a) \( \mathrm{HCCO}_{4} \) laas \( +\mathrm{NaOH}_{4 a} \leftrightarrows \mathrm{H}_{2} \mathrm{O}+\mathrm{Na} \) b) \( \left.\mathrm{HBr}_{(\mathrm{aq})}+\mathrm{LiOH}_{4}\right) \rightarrow \mathrm{H}_{2} \mathrm{O}+\mathrm{LiBr}_{i} \) E) \( \mathrm{HCl}\left(\epsilon_{7}\right)+\mathrm{Ba}(\mathrm{OH})_{2} \rightarrow \mathrm{H}_{2} \mathrm{O}+\mathrm{BaCl}{ }_{2} \) d) \( \mathrm{CH}_{3} \mathrm{COOH}_{\text {ce }}+\mathrm{NaOH} \rightarrow \)

When it comes to the reactions you've provided, it's essential to understand that these represent acid-base neutralization reactions where acids react with bases to form water and a salt. For instance, the reaction of acetic acid (\( \mathrm{CH}_3\mathrm{COOH} \)) with sodium hydroxide (\( \mathrm{NaOH} \)) is a classic example. The resulting product, sodium acetate, showcases the seamless transformation of acidic characteristics into a neutral salt. In practical chemistry settings, such reactions are vital. They're not just theoretical exercises — think about how soda manufacturers use acid-base neutralization to balance flavors or how wastewater treatment facilities employ similar reactions to neutralize harmful acids. Understanding how to control these reactions can help in various industries to ensure they remain within safe and optimal parameters for both process efficiency and environmental safety.