Iron Removal From Ground Water Biology Essay

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Iron is a troublesome element in water supplies. Making up at least 5% of the earths crust, iron is one of the earth's resources. Rainwater, as it infiltrates soil, the underlying geologic formations, dissolves iron, causing it to seep into the aquifers that serve as source of groundwater for wells. However, a little amount i.e. 0.3 ppm can cause water to turn a reddish brown color. Thus, the removal of iron is necessary to avoid health risk.

In the present dissertation work, adsorption of Iron has been studied for two types of grain sizes of sand in waste water. The effect of contact time, dose, and pH is studied in column experiment.

It has been found that when initial iron concentration is 5mg/L, 20gm/L dose of 0.5 mm sand can remove 97.6% iron at contact time of 2 hrs and 20gm/l dose of 1.0 mm sand can remove 95% iron contact time of 2 hrs at initial pH is 7.5. On comparing two grain size (i.e. GMS (geometric mean size) equal to 0.5 mm and 1 mm) of sand it is observed that 0.5 mm sand is much better adsorbent than other grain size of sand for iron removal. Finally, it has been concluded that low cost adsorbents has been found successfully removal of iron from waste water.




Iron (Fe) is metal that occur naturally in soils, rocks and minerals. In the aquifer, groundwater comes in contact with solid material dissolving them, releasing their constituents (Fe) to the water. At concentrations approaching 0.3 mg/L of Fe, the water's usefulness may become seriously impacted, e.g., there may be a metallic taste to the water and staining of plumbing fixtures may become common. At these concentrations, however, the health risk of dissolved Fe in drinking water is insignificant. The extent to which Fe dissolved in groundwater depends on the amount of oxygen in the water and, to a lesser extent, upon its degree of acidity, i.e., its pH. Iron can occur in two forms: as Fe2+ and as Fe3+. When levels of dissolved oxygen in groundwater are greater than 1-2 mg/L, iron occurs as Fe3+, while at lower dissolved oxygen levels, the iron occurs as Fe2+. Although Fe2+ is very soluble, Fe3+ will not dissolve appreciably. If the groundwater contains less oxygen then, iron will dissolve more readily, particularly if the pH of the water is on the low side (slightly more acidic). Dissolved oxygen content is typically low in deep aquifers, particularly if the aquifer contains organic matter. Decomposition of the organic matter depletes the oxygen in the water and the iron dissolves as Fe2+. When this water is pumped to the surface, the dissolved iron reacts with the oxygen in the atmosphere, changes to Fe3+ and forms rust-colored iron minerals. Treatment for dissolved iron takes advantage of the natural process of oxidation, through the use of aeration, i.e., injecting air into the water prior to the tap to precipitate iron from the water. Chlorine is also an effective oxidizer and will cause iron to precipitate, plus it provides protection from microbial contaminants. Usually a physical filter follows the treatment so that the particles will not exit through the tap. Additional treatment methods include greensand filters and water softeners. Local suppliers of water treatment devices should be consulted in order to select the best system for a given water supply. The amount of dissolved iron in groundwater may vary seasonally for a given well. Usually this is associated with an influx of oxygenated water from the surface during periods of high recharge. This oxygenated water will prevent the iron from dissolving and the water pumped from the well will have low concentrations of these metals. After the oxygen in the recharge water has been consumed, iron will again be dissolved and the water will have dissolved iron characteristics.

A final note is that even though treating the water for dissolved iron after it leaves the well will make the water more palatable, high concentrations of dissolved iron within the well bore may lead to growth of iron bacteria. These bacteria may coat the inside of the casing or any other submerged part of the plumbing in the well and may cause problems. In areas where elevated iron is common, it may be worth while to periodically disinfect the well to keep iron bacterial growth in check.

According to the W.H.O. drinking water standard recommended limit for iron in public water supplies is 0.3 mg/L. Even though this limit is not based on physiological consideration, the presence of iron in domestic and industrial water supplies has long plagued the householder and the manufacturer. When iron is present in a water supply at concentration exceeding 0.3 mg/l, it is undesirable for the following reason:

Iron precipitates give water a reddish colour when exposed to air.

Iron gives water a metallic taste.

Home softness becomes clogged by iron precipitates.

