The statement of the common ion effect can be written as follows – in a solution wherein there are several species associating with each other via a chemical equilibrium process, an increase in the concentration of one of the ions dissociated in the solution by the addition of another species containing the same ion will lead to an increase in the degree of association of ions.
An example of the common ion effect can be observed when gaseous hydrogen chloride is passed through a sodium chloride solution, leading to the precipitation of the NaCl due to the excess of chloride ions in the solution (brought on by the dissociation of HCl).
This effect cannot be observed in the compounds of transition metals. This is because the d-block elements tend to form complex ions. This can be observed in the compound cuprous chloride, which is insoluble in water. This compound can be dissolved in water by the addition of chloride ions leading to the formation of the CuCl2– complex ion, which is soluble in water.
Effect on Solubility:
How the solubility of a salt in a solution is affected by the addition of a common ion is discussed in this subsection.
The common ion effect can be used to obtain drinking water from aquifers (underground layers of water mixed with permeable rocks or other unconsolidated materials) containing chalk or limestone. Sodium carbonate (chemical formula Na2CO3) is added to the water to decrease the hardness of the water.
In the treatment of water, the common ion effect is used to precipitate out the calcium carbonate (which is sparingly soluble) from the water via the addition of sodium carbonate, which is highly soluble.
A finely divided calcium carbonate precipitate of a very pure composition is obtained from this addition of sodium carbonate. The CaCO3 precipitate is, therefore, a valuable by-product that can be used in the process of manufacturing toothpaste.
Since soaps are the sodium salts of carboxylic acids containing a long aliphatic chain (fatty acids), the common ion effect can be observed in the salting-out process which is used in the manufacturing of soaps. The soaps are precipitated out by adding sodium chloride to the soap solution to reduce its solubility.
However, it can be noted that water containing a respectable amount of Na+ ions, such as seawater and brackish water, can hinder the action of soaps by reducing their solubility and therefore their effectiveness.
This document discusses aromatic compounds and Hückel's rule for aromaticity. It defines aromatic compounds as cyclic, planar and fully conjugated compounds that have 4n + 2 π electrons according to Hückel's rule. These compounds are highly stable due to delocalization of π electrons over the whole ring. They undergo substitution rather than addition reactions and have intermediate bond lengths and diatropic NMR properties. Anti-aromatic compounds have 4n π electrons and show the opposite NMR characteristics. Molecular orbital theory is used to explain the stability and properties of aromatic compounds.
Hybridization is the idea that atomic orbitals fuse to form newly hybridized orbitals, which in turn, influences molecular geometry and bonding properties. Hybridization is also an expansion of the valence bond theory.
This presentation describes the concept of Hyperconjugation in simple words, gives definition of hyperconjugation, explains why it is called as 'No bond Resonance' and gives the effects of hyperconjugation on the chemical properties of compounds: alkyl cations and their relative stability, alkyl radicals and their relative stability, alkenes and their relative stability, bond length, anomeric effect and Baker - Nathan effect.
What is Gravimetric analysis, stepes invloved in gravimetry, Filteration medium in gravimetry, gravimetric factor, application, organic and inorganic prepecating agents
This document discusses Hard and Soft Acids and Bases (HSAB) theory presented by Dr. Satish S. Kola. It defines characteristics of hard vs soft acids and bases, with hard acids/bases being small with high oxidation states and no d-electrons, while soft acids/bases are large with low oxidation states and many d-electrons. Applications of HSAB principles are discussed, including predicting complex formation and metal catalyst poisoning. The theoretical basis involves concepts like pi-bonding, electrostatic interactions, and polarizability. Limitations are noted where inherent acid/base strength may override HSAB predictions.
This document discusses electronic displacement in organic compounds. It describes two types of electronic displacement: permanent displacement including inductive, resonance, and mesomeric effects; and temporary displacement through electromeric effects. Inductive effects are further broken down into +I effects where groups donate electron density and -I effects where groups withdraw electron density. Examples of inductive effects include their impact on acid/base strength, stability of carbocations/carbanions, and dipole moments.
This document provides an overview of redox reactions including:
- Redox reactions involve the transfer of electrons between chemical species, resulting in oxidation and reduction.
- Oxidizing agents gain electrons and are reduced, while reducing agents lose electrons and are oxidized.
- Latimer, Frost, and Pourbaix diagrams can be used to predict and understand redox reactions in aqueous solutions by showing the thermodynamic stability of different oxidation states.
- Key concepts like disproportionation, oxidizing/reducing abilities, and stable/unstable species can be determined from these types of diagrams.
This document defines aromaticity and its key rules. Aromatic molecules have complete delocalization of pi electrons in a planar, conjugated ring system containing 4n+2 pi electrons, as defined by Huckel's rule. Examples of aromatic molecules include benzene and naphthalene. Anti-aromatic molecules have higher energy and contain a 4n pi electron system that is not fully delocalized.
This document discusses aromatic compounds and Hückel's rule for aromaticity. It defines aromatic compounds as cyclic, planar and fully conjugated compounds that have 4n + 2 π electrons according to Hückel's rule. These compounds are highly stable due to delocalization of π electrons over the whole ring. They undergo substitution rather than addition reactions and have intermediate bond lengths and diatropic NMR properties. Anti-aromatic compounds have 4n π electrons and show the opposite NMR characteristics. Molecular orbital theory is used to explain the stability and properties of aromatic compounds.
Hybridization is the idea that atomic orbitals fuse to form newly hybridized orbitals, which in turn, influences molecular geometry and bonding properties. Hybridization is also an expansion of the valence bond theory.
This presentation describes the concept of Hyperconjugation in simple words, gives definition of hyperconjugation, explains why it is called as 'No bond Resonance' and gives the effects of hyperconjugation on the chemical properties of compounds: alkyl cations and their relative stability, alkyl radicals and their relative stability, alkenes and their relative stability, bond length, anomeric effect and Baker - Nathan effect.
What is Gravimetric analysis, stepes invloved in gravimetry, Filteration medium in gravimetry, gravimetric factor, application, organic and inorganic prepecating agents
This document discusses Hard and Soft Acids and Bases (HSAB) theory presented by Dr. Satish S. Kola. It defines characteristics of hard vs soft acids and bases, with hard acids/bases being small with high oxidation states and no d-electrons, while soft acids/bases are large with low oxidation states and many d-electrons. Applications of HSAB principles are discussed, including predicting complex formation and metal catalyst poisoning. The theoretical basis involves concepts like pi-bonding, electrostatic interactions, and polarizability. Limitations are noted where inherent acid/base strength may override HSAB predictions.
