Hydrogen

Typical wastewaters

• hydrofluoric acid, ammonia fertiliser, primary beryllium and semiconductor plant effluent free mineral acidity as H+ (or alkalinity as OH-) in process wastewater, neutralised to the 6.0 to 9.0 pH band that footnote 1 of each US effluent guideline table sets
• acid mine drainage and metal pickling H+ as free acidity with sulfate or chloride, neutralised with lime, caustic soda or magnesium hydroxide to the permit band; metals precipitate as hydroxides in the same step and gypsum sludge forms with lime and sulfuric acid no acid effluent concentrations were read
• alkaline discharges (lime, caustic soda, concrete washwater) OH- alkalinity at high pH, neutralised with sulfuric acid or carbon dioxide (CO₂ cannot overshoot below about pH 6 in buffered water)
• municipal sewage (secondary treatment effluent) pH held in the 6.0 to 9.0 band by the biological process itself; the band US secondary treatment effluent must stay within
• textile dyeing and finishing and leather tanning pH measured on site, 6 to 9 at every ZDHC level; the same band at Foundational, Progressive and Aspirational

1 · Identity

Symbol, number

H, 1

Oxidation states in water

+1 in water itself, in the hydrated proton (written H+ here, H₃O⁺ in reality) whose activity pH measures, in every acid, and bonded to carbon, nitrogen and oxygen in dissolved organic matter and ammonia; 0 as dissolved hydrogen gas H₂, released by electrolysis, by the cathodic reaction of corrosion in acid or anoxic water and by fermentation, and consumed by hydrogenotrophic bacteria; -1 (hydride) does not exist in water, hydrides react with it to give H₂ (book entry).

Note

The element entry already carries combustion, the H₂ and Cl₂ chain reaction and the hydride chemistry. This chapter is about pH and about H₂ as a reagent and by-product. Hydrogen sulfide belongs to the sulfur chapter and hydrogen peroxide to the oxygen chapter; the carbonate buffer that fixes the pH of most natural water is written out in the carbon chapter and only referenced here.

2 · Occurrence in water

Natural sources

The pH of most natural and drinking water is controlled by the carbon dioxide, bicarbonate and carbonate equilibrium: more dissolved CO₂ lowers pH, less raises it, and temperature shifts the equilibria (pure water falls about 0.45 pH unit for a 25 C rise). Acid rain gives lower pH, limestone catchments higher (WHO background document). Dissolved H₂ is produced by fermentation in anoxic sediments and anaerobic digesters, where it is an intermediate consumed by hydrogenotrophic methanogens (Metcalf and Eddy chapter 10).

Anthropogenic sources

Acid discharges (mine drainage, metal pickling, acid plants) and alkaline discharges (lime, caustic soda, concrete washwater, insufficiently cured cement mortar pipe linings, which WHO names as a cause of extreme pH); treatment chemicals that consume or release protons, the acid coagulants and chlorine gas among them; H₂ from water electrolysers and from on-site electrolytic hypochlorite generation; H₂ evolved on steel in anoxic water and on cathodically protected structures.

3 · Speciation

pH is the negative common logarithm of the hydrogen ion activity, which equals the concentration in dilute solution (WHO background document). Water dissociates to H+ and OH- with an ion product of 10⁻14 at 25 C, so pH 7 is neutral only at that temperature; the ion product rises with temperature and neutral pH falls (Stumm and Morgan chapter 3). In almost every natural water the free proton is a minor species and pH is set by the weak acid systems that buffer it: carbonic acid and bicarbonate between pH 6 and 9, ammonia above 9, hydroxide above 10 (carbon and nitrogen chapters). Dissolved H₂ is a sparingly soluble, unreactive gas at room temperature that only bacteria and electrodes use.

Solubility

Not applicable to H+; the solubility of H₂ in water is not quoted because no source read prints it (it is a sparingly soluble gas).

Hydrolysis

Water is the hydrolysis: H₂O splits to H+ and OH- and every other hydrolysis in the book (Fe³⁺, Al³⁺, chlorine, coagulants) is the transfer of one of those ions.

