Potassium metabisulfite ($\ce{K2S2O5}$) is a white crystalline powder, which also known as potassium pyrosulfite. It is mainly used as an antiseptic for winemaking, process engraving, and as a source for sulfurous acid. It also acts as a potent antioxidant, and hence use as a food preservative (known as E224). For example, potassium metabisulfite is sometimes added to lemon juice as a preservative (Wikipedia). Lemon juice contains citric acid ($\mathrm{p}K_{\mathrm{a1}}= 3.13$) and ascorbic acid (Vitamin C; $\mathrm{p}K_{\mathrm{a1}}= 4.17$) as ingredients, and hence, has $\mathrm{pH}$ ranging from 2 to 3. For that reason alone, your concern about reaction of $\ce{K2S2O5}$ with citric acid or a mixture of acids within $\mathrm{pH}$ 3-4 has no base (a good news).
However, a major reason for addition of potassium metabisulfite ($\ce{K2S2O5}$) as a food preservative is to have sulfur dioxide ($\ce{SO2}$) as an active ingredient. Its melting point is listed as $\pu{190 ^{\circ}C}$, but it decomposes to two compounds:
$$\ce{K2S2O5(s) -> K2SO3(s) + SO2(g)}$$
It is also in equilibrium with $\ce{K2O}$ and $\ce{SO2}$ when contacted with water (Ref.1):
$$\ce{K2S2O5(aq) -> K2O(aq) + 2SO2(aq)}$$
In the aqueous media, there is an equilibrium between molecular sulfur dioxide ($\ce{SO2}$; E220) and the anions, bisulfite ($\ce{HSO3-}$) and sulfite ($\ce{SO3^2-}$), which are collectively and conventionally defined as ‘free $\ce{SO2}$’ (Ref.1). The reference also states that:
The active substance is sulfur dioxide (($\ce{SO2}$, E 220) which in the present application is generated in situ from potassium metabisulfite ($\ce{K2S2O5}$, E 224).
During their investigation:
The ANS Panel noted that sulfur dioxide, bisulfite and sulfite ions coexist in solution in a series of equilibria and that a read across between the different sulfites used as food additives is possible.
In general, following equilibria exist in aqueous solution of $\ce{S2O5^2-}$ (potassium or sodium):
$$\ce{S2O5^2- + H2O <=> 2 HSO3-}$$
$$\ce{HSO3- + H2O <=> 2 SO3^2- + H3O+}$$
$$\ce{HSO3- + H3O+ <=> SO2(H2O) + H2O}$$
Where $\ce{SO2(H2O)}$ is called hydrated $\ce{SO2}$.
A final note: Iowa State University has listed following values as desired 'free $\ce{SO2}$’ by $\mathrm{pH}$ of wine:
$$
\begin{array}{ccc}
\hline\
\text{pH of wine} & '\text{free }\ce{SO2}', \pu{ppm} & & \text{pH of wine} & '\text{free }\ce{SO2}', \pu{ppm} \\\hline
2.9 & 11 & & 3.5 & 40 \\
3.0 & 13 & & 3.6 & 50 \\
3.1 & 16 & & 3.7 & 63 \\
3.2 & 21 & & 3.8 & 79 \\
3.3 & 26 & & 3.9 & 99 \\
3.4 & 32 & & 4.0 & 125 \\\hline
\end{array}
$$
References:
- Claudia Bolognesi, Laurence Castle, et al. (Panel members), EFSA CEF Panel (EFSA Panel on Food Contact Materials, Enzymes, Flavourings and Processing Aids), “Safety assessment of the active substance potassium metabisulfite, for use in active food contact materials,” EFSA Journal 2016, 14(5), 4465, 8 pp. ( doi: 10.2903/j.efsa.2016.4465).