This document discusses electronic displacement in organic compounds. It describes two types of electronic displacement: permanent displacement including inductive, resonance, and mesomeric effects; and temporary displacement through electromeric effects. Inductive effects are further broken down into +I effects where groups donate electron density and -I effects where groups withdraw electron density. Examples of inductive effects include their impact on acid/base strength, stability of carbocations/carbanions, and dipole moments.
This document provides an overview of redox reactions including:
- Redox reactions involve the transfer of electrons between chemical species, resulting in oxidation and reduction.
- Oxidizing agents gain electrons and are reduced, while reducing agents lose electrons and are oxidized.
- Latimer, Frost, and Pourbaix diagrams can be used to predict and understand redox reactions in aqueous solutions by showing the thermodynamic stability of different oxidation states.
- Key concepts like disproportionation, oxidizing/reducing abilities, and stable/unstable species can be determined from these types of diagrams.
This document defines aromaticity and its key rules. Aromatic molecules have complete delocalization of pi electrons in a planar, conjugated ring system containing 4n+2 pi electrons, as defined by Huckel's rule. Examples of aromatic molecules include benzene and naphthalene. Anti-aromatic molecules have higher energy and contain a 4n pi electron system that is not fully delocalized.
The document discusses hybridization, which is the mixing of atomic orbitals to form new hybrid orbitals. It provides rules for hybridization, including that orbitals with similar energies can hybridize to form equivalent hybrid orbitals. The number of hybrid orbitals formed equals the number participating in hybridization. Different types of hybridization, including sp, sp2, sp3, sp3d, sp3d2, and sp3d3 are covered along with their orbital and molecular geometries. Examples of beryllium chloride (sp hybridization) and boron trifluoride (sp2 hybridization) are given.
The document discusses inductive effect and resonance effect. Inductive effect refers to polarization of a sigma bond due to electron withdrawing or donating groups. Electron withdrawing groups have a negative inductive effect (-I) while electron donating groups have a positive inductive effect (+I). Resonance effect refers to delocalization of pi electrons or a lone pair. Electron withdrawing groups have a negative resonance effect (-R) while electron donating groups have a positive resonance effect (+R). In most cases, resonance effect is stronger than inductive effect. The document provides examples of how these effects influence acidity, reactivity, and stability.
This document discusses colligative properties of solutions and ways of expressing concentration. It begins by defining key terms like solute, solvent, concentration, dilute and concentrated solutions. It then describes various ways of expressing concentration including percentage by weight, mole fraction, molarity, molality, normality and parts per million. The document also discusses colligative properties like lowering of vapor pressure, elevation of boiling point, depression of freezing point and osmotic pressure. It provides equations and experimental methods for determining these properties and using them to calculate molecular masses. The concept of abnormal molar masses from association or dissociation in solution is introduced along with the van't Hoff factor.
Crystal Field Theory explains the colors of transition metal complexes based on ligand-metal interactions. The electrostatic interaction between ligands and metal d-orbitals splits the d-orbital energies. For an octahedral complex, the d-orbitals point directly at ligands have higher energy than those that bisect ligands. This splitting pattern determines if the complex is high or low spin, which then dictates its color and magnetic properties. The spectrochemical series orders ligands by their ability to cause crystal field splitting, correlating ligand type with complex color.
The document summarizes the hard and soft acid and base (HSAB) theory introduced by Ralph Pearson. It defines hard and soft acids and bases based on their polarizability, with hard species being less polarizable and soft species being more polarizable. The key principle of HSAB theory is that hard acids prefer to interact with hard bases via ionic interactions, while soft acids prefer interacting with soft bases via covalent interactions. Examples of interactions between different combinations of hard/soft acids and bases are provided. Limitations of Pearson's original HSAB model are also outlined.
This presentation consists of three topics that are:
1. conductance of electrolytic solution
2. Specific Conductance, Molar Conductance & Equivalent Conductance
3. Kohlrausch's Law
This document provides an overview of chapter 15 on acids and bases. It begins with an outline of the key concepts and sections to be covered, including the Arrhenius, Brønsted-Lowry, and Lewis definitions of acids and bases. It then discusses acid and base strengths, factors that influence strength, and the concepts of conjugate acid-base pairs and amphiprotic species. Another section covers self-ionization of water and how pH is defined and calculated. Learning objectives are provided for each major topic.
The document discusses four important techniques used in organic chemistry: recrystallization, distillation, sublimation, and chromatography. Recrystallization involves dissolving an impure solid in a hot solvent and slowly cooling the solution to form purer crystals. Distillation separates liquids with different boiling points. Sublimation transitions a substance directly from solid to gas without an intermediate liquid phase. Chromatography separates mixtures by differential absorption of compounds onto a stationary phase as they are carried through by a mobile phase.
Complexometric titration involves the titration of a metal ion with EDTA (ethylene diamine tetraacetic acid) where a colored complex is formed at the endpoint. There are several types of complexometric titrations including direct titration where the metal ion is directly titrated with EDTA, back titration where excess EDTA is added and then titrated with another metal ion, and replacement titration where the metal ion displaces another metal ion from an EDTA complex which is then titrated. Complexometric titrations can be used to determine mixtures of metal ions and are useful because EDTA forms very stable complexes with most metal ions.
The document summarizes Arrhenius theory of acids and bases as proposed by Svante Arrhenius in 1887. The key points are:
1) Arrhenius proposed that acids donate H+ ions and bases donate OH- ions when dissolved in water.
2) Electrolytes conduct electricity in solution by dissociating into ions, while nonelectrolytes do not dissociate or conduct.
3) The theory explained several phenomena like heat of neutralization and differentiation of strong/weak electrolytes but had limitations like not explaining ion formation mechanisms.
RAOULT'S LAW ( Physical & Analytical Chemistry)Hasnaın Sheıkh
Name; Hasnain Nawaz
Surname : Shaikh
ROLL NO: 16 CH 42
B.E: Chemical Engineering (In Progress).
Mehran University of Engineering and Technology
Jamshore, ISO 9001 Certified.
