Can Molecules Without OH Groups Hydrogen Bond with Water?

A molecule does **not** need an OH group to form hydrogen bonds with water. A molecule such as acetone, which contains a carbonyl group (C=O) but no OH group, can still hydrogen bond with water.

The key is to look at the whole molecular structure, not just search for an OH group. Water can donate a hydrogen bond to a suitable lone pair on another molecule, even when that molecule has no hydrogen attached to oxygen, nitrogen, or fluorine.

This is an important distinction in HSC Chemistry: a molecule can accept hydrogen bonds from water without being able to donate hydrogen bonds itself.

## 1. Can a molecule without an OH group hydrogen bond with water?

Yes. To understand why, think about what actually makes a hydrogen bond possible.

Water has the molecular structure H-O-H. Oxygen is highly electronegative, so it attracts the shared electrons in both O-H bonds towards itself.

This gives water:
- Two partially positive hydrogen atoms, written as \(\delta^+\).
- Two lone pairs of electrons on oxygen.
- A partially negative oxygen atom, written as \(\delta^-\).

Water can therefore play two roles:
1. It can donate a hydrogen bond through one of its partially positive hydrogen atoms.
2. It can accept a hydrogen bond through a lone pair on its oxygen atom.

The second role is the one students often overlook.

[IMAGE: Annotated molecular illustration of two water molecules and an acetone molecule. Show water's partially positive hydrogen atoms and oxygen lone pairs. Show a dotted hydrogen bond from a water hydrogen to the carbonyl oxygen of acetone. Include the C=O bond and clearly distinguish dotted intermolecular attractions from solid covalent bonds.]

Consider acetone, \(\ce{CH3COCH3}\). It contains a carbonyl group, \(\ce{C=O}\), but no OH group.

The carbonyl oxygen has lone pairs of electrons. A partially positive hydrogen atom from a water molecule is attracted to one of these lone pairs.

The resulting attraction is a hydrogen bond:

\[
\ce{H-O-H \bond{...} O=C(CH3)2}
\]

The dotted interaction represents an intermolecular hydrogen bond, not a new covalent bond.

**The important rule:** A molecule does not need an OH group to accept hydrogen bonds from water. It needs an appropriate electron-rich site, such as a lone pair on an oxygen, nitrogen, or fluorine atom.

## 2. Hydrogen bond donors and acceptors: two different jobs

There are two parts to a hydrogen bond:

- **Hydrogen bond donor:** provides a partially positive hydrogen atom covalently bonded to N, O, or F.
- **Hydrogen bond acceptor:** provides a lone pair of electrons that attracts the partially positive hydrogen.

These roles do not have to belong to the same molecule.

Water can donate its hydrogen atoms to an electron-rich site on another molecule. That other molecule does not need an OH group.

At the HSC level, use this decision table:

| Molecule | Hydrogen bond donor? | Hydrogen bond acceptor? | Can hydrogen bond with water? |
|---|---|---|---|
| Ethanol, \(\ce{CH3CH2OH}\) | Yes | Yes | Yes |
| Acetone, \(\ce{CH3COCH3}\) | No | Yes | Yes |
| Dimethyl ether, \(\ce{CH3OCH3}\) | No | Yes | Yes |
| Methylamine, \(\ce{CH3NH2}\) | Yes | Yes | Yes |
| Propane, \(\ce{CH3CH2CH3}\) | No | No | No, under the usual HSC hydrogen-bonding model |

Notice that acetone and dimethyl ether can accept hydrogen bonds despite having no OH groups.

Propane provides a useful contrast. It has neither a suitable hydrogen bond donor nor an appropriate N, O, or F lone-pair acceptor site.

[IMAGE: Comparison of five molecular structures: ethanol, acetone, dimethyl ether, methylamine, and propane. Annotate the donor and acceptor sites on each structure. Show a separate water molecule donating a hydrogen bond to the acceptor oxygen of acetone and dimethyl ether. Use solid lines for covalent bonds and dotted lines for intermolecular hydrogen bonds.]

