16.2 Binary Acids, Oxoacids, & Polyprotic Acids
PRE-HEALTH PREP
Acid Strength Trends
Binary Acids, Oxoacids & Polyprotic Acids
General Chemistry 16.2
Lesson Overview
Acid strength follows predictable periodic trends that differ depending on acid type: binary acids grow stronger down a group due to increasing bond length and across a period due to increasing electronegativity, while oxoacids grow stronger with more oxygen atoms (resonance stabilization) and more electronegative heteroatoms (inductive stabilization). The inverse relationship between acid strength and conjugate base stability underlies all of these trends — a more stable conjugate base always corresponds to a stronger parent acid. Polyprotic acids donate multiple protons in successive steps of decreasing strength, producing amphoteric intermediate species capable of acting as either acids or bases.
What Are Binary Acids?
A binary acid is an acid composed of exactly two elements: hydrogen and one other element. The prefix "bi-" means two, which makes the name easy to remember. Understanding binary acid structure is the first step toward predicting binary acid strength trends on the periodic table.
Definition: Binary Acid
A binary acid contains hydrogen bonded to one other element. The formula takes the general form HnX, where X is typically a nonmetal. The most common binary acids are the hydrohalic acids: HF, HCl, HBr, and HI.
Water (H2O) can technically be classified as a binary acid, which illustrates an important point: "binary" refers to having two elements, not necessarily two atoms. A binary acid may contain multiple hydrogen atoms alongside that second element.
How Does Binary Acid Strength Change Going Down a Group?
Binary acids increase in strength as you move down a group on the periodic table, and the reason comes down to size. Size is the single most important factor for determining binary acid strength going down a group.
The connection lies in bond length — the nucleus-to-nucleus distance between the hydrogen atom and the other atom. As you move down a group, atomic radii increase, which means the bond to hydrogen gets longer. A longer bond is a weaker bond, and a weaker bond dissociates more readily.
The Hydrohalic Acid Series: Acid Strength Down Group 17
Consider the hydrohalic acids going down Group 17:
HF → HCl → HBr → HI
Acid strength increases in this direction. HF is a weak acid, while HCl, HBr, and HI are all strong acids. The H–I bond is the longest and weakest of the four, so HI dissociates most readily and is the strongest binary acid in this series.
Recall that a strong acid dissociates 100% in water, whereas a weak acid dissociates only partially — often less than 5%. The more an acid dissociates, the stronger it is. Longer bonds lead to greater dissociation, which is why binary acid strength increases down a group.
How Does Binary Acid Strength Change Across a Period?
Binary acid strength also increases from left to right across a period, but the reasoning is different. Within the same period, atomic sizes do not vary as dramatically as they do going down a group. When size differences are minimal, the second most important factor takes over: electronegativity.
A more electronegative atom bonded to hydrogen produces a stronger acid. This trend in acid strength across a period becomes clear when you compare Period 2 hydrides across the periodic table.
Comparing Period 2 Hydrides: Acidity Trend Across a Period
Moving from left to right across Period 2:
NH3 (weak base) → H2O (neutral) → HF (weak acid)
Acidity clearly increases as electronegativity increases from nitrogen to oxygen to fluorine. Even though none of these are strong acids, the trend in relative acidity is unmistakable.
If you have your strong acids memorized from the previous lesson, these trends should feel familiar. Going down a group, HF is weak while HCl, HBr, and HI are strong — a clear increase in binary acid strength. Going across a period, you move from a weak base (NH3) to neutral (H2O) to a weak acid (HF) — acidity is definitely increasing to the right.
What Is the Relationship Between Acid Strength and Conjugate Base Stability?
The trend across a period connects to a deeper principle: conjugate base stability. When a binary acid donates a proton, it produces a conjugate base carrying a negative charge. That negative charge is more stable on a more electronegative atom — the same principle you applied when comparing resonance structures in Lewis structures.
The Inverse Relationship: Acid Strength vs. Conjugate Base Strength
There is always an inverse relationship between acid strength and conjugate base strength:
- The stronger the acid → the weaker its conjugate base
- The stronger the base → the weaker its conjugate acid
A more stable conjugate base is a weaker conjugate base, and a weaker conjugate base comes from a stronger acid.