Alkyl halides are derivatives of alkanes where one or more hydrogen atoms are replaced by halogen atoms such as fluorine, chlorine, bromine, or iodine. They are represented by R-X, where R is an alkyl group and X is a halogen. Common methods for preparing alkyl halides include direct halogenation of alkanes, addition of hydrogen halides to alkenes and alkynes, and reactions of hydrogen halides, phosphorus halides, or thionyl chloride with alcohols. Alkyl halides undergo nucleophilic substitution and elimination reactions. They can be reduced to alkanes or used to form Grignard reagents. Common uses
This document discusses oxidation-reduction (redox) reactions and concepts including definitions of oxidation and reduction in terms of gaining or losing electrons, oxygen, and hydrogen. It provides examples of redox reactions and identifies the oxidizing agent and reducing agent in reactions. It also discusses oxidation numbers and how to balance redox equations using the oxidation number change method. Finally, it discusses redox titrations and the specific methods of iodimetry and iodometry which involve the use of iodine as the titrant or analyte.
This document discusses acid-base theories and titration. It covers:
1) Arrhenius, Bronsted-Lowry, and Lewis acid-base theories.
2) Types of acids and bases as strong or weak.
3) The law of mass action and dissociation constants.
4) Neutralization curves for different types of acid-base titrations and the pH at equivalence points.
5) Choice of indicators for different titrations and mixed indicators.
The document discusses various factors that affect the stability of metal complexes. It explains that complexes formed with ligands having higher charge and smaller size are generally more stable. It also discusses the Irving-Williams order of stability and the factors of charge to radius ratio, electronegativity, and basicity of ligands. The chelate effect is described as an important ligand effect where multidentate ligands form more stable complexes due to entropy gains. Kinetic and thermodynamic stability are distinguished from reactivity concepts of labile and inert complexes.
1) The document describes acid-base titration techniques, including defining terms like equivalence point and end point.
2) It discusses different types of titrations including strong acid-strong base, weak acid-strong base, and constructing titration curves.
3) Key points are made about calculating pH values before, at, and after the equivalence point for different titration scenarios. The document provides examples of constructing titration curves step-by-step.
Resonance structures represent different arrangements of electrons in a molecule that have the same positions of nuclei but different bonding patterns. Resonance contributes to the stability of molecules like benzene by delocalizing electrons across multiple equivalent structures. The actual structure of a molecule represented by resonance is a hybrid of the contributing structures, with bond lengths intermediate between single and double bonds. Delocalization of electrons is depicted using curved arrows between resonance structures.
Electrophilic additions involve reactions of alkenes where the pi electrons in the double bond attack an electrophile. There are several types of additions including addition of HX, halogens, water, alcohols, and hydroboration. The mechanism typically involves formation of a carbocation intermediate that is then attacked by the nucleophile. Addition occurs regioselectively according to Markovnikov's rule, favoring the most stable carbocation. Exceptions include free radical additions, which give the anti-Markovnikov product. Oxymercuration-demercuration and hydroboration allow for Markovnikov addition without rearrangements.
This document discusses the distribution law, also known as the partition coefficient, which describes how a solute will distribute between two immiscible solvents at equilibrium. It provides the equation that the concentration of the solute in the first solvent (C1) divided by the concentration in the second solvent (C2) equals the distribution coefficient (KD). Several applications of the distribution law are described, including solvent extraction, partition chromatography, and determining solubility, dissociation, and association.
Common Ion Effect: The shift in equilibrium that occurs because of the addition of an ion already involved in the equilibrium reaction. AgCl(s) « Ag+(aq) + Cl-(aq) adding NaCl(aq) shifts equilibrium position.
Weak Acid - A stoichiometry problem - reaction is assumed to run to completion - then determine remaining species. Step 2 - An equilibrium problem - determine the position of weak acid equilibrium and calculate the pH
The document discusses hybridization, which is the mixing of atomic orbitals to form new hybrid orbitals. It provides rules for hybridization, including that orbitals with similar energies can hybridize to form equivalent hybrid orbitals. The number of hybrid orbitals formed equals the number participating in hybridization. Different types of hybridization, including sp, sp2, sp3, sp3d, sp3d2, and sp3d3 are covered along with their orbital and molecular geometries. Examples of beryllium chloride (sp hybridization) and boron trifluoride (sp2 hybridization) are given.
The document discusses inductive effect and resonance effect. Inductive effect refers to polarization of a sigma bond due to electron withdrawing or donating groups. Electron withdrawing groups have a negative inductive effect (-I) while electron donating groups have a positive inductive effect (+I). Resonance effect refers to delocalization of pi electrons or a lone pair. Electron withdrawing groups have a negative resonance effect (-R) while electron donating groups have a positive resonance effect (+R). In most cases, resonance effect is stronger than inductive effect. The document provides examples of how these effects influence acidity, reactivity, and stability.
This document discusses colligative properties of solutions and ways of expressing concentration. It begins by defining key terms like solute, solvent, concentration, dilute and concentrated solutions. It then describes various ways of expressing concentration including percentage by weight, mole fraction, molarity, molality, normality and parts per million. The document also discusses colligative properties like lowering of vapor pressure, elevation of boiling point, depression of freezing point and osmotic pressure. It provides equations and experimental methods for determining these properties and using them to calculate molecular masses. The concept of abnormal molar masses from association or dissociation in solution is introduced along with the van't Hoff factor.
Crystal Field Theory explains the colors of transition metal complexes based on ligand-metal interactions. The electrostatic interaction between ligands and metal d-orbitals splits the d-orbital energies. For an octahedral complex, the d-orbitals point directly at ligands have higher energy than those that bisect ligands. This splitting pattern determines if the complex is high or low spin, which then dictates its color and magnetic properties. The spectrochemical series orders ligands by their ability to cause crystal field splitting, correlating ligand type with complex color.
The document summarizes the hard and soft acid and base (HSAB) theory introduced by Ralph Pearson. It defines hard and soft acids and bases based on their polarizability, with hard species being less polarizable and soft species being more polarizable. The key principle of HSAB theory is that hard acids prefer to interact with hard bases via ionic interactions, while soft acids prefer interacting with soft bases via covalent interactions. Examples of interactions between different combinations of hard/soft acids and bases are provided. Limitations of Pearson's original HSAB model are also outlined.