### Can a molecule accept hydrogen bonds but not donate them?

Absolutely.

Dimethyl ether is a particularly useful example:

\[
\ce{CH3-O-CH3}
\]

Its oxygen has lone pairs, so it can accept hydrogen bonds from water.

However, it has no hydrogen directly bonded to oxygen, nitrogen, or fluorine. Its hydrogen atoms are bonded to carbon instead.

Under the standard HSC model, those carbon-bonded hydrogen atoms do not make it a hydrogen bond donor.

This means:
- A water molecule can donate a hydrogen bond to dimethyl ether.
- A dimethyl ether molecule cannot donate a conventional hydrogen bond to another dimethyl ether molecule.

The distinction is about the arrangement of atoms, not simply whether oxygen or nitrogen appears somewhere in the molecular formula.

## 3. Compare the whole structure, not just the functional group

Let's compare two molecules containing oxygen:

**Ethanal:**

\[
\ce{CH3-CHO}
\]

**Ethanol:**

\[
\ce{CH3-CH2-OH}
\]

Both contain oxygen, but their structures give them different hydrogen-bonding abilities.

Ethanal contains a carbonyl group, \(\ce{C=O}\). Its oxygen has lone pairs and can accept a hydrogen bond from water.

Ethanol contains an OH group. Its oxygen has lone pairs, allowing it to accept hydrogen bonds, while its O-H bond provides a partially positive hydrogen that can donate hydrogen bonds.

[IMAGE: Side-by-side structural comparison of ethanal and ethanol interacting with water. Show water donating a hydrogen bond to the carbonyl oxygen of ethanal. For ethanol, show both water donating to the ethanol oxygen and ethanol donating its OH hydrogen to a water oxygen. Keep all covalent bonds solid and hydrogen bonds dotted.]

The important difference is:

| Feature | Ethanal | Ethanol |
|---|---|---|
| Carbonyl group | Yes | No |
| OH group | No | Yes |
| Can accept hydrogen bonds from water | Yes | Yes |
| Can donate hydrogen bonds to water | No | Yes |

This is why identifying the correct functional group matters. You cannot determine hydrogen-bonding behaviour by spotting an oxygen atom and stopping there.

If you need a refresher on recognising carbonyl, hydroxyl, and other organic functional groups, see [HSC Chemistry: How to Identify Organic Functional Groups](https://tutorgum.com/learn/resources/hsc-chemistry-structures-of-organic-acids-and-bases/).

### One important limitation: hydrogen bonding does not guarantee solubility

You might now predict that any molecule containing an oxygen atom will dissolve easily in water.

Be careful.

Hydrogen bonding is one contribution to solubility, not the only one. The size of the non-polar hydrocarbon region also matters, as do other intermolecular attractions and the temperature.

For example, a long-chain alcohol still has an OH group that can hydrogen bond with water. However, its large non-polar carbon chain makes it less soluble than a small alcohol.

So, when predicting solubility, consider the entire molecule rather than counting its oxygen atoms.

## 4. How can an experiment provide evidence of hydrogen bonding?

Drawing a dotted line between water and acetone is a useful model. But how could we experimentally investigate whether the two molecules interact through hydrogen bonding?

Infrared spectroscopy provides a useful approach.

### The experimental question

Acetone has a carbonyl group, \(\ce{C=O}\), but no OH group.

If water hydrogen bonds with the carbonyl oxygen, the electronic environment around the carbonyl group should change. This can affect the vibration of the C=O bond.

We can investigate that change using an infrared (IR) spectrometer.

### How the apparatus works

In a simplified IR experiment:

1. An infrared source produces radiation across a range of frequencies.
2. The radiation passes through a sample held in a suitable cell.
3. The sample absorbs particular frequencies associated with molecular vibrations.
4. A detector measures the radiation that passes through.
5. The instrument produces a spectrum showing absorbance against wavenumber.