Consider the three Period 2 hydrides: when HF, H2O, or NH3 donate a proton, the conjugate base carries a negative charge on fluorine, oxygen, or nitrogen, respectively. Fluorine is the most electronegative of the three, so fluoride (F⁻) is the most stable conjugate base. A more stable base is a weaker base, and a weaker base comes from a stronger conjugate acid — confirming that HF is the strongest acid of the three.
This inverse relationship reflects a fundamental principle: stability and reactivity are always inversely related. If a species is stable, it has no driving force to change. Greater stability always corresponds to lower reactivity. If you are stable the way you are, you simply do not want to react.
Ranking Conjugate Base Strength for the Halide Ions
Given the halide ions I⁻, Br⁻, Cl⁻, and F⁻ — which is the strongest conjugate base?
Since acid strength increases going down the group (HF < HCl < HBr < HI), the conjugate base trend runs in the opposite direction. The weakest acid (HF) produces the strongest conjugate base (F⁻).
Conjugate base strength: F⁻ > Cl⁻ > Br⁻ > I⁻
What Are Oxoacids (Oxyacids)?
Oxoacids (also called oxyacids) are acids composed of three elements: hydrogen, oxygen, and one additional element known as the heteroatom. Two distinct rules govern their acid strength, and those rules differ fundamentally from the rules for binary acids.
Definition: Oxoacid (Oxyacid)
An oxoacid has a structure where hydrogen is bonded to oxygen, which is bonded to a central heteroatom. Common examples include HClO4 (perchloric acid), H2SO4 (sulfuric acid), and HNO3 (nitric acid). The key structural feature: the acidic hydrogen is always bonded to oxygen, never directly to the heteroatom.
2 Rules That Determine Oxoacid Strength
Oxoacid Rule 1: The more oxygen atoms an oxoacid contains, the stronger the acid. This is typically the most important rule when comparing two oxoacids with the same heteroatom.
Unlike binary acids, bond length never comes into play for oxoacids. The acidic hydrogen is always bonded to an oxygen, so the H–O bond length is roughly the same in every oxoacid. Instead, we determine oxoacid strength by examining conjugate base stability through resonance.
Oxoacid Strength Example: Perchloric Acid vs. Chloric Acid
Compare HClO4 (perchloric acid) with HClO3 (chloric acid):
- ClO4⁻ (perchlorate) has 4 resonance structures, spreading the negative charge across four oxygens. Each oxygen bears roughly −¼ of the charge.
- ClO3⁻ (chlorate) has 3 resonance structures, spreading the charge across three oxygens. Each oxygen bears roughly −⅓ of the charge.
Perchlorate distributes the charge more effectively, making it the more stable (and therefore weaker) conjugate base. A weaker conjugate base comes from a stronger acid, so HClO4 is the stronger acid.
The principle is straightforward: a negative charge means excess electron density, and stability favors being as close to neutral as possible. The more atoms that share the negative charge, the less burden each atom carries. This is resonance stabilization of the conjugate base.
Oxoacid Rule 2 applies when two oxoacids have the same number of oxygens. See the next section on the inductive effect for details.
How Does the Inductive Effect Determine Oxoacid Strength?
Oxoacid Rule 2: When two oxoacids have the same number of oxygens, the one with the more electronegative heteroatom is the stronger acid. This rule applies when Rule 1 cannot distinguish between the two.
Inductive Effect Example: Perchloric Acid vs. Perbromic Acid
Both HClO4 and HBrO4 have four oxygens, so Rule 1 provides no distinction — both conjugate bases have four resonance structures spreading the charge across four oxygens.
The tiebreaker is electronegativity. Chlorine is more electronegative than bromine, so it pulls electron density away from those oxygens more effectively through the bonds. By drawing some of that negative charge toward itself, chlorine makes the four oxygens slightly less negative — spreading the charge even further.
HClO4 is the stronger acid — and you likely already know this since perchloric acid appears on the strong acids list while perbromic acid does not.
