Reactivity Series

The tendency to form a cation depends on the atomic structure of the metal.

• Metals higher in the series have a greater atomic radius and more electronic shielding from inner electron shells.
• This reduces the electrostatic attraction between the positively charged nucleus and the negatively charged outer electrons.
• As a result, the outer electrons are lost much more easily, making the metal highly reactive.

The general half-equation for the oxidation of a metal ($M$) forming an $n+$ cation is:

$$M \rightarrow M^{n+} + ne^-$$

Reactions of Metals with Water

Dropping a piece of potassium into water looks more like a firework display than a typical chemistry experiment.

• When a reactive metal reacts with cold water, it produces a metal hydroxide and hydrogen gas.
• The visible observations depend heavily on the metal's position in the reactivity series.
• Potassium reacts violently, melting into a ball, darting across the surface, and burning with a distinctive lilac flame.
• Sodium reacts vigorously by melting into a shiny ball with rapid effervescence (fizzing), while lithium fizzes more steadily and shrinks until it disappears.

Calcium reacts moderately to produce steady bubbles, and the water turns cloudy because the calcium hydroxide formed is only slightly soluble.

• Metals below calcium in the series do not react noticeably with cold water.
• However, magnesium will react vigorously with steam to produce a metal oxide (magnesium oxide, a white powder) instead of a hydroxide, alongside hydrogen gas.

The balanced equation for magnesium reacting with steam is:

$$Mg(s) + H_2O(g) \rightarrow MgO(s) + H_2(g)$$

Reactions of Metals with Dilute Acids

Understanding exactly how metals behave in acid allows chemists to safely predict and control industrial reactions.

• Metals positioned above hydrogen in the reactivity series react with dilute acids to produce a salt and hydrogen gas.
• These reactions are exothermic, meaning they release heat energy to the surroundings, causing the reaction vessel to become hot.
• You can confirm the presence of the hydrogen gas produced by using the squeaky pop test, where a lit splint held near the mouth of the test tube creates a distinct popping sound.

The vigor of the acid reaction perfectly mirrors the reactivity series.

• Potassium and sodium are never reacted with dilute acids in a school laboratory because the reactions are dangerously explosive.
• Magnesium reacts very vigorously with rapid effervescence, zinc reacts steadily with moderate bubbling, and iron reacts very slowly.
• Metals below hydrogen, such as copper, silver, and gold, simply do not react with dilute acids at all.
• Aluminium is an anomaly; it often appears unreactive because it is coated in a tough, protective oxide layer ($Al_2O_3$) that prevents the acid from reaching the underlying metal.

The balanced equation for zinc reacting with dilute sulfuric acid is:

$$Zn(s) + H_2SO_4(aq) \rightarrow ZnSO_4(aq) + H_2(g)$$

Investigating Reactivity Experimentally

Every time you compare two chemical reactions, you have to ensure the playing field is completely level.

• To scientifically prove the order of the reactivity series, you can quantify the reactions by measuring specific variables.
• You can measure the rate of effervescence, track the maximum temperature change using a polystyrene cup (which acts as an insulator), or measure the total volume of hydrogen gas produced over time using a gas syringe.
• Another powerful experimental tool is the displacement reaction, where a more reactive metal forces a less reactive metal out of a compound (e.g., adding magnesium to copper sulfate solution will displace the copper).

To ensure your acid-metal investigation is a fair test, you must strictly control specific variables.

• You must use the same concentration and volume of the dilute acid.
• You must also control the mass (or moles) of the metal and its surface area (for instance, ensuring you always use uniformly sized metal powders rather than mixing powders and solid ribbons).