H2O. The spectator ions, such as Na+ and Cl- in HCl + NaOH, are left out of the net equation.
10What pH does the equivalence point have when a weak acid is titrated with a strong base, and why?
The pH is above 7, because the salt formed from a weak acid and a strong base produces a basic solution.
11What is a buffer, and how does it resist changes in pH?
A buffer is a weak acid plus its conjugate base, such as acetic acid with sodium acetate. The base part absorbs added H+, and the acid part neutralizes added OH-, so the pH changes little.
12Why does a good indicator matter in a titration?
A good indicator changes color at a pH close to the equivalence point, so the endpoint you observe matches the point where moles of H+ equal moles of OH-.

Practice quiz

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1What is the pH of 0.0010 M NaOH at 25 °C?
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A. 11.00 NaOH gives [OH-] = 1.0 × 10^-3 M, so pOH = 3.00. Then pH = 14.00 - 3.00 = 11.00. The value 3.00 is the pOH, not the pH.
2In the reaction HCl + H2O → H3O+ + Cl-, which pairs are conjugate acid-base pairs?
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D. HCl/Cl- and H3O+/H2O A conjugate pair differs by exactly one H+. HCl loses H+ to become Cl-, and H2O gains H+ to become H3O+.
3What is the pH at the equivalence point of a weak acid titrated with a strong base?
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C. Above 7 The salt formed from a weak acid and a strong base hydrolyzes to produce OH-, so the solution at the equivalence point is basic, with pH above 7.
4What is the pH of 0.050 M Ba(OH)2 at 25 °C?
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A. 13.00 Each Ba(OH)2 releases 2 OH-, so [OH-] = 0.10 M. pOH = 1.00, and pH = 14.00 - 1.00 = 13.00. The value 1.00 is the pOH.
520.0 mL of 0.100 M NaOH exactly neutralizes 25.0 mL of HCl. What is the concentration of the HCl?
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B. 0.0800 M Moles NaOH = 0.100 × 0.0200 L = 0.00200 mol. This equals moles HCl in a 1:1 reaction, so [HCl] = 0.00200 mol ÷ 0.0250 L = 0.0800 M.
6Which mixture is a buffer?
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B. Acetic acid and sodium acetate A buffer needs a weak acid plus its conjugate base. Acetic acid is a weak acid and acetate is its conjugate base. HCl and NaCl have no weak acid, and HCl with NaOH would neutralize rather than buffer.

Test traps

What teachers like to test on this, and the mistakes students make.

Mixing up pH and pOH in the final step. Students often report the pOH as the answer for a basic solution, when they need to subtract pOH from 14 to get pH.
Forgetting to multiply by the number of OH- units for bases like Ba(OH)2. Writing [OH-] = 0.050 M instead of 0.10 M gives the wrong pH.
Assuming pH 7 is the equivalence point for every titration. It is only true for a strong acid with a strong base; weak acid with strong base gives above 7, and strong acid with weak base gives below 7.
Confusing strong with concentrated. A dilute HCl solution is still a strong acid, because strength refers to how completely it ionizes.
Using pH + pOH = 14 at any temperature. This relationship holds at 25 °C because Kw changes with temperature.
Listing H3O+ and OH- as a conjugate pair. They are not; H3O+ and H2O are the conjugate pair.

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Summary

Arrhenius defines acids as substances that release H+ in water and bases as substances that release OH-, while Bronsted-Lowry defines acids as proton donors and bases as proton acceptors.
Bronsted-Lowry is the broader definition because it includes bases like NH3 that have no OH- in them.
Water is amphoteric, so it can act as an acid or as a base depending on the reaction.
pH = -log[H+] and pOH = -log[OH-], and at 25 °C pH + pOH = 14 because [H+][OH-] = 1.0 × 10^-14.
Each pH unit is a tenfold change in [H+], so pH 3 has 100 times more H+ than pH 5.
Strong acids and bases ionize essentially completely, while weak ones reach an equilibrium where only part of the substance has ionized.
In a titration, the endpoint should match the equivalence point, where moles of H+ equal moles of OH-, and for a 1:1 reaction Ma × Va = Mb × Vb.
A buffer is a weak acid plus its conjugate base, and it resists pH change when a small amount of acid or base is added.

Key terms

Arrhenius acid: A substance that releases H+ (present as H3O+) in water.
Bronsted-Lowry acid: A proton (H+) donor.
Bronsted-Lowry base: A proton (H+) acceptor. Ammonia, NH3, is a base under this definition even though it has no OH- in its formula.
Conjugate pair: Two substances that differ by exactly one H+, such as HCl and Cl-, or H3O+ and H2O.
Amphoteric: Able to act as either an acid or a base. Water is amphoteric.
Kw: The ion product of water, 1.0 × 10^-14 at 25 °C, equal to [H+][OH-].
Equivalence point: The point in a titration where moles of H+ equal moles of OH-.
Buffer: A solution of a weak acid and its conjugate base that resists pH change when a little acid or base is added.

