Determination of ca2+, mg2+ and fe2+ in water subject analytical chemistry

VIETNAM NATIONAL UNIVERSITY – HO CHI MINH CITY HO CHI MINH CITY UNIVERSITY OF TECHNOLOGY DEPARTMENT OF CHEMICAL ENGINEERING DETERMINATION OF Ca2+, Mg2+ AND Fe2+ IN WATER SUBJECT ANALYTICAL CHEMISTRY G[.]

HO CHI MINH CITY UNIVERSITY OF TECHNOLOGY DEPARTMENT OF CHEMICAL ENGINEERING

IN WATER SUBJECT: ANALYTICAL CHEMISTRY

GROUP 10_CH2113_ CC01

INSTRUCTOR: DR TRAN THI KIEU ANH

NGUYEN MINH HOANG YEN – 25% – 2052335

1 INTRODUCTION

The safety of drinking water is a very important public health issue The World Health Organization have established well-defined standards for drinking water purity Nowadays, almost the water source is mixed with ionic impurities that we called water hardness Water ‘hardness’ is a measure of the amount of hard water cations in water These hard water cations include calcium, magnesium, iron, zinc and the other polyvalent metal ions In most water samples, calcium and magnesium are the chief contributors to water hardness It

is unfit for drinking, bathing, washing and it also forms scales in boilers Hence

it is necessary to estimate the amount of hardness producing substances present

in the water sample Once it is estimated, the amount of chemicals required for the treatment of water can be calculated In this report we will discuss about the content of Ca2+¿, Mg2+¿∧Fe2+¿ ¿¿¿ in water by gravimetric method of analysis, titrimetric method of analysis and spectroscopic method of analysis

2 THEORY

2.1 Gravimetric methods of analysis

2.1.1 Principle

Gravimetric methods are quantitative methods based on measurement of weight of an analyte or a compound containing the analyte

Types of analysis:

 Precipitation

 Electrogravimetry

 Volatilization gravimetry (indirect method)

 Particulate gravimetry (direct method)

 In this report, we only focus on Precipitation Gravimetry as the application for determination of Ca2+ and Mg2+

2.1.2 Precipitation Gravimetry

Gravimetry is among the most accurate and precise methods of macroquantitative analysis

The analyte is selectively converted to an insoluble form

The separated precipitate is dried or ignited, possibly to another form, and

is accurately weighed

The weight of analyte in the desired form is calculated from the weight of the precipitate and a knowlegde of its chemical composition (stoichiometric mole relationships)

The steps required for Precipitation Gravimetry:

 Preparation of the solution: This involves several steps

including adjustment of the pH of the solution in order for the precipitate

to occur quantitatively and get a precipitate of desired properties, removing interferences, adjusting the volume of the sample to suit the amount of precipitating agent to be added

 Precipitation:

- This requires addition of a precipitating agent solution to the sample solution Upon addition of the first drops of the precipitating agent, supersaturation occurs, then nucleation starts to occur where every few molecules of precipitate aggregate together forming a nucleus At this point, addition of extra precipitating agent will either form new nuclei or will build up on existing nuclei to give a precipitate Relative supersaturation ¿Q – S S

- The Q is the concentration of reactants before precipitation, S

is the solubility of precipitate in the medium from which it is being precipitated Therefore, to get particle growth instead of further nucleation

we must make the relative supersaturation ratio as small as possible The optimum conditions for precipitation which make the supersaturation low are:

+ Precipitation using dilute solutions to decrease Q

+ Slow addition of precipitating agent to keep Q as low as possible + Stirring the solution during addition of precipitating agent to avoid concentration sites and keep Q low

+ Increase solubility by precipitation from hot solution

+ Adjust the pH to increase S, but not too much increase np as we do not want to lose precipitate by dissolution

+ Usually add a little excess of the precipitating agent for quantitative precipitation and check for completeness of the precipitation

 Digestion of the precipitate: The precipitate is left hot (below

boiling) for 30 min to one hour for the particles to be digested Digestion involves dissolution of small particles and reprecipitation on larger ones resulting in particle growth and better precipitate characteristics This process is called Ostwald ripening An important advantage of digestion is

observed for colloidal precipitates where large amounts of adsorbed ions cover the huge area of the precipitate Digestion forces the small colloidal particles to agglomerate which decreases their surface area and thus adsorption You should know that adsorption is a major problem in gravimetry in case of colloidal precipitate since a precipitate tends to adsorb its own ions present in excess, Therefore, forming what is called a primary ion layer which attracts ions from solution forming a secondary

or counter ion layer Individual particles repel each other keeping the colloidal properties of the precipitate Particle coagulation can be forced

by either digestion or addition of a high concentration of a diverse ions strong electrolytic solution in order to shield the charges on colloidal particles and force agglomeration Usually, coagulated particles return to the colloidal state if washed with water, a process called peptization