The Inductive Effect in Oxoacids
The ability of an electronegative atom to pull electron density through bonds is called the inductive effect (or simply induction). The two oxoacid strength rules rely on different stabilization mechanisms:
- Rule 1 — resonance stabilization of the conjugate base (more oxygens = more resonance structures = stronger oxoacid)
- Rule 2 — inductive stabilization of the conjugate base (more electronegative heteroatom = better electron withdrawal = stronger oxoacid)
A critical point for oxoacids: the heteroatom's size is completely irrelevant. It does not matter whether you are comparing elements in the same group or the same period. Oxoacid strength depends purely on electronegativity — no ifs, ands, or buts.
Binary Acids vs. Oxoacids: 4 Key Differences
Students often find it confusing that binary acids get stronger going down a group while oxoacids get stronger going up a group. The difference comes down to what the hydrogen is bonded to.
| Property | Binary Acids | Oxoacids |
|---|---|---|
| H bonded to | Directly to the heteroatom (X) | Always to oxygen (O) |
| Primary factor | Size / bond length | Electronegativity / resonance |
| Trend down a group | Acid strength increases | Acid strength decreases |
| Role of size | Dominant factor | Irrelevant |
| Why | Longer H–X bond → weaker bond → more dissociation | Less electronegative heteroatom → less inductive stabilization → weaker acid |
The Key Distinction Between Binary Acids and Oxoacids
In a binary acid, hydrogen bonds directly to the heteroatom, so the H–X bond length (determined by size) controls acid strength. In an oxoacid, hydrogen always bonds to oxygen, making the H–O bond length roughly constant. Since bond length cannot explain the difference, electronegativity becomes the deciding factor.
This is a classic exam question. If someone tells you HBr is stronger than HCl because of size, that reasoning is correct — for binary acids. But if they try to use size to compare HClO4 and HBrO4, that reasoning fails because the hydrogen in an oxoacid never bonds to chlorine or bromine. Always identify the acid type before applying an acid strength trend.
What Are Polyprotic Acids?
A polyprotic acid is an acid that can donate more than one proton (H⁺). Most common acids are monoprotic, meaning they donate a single proton. Polyprotic acids donate two or more, producing a series of successive dissociation steps each with its own acid dissociation constant (Ka).
3 Common Polyprotic Acids
- Phosphoric acid (H3PO4) — triprotic (3 acidic protons)
- Sulfuric acid (H2SO4) — diprotic (2 acidic protons), and the only strong acid that is also polyprotic
- Carbonic acid (H2CO3) — diprotic (2 acidic protons)
Why Does Each Successive Proton Donation Get Harder?
The defining trend of polyprotic acids: each successive proton donation is more difficult than the last, meaning Ka1 > Ka2 > Ka3. After the first proton leaves, the remaining species carries a negative charge. That negative charge attracts the remaining H⁺ ions more strongly, making them harder to remove.
Phosphoric Acid Dissociation Steps and Ka Values
H3PO4 ⇌ H⁺ + H2PO4⁻ (Ka1 — largest)
H2PO4⁻ ⇌ H⁺ + HPO4²⁻ (Ka2 — smaller)
HPO4²⁻ ⇌ H⁺ + PO4³⁻ (Ka3 — smallest)
Each equilibrium constant (Ka) is progressively smaller, reflecting the decreasing tendency to donate a proton at each stage.
The logic is intuitive: after losing the first proton, the species becomes negatively charged. A −1 species is less inclined to release another positive charge than a neutral species was. A −2 species is even more reluctant. The increasing negative charge creates an ever-stronger electrostatic attraction to the remaining protons, which is why Ka decreases at each successive step.
5 Species in a Phosphoric Acid Solution (Ranked by Concentration)
Because H3PO4 is a weak polyprotic acid, only a small fraction dissociates — perhaps around 1%. The vast majority of the acid remains in its undissociated form, making H3PO4 the most abundant species in solution.
