Hydroponic Equipment . com - Introduction to pH

Introduction to pH

Definitions
pH is the scale of measurement of acidity or alkalinity in aqueous (water based) solutions. A neutral solution such as pure water has a pH of 7. Solutions with a lower pH are termed acidic and solutions with a higher pH are termed alkaline. pH ranges from highly acidic pH 0 to highly alkaline pH 14. The pH of the human eye is about 7.5.

pH was introduced by a Danish biochemist Soren Peter Lauritz Sorensen in 1909 to measure the acidity of water in the brewing of beer. The letters pH are an abbreviate for pondus hydrogenii (translated as potential hydrogen) meaning hydrogen power as acidity is caused by a predominance of hydrogen ions (H+).

Water is H2O, this is two hydrogen atoms and one oxygen atom. Water is a polar molecular shaped like a boomerang with the oxygen atom in the middle and the hydrogen atoms at the tips. This shape comes about because of the way the 8 negatively charged electrons in oxygen pack themselves in close to the 8 positively charged nucleus of the oxygen atom. The nucleus of an atom is very small like a sparrow in a cloud (of electrons). The Hydrogen atoms at the tips have more positive charge than negative charge and the oxygen atom has more negative charge than positive charge. This causes the molecule to be polar, that is to have an electrical field difference along it. The tips of the boomerang are positively charged and the middle bend is negatively charged.

Water dissociates (breaks up) into H+ and OH- ( H2O < == > OH- + H+). The energy of the atoms bumping into each other cause the water molecules to split up. Because water is a polar molecule the H+ will attract the negatively charged Oxygen atoms in other water molecules and will “float” away. The OH- will attract the Hydrogen in other water molecules and float away on its own. Eventually a H+ and a OH- will bump into each other and join up to become water again. This process of dissociation and association is dynamic.

In pure water, for every one H+ There are 10,000,000 normal water molecules. This is 1 in107 and the pH of pure water is 7 meaning for every 107 molecules of water there is a H+ ion. Said in another way the H+ concentration is 10-7 parts (pH of 7), that is one part in 10Million parts. Likewise with the OH-. If the amount of H+ increases by adding acidic compounds (hydrochloric acid HCl < == > H+ and Cl-) then there are more H+'s to combine with OH-'s so the concentration of OH- is reduced. Likewise adding OH- by adding an alkaline compound (Caustic Soda NaOH < == > Na+ and OH-) reduces the concentration of H+. A Chemical rule is that the concentration of H+ times the concentration of OH- always equals 10-14.

Low pH (acidic) conditions cause water to be corrosive. Acids will cause pitting of concrete, dissolve metals, wrinkle vinyl, irritate skin and eye and stains pool and spa walls and pipes.

High pH (caustic) conditions cause scaling - minerals precipitate out of the water (calcium, copper, iron etc. and those minerals will block filters and pipes.

The numerical value of the pH is the negative of the exponent of the molar concentration. A mole is defined as the weight in grams that corresponds to the summed atomic weights of all the atoms of a molecule - its molecular mass. In the case of H+ the molecular mass equals 1, and a mole of H+ ions would weigh 1 gram. In the case of water (H2O) the molecular mass equals 34, and a mole of water ions would weigh 18 grams (The sum of the atomic weights).

Thus low pH values indicate high concentrations of H+ ions (acid), and high pH values indicate low concentrations. Each pH unit downward represents, therefore, a tenfold increase in the H+ concentration. A pH of 3, for example, indicates a 10-3 molar concentration of hydrogen ions. To put it simply the pH scale is logarithmic, meaning that every pH unit means "10 times". Therefore a pH of 6 is 10 times more acidic than a pH of 7, and a pH of 3 is 1000 times more acidic than 6 and 10,000 times more acidic than 7.

To modify pH use one of the specially designed pH adjusters on our nutrients page.

Combining an acidic compound with an alkaline compound produces a salt. Sodium Hydroxide plus Hydrochloric acid when mixed will become very hot and result in a salt - Sodium Chloride. Salts by themselves can increase the rate of corrosion and mineral precipitates.

Some salts give water the ability to resist changes in pH or buffer the water from wild pH swings. In water that contains no buffering ability, the pH can wander dramatically with the addition of small amounts of acids or bases (alkali), or other pH altering agents like chlorine or bromine. With a buffer the pH will hold steady and pH bounce will be eliminated for a while until so much acid or alkaline has been added that the buffering effect is overwhelmed.

Why pH Varies
The ratio in uptake of anions (negatively charged nutrients) and cations (positively charged nutrients) by plants may cause substantial shifts in pH. In general, an excess of cation over anion leads to a decrease in pH, whereas an excess of anion over cation uptake leads to an increase in pH. As nitrogen (an element required in large quantities for healthy plant growth) may be supplied either as a cation (ammonium - NH4+) or an anion (nitrate - NO3), the ratio of these two forms of nitrogen in the nutrient solution can have large effects on both the rate and direction of pH changes with time. This shift in pH can be surprisingly fast.