 Washing and Filtering the precipitate: It is crucial to wash the

precipitate thoroughly to remove all adsorbed species that would add to the weight of the precipitate One should be careful to use too much water since part of the precipitate may be lost Also, in case of colloidal precipitates we should not use water as a washing solution since peptization would occur In such situations dilute nitric acid, ammonium nitrate, or dilute acetic acid may be used Usually, it is a good practice to check for the presence of precipitating agent in the filtrate of the final washing solution The presence of precipitating agent means that extra washing is required Filtration should be done in appropriately sized ignition filter paper

 Drying and Ignition: The purpose of drying (heating at about

120 - 150 in an oven) or ignition in a muffle furnace at temperatures℃ ranging from 600 - 1200 ℃ is to get a material with exactly known chemical structure so that the amount of analyte can be accurately determined

 Precipitation from Homogeneous Solution: To make Q

minimum we can, in some situations, generate the precipitating agent in the precipitation medium rather than adding it For example, to precipitate iron as the hydroxide, we dissolve urea in the sample Heating of the solution generates hydroxide ions from the hydrolysis of urea Hydroxide ions are generated at all points in solution and thus there are no sites of concentration We can also adjust the rate of urea hydrolysis and thus control the hydroxide generation rate This type of procedure can be very advantageous in case of colloidal precipitates The identification of the quantity of an unknown solute in a solution by making it the limiting

reagent of precipitation reaction, weighting the mass of the resulting precipitate and then using the stoichiometry of the precipitation reaction

to determine the moles of the unknown This technique takes advantage of the solubility rules, and the strategy is to use a double displacement reaction where where one of the products is soluble, while the other forms

a precipitate

2.1.3 Calculation of results - Precipitation Gravimetry

- Solid sample:

m

- Solid sample (dilution):

v × 100 m

- Liquid sample:

X (g/l)=m ' × F × V

v1× 1000 v

2.2 Titrimetric methods of analysis

2.2.1 Principle

Titrimetric methods include a large and powerful group of quantitative procedures based on measuring the amount of a reagent of known concentration that is consumed by an analyte

Titrimetric methods are classified into four groups based on the type of

reaction involved These groups are acid–base titrations, in which an acidic or basic titrant reacts with an analyte that is a base or an acid; complexometric

titrations involving a metal–ligand complexation reaction; redox titrations,

where the titrant is an oxidizing or reducing agent; and precipitation titrations,

in which the analyte and titrant react to form a precipitate

 In this report, we only focus on Complexometric Titrations as the application for determination of Ca2+ and Mg2+

2.2.2 Complexometric Titration

Complexation titration is a titration in which the reaction between the analyte and titrant is a complexation reaction

The most widely used of new ligands was ethylenediaminetetraacetic acid, EDTA, which forms strong 1:1 complexes with many metal ions The first use

of EDTA as a titrant occurred in 1946, when Schwarzenbach introduced metallochromic dyes as visual indicators for signaling the end point of a complexation titration

Figure 1 Structure of EDTA

The equivalence point of a complexation titration occurs when stoichiometrically equivalent amounts of analyte and titrant have reacted For titrations involving metal ions and EDTA, the equivalence point occurs when

CM and CEDTA are equal and may be located visually by looking for the titration curve’s inflection point

As with acid–base titrations, the equivalence point of a complexation titration is estimated by an experimental end point A variety of methods have been used to find the end point, including visual indicators and sensors that respond to a change in the solution conditions For determining water hardness, finding the end point with a visual indicator is chosen as the method used

Finding the End Point with a Visual Indicator: Most indicators for complexation titrations are organic dyes that form stable complexes with metal ions To function as an indicator for an EDTA titration, the metal–indicator complex must possess a color different from that of the uncomplexed indicator Furthermore, the formation constant for the metal–indicator complex must be less favorable than that for the metal–EDTA complex The indicator, Inm–, is added to the solution of analyte, forming a colored metal–indicator complex, MInn-m As EDTA is added, it reacts first with the free analyte, and then

displaces the analyte from the metal–indicator complex, affecting a change in the solution’s color

2.2.3 Calculation of results

- Liquid sample:

X(g/l)=C C × 10−3×V C ×equivalent X × V1

V

- Solid sample:

V X × 100 m

C(standard):the standard solution with knownconcentration

X(analyte):the solution X withknown concentration

2.3 Spectroscopic methods of analysis

2.3.1 Principle

Colorimetric analysis is based on the change in the intensity of the color of

a solution with variations in concentration Colorimetric methods represent the simplest form of absorption analysis The human eye is used to compare the color of the sample solution with a set of standards until a match is found