How to study it

Acid-base definitions and conjugate pairs

Learn the difference between the Arrhenius definition (H+ or OH- in water) and the Bronsted-Lowry definition (proton donor or acceptor).
For any acid-base reaction, identify which species donates H+ and which accepts it, then pair each acid with its conjugate base and each base with its conjugate acid.
Check your pairs by confirming that each pair differs by exactly one H+.

pH, pOH and calculations

Use pH = -log[H+] and pOH = -log[OH-], and remember that pH + pOH = 14 only at 25 °C.
For a strong base, find [OH-] first (multiply by the number of OH- units for Ba(OH)2), then find pOH, then convert to pH.
Practice converting between [H+] and pH in both directions, including values like 2.5 × 10^-4 M where you use pH = 4 - log(2.5).

Titration and buffers

In a titration, write the neutralization equation first, then use moles = M × V to find moles of the known solution, then use the mole ratio to find moles of the unknown.
Know which pH the equivalence point has for each combination of strong and weak acid or base.
Be able to explain why a buffer needs both the weak acid and its conjugate base to work.

The notes this pack came from

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Definitions

Arrhenius: an acid releases H+ in water (present as hydronium, H3O+); a base releases OH- in water. Examples: HCl → H+ + Cl- and NaOH → Na+ + OH-.
Bronsted-Lowry: an acid is a proton (H+) donor; a base is a proton acceptor. It covers bases like NH3, which has no OH-.
Water is amphoteric: it can act as an acid or a base.

Conjugate pairs

An acid that gives away H+ becomes its conjugate base; a base that accepts H+ becomes its conjugate acid. A pair differs by exactly one H+.
HCl + H2O → H3O+ + Cl- has the pairs HCl/Cl- and H3O+/H2O.
The stronger the acid, the weaker its conjugate base.

pH and pOH

pH = -log[H+] and pOH = -log[OH-].
At 25 °C, [H+][OH-] = 1.0 × 10^-14 (called Kw), so pH + pOH = 14. Below 7 is acidic, 7 is neutral, above 7 is basic.
Each pH unit is a tenfold change in [H+]: pH 3 has 100 times more H+ than pH 5.

Worked pH examples

0.010 M HCl (strong): [H+] = 1.0 × 10^-2 M, so pH = 2.00 and pOH = 12.00.
0.0010 M NaOH: [OH-] = 1.0 × 10^-3 M, so pOH = 3.00 and pH = 11.00.
0.050 M Ba(OH)2 gives 2 OH- per unit: [OH-] = 0.10 M, pOH = 1.00, pH = 13.00.
[H+] = 2.5 × 10^-4 M: pH = 4 - log(2.5) = 4 - 0.40 = 3.60.

Strong vs weak

Strong acids ionize essentially 100 percent. The six usually memorized: HCl, HBr, HI, HNO3, HClO4 and H2SO4 (first H+ only).
Strong bases: group 1 hydroxides (NaOH, KOH) and heavier group 2 ones like Ba(OH)2.
Weak acids and bases ionize only partly and reach equilibrium: acetic acid (CH3COOH), HF and the weak base NH3.
A larger Ka (smaller pKa) means a stronger acid.

Neutralization

Acid + base → salt + water: HCl + NaOH → NaCl + H2O, with the net ionic equation H+ + OH- → H2O.

Indicators

Litmus: red in acid, blue in base.
Phenolphthalein: colorless in acid, turning pink from about pH 8.2 to 10.
Bromothymol blue: yellow in acid, blue in base, changing near pH 7.

Titration basics

Put a measured volume of the unknown (analyte) in a flask with indicator.
Add a solution of known concentration (titrant) from a buret.
Stop at the endpoint, when the color changes. A good indicator makes this match the equivalence point, where moles of H+ equal moles of OH-.

For a 1:1 reaction: Ma × Va = Mb × Vb.
Example: 20.0 mL of 0.100 M NaOH neutralizes 25.0 mL of HCl. Moles NaOH = 0.100 × 0.0200 L = 0.00200 mol = moles HCl, so [HCl] = 0.00200 mol ÷ 0.0250 L = 0.0800 M.
pH at equivalence: 7 for strong acid with strong base, above 7 for weak acid with strong base, below 7 for strong acid with weak base.

Buffers in brief

A buffer resists pH change when a little acid or base is added.
It is a weak acid plus its conjugate base, such as acetic acid with sodium acetate. The base part absorbs added H+; the acid part neutralizes added OH-.
Blood uses carbonic acid and bicarbonate to stay near pH 7.4.

Common mix-ups

Strong does not mean concentrated: dilute HCl is still strong.
pH + pOH = 14 only at 25 °C, because Kw changes with temperature.
The equivalence point is pH 7 only for a strong acid with a strong base.
H3O+ and H2O are a conjugate pair; H3O+ and OH- are not.

Questions

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