Daylight photosynthesis produces hydrogen ions which can cause the nutrient acidity to increase (lowering the pH). At dusk photosynthesis stops and the plants increase their rate of respiration and this coupled with the respiration of micro organisms and the decomposition of organic matter uses up the hydrogen ions so the acidity of the solution tends to decrease ( pH rises )

Most varieties of vegetables grow at their best in a nutrient solution having a pH between 6.0 and 7.5 and a nutrient temperature between 20 and 22 degrees celcius

In low light ( overcast days or indoor growing environments) plants take up more potassium and phosphorous from the nutrient solution so the acidity increases (pH drops). In strong intense light (clear sunny days) plants take up more nitrogen from the nutrient solution so the acidity decreases (pH rises). pH can be controlled in two ways.

How pH Affects Plant Growth

pH can affect the availability of nutrients.
pH can affect the absorption of nutrients by plant roots
pH values above 7.5 cause iron, manganese, copper, zinc and boron ions to be less available to plants.
pH values below 6 cause the solubility of phosphoric acid, calcium and magnesium to drop.
pH values between 3 and 5 and temperatures above 26 degrees Celcius encourage the development of fungal diseases.
For plant roots to be able to absorb nutrients, the nutrients must be dissolved in solution. The process of precipitation (the reverse of dissolving) results in the formation of solids in the nutrient solution, making nutrients unavailable to plants. Not all precipitation settles to the bottom of the tanks, some precipitates occur as very fine suspension invisible to the naked eye.

Plants can tell us their problems through leaf symptoms (e.g. iron [Fe] deficiency) when it's too late. Iron (Fe) is one essential plant nutrient whose solubility is affected by pH which is why it is added in a chelated form (or daily), Fe deficiency symptoms occur readily. At pH values over 7, less than 50% of the Fe is available to plants. At pH 8.0, no Fe is left in solution due to iron hydroxide precipitation (Fe(OH)3 - which eventually converts to rust). As long as the pH is kept below 6.5, over 90% of the Fe is available to plants.

Varying pH of summer lettuce nutrient solutions also affects the solubility of calcium (Ca) and phosphorus (P). Due to calcium phosphate precipitation (Ca3(PO4)2) the availability of Ca and P decreases at pH values above 6.0. All other nutrients stay in solution and do not precipitate over a wide pH range. Poor water quality could exacerbate any precipitation reactions that may occur.

Generally in the pH range 4.0 to 6.0, all nutrients are available to plants. Precipitation reduces Fe, Ca and P availability at pH 6.0 and over.

Measuring pH
Two parts are involved in a pH electrode system. The first is a pH sensor, which consists of a special pH sensitive glass and whose voltage output is proportional to the pH. The second is a reference sensor which provides a stable and constant reference point. Electrical contact is made with the solution using a saturated salt solution (usually KCl) which leaks slowly out of a porous junction. These two sensors are usually built into the same housing, hence the term "combination" pH electrode.

A pH meter is simply a device which measures the voltage from the electrodes and converts it to a pH reading on a display.

In theory, the output of a pH electrode is zero millivolts (mV) at pH7. At 25 degrees Celsius the output of the electrode changes by 59 mV per pH. At pH4 the output is +177mV, and at pH 10 the output is -177mV.

In practice, the output varies from electrode to electrode. The "zero" output (called the "asymmetry") may be slightly away from pH7, and the mV change per pH (called the "slope" or "span") may be slightly lower than 59 (especially as the probe ages). Therefore, pH meters must have two calibration controls, which are adjusted when the pH electrode is immersed in pH buffers (solutions of a known pH which are resistant to pH change).

The first control, usually labelled CALIBRATE (also ISO, ZERO, ASYMMETRY), is used to calibrate the zero point of the electrode. This is done in a pH buffer at or near pH7. Instruments Australia supplies pH6.88 buffer. This calibration MUST be done FIRST.

The second calibration point calibrates the slope of the electrode. This is usually labelled SLOPE (also SPAN). In order for this calibration step to be effective, it must be done at least 2 to 3 pH away from 7. Instruments Australia supplies pH4.00 buffer. This calibration MUST be done SECOND.

Micro-processor based instruments calculate the asymmetry and slope regardless of which buffer they are calibrated in, so it is not necessary to calibrate at 7 and then at 4.

The frequency of calibration varies from application to application. As a guideline, calibration should be checked at least weekly.

Factors that affect the output of pH electrodes include temperature changes, a blocked reference junction, and coatings on the pH glass.

Temperature effects can be compensated for in two ways. Firstly, the pH meter can have a manual control which can be set to the temperature of the solution. Secondly, the pH meter can be fitted with an Automatic Temperature Compensation ("ATC") probe, which, as the name suggests, measures the temperature and automatically compensates for the temperature effect on pH electrode performance.