Complexation

H+ competes with every metal for ligands, so lowering pH frees metals from hydroxide, carbonate and organic complexes; that is why acidification preserves metal samples and why acid water carries metals (WHO pH document: the lower the pH, the higher the potential corrosion).

Precipitates

None of hydrogen itself. pH decides which hydroxides and carbonates precipitate and at what pH, treated under each metal.

4 · Role in treatment

5 · Removal and control

6 · Analytics

Sampling pitfalls

Measure pH on site: loss or gain of carbon dioxide in the bottle shifts it, and temperature changes both the equilibrium and the electrode response (WHO background document, APHA). Calibrate with two buffers bracketing the sample. In low ionic strength water the glass electrode drifts; use a low ionic strength electrode or add a neutral salt. Report the temperature with every pH.

7 · Regulatory limits

Limits change, and many are set locally. Treat these as the published values to start from, not as your compliance target: check the standard in force at your site and the numbers written into your own permit.

8 · Health and environmental effects

Toxicity

pH alone is not the primary determinant of adverse effects; acids and alkalis in water are extremely dilute and stomach fluid is pH 1 to 3.5; any health effect is indirect, through metals leached from plumbing and inadequate disinfection (WHO background document). Hydrogen gas is not toxic and acts only as an asphyxiant and an explosion hazard (element entry).

Bioaccumulation

Not applicable.

Ecotoxicity

US EPA national recommended criteria: freshwater pH 6.5 to 9, saltwater 6.5 to 8.5, and open ocean pH not changed more than 0.2 unit from natural variation (1986 Gold Book, cited in the criteria table). pH also governs the toxicity of ammonia, aluminium and metals, which is why their criteria are pH dependent.

Flags

• The pH ranges given as concentrations are regulatory bands and a WHO generalisation, not surveys.
• The on-site hypochlorite and hydrogenotrophic denitrification equations are written from the electrochemistry and electron balance; the EPA manual and Metcalf and Eddy name the processes without printing them.
• Kw and the temperature dependence are cited to Stumm and Morgan chapter 3 from memory of the text, not re-read this session.
• The 40 CFR industrial pH band is cited to the four sections read; other categories were not checked.
• Abu Dhabi values cover two media (marine outfall 6.0 to 9.0, sewer above 6 and below 9); other GCC states were not read.

Gaps

• No source read gives dissolved hydrogen concentrations in any water matrix.
• The solubility of H₂ and the temperature dependence of Kw are not quoted as numbers; Stumm and Morgan has them but was not re-read.
• No survey of pH in municipal or industrial wastewater was read; permit bands stand in for it.
• The 2016 CWW BREF text on pH neutralisation techniques was not re-read.
• Other GCC discharge standards (Saudi, Oman, Qatar) were not read.

Sources

Identity

Name and symbol

Hydrogen, H

Atomic number

1 protons

Position

group 1 · period 1 · s-block · diatomic nonmetal

CAS number

12385-13-6

Atomic structure

Atomic mass

1.0079 u

Electron configuration

1s¹
1s¹

Electrons per shell

1

Valence electrons

1 outer shell

Physical properties

State at room temperature

Gas

Melting point

13.81 K (-259.34 °C)

Boiling point

20.28 K (-252.87 °C)

Density

0.0001 g/cm3 (gas at STP, so 0.0899 g/L)

Appearance

Colorless gas

Thermal conductivity

0.1805 W/(m·K)

Electrical resistivity

not in sources

Electrical conductivity

not in sources

Crystal structure

hexagonal

Molar heat capacity

28.836 J/(mol·K)

Chemical properties

Oxidation states

+1, -1

Electronegativity

2.2 (Pauling Scale)

Ionisation energy

13.598 eV
1st 1,312 kJ/mol

Electron affinity

0.754 eV

Atomic radius

empirical 31, covalent 31, van der Waals 120 pm

Ionic radius

H⁻ 140; H⁺ -18 (2-coordinate) pm

Reactivity

Molecular hydrogen is only moderately reactive at room temperature because the H-H bond is strong (about 436 kJ/mol), but once started it combines with most nonmetals and with the reactive metals, sometimes explosively.