Deposition of iron precipitates in the distribution system can reduce the effective pipe diameter and eventually clog the pipe.

Iron is a substrate for the growth of bacteria in the water mains, when iron bacteria die and slough off, bad odors and unpleasant tastes may be caused.

In paper industries, iron causes discoloration of pulp and paper. Iron is responsible for spot and discoloration of leather goods. Distilleries require clear, potable water free from iron. Deposits are formed in boiler operating at high pressure due to presence of iron in the feed water. Since these deposits can cause tube failures, the feed water in high - pressure boiler system should be free from iron.

In the present dissertation work, adsorption of Iron has been studied for two types of grain sizes of sand in waste water. The effect of contact time, dose, and pH is studied in column experiment.



2.1 Introduction

Iron is a lustrous, ductile, malleable, silver-gray metal (group VIII of the periodic table). It is known to exist in four distinct crystalline forms. Iron rusts in dump air, but not in dry air. It dissolves readily in dilute acids. Iron is chemically active and forms two major series of chemical compounds, the bivalent iron (II), or ferrous, compounds and the trivalent iron (III), or ferric, compounds. The chemical properties of iron are given below:

Atomic number


Atomic mass

55.85 g.mol-1

Electronegativity according to Pauling



7.8 20°C

Melting point

1536 °C

Boiling point

2861 °C

Vanderwaals radius

0.126 nm

Ionic radius

0.076 nm (+2) ; 0.064 nm (+3)



Electronic shell

[Ar ] 3d64s2

Energy of first ionisation

761 kJ.mol-1

Energy of second ionisation

1556.5 kJ.mol-1

Energy of third ionisation

2951 kJ.mol-1

Standard potential

- o.44 V (Fe2+/ Fe ) ; 0.77 V ( Fe3+/ Fe2+)


Iron is the most used of all the metals, including 95 % of all the metal tonnage produced worldwide. Thanks to the combination of low cost and high strength it is indispensable. Its applications go from food containers to family cars, from screwdrivers to washing machines, from cargo ships to paper staples. Steel is the best known alloy of iron, and some of the forms that iron takes include: pig iron, cast iron, carbon steel, wrought iron, alloy steels, iron oxides.

2.1.1 Health effects of iron

Iron can be found in meat, whole meal products, potatoes and vegetables. The human body absorbs iron in animal products faster than iron in plant products. Iron is an essential part of hemoglobin; the red colouring agent of the blood that transports oxygen through our bodies.

Iron may cause conjunctivitis, choroiditis, and retinitis if it contacts and remains in the tissues. Chronic inhalation of excessive concentrations of iron oxide fumes or dusts may result in development of a benign pneumoconiosis, called siderosis, which is observable as an x-ray change. No physical impairment of lung function has been associated with siderosis. Inhalation of excessive concentrations of iron oxide may enhance the risk of lung cancer development in workers exposed to pulmonary carcinogens. LD50 (oral, rat) =30 gm/kg. (LD50: Lethal dose 50. Single dose of a substance that causes the death of 50% of an animal population from exposure to the substance by any route other than inhalation. A more common problem for humans is iron deficiency, which leads to anaemia. A man needs an average daily intake 7 mg of iron and a woman 11 mg; a normal diet will generally provided all that is needed.

2.1.2 Iron and water

Seawater contains approximately 1-3 ppb of iron. The amount varies strongly, and is different in the Atlantic and the Pacific Ocean. Rivers contain approximately 0.5-1 ppm of iron, and groundwater contains 100 ppm. Drinking water may not contain more than 200 ppb of iron. Most algae contain between 20 and 200 ppm of iron, and some brown algae may accumulate up to 4000 ppm. The bio concentration factor of algae in seawater is approximately 104 - 105. Sea fish contain approximately 10-90 ppm and oyster tissue contains approximately 195 ppm of iron (all are dry mass). Dissolved iron is mainly present as Fe(OH)2+ under acidic and neutral, oxygen rich conditions. Under oxygen-poor conditions it mainly occurs as binary iron. Iron is part of many organic and inorganic chelation complexes that are generally water soluble.