A wavenumber is measured in \(\mathrm{cm^{-1}}\). Different bonds absorb infrared radiation at characteristic wavenumbers.

[IMAGE: Scientifically accurate experimental apparatus diagram showing an infrared source on the left, an infrared beam passing through a sample cell containing acetone solution, and a detector on the right connected to a spectrum display. Label the direction of radiation, sample cell, detector, and the carbonyl absorption region on a simple spectrum.]

### What should the experiment show?

Prepare two samples:

- Reference sample: acetone dissolved in a dry solvent that does not donate hydrogen bonds.
- Test sample: acetone dissolved in water.

Keep the acetone concentration, temperature, and sample-cell path length controlled.

If water forms hydrogen bonds with the carbonyl oxygen, the C=O bond experiences a different electronic environment.

Hydrogen bonding to the carbonyl oxygen generally weakens the effective C=O stretching interaction. The carbonyl stretching absorption therefore shifts towards a lower wavenumber.

This is evidence consistent with hydrogen bonding between water and acetone.

### What could go wrong with the original experiment?

Suppose a student simply compares pure acetone with an acetone-water mixture.

They observe a change in the carbonyl absorption peak and immediately conclude that hydrogen bonding must be responsible.

That conclusion is too quick.

The samples have different solvent environments and potentially different concentrations. Solvent polarity can also affect vibrational frequencies. Water itself absorbs infrared radiation, which must be accounted for.

A better experiment uses a suitable reference solvent, matches the acetone concentrations and temperatures, and accounts for background absorption.

Even then, a spectral shift is evidence of an interaction, rather than a photograph of an individual hydrogen bond. Other experimental evidence and appropriate controls strengthen the interpretation.

### Worked measurement: Interpreting an IR shift

Consider the following illustrative data. These values are hypothetical and are provided to demonstrate the calculation, not as reported experimental measurements.

| Sample | C=O stretching absorption |
|---|---:|
| Acetone in dry reference solvent | \(1715\ \mathrm{cm^{-1}}\) |
| Acetone in water | \(1705\ \mathrm{cm^{-1}}\) |

Calculate the change in wavenumber.

**Step 1**: Identify the initial and final measurements.

Initial wavenumber:

\[
\tilde{\nu}_i=1715\ \mathrm{cm^{-1}}
\]

Final wavenumber:

\[
\tilde{\nu}_f=1705\ \mathrm{cm^{-1}}
\]

Here, \(\tilde{\nu}\) represents the wavenumber.

**Step 2**: Calculate the change.

\[
\begin{aligned}
\Delta\tilde{\nu}
&=\tilde{\nu}_f-\tilde{\nu}_i\\
&=1705-1715\\
&=-10\ \mathrm{cm^{-1}}
\end{aligned}
\]

**Step 3**: Interpret the result.

The negative value indicates a shift towards a lower wavenumber.

This is consistent with hydrogen bonding between water and the carbonyl oxygen of acetone.

Notice what has changed: the vibrational behaviour of the C=O bond.

The experiment does not indicate that the C=O covalent bond has broken. Acetone remains acetone.

## 5. Hydrogen bonding, boiling, and covalent bond breaking

There is another common misconception worth fixing.

Students sometimes hear that hydrogen bonds must be overcome during boiling and assume that covalent bonds within molecules are also being broken.

These are two very different processes.

**Boiling involves overcoming intermolecular attractions, not breaking the covalent bonds within molecules.**

Consider liquid water.

Water molecules attract one another through hydrogen bonds. When water boils, energy is supplied to allow molecules to separate and enter the gas phase.

The O-H covalent bonds inside each water molecule remain intact.

