Press Start to drop the magnesium into the acid.
Not to time scale. Water molecules are not shown. The share of successful collisions is exaggerated so you can see it; in reality only a tiny fraction succeed. More concentrated acid = more H⁺ ions in the same space = more hits on the magnesium every second.
| # | HCl (M) | Start T (°C) | Acid (cm³) | Magnesium | Time (s) | 1/time (s⁻¹) | Initial rate (cm³/s) | H₂ (cm³) |
|---|
Readings as a student would write them: the syringe read to the nearest 1 cm³, the thermometer to 0.5 °C. Use Copy readings to paste them into a spreadsheet and draw your own graph.
Click any piece of equipment in the 3D lab, or a button below, to find out what it does. The strip held by the tweezers is the magnesium ribbon.
Settings are locked while a run is going. Press New run to change them.
Some reactions are over in a flash (fireworks), others take years (rusting). The rate of reaction tells you how fast reactants are used up or products are made. Drop a piece of magnesium into hydrochloric acid and it fizzes, because bubbles of hydrogen gas are being made, and the shiny metal slowly disappears. The faster the fizzing and the sooner the magnesium is gone, the higher the rate. Think of a crowded school corridor: the more people squeezed into it, the more often someone bumps into the lockers.
The word equation is magnesium + hydrochloric acid → magnesium chloride + hydrogen, and the balanced symbol equation is $\mathrm{Mg(s) + 2HCl(aq) \rightarrow MgCl_2(aq) + H_2(g)}$. In the acid the particles that actually react are the hydrogen ions: $\mathrm{Mg(s) + 2H^+(aq) \rightarrow Mg^{2+}(aq) + H_2(g)}$. The chloride ions are spectator ions. Collision theory says particles can only react when they collide, and only collisions with at least a minimum amount of energy, the activation energy, succeed. A more concentrated acid has more H⁺ ions in the same volume, so more of them hit the magnesium every second. A hotter acid makes the ions move faster, so they collide more often and a much larger share of the collisions have enough energy.
Collect the hydrogen in a gas syringe and read the volume every 10 seconds, or simply time how long the magnesium takes to disappear. The mean rate is the amount made divided by the time:
On a volume–time graph the rate at any moment is the gradient. The curve is steepest at the start, when there is most metal surface, then gets less steep as the ribbon thins, pits and breaks up, and goes flat when the last piece has gone. The initial rate is the gradient of the tangent at t = 0 (the dashed white line in the lab). If the same mass of magnesium is used every time, a shorter time means a higher rate, so 1/time is a quick measure of rate. The final volume is the same for every concentration, because the same amount of magnesium is used and the acid is in excess.
At A-level/AP the rate is written as a rate equation, $\text{rate} = k[\mathrm{H^+}]^n$, where $n$ is the order of reaction with respect to H⁺. If $n = 1$, doubling the concentration doubles the initial rate. To find $n$, plot $\log(\text{rate})$ against $\log[\mathrm{HCl}]$: the gradient is $n$ (tick log–log). The rate constant $k$ depends on temperature through the Arrhenius equation, $k = A\,e^{-E_a/RT}$, so a graph of $\ln(\text{rate})$ against $1/T$ is a straight line of gradient $-E_a/R$ (choose vs temperature and tick Arrhenius plot). The reaction is exothermic (about $-467$ kJ per mole of Mg), so the acid warms by a few degrees during a run and speeds itself up a little, which is one reason real class results are never perfectly tidy.
(1) Press Run the full investigation and compare the curves: all end at the same volume, but the steep start gets steeper as the acid gets stronger. (2) Press Temperature series and watch the time fall from about 1.5 minutes at 20 °C to about 13 s at 60 °C. (3) Turn on Real-lab variation and repeat 1.0 M three times: the bung goes in late, some gas escapes, the plunger sticks and jumps, and every time is a little different, so find the mean. (4) Try the same ribbon cut into 6 pieces (almost no change) and then powder (very fast). (5) Switch to Mass loss and see how little the balance changes, because hydrogen is so light.
| Variable | Concentration investigation | Temperature investigation |
|---|---|---|
| Independent | Concentration of HCl | Starting temperature of the acid |
| Dependent | Volume of hydrogen over time (initial rate), or time for the magnesium to disappear | |
| Control | Temperature, volume of acid, length and cleanliness of the ribbon | Concentration, volume of acid, length and cleanliness of the ribbon |
Safety: wear eye protection; hydrochloric acid of 2 M or more is an irritant; hydrogen is flammable, so keep flames away; do not heat the acid above about 60 °C; magnesium powder is a highly flammable solid, so use only small amounts, keep it away from flames and add it slowly.
