Provide the formula for each acid. a. hydrofluoric acid b. hydrobromic acid c. sulfurous acid

When studying hydrofluoric acid, it's crucial to recognize its unique properties and risks associated with handling it. Hydrofluoric acid, with the chemical formula HF, is not only used for etching glass but also in the production of refrigerants and herbicides.

The formation of HF is relatively simple. One hydrogen atom (H) bonds with one fluorine atom (F), because fluorine is a halogen that typically forms a single negative ion, denoted as F^-. This results in a stable diatomic molecule, HF, which, despite being a weak acid, is highly corrosive and can pose significant health hazards such as tissue damage upon contact.

Students should handle advanced exercises involving HF with care, keeping in mind that despite its simple formula, hydrofluoric acid requires a complex understanding of its properties and safe usage in practical applications.

Hydrobromic acid is another hydrogen halide, akin to hydrofluoric acid, but instead features bromine. This acid, symbolized as HBr, is widely utilized in industry for the synthesis of bromine compounds and as a catalyst in certain reactions.

Its construction is straightforward: a single proton, or hydrogen ion (H), combines with a bromide ion (Br^-), which is the form bromine takes when it gains an electron. This bond creates a strong, corrosive acid that is much stronger than hydrofluoric acid. However, HBr is also a gas at room temperature and standard atmospheric pressure, which means that the acid is usually found in solution form.

Focus on exercises that encourage the safe handling and application of hydrobromic acid, emphasizing the need to understand the concentration of the solution, as this affects its acidity and potential uses in the lab and industry.

Sulfurous acid's formula, H₂SO₃, reveals a compound different from the previous halogen acids. It involves sulfur, a non-metal element that forms various oxoacids. Sulfurous refers to the presence of sulfur in a lower oxidation state, combined with oxygen. Here, the focus should be on the sulfite ion (SO₃^2-), which pairs with two hydrogen ions to neutralize the charge.

The inclusion of oxygen changes the nature of the acid, creating what is known as an oxyacid. Sulfurous acid is less common and more unstable than sulfuric acid (H₂SO₄), and it exists mostly in solution. It's vital to approach exercises with an understanding of the delicate balance between oxygen, sulfur, and hydrogen which dictates its reactivity and uses. Practical applications include acting as an intermediate in the synthesis of other chemicals and as a reducing agent in industry.

In conclusion, understanding these acids' formulas is the foundational step towards appreciating their unique chemical behaviors, risks, and roles in various industrial processes.