[IMAGE: Side-by-side molecular representation of liquid water and water vapour. In the liquid, show several intact H-O-H molecules connected by dotted intermolecular hydrogen bonds. In the vapour, show separated intact H-O-H molecules with no continuous hydrogen-bond network. Use solid lines for the O-H covalent bonds and dotted lines for intermolecular attractions.]

The same idea applies to an acetone-water mixture.

Hydrogen bonds between water and acetone contribute to the attractions in the liquid. During vaporisation, molecules separate as these intermolecular attractions are overcome.

The covalent C=O bond inside each acetone molecule does not need to break.

If covalent bonds were being broken, you would be describing a chemical reaction or decomposition, not ordinary boiling.

### A useful comparison: ethanol and dimethyl ether

These two molecules make the distinction particularly clear.

Both have the molecular formula:

\[
\ce{C2H6O}
\]

Both have a molar mass of approximately \(46.07\ \mathrm{g\,mol^{-1}}\).

However, their structures are different.

- Ethanol: \(\ce{CH3CH2OH}\)
- Dimethyl ether: \(\ce{CH3OCH3}\)

Their normal boiling points are also very different:

| Property | Ethanol | Dimethyl ether |
|---|---:|---:|
| Molecular formula | \(\ce{C2H6O}\) | \(\ce{C2H6O}\) |
| OH group | Yes | No |
| Hydrogen bond donor | Yes | No |
| Hydrogen bond acceptor | Yes | Yes |
| Normal boiling point | Approximately \(78.4^\circ\mathrm{C}\) | Approximately \(-24.8^\circ\mathrm{C}\) |

Why is the difference so large?

In pure liquid ethanol, molecules can form hydrogen bonds with other ethanol molecules. The OH group provides the hydrogen bond donor, while the oxygen provides an acceptor site.

In pure dimethyl ether, the oxygen can accept hydrogen bonds, but the molecules do not have a suitable hydrogen bond donor. They therefore cannot form the same conventional hydrogen-bond network with one another.

Ethanol has stronger overall intermolecular attractions, particularly because of hydrogen bonding, so more energy is needed to separate its molecules during boiling.

But here's the important twist:

Dimethyl ether can still accept hydrogen bonds from water.

Water supplies the partially positive hydrogen that dimethyl ether itself cannot provide.

[IMAGE: Molecular comparison of liquid ethanol and liquid dimethyl ether. Show ethanol molecules connected through intermolecular O-H...O hydrogen bonds. Show dimethyl ether molecules separated without conventional hydrogen bonds between them. Beside them, show a water molecule donating a hydrogen bond to the oxygen of dimethyl ether. Clearly distinguish pure-liquid interactions from water-solute interactions.]

The boiling-point comparison tells us about attractions between molecules of the pure substances. It does not tell us that dimethyl ether cannot hydrogen bond with a different substance.

Also, boiling points depend on all intermolecular attractions, including dispersion and dipole-dipole forces. Hydrogen bonding is the major distinction in this comparison, not the only factor affecting boiling point.

## 6. Worked examples

### Worked example: Can dimethyl ether hydrogen bond with water?

Dimethyl ether has the structural formula \(\ce{CH3OCH3}\).

Determine whether it can hydrogen bond with water. If it can, identify which molecule donates the hydrogen bond.

**Step 1**: Identify the relevant atom.

Dimethyl ether contains an oxygen atom bonded to two carbon atoms.

The oxygen has lone pairs of electrons, making it a suitable hydrogen bond acceptor.

**Step 2**: Check for a hydrogen bond donor.

There is no hydrogen directly bonded to oxygen, nitrogen, or fluorine.

Therefore, dimethyl ether cannot donate a conventional hydrogen bond.

**Step 3**: Consider the interaction with water.

Water has partially positive hydrogen atoms that can be attracted to the lone pairs on the ether oxygen.

**Conclusion:** Dimethyl ether can hydrogen bond with water as a hydrogen bond acceptor. It does not need an OH group because water supplies the hydrogen bond donor.