2.1.3 Environmental effects of iron in water

Iron is a dietary requirement for most organisms, and plays an important role in natural processes in binary and tertiary form. Oxidized tertiary iron cannot be applied by organisms freely, except at very low pH values. Still, iron usually occurs in this generally water insoluble form. Adding soluble iron may rapidly increase productivity in oceanic surface layers. It might than play an important role in the carbon cycle. Iron is essential for nitrogen binding and nitrate reduction, and it may be a limiting factor for phytoplankton growth. Solubility in salt water is extremely low.The iron cycle means reduction of tertiary iron by organic ligands (a process that is photo catalysed in surface waters), and oxidation of binary iron. Iron forms chelation complexes that often play an important role in nature, such as haemoglobin, a red colouring agent in blood that binds and releases oxygen in breathing processes. Organisms take up higher amounts of binary iron than of tertiary iron, and uptake mainly depends on the degree of saturation of physical iron reserves.

Iron is often a limiting factor for water organisms in surface layers. When chelation ligands are absent, water insoluble tertiary iron hydroxides precipitate. This is not thought to be hazardous for aquatic life, because not much is known about hazards of water borne iron.

Mollusks have teeth of magnetite of goethite. Green plants apply iron for energy transformation processes. Plants that are applied as animal feed may contain up to 1000 ppm of iron, but this amount is much lower in plants applied for human consumption. Generally plants contain between 20 and 300 ppm iron (dry mass), but lichens may consist up to 5.5% of iron. When soils contain little iron, or little water soluble iron, plants may experience growth problems. Plant uptake capacity strongly varies, and it does not only depend on soil iron concentrations, but also upon pH values, phosphate concentrations and competition between iron and other heavy metals. Limes soils are often iron deficit, even when sufficient amounts of iron are present. This is because of the generally high pH value, which leads to iron precipitation. Iron usually occurs in soils in tertiary form, but in water saturated soils it is converted to binary iron, thereby enabling plant iron uptake. Plants may take up water insoluble iron compounds by releasing H+ ions, causing it to dissolve. Micro organisms release iron siderochrome, which can be directly taken up by plants.

Iron may be harmful to plants at feed concentrations of between 5 and 200 ppm. These cannot be found in nature under normal conditions, when low amounts of soil water are present. A number of bacteria take up iron particles and convert them to magnetite, to apply this as a magnetic compass for orientation. Iron compounds may cause a much more serious environmental impact than the element itself. A number of values are known for rats (oral intake): iron (III) acetyl acetone 1872 mg/kg, iron (II) chloride 984 mg/kg, and iron pent carbonyl 25 mg/kg. There are four naturally occurring non-radioactive iron isotopes.

2.1.4 The health effects of iron in water

The total amount of iron in the human body is approximately 4 g, of which 70% is present in red blood colouring agents. Iron is a dietary requirement for humans, just as it is for many other organisms. Men require approximately 7 mg iron on a daily basis, whereas women require 11 mg. The difference is determined by menstrual cycles. When people feed normally these amounts can be obtained rapidly. The body absorbs approximately 25% of all iron present in food. When someone is iron deficit feed iron intake may be increased by means of vitamin C tablets, because this vitamin reduces tertiary iron to binary iron. Phosphates and phytates decrease the amount of binary iron. In food iron is present as binary iron bound to haemoglobin and myoglobin, or as tertiary iron. The body may particularly absorb the binary form of iron. Iron is a central component of haemoglobin. It binds oxygen and transports it from lungs to other body parts. It transports CO2 back to the lungs, where it can be breathed out. Oxygen storage also requires iron. Iron is a part of several essential enzymes, and is involved in DNA synthesis. Normal brain functions are iron dependent.

In the body iron is strongly bound to transferring, which enables exchange of the metal between cells. The compound is a strong antibiotic, and it prevents bacteria from growing on the vital element. When one is infected by bacteria, the body produces high amounts of transferring. When iron exceeds the required amount, it is stored in the liver. The bone marrow contains high amounts of iron, because it produces haemoglobin. Iron deficits lead to anaemia, causing tiredness, headaches and loss of concentration. The immune system is also affected. In young children this negatively affects mental development, leads to irritability, and causes concentration disorder. Young children, pregnant women and women in their period are often treated with iron (II) salts upon iron deficits.