with water

Does not react; water is instead a source of hydrogen, by electrolysis or by passing steam over hot carbon.

with oxygen, air

Burns in air or oxygen to water, and mixtures of 4 to 74 percent hydrogen in air explode: $\ce{2 H2 + O2 -> 2 H2O}$

with acids

Does not react; hydrogen is the gas released when reactive metals displace it from dilute acids.

with halogens

Explodes with fluorine even in the cold; with chlorine it reacts by a light-initiated chain reaction and with bromine on heating: $\ce{H2 + Cl2 -> 2 HCl}$

Typical compounds

• H₂O water the oxide of hydrogen and its main home on Earth
• NH₃ ammonia from nitrogen and hydrogen by the Haber process
• HCl hydrogen chloride from hydrogen and chlorine; hydrochloric acid in water
• CH₄ methane simplest hydrocarbon; steam reforming of natural gas gives hydrogen
• H₂O₂ hydrogen peroxide peroxide with oxygen in the minus one state
• LiH lithium hydride ionic hydride used to store hydrogen; gives H2 with water

Occurrence, production and use

Crustal abundance

1.40×103 milligrams per kilogram

Oceanic abundance

1.08×105 milligrams per liter

Occurrence and sources

Hydrogen is estimated to make up more than 90% of all the atoms three quarters of the mass of the universe! This element is found in the stars, and plays an important part in powering the universe through both the proton-proton reaction and carbon-nitrogen cycle. Stellar hydrogen fusion processes release massive amounts of energy by combining hydrogens to form helium.

Production of hydrogen in the U.S. alone amounts to about 3 billion cubic feet per year. Hydrogen is prepared by

▸ steam on heated carbon,

▸ decomposition of certain hydrocarbons with heat,

▸ reaction of sodium or potassium hydroxide on aluminum

▸ electrolysis of water, or

▸ displacement from acids by certain metals.

Liquid hydrogen is important in cryogenics and in the study of superconductivity, as its melting point is only 20 degrees above absolute zero.

Tritium is readily produced in nuclear reactors and is used in the production of the hydrogen bomb.

Hydrogen is the primary component of Jupiter and the other gas giant planets. At some depth in the planet's interior the pressure is so great that solid molecular hydrogen is converted to solid metallic hydrogen.

In 1973, a group of Russian experimenters may have produced metallic hydrogen at a pressure of 2.8 Mbar. At the transition the density changed from 1.08 to 1.3 g/cm3. Earlier, in 1972, at Livermore, California, a group also reported on a similar experiment in which they observed a pressure-volume point centered at 2 Mbar. Predictions say that metallic hydrogen may be metastable; others have predicted it would be a superconductor at room temperature.

• water (H2O) oceans, ice and groundwater; the greatest quantity of terrestrial hydrogen
• organic matter and fossil hydrocarbons living plants, petroleum, coal and natural gas; natural gas is the main industrial feedstock
• free H2 atmosphere, below 1 ppm by volume; escapes Earth's gravity

Extraction, production

Steam reforming of natural gas (primary reformer of an ammonia plant)

$\ce{CH4 + H2O <=> CO + 3 H2}$

Highly endothermic, delta H0 = 206 kJ/mol; the desulphurised gas is mixed with steam and enters the primary reformer at 400 to 600 degrees C; hydrocarbon conversion in a conventional plant is about 60 percent per pass (AAF BREF, PDF p68). Most hydrogen for ammonia is made this way, which is why ammonia plants cluster where gas is cheap.

Water-gas shift of the reformer gas

$\ce{CO + H2O <=> CO2 + H2}$

Delta H0 = -41 kJ/mol; the 12 to 15 percent CO in reformer gas is shifted in two steps with intermediate heat removal, first over an iron oxide/chromium oxide catalyst at about 350 degrees C (AAF BREF, PDF p69). Residual carbon oxides are then removed by methanation, $\ce{CO + 3 H2 -> CH4 + H2O}$, because they poison the ammonia catalyst (PDF p70).

Electrolysis of water

$\ce{2 H2O (l) -> 2 H2 (g) + O2 (g)}$

Stated as the second commercial route by the RSC; the sources describe it in words and print no equation.