### Worked example: Calculate and interpret a carbonyl shift

An infrared experiment gives the following illustrative measurements for an organic molecule containing a carbonyl group:

- In a dry reference solvent: \(1720\ \mathrm{cm^{-1}}\).
- In water: \(1712\ \mathrm{cm^{-1}}\).

Calculate the percentage decrease in wavenumber and explain what the result suggests.

**Step 1**: Calculate the change.

\[
\Delta\tilde{\nu}=1712-1720
\]

\[
\Delta\tilde{\nu}=-8\ \mathrm{cm^{-1}}
\]

**Step 2**: Calculate the percentage decrease.

\[
\begin{aligned}
\%\text{ decrease}
&=\frac{8}{1720}\times100\\
&=0.47\%
\end{aligned}
\]

**Step 3**: Interpret the measurement.

The carbonyl stretching absorption has shifted to a lower wavenumber by approximately \(0.47\%\).

This is consistent with hydrogen bonding between water and the carbonyl oxygen, provided other solvent effects have been controlled.

The carbonyl group is still present. Its covalent bond has not been broken.

## Questions and solutions

### Question 1

Acetone has the structural formula:

\[
\ce{CH3COCH3}
\]

Can acetone form hydrogen bonds with water? Identify whether it acts as a donor, an acceptor, or both.

### Solution 1

Acetone can form hydrogen bonds with water, but it acts only as a hydrogen bond acceptor.

Its carbonyl oxygen has lone pairs of electrons that can attract a partially positive hydrogen atom from water.

However, acetone has no hydrogen directly bonded to oxygen, nitrogen, or fluorine. It therefore cannot donate a conventional hydrogen bond.

The interaction is between a hydrogen atom on water and the carbonyl oxygen on acetone.

The absence of an OH group does not prevent hydrogen bonding.

### Question 2

Consider two molecules:

- Molecule A: ethanol, \(\ce{CH3CH2OH}\).
- Molecule B: ethanal, \(\ce{CH3CHO}\).

Both are added separately to water.

**(a)** Identify the hydrogen bond donor and acceptor sites in each molecule.

**(b)** Explain why both molecules can hydrogen bond with water, despite having different functional groups.

**(c)** Which molecule can donate hydrogen bonds to water? Explain.

### Solution 2

**(a)** Ethanol has both a hydrogen bond donor site and an acceptor site. Its OH group provides the donor hydrogen, while the oxygen has lone pairs that can accept hydrogen bonds.

Ethanal has an acceptor site at its carbonyl oxygen, but it has no suitable hydrogen bond donor.

**(b)** Both molecules contain an oxygen atom with available lone pairs.

In ethanol, the oxygen in the OH group can accept a hydrogen bond from water.

In ethanal, the carbonyl oxygen can also accept a hydrogen bond from water.

The oxygen does not need to be part of an OH group to act as an acceptor.

**(c)** Ethanol can donate hydrogen bonds to water, while ethanal cannot under the standard HSC model.

The hydrogen in ethanol's O-H bond is partially positive because oxygen attracts the shared bonding electrons.

Ethanal has hydrogen atoms bonded to carbon, not to an electronegative N, O, or F atom. These are not conventional hydrogen bond donor sites.

### Question 3

Methylamine, \(\ce{CH3NH2}\), and dimethyl ether, \(\ce{CH3OCH3}\), are considered separately.

**(a)** Explain which molecule can both donate and accept hydrogen bonds with water.

**(b)** Predict which molecule can form conventional hydrogen bonds between its own molecules in a pure sample.

**(c)** Explain why dimethyl ether can interact through hydrogen bonding with water but cannot form the same type of hydrogen-bond network with itself.

### Solution 3

**(a)** Methylamine can both donate and accept hydrogen bonds with water.

Its N-H bonds contain partially positive hydrogen atoms, allowing it to donate hydrogen bonds. Its nitrogen also has a lone pair that can accept hydrogen bonds.