When high concentrations of iron are absorbed, for example by haemochromatose patients, iron is stored in the pancreas, the liver, the spleen and the heart. This may damage these vital organs. Healthy people are generally not affected by iron overdose, which is also generally rare. It may occur when one drinks water with iron concentrations over 200 ppm.

Iron compounds may have a more serious effect upon health than the relatively harmless element itself. Water soluble binary iron compounds such as FeCl2 and FeSO4 may cause toxic effects upon concentrations exceeding 200 mg, and are lethal for adults upon doses of 10-50 g. A number of iron chelates may be toxic, and the nerve toxin iron penta carbonyl is known for its strong toxic mechanism. Iron dust may cause lung disease.

About 80% of communicable diseases in the world are due to water. Various undesirable and naturally occurring pollutants in water such as coli form bacteria, iron, fluoride and arsenic are very important as theses pose severe health problems (Joshi and Chaudhuri, 1996). Iron comprises 5 percent of the earth's crust. Iron ores include haemetite (Fe2O4 ), magnetite (Fe3O4 ), limonite (Fe2O33H2O), siderite (FeCO3) and pyrite (FeS2) .Fe+3 minerals are virtually in soluble in water, but siderite has solubility of 65 mg/l. The solubility of siderite (FeCO2) is highly increased by the presence of carbon dioxide or carbonic acid (Das et al. 2007).

Iron is also present as ferrous sulphate in river, lakes or reservoirs containing acid wastes or in ground water containing sulphur particularly, hydrogen sulphide. Commonly referred to as organically bound iron is found in ground and surface supplies. High concentration of Fe+3 occurs in the hypolimniteic zones of entropic lakes and reservoirs when the dissolved oxygen in these segment zones is depleted and the bottom mud contain iron materials. When such waters are used for water supplies, they come in contact with the atmospheric oxygen and Fe+2 is oxidized to Fe+3 forming yellow or red precipitates of ferric hydroxide. Iron containing color compound are stable and are not usually regarded as iron sources, though there may be practical or aesthetic objection to the use of coloured water. Red water owes its appearance and name to suspended insoluble ferric hydroxide formed by the corrosion of the ferrous metal of mains, piping and toxic and oxidation of ferrous carbonate.


Oxidation in Iron Removal

There are several methods for removal of iron from drinking water like ion exchange and water softening (Vaaramaa and Lehto, 2003), activated carbon and other filtration materials (Munter et al. 2005), bioremediation (Berbenni et al. 2000). Oxidation by aeration, chlorination, ozonation is followed by filtration (Ellis et al. 2000). However, oxidation processes are generally used to remove soluble iron from ground water. Actually this is a reversal of the natural process whereby iron is rendered soluble. The most commonly used oxidizing agent is oxygen, which is added to the water by means of aeration. When water is aerated, CO2 is removed from water resulting in an increase in the pH. The rate of oxidation increases rapidly at a pH of 7.0 or more. Therefore any factor which tends to displace the equilibrium:

H2O + CO2 H+ + HCO3- 2H+ + CO3-2

May determine the course of reactions involved in iron removal. The oxidation is usually accomplished by:

(i) Open devices over which water flows by gravity. With or without counter current

forced draft, e.g. open cake tray aerator, open salt tray aerator, closed forced

draft aerator etc.

(ii) Spray devices which spray the water in to the air.

(iii) Diffused air aeration.

(iv) Aspiration devices e.g. venture devices. The most common method of aeration

is cake tray aerator.

Ferrous iron can be oxidized by using other oxidizing agents also. It has been reported that iron in water can be removed almost completely by a free residual chlorine of approximately 0.5 mg/l at normal pH values without using elaborate treatment facilities. Other oxidizing agents usually are ozone potassium permanganate etc.

2.3.2 Unit Processes in Iron Removal

The most common method of iron removal from ground water involve four basic unit processes, viz, adsorption , oxidation , settling and filtration. The precipitated iron is removed partially by sedimentation and partially by filtration. Rapid sand filter or pressure filter are used. Vander Wal (1952) has reported that iron and manganese are too difficult to remove from water without prior flock formation. Whatever treatment method is employed, difficulties of incomplete iron removal are often encountered and in some plants reduction of iron from the ferric state to the ferrous state during filtration has been reported. This is always associated with a marked depletion of dissolved oxygen and a considerable growth of biological slime on the filter. It has been postulated that the chemical reduction of iron is mediated by the bacterial growth in the filter.