This presentation consists of three topics that are:
1. conductance of electrolytic solution
2. Specific Conductance, Molar Conductance & Equivalent Conductance
3. Kohlrausch's Law
This document provides an overview of chapter 15 on acids and bases. It begins with an outline of the key concepts and sections to be covered, including the Arrhenius, Brønsted-Lowry, and Lewis definitions of acids and bases. It then discusses acid and base strengths, factors that influence strength, and the concepts of conjugate acid-base pairs and amphiprotic species. Another section covers self-ionization of water and how pH is defined and calculated. Learning objectives are provided for each major topic.
The document discusses four important techniques used in organic chemistry: recrystallization, distillation, sublimation, and chromatography. Recrystallization involves dissolving an impure solid in a hot solvent and slowly cooling the solution to form purer crystals. Distillation separates liquids with different boiling points. Sublimation transitions a substance directly from solid to gas without an intermediate liquid phase. Chromatography separates mixtures by differential absorption of compounds onto a stationary phase as they are carried through by a mobile phase.
Complexometric titration involves the titration of a metal ion with EDTA (ethylene diamine tetraacetic acid) where a colored complex is formed at the endpoint. There are several types of complexometric titrations including direct titration where the metal ion is directly titrated with EDTA, back titration where excess EDTA is added and then titrated with another metal ion, and replacement titration where the metal ion displaces another metal ion from an EDTA complex which is then titrated. Complexometric titrations can be used to determine mixtures of metal ions and are useful because EDTA forms very stable complexes with most metal ions.
The document summarizes Arrhenius theory of acids and bases as proposed by Svante Arrhenius in 1887. The key points are:
1) Arrhenius proposed that acids donate H+ ions and bases donate OH- ions when dissolved in water.
2) Electrolytes conduct electricity in solution by dissociating into ions, while nonelectrolytes do not dissociate or conduct.
3) The theory explained several phenomena like heat of neutralization and differentiation of strong/weak electrolytes but had limitations like not explaining ion formation mechanisms.
RAOULT'S LAW ( Physical & Analytical Chemistry)Hasnaın Sheıkh
Name; Hasnain Nawaz
Surname : Shaikh
ROLL NO: 16 CH 42
B.E: Chemical Engineering (In Progress).
Mehran University of Engineering and Technology
Jamshore, ISO 9001 Certified.
Alkyl halides are derivatives of alkanes where one or more hydrogen atoms are replaced by halogen atoms such as fluorine, chlorine, bromine, or iodine. They are represented by R-X, where R is an alkyl group and X is a halogen. Common methods for preparing alkyl halides include direct halogenation of alkanes, addition of hydrogen halides to alkenes and alkynes, and reactions of hydrogen halides, phosphorus halides, or thionyl chloride with alcohols. Alkyl halides undergo nucleophilic substitution and elimination reactions. They can be reduced to alkanes or used to form Grignard reagents. Common uses
This document discusses oxidation-reduction (redox) reactions and concepts including definitions of oxidation and reduction in terms of gaining or losing electrons, oxygen, and hydrogen. It provides examples of redox reactions and identifies the oxidizing agent and reducing agent in reactions. It also discusses oxidation numbers and how to balance redox equations using the oxidation number change method. Finally, it discusses redox titrations and the specific methods of iodimetry and iodometry which involve the use of iodine as the titrant or analyte.
This document discusses acid-base theories and titration. It covers:
1) Arrhenius, Bronsted-Lowry, and Lewis acid-base theories.
2) Types of acids and bases as strong or weak.
3) The law of mass action and dissociation constants.
4) Neutralization curves for different types of acid-base titrations and the pH at equivalence points.
5) Choice of indicators for different titrations and mixed indicators.
The document discusses various factors that affect the stability of metal complexes. It explains that complexes formed with ligands having higher charge and smaller size are generally more stable. It also discusses the Irving-Williams order of stability and the factors of charge to radius ratio, electronegativity, and basicity of ligands. The chelate effect is described as an important ligand effect where multidentate ligands form more stable complexes due to entropy gains. Kinetic and thermodynamic stability are distinguished from reactivity concepts of labile and inert complexes.
1) The document describes acid-base titration techniques, including defining terms like equivalence point and end point.
2) It discusses different types of titrations including strong acid-strong base, weak acid-strong base, and constructing titration curves.
3) Key points are made about calculating pH values before, at, and after the equivalence point for different titration scenarios. The document provides examples of constructing titration curves step-by-step.
Resonance structures represent different arrangements of electrons in a molecule that have the same positions of nuclei but different bonding patterns. Resonance contributes to the stability of molecules like benzene by delocalizing electrons across multiple equivalent structures. The actual structure of a molecule represented by resonance is a hybrid of the contributing structures, with bond lengths intermediate between single and double bonds. Delocalization of electrons is depicted using curved arrows between resonance structures.
Electrophilic additions involve reactions of alkenes where the pi electrons in the double bond attack an electrophile. There are several types of additions including addition of HX, halogens, water, alcohols, and hydroboration. The mechanism typically involves formation of a carbocation intermediate that is then attacked by the nucleophile. Addition occurs regioselectively according to Markovnikov's rule, favoring the most stable carbocation. Exceptions include free radical additions, which give the anti-Markovnikov product. Oxymercuration-demercuration and hydroboration allow for Markovnikov addition without rearrangements.
This document discusses the distribution law, also known as the partition coefficient, which describes how a solute will distribute between two immiscible solvents at equilibrium. It provides the equation that the concentration of the solute in the first solvent (C1) divided by the concentration in the second solvent (C2) equals the distribution coefficient (KD). Several applications of the distribution law are described, including solvent extraction, partition chromatography, and determining solubility, dissociation, and association.
Common Ion Effect: The shift in equilibrium that occurs because of the addition of an ion already involved in the equilibrium reaction. AgCl(s) « Ag+(aq) + Cl-(aq) adding NaCl(aq) shifts equilibrium position.
Weak Acid - A stoichiometry problem - reaction is assumed to run to completion - then determine remaining species. Step 2 - An equilibrium problem - determine the position of weak acid equilibrium and calculate the pH
Solubility is the amount of solute that will dissolve in a given amount of solution at a particular temperature (in grams or moles)
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The molar solubility (mol/L) is the number of moles of solute that will dissolve in 1L of a saturated solution.