Dimethyl ether can accept hydrogen bonds through its oxygen lone pairs, but it cannot donate conventional hydrogen bonds.

**(b)** Methylamine can form conventional hydrogen bonds between its own molecules. Dimethyl ether cannot do so under the standard HSC model.

**(c)** The difference is the availability of a suitable hydrogen bond donor.

Methylamine has both N-H bonds and a lone pair on nitrogen. One molecule can therefore donate a hydrogen bond to the nitrogen of another methylamine molecule.

Dimethyl ether has oxygen lone pairs, but no hydrogen bonded to N, O, or F. Two dimethyl ether molecules therefore lack the necessary donor-acceptor combination for conventional hydrogen bonding between themselves.

Water changes the situation because it supplies the missing hydrogen bond donor.

A water hydrogen can interact with the oxygen lone pair on dimethyl ether.

This demonstrates why hydrogen bonding depends on the molecules interacting, not simply on the properties of one molecule considered in isolation.

### Question 4

An experiment investigates whether acetone can hydrogen bond with water using infrared spectroscopy.

The following illustrative results are obtained:

| Sample | Carbonyl stretching wavenumber |
|---|---:|
| Acetone in a suitable dry reference solvent | \(1718\ \mathrm{cm^{-1}}\) |
| Acetone in water | \(1709\ \mathrm{cm^{-1}}\) |

**(a)** Calculate the change in wavenumber.

**(b)** Calculate the percentage decrease.

**(c)** Explain how the result supports the possibility of hydrogen bonding between water and acetone.

**(d)** A student concludes that the carbonyl covalent bond has broken because its absorption frequency changed. Explain why this conclusion is incorrect.

### Solution 4

**(a)** The carbonyl absorption shifts down by \(9\ \mathrm{cm^{-1}}\).

Using:

\[
\Delta\tilde{\nu}
=\tilde{\nu}_f-\tilde{\nu}_i
\]

Substitute:

\[
\begin{aligned}
\Delta\tilde{\nu}
&=1709-1718\\
&=-9\ \mathrm{cm^{-1}}
\end{aligned}
\]

The negative value indicates a shift to a lower wavenumber.

**(b)** The percentage decrease is approximately \(0.52\%\).

\[
\begin{aligned}
\%\text{ decrease}
&=\frac{9}{1718}\times100\\
&=0.52\%
\end{aligned}
\]

**(c)** The shift is consistent with hydrogen bonding between water and the carbonyl oxygen of acetone.

The partially positive hydrogen on water interacts with the oxygen lone pair. This changes the electronic environment of the carbonyl group and affects its stretching vibration.

The observation supports the hydrogen-bonding model, although solvent effects must be controlled before attributing the entire shift to hydrogen bonding.

**(d)** The covalent carbonyl bond has not necessarily broken.

An infrared absorption frequency depends on the vibrational properties of a bond, which are influenced by its electronic environment.

An intermolecular attraction can change the bond's vibrational frequency without destroying the covalent bond.

Acetone remains a molecule containing a C=O group. The experiment detects a change in its vibrational behaviour, not the breaking of that bond.

### Question 5

Two organic molecules, A and B, have the same molecular formula, \(\ce{C2H6O}\), and the same molar mass.

Their structures and normal boiling points are:

| Property | Molecule A | Molecule B |
|---|---|---|
| Structure | \(\ce{CH3CH2OH}\) | \(\ce{CH3OCH3}\) |
| Name | Ethanol | Dimethyl ether |
| Normal boiling point | \(78.4^\circ\mathrm{C}\) | \(-24.8^\circ\mathrm{C}\) |

Separate aqueous samples of both molecules are investigated using infrared spectroscopy. A controlled experiment shows changes in the absorption associated with water's O-H vibrations when either molecule is added.