Species Concentrations in H3PO4 Solution (Greatest to Least)
- H3PO4 — most abundant (undissociated, ~99%)
- H⁺ — produced in the first dissociation, boosted slightly by subsequent steps
- H2PO4⁻ — slightly less than H⁺ because some dissociates further
- HPO4²⁻ — much less (second dissociation is far weaker)
- PO4³⁻ — least abundant (third dissociation is weakest of all)
An important subtlety: you might expect equal concentrations of H⁺ and H2PO4⁻ from the first dissociation since they form in a 1:1 ratio. However, the second dissociation consumes some H2PO4⁻ while producing additional H⁺, so the H⁺ concentration ends up slightly higher. Each successive Ka is dramatically smaller — the decrease in acidity at each step is often far larger than you might expect.
4 Species in a Sulfuric Acid Solution (Ranked by Concentration)
Sulfuric acid (H2SO4) requires special attention because it is a strong polyprotic acid. Unlike phosphoric acid, the first dissociation of H2SO4 goes to completion — 100%.
Sulfuric Acid: The Only Strong Polyprotic Acid
H2SO4 → H⁺ + HSO4⁻ (100% dissociation — strong acid)
HSO4⁻ ⇌ H⁺ + SO4²⁻ (partial dissociation — weak acid)
Notice the first step uses a single arrow (→) because the dissociation is complete, while the second step uses equilibrium arrows (⇌) because it is only partial.
Species Concentrations in H2SO4 Solution (Greatest to Least)
- H⁺ — most abundant (produced by both dissociations)
- HSO4⁻ — second most abundant
- SO4²⁻ — third (from partial second dissociation)
- H2SO4 — effectively zero (100% dissociated)
This is the key difference from a weak polyprotic acid: for H2SO4, the undissociated form is the least abundant species rather than the most. If an exam asks you to rank species concentrations, do not forget that H2SO4 is a strong acid.
What Are Amphoteric Species in Polyprotic Acids?
Look at the middle species in any polyprotic acid series and you will find amphoteric compounds — species that can act as either an acid or a base depending on the circumstances. Recognizing amphoteric species is a common general chemistry exam skill.
Definition: Amphoteric Species
An amphoteric species can donate a proton (acting as an acid) or accept a proton (acting as a base). Amphoteric species always appear in the intermediate steps of a polyprotic acid dissociation series.
Amphoteric Species in Phosphoric Acid and Sulfuric Acid Series
| Polyprotic Acid | Amphoteric Species | As Acid (donates H⁺) | As Base (accepts H⁺) |
|---|---|---|---|
| H3PO4 | H2PO4⁻, HPO4²⁻ | Donates H⁺ to form the next species in the series | Accepts H⁺ to reform the previous species |
| H2SO4 | HSO4⁻ | Donates H⁺ to form SO4²⁻ | Accepts H⁺ to reform H2SO4 |
For example, H2PO4⁻ can donate a proton to form HPO4²⁻ (acting as an acid) or accept a proton to reform H3PO4 (acting as a base). Any species that sits between the fully protonated and fully deprotonated forms is amphoteric.
10 Key Takeaways: Acid Strength Trends
- Binary acids contain hydrogen and one other element. Their strength depends primarily on bond length (size) when comparing down a group and on electronegativity when comparing across a period.
- Binary acids increase in strength going down a group because larger atoms form longer, weaker H–X bonds that dissociate more readily.
- Binary acids increase in strength going across a period because more electronegative atoms produce more stable conjugate bases.
- The stronger the acid, the weaker its conjugate base — and vice versa. Stability and reactivity are always inversely related.
- Oxoacids contain hydrogen, oxygen, and a heteroatom. More oxygens means greater resonance stabilization of the conjugate base, producing a stronger oxoacid.
- When oxoacids have equal numbers of oxygens, a more electronegative heteroatom provides greater inductive stabilization, producing a stronger acid.
- Size matters for binary acids (H bonded to heteroatom) but not for oxoacids (H always bonded to oxygen) — this is why the acid strength trends point in opposite directions down a group.
- Polyprotic acids donate multiple protons, but each successive Ka is smaller because the increasing negative charge holds remaining protons more tightly (Ka1 > Ka2 > Ka3).
- In weak polyprotic acid solutions, the undissociated form is most abundant. In H2SO4 (the only strong polyprotic acid), the undissociated form is effectively absent.
- Amphoteric species appear in the middle of polyprotic acid series and can act as either acids or bases.