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The molarity of the dissolved solute in a saturated solution.
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Solubility (g/L) is the number of grams of solute dissolved in 1 L of a saturated solution.
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A saturated solution contains the maximum amount of solute possible at a given temperature in equilibrium with an undissolved excess of the substance.
The solubility product is a kind of equilibrium constant and its value depend...RidhaTOUATI1
The solubility product is a kind of equilibrium constant and its value depends on temperature. Ksp usually increases with an increase in temperature due to increased solubility. Solubility is defined as a property of a substance called solute to get dissolved in a solvent in order to form a solution
Adding AgNO3 to a saturated solution of AgCl increases the concentration of Ag+ ions, making the ion product (Qsp) greater than the solubility product (Ksp) and causing the solution to be out of equilibrium. The solution will re-equilibrate by precipitating excess AgCl, resulting in a higher Ag+ concentration than Cl- concentration. Similarly, adding NaCl increases the Cl- concentration, making Qsp > Ksp and causing re-equilibration with a lower Ag+ concentration than Cl-. The concentrations of ions in equilibrium solutions fall along a curved line where the ion product equals Ksp.
This document summarizes Chapter 19 from a general chemistry textbook. The chapter covers solubility product constants, solubility equilibria, common ion effects, limitations of Ksp, precipitation criteria, fractional precipitation, solubility and pH, and equilibria involving complex ions. It provides examples and explanations of these concepts. Qualitative cation analysis is also discussed, describing how cations can be selectively precipitated into groups for identification purposes.
This chapter of the general chemistry textbook discusses solubility and complex ion equilibria. It covers topics such as the solubility product constant Ksp, the common ion effect, limitations of Ksp, criteria for precipitation, fractional precipitation, effects of pH on solubility, and equilibria involving complex ions. It also describes the process of qualitative cation analysis using selective precipitation of cations in different solubility groups.
This document discusses various concepts related to gravimetric analysis methods. It covers three key points:
1) Gravimetric analysis involves selectively precipitating the analyte of interest and weighing the precipitate to determine the amount of analyte. Factors like solubility products (Ksp), common ion effects, and pH can impact precipitation.
2) Key steps in gravimetric analysis are discussed, including filtering, drying, and transferring precipitates. Equipment like filters, crucibles, and drying ovens are also mentioned.
3) Solubility is impacted by various equilibrium concepts like Ksp values, common ion effects, salt effects, pH, complexation, and temperature. These concepts are illustrated through
1. Precipitation involves combining two ionic species to form an insoluble product, forcing the reaction to completion. For precipitation reactions to be useful in titrimetric analysis, the precipitate must be insoluble and form rapidly and quantitatively without interference from adsorption effects. The equivalence point must also be detectable.
2. Factors that affect solubility include common ion effect, temperature, solvent, and pH. Solubility product constants describe the solubility equilibrium of slightly soluble salts. Fractional precipitation determines which salt will precipitate first when multiple anions are present.
3. Direct and indirect precipitation titrations can be used. Direct titrations detect the equivalence point by formation of a colored
This document discusses solubility products (Ksp) and how they relate to the solubility of various salts such as AgCl, Hg2Cl2, and PbCl2. It also discusses how to use Ksp values to determine if a precipitate will form when solutions are mixed, and describes experiments performed with the various salts including forming complexes and disproportionation reactions.
This document discusses solubility and solubility products. It defines solubility as the amount of a substance that will dissolve in a given amount of solvent. The solubility product (Ksp) is the mathematical product of the dissolved ion concentrations raised to their stoichiometric coefficients. Ksp is a constant for a given ionic compound at a given temperature. Examples are provided to demonstrate how to calculate Ksp from solubility data and vice versa. Factors that affect solubility like temperature, common ions, pH, and complex ion formation are also mentioned. Finally, different types of solutions like unsaturated, saturated, and supersaturated solutions are defined.
Solubility is defined as the maximum amount of solute that can dissolve in a solvent at equilibrium. The solubility product is the maximum product of ion concentrations of an electrolyte that can be in equilibrium with the undissolved solid phase. When an ionic solid dissolves, it dissociates into separate cations and anions. At equilibrium, the rates of dissolution and precipitation are equal. The solubility product principle states that the product of ion concentrations in a saturated solution is constant at a given temperature. The solubility product (Ksp) can be used to calculate solubility from ion concentrations or vice versa. Precipitation will occur if the ionic product (Q) is greater than Ksp when
1. Solubility is defined as the amount of a substance that dissolves in water completely to give free ions, and can be expressed in units of g/L or mol/dm3. Molar solubility refers to the amount of solute in moles that dissolves in a given volume of solvent.
2. The solubility product constant (Ksp) represents the equilibrium between ions in solution and undissolved solid for sparingly soluble salts. Ksp is the product of the concentrations of ions raised to their stoichiometric coefficients at the point of saturation.
3. The molar solubility of a salt can be calculated from its Ksp value by setting up an ionic
This document discusses heterogeneous equilibria involving precipitation reactions and solubility products (Ksp). It defines Ksp as the equilibrium constant for the formation of a precipitate from a saturated solution. For unsaturated solutions, the ion product (Qsp) is used instead of Ksp. Examples are provided to demonstrate calculating Ksp from molar solubility and vice versa. The relationship between Qsp, Ksp and saturation is explained. Factors that affect solubility like common ions are also discussed.
3rd Lecture on Ionic Equilibria | Chemistry Part I | 12th StdAnsari Usama
Buffer solutions resist changes in pH when small amounts of acid or base are added. There are two types: acidic buffers containing a weak acid and its salt, and basic buffers containing a weak base and its salt. The pH of a buffer solution is related to the ratio of the salt and acid or base concentrations. Buffer solutions have important applications in biological systems, agriculture, medicine, and analytical chemistry. They maintain pH despite additions of H+ or OH-.
1. Precipitation titrations involve the titration of an analyte solution with a reagent that causes the formation of an insoluble precipitate.
2. The Mohr method is commonly used for chloride determination and involves titrating a chloride solution with silver nitrate using potassium chromate as an indicator.
3. At the endpoint, excess silver nitrate reacts with potassium chromate to form a reddish brown silver chromate precipitate, indicating that all the chloride ions have been precipitated out as silver chloride.
Discusses the chemical of slightly soluble compounds. Ksp and factors affecting solubility are included as well as solved problems.