Owing to the increased demand for better quality water, various modifications for the conventional method have been suggested. The uses of diatomaceous earth filter for iron removal have been found feasible both practically and economically. The use of precipitators relieves the sand filter of the tremendous load they have to carry in iron removal plants of the aeration, adsorption, and sedimentation and filtration type. Approximately 95% iron can be removed by the Precipitators, when proper pH is maintained. Iron can also be removed during ion exchange softening processes. A sodium zeolite bed or a manganese zeolite bed can be used.


2.4.1 Mechanism of Filtration

The development of the sand filter for water purification took place in England in the mid-nineteenth century. Those filters were developed in the United States to operate at higher filtration rates. The higher rate means less filter area and less capital investment to achieve the desired capacity. A lot of investigations have been conducted to study the mechanism of filtration. According to Cleasby et al. (1963) a very delicate balance exists between those forces tending to deposit and hold to particles. The removal of suspended particles in a filter is believed to be achieved in two steps, transport step followed by an attachment step O'Melia and Stumm, 1967). Particle transport is a physico hydraulic process and is principally affected by those parameters which govern mass transfer. Particle attachment is basically physicochemical process and is influenced by both physical and chemical parameters. In actual filtration practice, removal results from a combination of these mechanisms (Cleasby, 1969). The major transport mechanisms include straining, bulk flow or convective flux (which promotes interception), gravity settling, Brownian diffusion and hydrodynamic action (Ives, 1971) and are affected by such physical characteristics as medium size, filtration rate ,fluid temperature and the density and size of the suspended particle. As the particle approaches the surface of the medium or previously deposited solids on the medium attachment mechanisms are required to retain the particle which are believed to be dependent upon London forces, electrical double layer interaction (electrokinetic's forces) and chemical bridging or specific adsorption. The attachment forces are affected by the coagulants applied in the pre-treatment, and the chemical characteristics of the water and the filter medium. According to Conley Hsiung (1969) flocculation within filter pores may play an important role and the effect of these needs to be re-evaluated.

A variety of materials can be used as effective filtering media. The most durable and perhaps, the cheapest material available is sand, Rapid sand filter are usually operated at filtration rate 94lpm/m2. Other materials like anthracite also are used as filter medium in countries like United State. These filters can be operated at rate much higher than the sand filter.


Many investigations have been made to study the use of materials others than sand as filter media. These investigations have been necessitated due to the comparatively low flow rate in rapid sand filter and increased demand of water. Anthrathracively is used as a filter medium in the United States. Coal is also used the filter medium. One of the most valuable properties of coal is its ability to adsorb from solution many of the dissolved contaminants. Gr. Eskenazy (1970) has reported that beryllium can be adsorbed on peat and coal. It is stated that selected type of dissolved inorganic matter can be removed from water solution by coal either by adsorption or by mechanisms yet to be defined.

The close relationship between the composition of coal and active carbon would indicate that a mean for using coals as an adsorbent, though perhaps in greater quantity, might economically provide comparable results with active carbon. Use of coal has advantages over use of active carbon because of their availability, lower cost and recovery of fuel value after exhaustion.



Adsorption is recognized as a significant phenomenon in most natural physical, biological and chemical processes. Sorption has widely used operation for purification of waters and waste waters. Some important definition used in adsorption is given as (Weber, 1972):

Adsorbate: It is material getting adsorbed or being concentrated (or being removed from one phase).

Adsorbent: It is the adsorbing phase on which the material is getting adsorbed.

Sorption: Include both adsorption and absorption and it is general expression for a process in which a component moves from one phase to be accumulated in another phase, particularly for the cases when second phase is liquid.

Adsorption Equilibrium: In a solid liquid system during adsorption process, the solute from the solution is adsorbed at the surface of solid. This positive adsorption continues till the time, when the concentration of solute remaining in the solution achieves a dynamic equilibrium with the solute concentrated at the surface. However equilibrium is a function of many factor live concentration of solute of nature of solution, temperature of the solution.