3 cm of clean ribbon (0.031 g) in 50 cm³ of acid, model order 1, no real-lab errors. Your own results will differ a little, as real ones do.
| [HCl] (M) at 20 °C | Time for Mg to disappear (s) | 1/time (s⁻¹) | Initial rate (cm³/s) |
|---|---|---|---|
| 2.0 | 44 | 0.023 | 1.8 |
| 1.5 | 60 | 0.017 | 1.4 |
| 1.0 | 91 | 0.011 | 0.95 |
| 0.5 | 191 | 0.0052 | 0.48 |
| 0.25 | 424 | 0.0024 | 0.24 |
Halving the concentration roughly halves the initial rate and doubles the time, and a graph of rate against concentration is a straight line through the origin: the rate is directly proportional to the concentration (first order).
| Start temperature, 1.0 M | Time (s) | 1/time (s⁻¹) | Initial rate (cm³/s) |
|---|---|---|---|
| 20 °C | 91 | 0.011 | 0.95 |
| 30 °C | 55 | 0.018 | 1.6 |
| 40 °C | 33 | 0.030 | 2.5 |
| 50 °C | 21 | 0.048 | 3.8 |
| 60 °C | 13 | 0.076 | 5.5 |
Each 10 °C rise makes the reaction about 1.4–1.7 times faster here (the factor gets smaller at higher temperatures). The exact factor depends on the activation energy, which has to be measured.
| Error | Effect | How to reduce it |
|---|---|---|
| Bung put back late | Gas escapes at the start, so early volumes are too low | Practise; use a side-arm flask or a tube to drop the ribbon in |
| Bung pushes air in | The syringe shows 1–3 cm³ before any gas is made | Note the starting reading and subtract it |
| Sticky plunger | Volume goes up in jumps, readings lag | Clean, dry syringe; tap it gently before reading |
| Reaction time on the stopwatch | ±0.2–0.5 s on every time | Repeat and take a mean; use longer runs |
| Ribbon not identical | Thickness and oxide vary between pieces | Cut all pieces from the same roll; sand every piece |
| Solution warms up | The rate increases during the run | Use more acid or less magnesium |
Turn on Real-lab variation to see all of these in the simulation.
| Factor | Change | Why (collision theory) |
|---|---|---|
| Concentration | Higher → faster | More H⁺ ions per cm³, so more collisions with the metal each second |
| Temperature | Higher → much faster | Particles move faster: more collisions, and a much larger share have the activation energy |
| Surface area | Larger → faster | More magnesium atoms exposed to the acid. Powder reacts far faster than ribbon |
| Catalyst | Faster | Gives a route with lower activation energy (not normally used for this reaction) |
The magnesium dissolves from its surface at a speed proportional to $[\mathrm{H^+}]^n$. As it reacts the exposed area shrinks with the mass, $A \propto m^{2/3}$ (the ribbon thins, pits and breaks up), which gives the curved volume–time graph. The speed is set so that 3 cm of clean ribbon in 1.0 M acid at 20 °C is used up in about 1.5 minutes; with 6 cm this gives about 26, 43 and 60 cm³ at 15, 30 and 60 s, close to an example data set used in A-level revision material (25, 41 and 60 cm³). Temperature works through the Arrhenius equation with an assumed activation energy of 40 kJ/mol (measured values for this reaction vary between experiments). The heat of reaction warms the acid and flask, and heat is lost to the room. The oxide-layer delay and the real-lab errors are simple illustrations, not measured data. The 3D bubbles are drawn at a rate that follows the reaction rate but are not counted one-for-one.