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This document discusses solubility equilibria and the formation of precipitates. It defines key terms like solubility product constant (Ksp), explains how to calculate Ksp values from molar solubility and vice versa, and shows examples of using Ksp to determine whether precipitates will form when solutions are mixed. The key points are that solubility is dependent on equilibrium, saturated solutions have concentrations where Ksp = Q, and precipitates form when mixing produces Q > Ksp (supersaturation).
This document discusses the solubility product expression for silver chloride (AgCl). It explains that AgCl dissociates into silver ions (Ag+) and chloride ions (Cl-) in water according to the equation AgCl(s) ⇌ Ag+(aq) + Cl-(aq). The solubility product expression, Ksp, is defined as the product of the molar concentrations of Ag+ and Cl- ions at equilibrium. The value of Ksp is related to the solubility of AgCl. Solubility product calculations can be used to determine the solubility of salts like silver bromide (AgBr) in water.
1. IMPORTANT FEATURES • Chi0nous cell walls • Hyphae with septa (cross‐walls) – Except for Yeasts (non‐hypha) • Ability of soma0c assimila0ve hypha to fuse with one another and to exchange nuclei (anastomosis) • Occurrence in their life cycle of a dikaryon (diploid)
2. THE MEIOSPORANGIA OR ASCUS
3. IMPORTANT CHARACTERS IN AN ASCOMYCETE LIFE CYCLE HOLOMORPH = ANAMORPH + TELEOMORPH WHOLE FUNGUS = ASEXUAL REPRODUCTION + SEXUAL REPRODUCTION
4. THE TELEOMORPH (SEXUAL) • When an ascospore germinates, it establishes a haploid mycelium. • In heterothallic ascomycetes, this can't undergo sexual reproduc0on un0l it meets another compa0ble haploid mycelium. • When this rare event takes place, the fungus cleverly maximizes the ensuing poten0al for gene0c recombina0on. • One would expect a single sexual fusion, resul0ng in a single zygote
5. THE ASCOMA OR ASCOMATA
6. ASCOMYCETE FRUITING STRUCTURES • The mul0cellular structures (ascomata) that produce the asci, and act as the plaMorms from which the spores are launched • Four Types – Apothecial – Perithecial – Pseudothecial – Cleistothecial
7. APOTHECIAL ASCOMA • The construc0on of the ascoma may allow several or many asci to discharge simultaneously because the en0re fer0le layer or hymenium is exposed...
8. PERITHECIAL ASCOMATA • contain unitunicate‐ inoperculate asci
9. PSEUDOTHECIAL ASCOMATA • contain bitunicate asci
10. CLEISTOTHECIAL ASCOMATA • asci oPen spherical and no longer shoot their spores: the fungus has evolved a new dispersal strategy • the fungus fruits in a confined space (for example, under bark, or below the surface of the ground) where airborne dispersal cannot operate • asci are clearly visible and not arranged in a layer or hymenium (as they were in the other three kinds of ascoma)
11. THE ASCUS/ASCI The Ascomycete Spores
12. FOUR TYPES OF ASCUS/ASCI • Unitunicate Operculate • Unitunicate Inoperculate • Protunicate • Bitunicate
13. UNITUNICATE OPERCULATE ASCI • have a single wall • have a built‐in lid or operculum which pops open so that the spores can be ejected • found only in apothecial ascomata
14. UNITUNICATE INOPERCULATE ASCI • have no operculum, but have a special elas0c ring mechanism built into their 0p (apical ring) • this is a pre‐set pressure release valve, or sphincter, and the ring eventually stretches momentarily, or turns inside out, to let the spores shoot through • found in perithecial and some apothecial ascomata
Quality Control
This part of the proposal which has three names: Quality Control, Warranty, and Evaluation procedure, states how the proponent will determine or test the quality, effectiveness, or greatness of the accomplished project. The following are some ways by which the proponent may reveal to the reader his method of determining the quality or effectiveness of his offer:
1. Submitting progress reports to assess any improvements or advancement of the project vis-à-vis the objectives of the plan.
2. Conducting a pre-test and post-test questionnaire about salient aspects of the project at a certain interval of time like semestral, annual, or biennial.
3. Giving a warranty or guarantee for a certain period.
4. Honest disclosure or admittance of inevitable limitations, defects, or drawbacks needing replacements, changes, exceptions, and recommendations.
Quality control, or QC for short, is a process by which entities review the quality of all factors involved in production. This approach emphasizes three aspects: 1. Elements such as controls, job management, defined and well-managed processes, performance and integrity criteria, and identification of records 2. Competence, such as knowledge, skills, experience, and qualifications 3. Soft elements, such as personnel, integrity, confidence, organizational culture, motivation, team spirit, and quality relationships
Volumetric flasks are used for the dilution of solutions to a specific volume. They have accurate volume. come in various sizes, from 1 mL to 2 L or more. A typical flask is shown in the figure. These flasks are designed to contain an accurate volume at the specified temperature (20 or 25◦ C) when the bottom of the meniscus (the concave curvature of the upper surface of the water in a column caused by capillary action) touches the line across the neck of the glass. These flasks are marked with “TC” to indicate “to contain.”
Fundamentals of Analytical Chemistry by Skoog
Advanced analytical chemistry By GD Christian
https://en.wikipedia.org/wiki/Volumetric_flask
They are directly marked on the face by the manufacturer as to the uncertainty of the container measurement; for example, a 250 mL volumetric flask is “±0.24 mL,” or roughly a 0.1% error. Initially, a small amount of diluent (usually distilled water) is added to the empty flask. Reagent chemicals should never be added directly to a dry glass surface, as glass is highly absorbent. When using a volumetric flask, a solution should be prepared stepwise.
Fundamentals of analytical Chemistry by Skoog
Advanced analytical chemistry By GD Christian
https://www.youtube.com/watch?v=hrvXuX0Ow3s
The Krebs cycle summarizes a circular series of reactions in the mitochondria to metabolize AcCoA to two molecules of CO2 with the resultant generation of one molecule of GTP, three molecules of NADH, and one molecule of FADH2. GTP is equivalent to ATP in terms of energy charge.
FUNCTION: The main function of the Krebs cycle is to produce energy, stored and transported as ATP or GTP. The cycle is also central to other biosynthetic reactions where the intermediates produced are required to make other molecules, such as amino acids, nucleotide bases, and cholesterol.
