Negative charged intermediate in acidic conditions - what's the problem?
Acidic conditions — whether you are using the Brønsted or the Arrhenius definition of acidity — is characterised by readily available protons. Remember the general dissolution ionisation of a Brønsted acid in water  as shown below:

$$\ce{HA + H2O &lt;=&gt; H3O+ + A-}\tag{1}$$

The proton is readily liberated and transferred to species which are more happy to accept it — for strong acids, that means it is transferred to water, a really weak base. A proton typically gladly attaches to anything that has a negative charge, because those species are typically more basic than water in this example. Therefore, postulating an enolate or any other oxygen-centred anion in acidic solution is not considered correct since there will be enough protons to quench it.

However, this does not mean that no anions can exist in aquaeous solution. For example, a typical method to transform alcohols into iodides is refluxing in aquaeous hydrogen iodide. Here, $\ce{HI}$ is a very strong acid that very readily dissociates to give iodide anions, $\ce{I-}$. These, even though negatively charged, do exist in acidic conditions and are able to attack nucleophilicly. That’s all due to the high acidity of $\ce{HI}$.

$$\ce{R-OH + H3O+ + I- &lt;&lt;=&gt; R-OH2+ + H2O + I- -&gt; R-I + 2 H2O}\tag{2}$$