A student makes the following claims:

1. Molecule B cannot hydrogen bond with water because it has no OH group.
2. Molecule B must contain an OH group because the water O-H absorption changes when B is added.
3. The much lower boiling point of B proves that it cannot form hydrogen bonds at all.

Evaluate all three claims.

Your answer should explain the molecular structures, the direction of hydrogen bonding, the boiling-point difference, and what the infrared evidence can establish.

[IMAGE: Question setup showing the structural formula of ethanol beside dimethyl ether, with a separate water molecule near each. Label the molecules A and B and show their supplied structural connections accurately. Do not draw hydrogen bonds, donor or acceptor arrows, or any completed interpretation.]

### Solution 5

All three claims are incorrect.

The central issue is that the student is confusing hydrogen bonding between molecules of a pure substance with hydrogen bonding between a solute and water.

**Claim 1: Molecule B cannot hydrogen bond with water because it has no OH group.**

This is incorrect.

Molecule B is dimethyl ether:

\[
\ce{CH3-O-CH3}
\]

Its oxygen has lone pairs that can accept hydrogen bonds.

Water supplies partially positive hydrogen atoms, which can interact with these lone pairs.

Therefore, dimethyl ether can accept hydrogen bonds from water even though it cannot donate conventional hydrogen bonds itself.

**Claim 2: Molecule B must contain an OH group because the water O-H absorption changes.**

This is also incorrect.

The change in the water O-H absorption is consistent with water molecules experiencing a different hydrogen-bonding environment.

In this case, water can donate hydrogen bonds to the oxygen in dimethyl ether.

The water O-H bond is involved in the interaction. Dimethyl ether does not need an O-H bond of its own.

The observed spectral change is therefore compatible with dimethyl ether acting as a hydrogen bond acceptor.

**Claim 3: The lower boiling point proves that molecule B cannot form hydrogen bonds at all.**

This confuses two different situations.

In pure liquid ethanol, molecules can donate and accept hydrogen bonds. Its OH group allows neighbouring ethanol molecules to form intermolecular hydrogen bonds.

In pure dimethyl ether, molecules have oxygen acceptor sites but lack conventional hydrogen bond donor sites. They cannot form the same hydrogen-bond network with one another.

Consequently, ethanol has stronger overall intermolecular attractions, particularly due to hydrogen bonding, and a much higher normal boiling point.

However, when dimethyl ether interacts with water, water provides the hydrogen bond donor.

The lower boiling point of pure dimethyl ether does not prevent it from accepting hydrogen bonds in an aqueous solution.

**The final distinction:**

- Ethanol can donate and accept hydrogen bonds with water.
- Dimethyl ether can accept hydrogen bonds from water but cannot donate conventional hydrogen bonds.
- The boiling-point difference reflects the different intermolecular attractions in the pure substances.
- The infrared changes provide evidence of altered molecular interactions, not evidence that new OH covalent bonds have appeared or that existing covalent bonds have broken.

This is the key reasoning: hydrogen bonding is an interaction between suitable donor and acceptor sites on neighbouring molecules. A molecule does not have to supply both sites itself.

## 7. The decision rule to remember

When an HSC Chemistry question asks whether an organic molecule can form hydrogen bonds with water, use this approach:

1. Find any N, O, or F atoms in the complete structural formula.
2. Check whether any hydrogen is directly bonded to N, O, or F. This identifies possible donor sites.
3. Check for suitable lone pairs on N, O, or F. These identify possible acceptor sites.
4. Consider water as both a possible donor and acceptor.
5. Identify the actual interaction between the two molecules.

The most important distinction is:

**An OH group is needed for an alcohol to donate hydrogen bonds through its hydroxyl hydrogen, but it is not needed for a molecule to accept hydrogen bonds from water.**

Once you can identify donors and acceptors, you can use the same reasoning to predict intermolecular attractions, compare boiling points, and begin making more informed predictions about solubility.