The pentose phosphate pathway is also called the Hexose monophosphate pathway/ HMP shunt/ Phosphogluconate pathway.
It is an alternative route for the metabolism of glucose.
It is a more complex pathway than glycolysis.
It is more anabolic.
It takes place in the cytosol.
The tissues such as the liver, adipose tissue, adrenal gland, erythrocytes, testes, and lactating mammary gland are highly active in the HMP shunt.
It is concerned with the biosynthesis of NADPH and pentoses.
Ch 13 biochemistry-Bioenergetics and Thermodynamics (Enthalpy, Entropy, Gibs ...IshaTariq8
Bioenergetics is the study of energy flow and transformation within living organisms. All living things require energy to grow and reproduce, and to maintain order and carry out life's processes. Energy enters living systems through photosynthesis or consumption of organic compounds, and is used to drive cellular processes like biosynthesis, movement, transport, and more through the release of energy from breaking down food molecules.
Histololgy of Female Reproductive System.pptxAyeshaZaid1
Dive into an in-depth exploration of the histological structure of female reproductive system with this comprehensive lecture. Presented by Dr. Ayesha Irfan, Assistant Professor of Anatomy, this presentation covers the Gross anatomy and functional histology of the female reproductive organs. Ideal for students, educators, and anyone interested in medical science, this lecture provides clear explanations, detailed diagrams, and valuable insights into female reproductive system. Enhance your knowledge and understanding of this essential aspect of human biology.
Local Advanced Lung Cancer: Artificial Intelligence, Synergetics, Complex Sys...Oleg Kshivets
Overall life span (LS) was 1671.7±1721.6 days and cumulative 5YS reached 62.4%, 10 years – 50.4%, 20 years – 44.6%. 94 LCP lived more than 5 years without cancer (LS=2958.6±1723.6 days), 22 – more than 10 years (LS=5571±1841.8 days). 67 LCP died because of LC (LS=471.9±344 days). AT significantly improved 5YS (68% vs. 53.7%) (P=0.028 by log-rank test). Cox modeling displayed that 5YS of LCP significantly depended on: N0-N12, T3-4, blood cell circuit, cell ratio factors (ratio between cancer cells-CC and blood cells subpopulations), LC cell dynamics, recalcification time, heparin tolerance, prothrombin index, protein, AT, procedure type (P=0.000-0.031). Neural networks, genetic algorithm selection and bootstrap simulation revealed relationships between 5YS and N0-12 (rank=1), thrombocytes/CC (rank=2), segmented neutrophils/CC (3), eosinophils/CC (4), erythrocytes/CC (5), healthy cells/CC (6), lymphocytes/CC (7), stick neutrophils/CC (8), leucocytes/CC (9), monocytes/CC (10). Correct prediction of 5YS was 100% by neural networks computing (error=0.000; area under ROC curve=1.0).
These lecture slides, by Dr Sidra Arshad, offer a simplified look into the mechanisms involved in the regulation of respiration:
Learning objectives:
1. Describe the organisation of respiratory center
2. Describe the nervous control of inspiration and respiratory rhythm
3. Describe the functions of the dorsal and respiratory groups of neurons
4. Describe the influences of the Pneumotaxic and Apneustic centers
5. Explain the role of Hering-Breur inflation reflex in regulation of inspiration
6. Explain the role of central chemoreceptors in regulation of respiration
7. Explain the role of peripheral chemoreceptors in regulation of respiration
8. Explain the regulation of respiration during exercise
9. Integrate the respiratory regulatory mechanisms
10. Describe the Cheyne-Stokes breathing
Study Resources:
1. Chapter 42, Guyton and Hall Textbook of Medical Physiology, 14th edition
2. Chapter 36, Ganong’s Review of Medical Physiology, 26th edition
3. Chapter 13, Human Physiology by Lauralee Sherwood, 9th edition
Osteoporosis - Definition , Evaluation and Management .pdfJim Jacob Roy
Osteoporosis is an increasing cause of morbidity among the elderly.
In this document , a brief outline of osteoporosis is given , including the risk factors of osteoporosis fractures , the indications for testing bone mineral density and the management of osteoporosis
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These lecture slides, by Dr Sidra Arshad, offer a quick overview of the physiological basis of a normal electrocardiogram.
Learning objectives:
1. Define an electrocardiogram (ECG) and electrocardiography
2. Describe how dipoles generated by the heart produce the waveforms of the ECG
3. Describe the components of a normal electrocardiogram of a typical bipolar lead (limb II)
4. Differentiate between intervals and segments
5. Enlist some common indications for obtaining an ECG
6. Describe the flow of current around the heart during the cardiac cycle
7. Discuss the placement and polarity of the leads of electrocardiograph
8. Describe the normal electrocardiograms recorded from the limb leads and explain the physiological basis of the different records that are obtained
9. Define mean electrical vector (axis) of the heart and give the normal range
10. Define the mean QRS vector
11. Describe the axes of leads (hexagonal reference system)
12. Comprehend the vectorial analysis of the normal ECG
13. Determine the mean electrical axis of the ventricular QRS and appreciate the mean axis deviation
14. Explain the concepts of current of injury, J point, and their significance
Study Resources:
1. Chapter 11, Guyton and Hall Textbook of Medical Physiology, 14th edition
2. Chapter 9, Human Physiology - From Cells to Systems, Lauralee Sherwood, 9th edition
3. Chapter 29, Ganong’s Review of Medical Physiology, 26th edition
4. Electrocardiogram, StatPearls - https://www.ncbi.nlm.nih.gov/books/NBK549803/
5. ECG in Medical Practice by ABM Abdullah, 4th edition
6. Chapter 3, Cardiology Explained, https://www.ncbi.nlm.nih.gov/books/NBK2214/
7. ECG Basics, http://www.nataliescasebook.com/tag/e-c-g-basics
Promoting Wellbeing - Applied Social Psychology - Psychology SuperNotesPsychoTech Services
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Muktapishti is a traditional Ayurvedic preparation made from Shoditha Mukta (Purified Pearl), is believed to help regulate thyroid function and reduce symptoms of hyperthyroidism due to its cooling and balancing properties. Clinical evidence on its efficacy remains limited, necessitating further research to validate its therapeutic benefits.
share - Lions, tigers, AI and health misinformation, oh my!.pptxTina Purnat
• Pitfalls and pivots needed to use AI effectively in public health
• Evidence-based strategies to address health misinformation effectively
• Building trust with communities online and offline
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3. The reduction of the degree of dissociation of a salt by the
addition of a common-ion is called the common-ion effect.