Coproduct of chlor-alkali brine electrolysis

$\ce{2 NaCl + 2 H2O -> 2 NaOH + H2 + Cl2}$

Cathode reaction in membrane and diaphragm cells: $\ce{2 Na+ + 2 e- + 2 H2O -> 2 NaOH + H2}$. The products come in a fixed ratio, 1,070 to 1,128 kg NaOH and about 28 kg H2 per tonne of Cl2 (CAK BREF, PDF p40).

Uses

Hydrogen is a commercially important element. Large amounts of hydrogen are combined with nitrogen from the air to produce ammonia (NH3) through a process called the Haber process. Hydrogen is also added to fats and oils, such as peanut oil, through a process called hydrogenation. Liquid hydrogen is used in the study of superconductors and, when combined with liquid oxygen, makes an excellent rocket fuel.

Hydrogen combines with other elements to form numerous compounds. Some of the common ones are: water (H2O), ammonia (NH3), methane (CH4), table sugar (C12H22O11), hydrogen peroxide (H2O2) and hydrochloric acid (HCl).

Hydrogen has three common isotopes. The simplest isotope, called protium, is just ordinary hydrogen. The second, a stable isotope called deuterium, was discovered in 1932. The third isotope, tritium, was discovered in 1934.

Great quantities of hydrogen are required commercially for nitrogen fixation using the Haber ammonia process, and for the hydrogenation of fats and oils. It is also used in large quantities in methanol production, in hydrodealkylation, hydrocracking, and hydrodesulfurization. Other uses include rocket fuel, welding, producing hydrochloric acid, reducing metallic ores, and filling balloons.

The lifting power of 1 cubic foot of hydrogen gas is about 0.07 lb at °C, 760 mm pressure.

The hydrogen fuel cell is a developing technology that will allow great amounts of electrical power to be obtained using a source of hydrogen gas.

Consideration is being given to an entire economy based on solar- and nuclear-generated hydrogen. Public acceptance, high capital investment, and the high cost of hydrogen with respect to today's fuels are but a few of the problems facing such an economy. Located in remote regions, power plants would electrolyze seawater; the hydrogen produced would travel to distant cities by pipelines. Pollution-free hydrogen could replace natural gas, gasoline, etc., and could serve as a reducing agent in metallurgy, chemical processing, refining, etc. It could also be used to convert trash into methane and ethylene.

• Chemicals: ammonia synthesis, $\ce{N2 + 3 H2 <=> 2 NH3}$, on an iron catalyst at 100 to 250 bar and 350 to 550 degrees C, delta H0 = -46 kJ/mol, with only 20 to 30 percent converted per pass so the unreacted gas is recycled (AAF BREF, PDF p71); methanol synthesis and cyclohexane, intermediates for plastics and pharmaceuticals about 80 percent of ammonia is used as the nitrogen source in fertilisers, the rest in plastics, fibres, explosives and other nitrogen chemicals (AAF BREF, PDF p63)
• Oil refining: hydrodesulfurisation, removing sulfur from fuels; hydrocracking and hydrodealkylation
• Food and beverage: hydrogenation of oils to solid fats, for example margarine; large quantities of hydrogen are used this way
• Electronics and glass: flushing gas in silicon chip manufacture; protective atmosphere for making flat glass sheets
• Energy and transport: fuel cells in buses and cars; hydrogen burns back to water; rocket fuel; reduction of metal ores and welding

Safety, toxicity

GHS classification, signal word Danger

• H226 Flammable liquid and vapor Flammable liquids
• H302 Harmful if swallowed Acute toxicity, oral
• H315 Causes skin irritation Skin corrosion/irritation
• H318 Causes serious eye damage Serious eye damage/eye irritation
• H335 May cause respiratory irritation Specific target organ toxicity, single exposure; Respiratory tract irritation
• H336 May cause drowsiness or dizziness Specific target organ toxicity, single exposure; Narcotic effects

Discovery and name

Discovered by

Henry Cavendish

Discovered

1766

First isolated

not in sources

Named by

Antoine Lavoisier

Origin of the name

name means 'water-former' in Greek