A more concentrated acid contains more hydrogen ions (H+) in the same volume. More of them hit the surface of the magnesium every second, so there are more successful collisions per second and the reaction is faster. The ions do not move faster; there are simply more of them to collide.
Key takeaway: more particles in the same space = more frequent collisions = faster reaction.Magnesium + hydrochloric acid → magnesium chloride + hydrogen. The balanced symbol equation is Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g). The ionic equation is Mg(s) + 2H+(aq) → Mg2+(aq) + H2(g); the chloride ions are spectator ions. It is a redox reaction: magnesium is oxidised and hydrogen ions are reduced.
Key takeaway: Mg + 2HCl → MgCl2 + H2.Collect the hydrogen in a gas syringe (or an upturned measuring cylinder full of water) and record the volume every 10 seconds, or time how long the magnesium takes to disappear. The initial rate is the gradient of the tangent to the volume–time curve at t = 0. If the same mass of magnesium is used every time, 1/time is proportional to the average rate, so it can be used as a quick measure. Keep the volume of acid, the length of ribbon and the temperature the same.
Key takeaway: gas volume vs time (initial rate from the tangent), or time for the ribbon to disappear.The graph is steepest at the start, when there is most magnesium surface. As the ribbon thins, develops holes and breaks into smaller pieces, there is less surface for the acid to attack, so the gradient falls, and the graph goes flat when the last piece has reacted. With the acid in excess, the magnesium is the limiting reactant, so the same mass of magnesium always makes the same volume of hydrogen (about 31 cm3 from 3 cm of ribbon). Higher concentration only makes it arrive sooner.
Key takeaway: concentration changes how fast, not how much (when Mg is limiting).Raising the temperature makes the reaction much faster. The hydrogen ions move faster, so they hit the magnesium more often, and, more importantly, a much larger fraction of the collisions have at least the activation energy. In this simulation 3 cm of ribbon in 1.0 M acid takes about 91 s at 20 °C but only about 13 s at 60 °C. Plotting ln(rate) against 1/T (an Arrhenius plot) gives a straight line whose gradient is −Ea/R.
Key takeaway: hotter = more frequent and more energetic collisions = much faster.Yes. About 467 kJ of heat is released for every mole of magnesium. With 0.03 g of magnesium in 50 cm3 of acid the temperature rises by only about 2 °C (some heat is lost to the flask and the room), but with more magnesium or less acid the flask gets noticeably warm. The warming speeds the reaction up a little, which is one reason real results do not fit a simple rate law perfectly.
Key takeaway: exothermic: the acid warms by a few degrees.Magnesium slowly reacts with air to form a thin, dull layer of magnesium oxide. The acid has to remove this layer before it reaches the metal, so an uncleaned ribbon starts slowly and gives a time that is too long. Sanding every piece the same way makes the runs a fair test. Untick “Rubbed clean” in the sim to see the delay.
Key takeaway: remove the oxide layer so every run starts the same way.Measure the rate at several concentrations and plot log(rate) against log[HCl]; the gradient is the order. In this simulation the default model is first order, so doubling the concentration doubles the rate. Class experiments often give a value a little different from 1, because the solution warms up, hydrogen bubbles cover part of the surface and real ribbon is uneven, so it is best treated as an experimental result rather than something you can read from the equation.
Key takeaway: order = gradient of the log–log graph; it cannot be read from the balanced equation.Hydrogen is the lightest gas (2 g per mole). The 31 cm3 of hydrogen from 3 cm of ribbon has a mass of only about 0.003 g, smaller than the resolution of most school balances (0.01 g). The mass-loss method works much better for reactions that give off a heavier gas, such as marble chips (calcium carbonate) with acid, which releases carbon dioxide.
Key takeaway: hydrogen is too light to weigh well; use a gas syringe.Hardly at all. Almost all of a ribbon's surface area is its two flat faces, and cutting it adds only a few tiny edges, so the rate barely changes. Grinding the same mass into a fine powder, however, exposes many times more surface, so powder reacts much faster. Try both in the simulation.
Key takeaway: cutting ribbon adds very little surface; powder adds a lot.