Surface tension of liquid is a major factor causing, adsorption of solute solid because adsorptions take place at boundary of liquid phase. Hence is caused as an effect of increased concentration of solute at the surface of liquid. The stages in the adsorption process are:

Film Diffusion: In this process the transport of adsorption molecules/ions through a surface film to the exterior of adsorbent takes place.

Pore Diffusion: In this process the molecules diffuse in to and through pore spaces of the adsorbent.

Interparticle Transport: It is adsorption of active sites on surface bounding the inner pore space of the adsorbent. The slowest of these transport/reaction steps controls the overall rate of uptake by adsorbate and depends on the method of contact. For 'Batch Process' pore diffusion controls the rate of reaction more. Also for sufficient turbulence, transport of adsorbate within the pores is likely to control the overall kinetics.


Adsorption capacity of an adsorbent is a complex function of many variables. Some known variables are being discussed here:

Nature of Adsorbate: Adsorption equilibrium is affected by solubility of solute. Decrease in the solubility of solute results in increased in adsorption. For effective adsorption solute solvent bond must be broken.

A polar solute is strongly adsorbed from a nonpolar solvent by a polar adsorbent (polar of an inorganic compound is a function of charge separation within the molecules). Water solubility is expected to increase with increasing polarity, hence adsorption decreases as polarity increases.

Molecular size of the solute influences the rate of adsorption and can be generalized with particular chemical class. This rate dependence on size is expected only for rapidly agitated batch reactors as it affects interparticle transport. Geometry of molecules may have only smaller effect on equilibrium conditions. Adsorption rate is decreased with increasing molecular weight.

Ionization affects the adsorption rate. Maximum adsorption capacity is for neutral species and structurally simple compounds. More compact molecules of a related pair are adsorbed more rapidly.

Nature of Adsorbent: Pore size distribution governs the rate of transport of adsorbed species from exterior surface to interior surface. It is also related to surface area, as the size distribution of molecular pores may determine what portion of the total surface area will be finally available for the adsorption of solute.

Adsorption capacity does increase with increase in pore size.

Particle size is another characteristic affecting the rate of adsorption, for adsorption on exterior surface of adsorbent.

Chemical nature of surface adsorbent has some effect on adsorbent surface of carbon can be considered as non-polar but usually is slightly polar due to interaction of oxygen with carbon. This causes so called surface acidity of carbon.

Concentration of Solute in Solution: Adsorption capacity depend on concentration of solute in solution phase at constant temperature and in a given system rate and capacity of adsorption increases with increasing with concentration of solute in solution phase.

Dose of Adsorbent: Rate and extent of adsorption vary with doses of adsorbent for a rangs so that no great difference in concentration of solute remains in bulk solution. Second variable is the concentration of adsorbate in the bulk and is created for large differences in concentration of residual solute.

Type of Control and Time of Contact: Batch mixing or continuous flow system may be adopted. Rate of adsorption may differ in the two cases. Mixing increases adsorption rate, but does not alter the adsorption capacity.

Temperature: Adsorption reactions are normally exothermic. Extent of adsorption increases with decreasing temperature but rate of adsorption decreases.

Competitive Interactions : Degree of mutual inhibition of compacting adsorbates is related to the relative sizes of molecules, relative adsorption affinities and relative concentration of solute Total adsorption capacity may increase with mixed solutes due to competitive inter factors.



3.1 Materials and method

Sand with geometric mean size, GMS 0.5 mm and 1.0 mm, has been used in adsorption study.

Preparation of Sample: Appropriate quantities of FeSO47H2O solution are added to tap water to make synthetic raw waters. Initial iron concentration has been taken as mean of 0, 1, 2 hr. concentration.

Table - 4 : Analysis of Tap Water

pH 7.5

Total Iron nil

Alkalinity 378 mg/l

Total Hardness 194 mg/l

Calcium 128 mg/l

Dissolved Oxygen 7. mg/l

Sulphate 320 mg/l

Chlorides 210 mg/l

Conductivity 430. micromoh/cm

3.2 IRON DETERMINATION: Total iron, ferrous iron and dissolved iron were determined by phenanthroline method (APHA et al. 1981) as explained below:

Total Iron:

Mix sample thoroughly and measure 50 ml in to a 125 ml volumetric flask.

