Reactions of Metals with Water and Acids
Chemical changes • Reactivity of metals
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Key concepts
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Reactivity and general definitions
Reactivity measures a metal’s tendency to lose electrons and form positive ions. Metals higher in the reactivity series lose electrons more readily and react with water or acids to produce hydrogen and a metal salt or hydroxide. Dilute acids provide H+ ions that accept electrons from metals, producing hydrogen gas and a soluble metal salt. Limiting factors include temperature, surface area, presence of oxide layers and acid concentration. Cold water reactions occur for the most reactive metals. Less reactive metals require steam or dilute acid to react. Some metals form protective oxide layers that prevent reaction with water or dilute acids.
Reactions with cold water
Very reactive metals undergo vigorous reactions with cold water to form hydroxides and hydrogen. Potassium, sodium and lithium react violently with cold water; calcium reacts less violently but still produces hydrogen and a slightly soluble alkaline hydroxide. Observations include rapid fizzing, heat release and, for potassium, possible flame from ignited hydrogen. Magnesium, zinc, iron and copper do not react appreciably with cold water under normal conditions. Magnesium reacts very slowly with cold water because a protective surface layer of hydroxide forms; significant reaction requires steam or higher temperatures.
Reactions with steam
Metals lower in the reactivity series react with steam rather than cold water. Steam supplies higher-energy water molecules that remove oxygen to form metal oxides and release hydrogen. Magnesium and zinc react with steam to form metal oxides and hydrogen gas; iron reacts with steam at high temperature to form iron(II,III) oxide and hydrogen. Equations for steam reactions: Mg + H2O (steam) -> MgO + H2; Zn + H2O (steam) -> ZnO + H2; Fe + H2O (steam) -> FeO + H2 (or mixtures of iron oxides depending on conditions).
Reactions with dilute acids
Dilute acids react with many metals to form soluble metal salts and hydrogen gas. Reaction occurs when the metal is more reactive than hydrogen in the reactivity series. Typical acid used in examples is dilute hydrochloric acid, producing chlorides and H2. Equations follow the pattern: metal + acid -> metal salt + hydrogen. Examples: Zn + 2HCl -> ZnCl2 + H2; Fe + 2HCl -> FeCl2 + H2. Copper does not react with dilute non-oxidising acids because copper lies below hydrogen in the reactivity series; copper requires oxidising acids to dissolve.
Specific metal behaviours — summary
Potassium, sodium and lithium react vigorously and often explosively with cold water, producing hydroxides and hydrogen. Observations include immediate fizzing, strong exotherm, melting of metal (sodium/potassium) and possible flame with potassium. Calcium reacts with cold water, producing Ca(OH)2 and H2 with slower evolution than alkali metals. Magnesium shows little reaction with cold water but reacts with steam to give MgO and H2; with dilute acid magnesium dissolves to form Mg2+ salts and hydrogen. Zinc and iron do not react with cold water; both react with dilute acids to produce salts and hydrogen. Copper shows almost no reaction with cold water or dilute acids.
Safety, limiting factors and experimental observations
Observation describes cause → effect: metal oxidation (cause) produces hydrogen gas and heat (effect). Hydrogen gas presence appears as effervescence and produces a squeaky pop when ignited. Reaction speed increases with higher temperature, higher acid concentration and greater metal surface area. Protective oxide or hydroxide layers reduce reaction rate or prevent reaction entirely. Safety consideration: highly reactive metals produce large volumes of hydrogen and heat; gas ignition and explosions can occur. Dilute acid concentration and small metal samples reduce risk and allow observation of characteristic reactions without extreme violence.
Key notes
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