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Nitrogen, fertiliser and freshwater quality

Covers AS91389

Why this is the New Zealand environmental issue

New Zealand's economy rests substantially on pastoral agriculture, and pastoral agriculture rests on nitrogen.
Nitrate in groundwater is the most contested environmental issue in the country, particularly in Canterbury, where intensive dairying sits above shallow, permeable gravel aquifers that supply drinking water.
This case study is unusual in that it can be argued from either of the standard's two routes: as an environmental issue (freshwater quality) or as a technology meeting a societal need (nitrogen fixation for food production). The strongest reports use both.

Step 1: nitrogen is abundant and useless

The atmosphere is 78% nitrogen, and almost none of it is available to plants.
The reason is the triple bond in N₂. Its bond enthalpy is about 945 kJ mol⁻¹ — one of the strongest bonds in chemistry.
Breaking it requires a large energy input, so N₂ is kinetically inert: it does not react under ordinary conditions despite the atmosphere being full of it.
This single fact — a very strong bond in a very abundant molecule — is why nitrogen fixation was one of the central chemical problems of the twentieth century.

Step 2: the Haber process

The Haber process converts atmospheric nitrogen into ammonia:

$\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g) \qquad \Delta_r H = -92 \text{ kJ mol}^{-1}$

The conditions are a compromise dictated entirely by the chemistry:

Condition	Value	Why
Temperature	~400–450 °C	the reaction is exothermic, so a low temperature favours a high yield — but the rate would be far too slow, so a compromise is used
Pressure	~200 atm	there are 4 mol of gas on the left and 2 on the right, so high pressure shifts the equilibrium right
Catalyst	iron	provides an alternative pathway of lower activation energy, increasing the rate without affecting the position of equilibrium
Recycling	unreacted gases returned	the single-pass yield is low, so recycling raises the overall conversion

New Zealand connection: the Ballance Agri-Nutrients plant at Kapuni, Taranaki, has made ammonia and urea from natural gas since 1982 — the country's only domestic nitrogen fertiliser production.

Step 3: ammonia becomes urea, and urea becomes nitrate

Most nitrogen fertiliser in New Zealand is applied as urea, made from ammonia and carbon dioxide:

$2\text{NH}_3(g) + \text{CO}_2(g) \rightarrow \text{CO(NH}_2)_2(s) + \text{H}_2\text{O}(l)$

In soil, urea is hydrolysed to ammonium:

$\text{CO(NH}_2)_2(s) + 2\text{H}_2\text{O}(l) \rightarrow 2\text{NH}_4^+(aq) + \text{CO}_3^{2-}(aq)$

Soil bacteria then oxidise ammonium to nitrate, a two-stage process called nitrification:

$2\text{NH}_4^+(aq) + 3\text{O}_2(g) \rightarrow 2\text{NO}_2^-(aq) + 4\text{H}^+(aq) + 2\text{H}_2\text{O}(l)$

$2\text{NO}_2^-(aq) + \text{O}_2(g) \rightarrow 2\text{NO}_3^-(aq)$

Nitrogen's oxidation number rises from −3 in ammonium, through +3 in nitrite, to +5 in nitrate. Nitrification is a redox process, and identifying it as one is the kind of integration this standard rewards.

Step 4: why nitrate leaches and ammonium does not

This is the pivotal chemistry, and it explains the entire problem.

Soil particles carry a net negative surface charge.
Ammonium, NH₄⁺, is a cation, so it is electrostatically attracted to soil particles and is held in place.
Nitrate, NO₃⁻, is an anion, so it is repelled by the negatively charged soil and stays dissolved in the soil water.
Nitrate therefore moves downwards with percolating rainwater into groundwater. Leaching is a direct consequence of the charge on the ion, and no amount of careful application changes that once the nitrogen has been nitrified.
The solubility rules make the same point: all nitrates are soluble, so nothing precipitates nitrate out on the way down.

Step 5: the consequences

In drinking water:

Nitrate itself has low toxicity, but it is reduced in the body to nitrite, which oxidises the iron in haemoglobin from Fe²⁺ to Fe³⁺, producing methaemoglobin, which cannot carry oxygen.
Infants are most at risk, and the New Zealand drinking-water standard sets a maximum acceptable value of 50 mg L⁻¹ of nitrate (11.3 mg L⁻¹ as nitrate-nitrogen).

In waterways — eutrophication:

Nitrate is a plant nutrient, so elevated nitrate stimulates growth of algae and aquatic plants.
When that biomass dies, decomposition consumes dissolved oxygen:

$\text{organic matter} + \text{O}_2(g) \rightarrow \text{CO}_2(g) + \text{H}_2\text{O}(l)$

The result is hypoxia — water with too little dissolved oxygen to support fish and invertebrates.

In the atmosphere:

Under low-oxygen conditions, bacteria denitrify nitrate, and one intermediate is nitrous oxide, N₂O:

$2\text{NO}_3^-(aq) + 2\text{H}^+(aq) \rightarrow \text{N}_2\text{O}(g) + \text{O}_2(g) + \text{H}_2\text{O}(l)$

N₂O is a greenhouse gas roughly 270 times more effective per molecule than CO₂, and it is a significant part of New Zealand's agricultural emissions profile.