Add 2 ml concentrate HCL and 1.0 ml hydroxylamine solution

Add a few glass beads and heat to boiling until volume is reduced to 15 to 20 ml.

Cool to room temperature and transfer to a 50 or 100 ml nessler tube.

Add 10 ml ammonium acetate buffer solution and 4 ml phenanthroline solution and dilute to mark with water.

Mix thoroughly and allow at least 10 to 15 min. for maximum colour development.

Ferrous Iron: Add 2.0 ml concentration HCL/100 ml sample at time of collection immediately with draw a 50 ml portion of acidified sample and add 20 ml phenanthroline solution and 10 ml ammonium acetate buffer solution with vigorous stirring. Dilute to 100 ml and measure colour intensity within 5 to 10 minute.


3.3.1 Adsorption:

The adsorption studies along with the effect of certain parameters like initial iron concentration, initial pH, contact time, dose of adsorption etc. For these studies, 0.5 mm and 1.0 mm geometric mean size i.e. GMS of sand are used of adsorbents preselected amount of adsorbent has been contacted with 100 ml solution of iron at various concentrations in a flask and agitated by stirred for desired time. After desired agitation, solution are allowed to settle for 30 minutes and then passed through the filter paper prior to iron determination.

3.3.2 Column Study:

The experiment has been conducted using a 5.0 cm diameter column as shown in Figure 3.1. A perforated glass plate fixed in position is used as the support for the filter media. The overhead influent reservoir had a capacity of about 100 liters and provided about 100 cm of available head.

Input solution


Output solution

Figure 3.1 Experimental set-up of column experiment.

All glass ware washed with concentration hydrochloric acid followed by rising in tap and distilled water. Appropriate quantity of the stock iron solutions is added to tap water and mixed to make the synthetic iron bearing water. The conductivity and pH of the water before and after adding the ferrous sulphate solution are noted. The filtration rate was adjusted manually and maintained constant during a run. Effluent and influent sample are taken at desired intervals and analyzed for iron.



Some of the important parameters which have great effect on the rate of adsorption, nature of adsorbent and adsorbate, contact time, adsorbent dose initial iron concentration, pH etc. The results of present work are presented in detail. Similarly, for finding the effect of contact time, grain size, and pH at different flow rates are also studied.


To find out the iron concentration at each stage of experiment, calibration of iron solution is necessary. A plot between iron concentration and absorbance at wave length of 500 nm is shown in Table 4.1 and Figure 4.1.

TABLE - 4.1

Calibration of Iron Solution: Absorbance Vs. Concentration of Iron (Wavelength = 500 nm and pH = 7. 5)

Sample Concentration (mg/L)


% Transmittance



















Figure 4.1. Plot between absorbance and concentration of iron.

The equation of line is

X = 7.57 y - 0.35 …………………… ………………(4.1)

Where x = concentration of iron in mg/l,

y = absorbance


The effect of adsorbent dose on iron removal efficiency varying from 1 gm/l to 20 gm/l has been shown in Tables 4.2 and 4.3 and Figures 4.2, 4.3 and 4.4.

Table - 4.2

Adsorption by 0.5 mm sand at initial Iron Concentration of 5 mg/l, initial pH = 7.5,

Contact time = 2 hrs and sample value = 100 ml

S. No.


Iron Conc. after Adsorption (Ce) (mg/l

Amount Adsorbed (mg/l)

Percentage removal of iron




































Figure 4.2. Plot between adsorbent dose and adsorbed iron.

Figure 4.3. Plot between adsorbent dose and percentage removal of iron.

Figure 4.2 indicates that as a dose increases, the adsorption efficiency increases very rapidly up to 8 gm/l but becomes slow beyond this dose. Figure 4.3 represents percentage removal of iron with respect to adsorbent dose. This may be attributed to the fact that with increase in adsorbent dose, more and more adsorbent surface is available for solute to adsorb and this increases the percentage removal. At higher dose, however the adsorption becomes slow, as the large number of particles on stirring had more contact with each other, where the adsorbed iron is partially removed.