Radiation is a form of energy and we are constantly reminded of its presence via our sense of sight and ability to feel radiant heat It may be considered in terms of a wave motion where the wavelength, λ, is the distance between two successive peaks The frequency, ν, is the number of peaks passing

a given point per second These terms are related so that:

c =νλ

where c is the velocity of light in a vacuum The full electromagnetic radiation spectrum is continuous and each region merges slowly into the next For spectroscopy purposes, we choose to characterize light in the ultraviolet and visible regions in terms of wavelength expressed in nanometers Other units which may be encountered, but whose use is now discouraged, are the Angstrom (Å) and the millimicron (mµ)

The energy of a photon absorbed or emitted during a transition from one molecular energy level to another is given by the equation

E=hν= hc/λ

h :Planck's constant

ν: frequency of the photon

Figure 2 Electromagnetic Spectrum- Regions

An increase in sensitivity and accuracy results when a spectrophotometer is used to measure the color intensity Basically, it measures the fraction of an incident beam of light which is transmitted by a sample at a particular wavelength You will use a Spectronic 21 in this experiment There are two ways to measure the difference in intensity of the light beam One is the percent transmittance, %T, which is defined as:

For any given compound, the amount of light absorbed depends upon (a) the concentration, (b) the path length, (c) the wavelength and (d) the solvent Absorbance is related to the concentration according to the Beer-Lambert law:

where ε is molar absorptivity ( M−1cm−1 ), b is the path length (cm) and c is the concentration (mol/L )

2.3.2 Instrument

Figure 3 Some kinds of Spectrophotometer 2.3.3 Performing a quantitative measurement:

Step 1: Prepare a quantitative sample on an analytical or semi-analytical

balance then dilute it with suitable solvent

Step 2: Turn on the UV-Vis spectrometer and allow the lamps to warm up

for an appropriate period of time (around 20 min) to stabilize them

Step 3: Fill a cuvette with the solvent for the sample and make sure the

outside is clean This will serve as a blank and help account for light losses due

to scattering or absorption by the solvent

Step 4: Place the cuvette in the spectrometer Make sure to align the

cuvette properly, as often the cuvette has two sides, which are meant for handling (may be grooved) and are not meant to shine light through

Step 5: Take a reading for the blank The absorbance should be minimal,

but any absorbance should be subtracted out from future samples Some instruments might store the blank data and perform the subtraction automatically

3 RESULT

Standard criteria:

Water hardness is calculated as the sum of Ca2+ and Mg2+ content If the hardness does not exceed 300 mg/l, it can be used for normal activities

TDS or Total Dissolved Solids means concentration of dissolved particles

or solids in water TDS for drinking water should be less than 300mg/liter and the maximum limit considered safe is 500mg/liter – this is ideal for healthy mineral-rich water

Figure 4 TDS range for water

Type of gravimetric methods: Precipitation gravimetry

 Ca 2+ :

Implementation process:

Step 1: Preparation of solution

An excess of oxalic acid, H2C2O4, is added to a measured, known volume

of water

Step 2: Precipitation

By adding a reagent, here ammonium oxalate¿¿, the calcium will precipitate as calcium oxalate (CaC2O4) The proper reagent, when added to aqueous solution, will produce highly insoluble precipitates from the positive and negative ions that would otherwise be soluble with their counterparts

The reaction is:

Formation of calcium oxalate: Ca2+(aq) + C2O42- (aq) → CaC2O4 (s)

Step 3: Washing-Drying-Igniting

The precipitate is collected, dried and ignited to high (red) heat which converts it entirely to calcium oxide (CaO)

The reaction is pure calcium oxide formed:

CaC2O4 (s) → CaO(s) + CO(g)+ CO2(g)

Step 4: Weighing

The pure precipitate is cooled, then measured by weighing, and the difference in weights before and after reveals the mass of analyte lost, in this case calcium oxide (CaO)

Calculation: Assume we prepare 100ml of water sample

Assume we meassure the weigh of CaO is 2.268mg = 2.268×10−3g

We have n CaO= m CaO

n C a2 + ¿=n CaC2O4¿ CaO =4.05 × 10−5(mol)¿

C

M Ca2+ ¿ = n C a2+ ¿

V water sample = 4.05× 10−5

100× 10−3=4.05×10−4 (M)=4.05×10−4 (mol/l) ¿¿

Water hadness(C a2+ ¿¿ =C M Ca2+ ¿× M CaCO3× 1 03

−4×100 × 103

 Mg 2+ :

Implementation process:

Step 1: Preparation of solution

An excess of acid phosphoric, H3PO4, is added to a measured, known volume of water

Step 2: Precipitation

By adding a reagent, here diamoni hydro phosphat ¿¿, the magnesium will precipitate as ammonium magnesium phosphate ¿¿ ) The proper reagent, when