Nitrogen, fertiliser and freshwater quality

Covers AS91389

Why this is the New Zealand environmental issue

New Zealand's economy rests substantially on pastoral agriculture, and pastoral agriculture rests on nitrogen.
Nitrate in groundwater is the most contested environmental issue in the country, particularly in Canterbury, where intensive dairying sits above shallow, permeable gravel aquifers that supply drinking water.
This case study is unusual in that it can be argued from either of the standard's two routes: as an environmental issue (freshwater quality) or as a technology meeting a societal need (nitrogen fixation for food production). The strongest reports use both.

Step 1: nitrogen is abundant and useless

The atmosphere is 78% nitrogen, and almost none of it is available to plants.
The reason is the triple bond in N₂. Its bond enthalpy is about 945 kJ mol⁻¹ — one of the strongest bonds in chemistry.
Breaking it requires a large energy input, so N₂ is kinetically inert: it does not react under ordinary conditions despite the atmosphere being full of it.
This single fact — a very strong bond in a very abundant molecule — is why nitrogen fixation was one of the central chemical problems of the twentieth century.

Step 2: the Haber process

The Haber process converts atmospheric nitrogen into ammonia:

$\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g) \qquad \Delta_r H = -92 \text{ kJ mol}^{-1}$

The conditions are a compromise dictated entirely by the chemistry:

Condition	Value	Why
Temperature	~400–450 °C	the reaction is exothermic, so a low temperature favours a high yield — but the rate would be far too slow, so a compromise is used
Pressure	~200 atm	there are 4 mol of gas on the left and 2 on the right, so high pressure shifts the equilibrium right
Catalyst	iron	provides an alternative pathway of lower activation energy, increasing the rate without affecting the position of equilibrium
Recycling	unreacted gases returned	the single-pass yield is low, so recycling raises the overall conversion

New Zealand connection: the Ballance Agri-Nutrients plant at Kapuni, Taranaki, has made ammonia and urea from natural gas since 1982 — the country's only domestic nitrogen fertiliser production.

Step 3: ammonia becomes urea, and urea becomes nitrate

Most nitrogen fertiliser in New Zealand is applied as urea, made from ammonia and carbon dioxide:

$2\text{NH}_3(g) + \text{CO}_2(g) \rightarrow \text{CO(NH}_2)_2(s) + \text{H}_2\text{O}(l)$

In soil, urea is hydrolysed to ammonium:

$\text{CO(NH}_2)_2(s) + 2\text{H}_2\text{O}(l) \rightarrow 2\text{NH}_4^+(aq) + \text{CO}_3^{2-}(aq)$

Soil bacteria then oxidise ammonium to nitrate, a two-stage process called nitrification:

$2\text{NH}_4^+(aq) + 3\text{O}_2(g) \rightarrow 2\text{NO}_2^-(aq) + 4\text{H}^+(aq) + 2\text{H}_2\text{O}(l)$

$2\text{NO}_2^-(aq) + \text{O}_2(g) \rightarrow 2\text{NO}_3^-(aq)$

Nitrogen's oxidation number rises from −3 in ammonium, through +3 in nitrite, to +5 in nitrate. Nitrification is a redox process, and identifying it as one is the kind of integration this standard rewards.

Step 4: why nitrate leaches and ammonium does not

This is the pivotal chemistry, and it explains the entire problem.

Soil particles carry a net negative surface charge.
Ammonium, NH₄⁺, is a cation, so it is electrostatically attracted to soil particles and is held in place.
Nitrate, NO₃⁻, is an anion, so it is repelled by the negatively charged soil and stays dissolved in the soil water.
Nitrate therefore moves downwards with percolating rainwater into groundwater. Leaching is a direct consequence of the charge on the ion, and no amount of careful application changes that once the nitrogen has been nitrified.
The solubility rules make the same point: all nitrates are soluble, so nothing precipitates nitrate out on the way down.

Step 5: the consequences

In drinking water:

Nitrate itself has low toxicity, but it is reduced in the body to nitrite, which oxidises the iron in haemoglobin from Fe²⁺ to Fe³⁺, producing methaemoglobin, which cannot carry oxygen.
Infants are most at risk, and the New Zealand drinking-water standard sets a maximum acceptable value of 50 mg L⁻¹ of nitrate (11.3 mg L⁻¹ as nitrate-nitrogen).

In waterways — eutrophication:

Nitrate is a plant nutrient, so elevated nitrate stimulates growth of algae and aquatic plants.
When that biomass dies, decomposition consumes dissolved oxygen:

$\text{organic matter} + \text{O}_2(g) \rightarrow \text{CO}_2(g) + \text{H}_2\text{O}(l)$

The result is hypoxia — water with too little dissolved oxygen to support fish and invertebrates.

In the atmosphere:

Under low-oxygen conditions, bacteria denitrify nitrate, and one intermediate is nitrous oxide, N₂O:

$2\text{NO}_3^-(aq) + 2\text{H}^+(aq) \rightarrow \text{N}_2\text{O}(g) + \text{O}_2(g) + \text{H}_2\text{O}(l)$

N₂O is a greenhouse gas roughly 270 times more effective per molecule than CO₂, and it is a significant part of New Zealand's agricultural emissions profile.