It can be further observed from the Figure 4.3 and 4.4 that at an initial concentration 5 mg/l of iron, the maximum efficiency of removal decreases with the geometric mean size (GMS) of sand as at 0.5 mm, it is 97.6%, at 1.0 mm GMS it is 95 %. The reason for this behavior can be explained as an increase in geometric mean size, GMS of the sand for the same mass leads to decrease in surface area available for adsorption and hence decrease in percentage removal.

Table - 4.3

Adsorption by 1.0 mm sand at initial concentration of 5 mg/l, initial pH = 7.5,

Contact time = 2 hrs, sample value = 100 ml.

S. No.


Iron Conc. after Adsorption (mg/L)

Amount Adsorbed (mg/L)

Percentage removal of iron




































Figure 4.4. Plot between adsorbent dose and percentage removal of iron.


The effect of contact time on iron removal efficiency has been shown in Tables 4.6, 4.7 and 4.8:

Table - 4.4

Effect of contact time on Percentage iron removal with pH = 7.5; GMS of sand = 0.5 mm, Initial conc. = 5 mg/l and Dose = 20gm/l

Time (min.)


Iron conc. after adsorption (mg/L)


% Removal (P)=

100*(initial conc.-(2))/initial conc

























Figure 4.5 Plot between time and percentage removal of iron.

Table - 4.5

Effect of contact time on percentage Iron removal with pH = 7.5, GMS of sand = 1.0 mm, Initial concentration = 5 mg/l and Dose = 20gmsl

Time (min.)


Iron conc. after adsorption (mg/L)


% Removal (P)=

100*(initial conc.-(2))/initial conc

























Figure 4.6 Plot between time and percentage of iron removal

Table - 4.6

Effect of concentration of contact time on percentage iron removal with pH =7.5; GMS of Sand = 2.0 mm; Initial concentration = 5 mg/l and Dose = 20gms/l

Time (min.)


Iron conc. after adsorption (mg/L)


% Removal (P)=

100*(initial conc.-(2))/initial conc.




























Figure 4.7 Plot between time and percentage of iron removal.

Tables 4.4 to 4.6 and corresponding Figures 4.5 to 4.7 indicate that as the contact time increases, the adsorption efficiency also increases very rapidly upto time of 60 minutes, but it become slow beyond this time for all grain sizes of sand. The reason for this behavior can be expressed as an increase in contact time up to 60 minute iron removal is faster but after 60 minutes contact time, adsorption and desorption becomes equal causing iron removal to be very slow. Also the surface area available for adsorption reduces as most of the adsorption sites are saturated.


The effect of pH varying from 4 to 10 has been shown in Table 4.7nand 4.8.

Table - 4.7

Effect of Ph percentage Iron Removal Using Sand as Adsorbent In Removal Using Sand as percentage Iron Removal Using Sand As Adsorbent Initial Contact = 5 Mg/l, Contact time = 2 hrs. , DOSE = 10 mg/l GMS of sand = 0.5 mm


After adsorption concentration (mg/L)

% Removal of Iron













Table 4.8

Geometric mean size, GMS of sand = 1.0 mm


After adsorption conc. Mg/l

% removal













Table 4.7 and 4.8 indicate that as the pH is increased, the adsorption efficiency also increases. This may be attributed to the fact that at lower pH below 4, the total iron in the initial iron concentration of 5 mg/l may be expected to be in the soluble state and the removal at lower pH become slow. Above pH 4 a part of the iron is oxides in the presence of dissolved oxygen and the higher percentage removal of iron is obtained. However, percentage removal of iron is higher in case of sand of GMS 0.5mm than compared to sand of GMS 1 mm.



Following conclusions are drawn from the present work.

Two types of geometrical mean size (GMS) of sand, 0.5 mm and 1.0 mm adsorb iron quite significantly but it has been seen that 0.5 mm sand is a better adsorbent.

At lower pH values of the iron solution, the adsorption is slower, while with increase in pH value up to 10 adsorption increases.

Contact time is also one of important parameter. After a particular contact time iron removal occurs at slower rate and smaller in quantity. It is different for different grain size of adsorbents.

On increasing the adsorbent dose, initially iron removal is faster, it becomes almost steady after a certain optimum dose.

Finally, adsorbent material (i.e. sand of geometric mean size 0.5 mm) is found successful for removal of iron up to 97.6% from waste water without using treatment plant.

The present method is very simple and economical.