Example
In a saturated solution of silver chloride, we have the
equilibrium
AgCl (s) Ag+
(aq) + Cl–
(aq)
When sodium chloride is added to the solution, the
concentration of Cl– ions will increase. The equilibrium
shown above will be shifted to the left to form more of
solid Ag Cl. Thus the solubility of AgCl, a typical sparingly
soluble salt, will decrease.
Common ion effect
4. Factors effecting degree of dissociation
Nature of solute
Nature of solvent
Concentration
Temperature
common ion effect
5. Example
Dissociation of hydrogen sulphide in presence of
hydrochloric acid
H2S 2H+ + S2-
By applying the law of mass action, we have
Ka = [H+]2 [S2-]/ [H2S]
To the above solution of H2S , if we add hydrochloric acid,
then it ionizes completely as
HCl H+ + Cl-
common ion effect
6. Solubility equilibria and solubility product
When an ionic solid substance dissolves in water, it
dissociates to give separate cations and anions. For
example, for a sparingly soluble salt, say Ag Cl, we can
write the equilibrium equations as follows :
AgCl(S) Ag+ + Cl-
common ion effect
7. Solubility equilibria and solubility product
According to law of mass action
K= [Ag+] [Cl–] / [AgCl]
The amount of Ag Cl in contact
with saturated solution does
not change with time and the
factor [Ag Cl] remains the same.
common ion effect
8. Solubility equilibria and solubility product
As [AgCl] is constant so
equilibrium expression is
Ksp= [Ag+] [Cl–]
where
[Ag+] and [Cl– ] are expressed in mol/L
The equilibrium constant in the new context is called the solubility
product constant (or simply the solubility product) and is denoted by
Ksp.
common ion effect
9. Solubility of barium iodate in presence of barium
nitrate
Barium iodate, Ba(IO3)2, has a solubility product
Ksp = [Ba2+][IO3
-]2 = 1.57 x 10-9
Its solubility in pure water is 7.32 x 10-4 M.
common ion effect
12. Activity
Calculate the solubility of silver chromate, Ag2CrO4, in a 0.100 M solution of
AgNO3. (Ksp for Ag2CrO4 = 9.0 × 10– 12)
Activity
Calculate the Ksp for Bismuth sulphide (Bi2S3), which has a solubility of 1.0 ×
10– 15 mol/L at 25°C.
Activity
The solubility of BaSO4 is 2.33 × 10– 4 g/ml at 20°C. Calculate the solubility
product of BaSO4 assuming that the salt is completely ionised.
common ion effect
13. Common ion effect on solubility
Adding a common ion decreases solubility, as the reaction
shifts toward the left to relieve the stress of the excess
product. Adding a common ion to a dissociation reaction
causes the equilibrium to shift left, toward the reactants,
causing precipitation.
Applications of common ion effect
14. Common ion effect on solubility
Example
AgCl(s) Ag+ + Cl–
If S be the solubility of AgCl, we have
Ksp = [S mol/l Ag+] [S mol/l Cl–]
Suppose 0.25 mol/L excess of HCl is added to the solution.
Then ion product (Q) will be
Q = [S mol/l Ag+] [(S + 0.25) mol/l Cl– ]
If Q > Ksp Precipitation
If Q = Ksp Saturated solution
If Q < Ksp No precipitation
Applications of common ion effect
15. Salting out of soap
RCOONa(aq.) RCOO-
(aq.) + Na+
(aq.)
NaCl(aq.) Na+
(aq.) + Cl-
(aq.)
Applications of common ion effect
17. Lime Softening
CO2 + Ca(OH)2 CaCO3 + H2O
Ca(HCO3)2 + Ca(OH)2 2CaCO3 + 2H2O
Mg(HCO3)2 + 2Ca(OH)2 2CaCO3 + 2H2O
Applications of common ion effect
18. Harris, B. C. and W.H. Harris. 2010. Quantitative
Chemical Analysis. 8th Edition. Freeman and Company
New York.
Christian, G.D. 2006. Analytical Chemistry. Sixth edition,
John Wiley and Sons, New York.
Skoog, D. A. and D.M. West 2005. Fundamentals of
Analytical Chemistry. Hot Reinehart Inc., London.
Kealey, D and P.J.Haines, 2002. Analytical Chemistry,
Bios Scientific Publishers Limited, Oxford, UK.
Sharma, B. K. 2004. Instrumental methods of chemical
analysis, In; Introduction to Analytical chemistry: Goel
Publishing House Meerut, 23th Edition.
Reilley, C. 1993. Laboratory Manual of Analytical
Chemistry. Allyn & Bacon, London.
Reference books
Editor's Notes
When a soluble salt (say A+C–) is added to a solution of another salt (A+B–) containing a common ion (A+), the dissociation of AB is suppressed. AB into A+ and B– . By the addition of the salt (AC), the concentration of A+ increases. Therefore, according to Le Chatelier’s principle, the equilibrium will shift to the left, thereby decreasing the concentration of A+ ions. Or that, the degree of dissociation of AB will be reduced. When solid NH4 Cl is added to NH4 OH solution, the equilibrium NH4 OH ----------NH4+ OH− shifts to the left. Thereby the equilibrium concentration of OH– decreases. This procedure of reducing the concentration of OH– ions is used in qualitative analysis
Nature of Solute The nature of solute is the chief factor which determines its degree of dissociation in solution. Strong acids and strong bases, and the salts obtained by their interaction are almost completely dissociated in solution. On the other hand, weak acids and weak bases and their salts are feebly dissociated. (2) Nature of the solvent The nature of the solvent affects dissociation to a marked degree. It weakens the electrostatic forces of attraction between the two ions and separates them. This effect of the solvent is measured by its ‘dielectric constant’. The dielectric constant of a solvent may be defined as its capacity to weaken the force of attraction between the electrical charges immersed in that solvent. The dielectric constant of any solvent is evaluated considering that of vaccum as unity. It is 4.1 in case of ether, 25 in case of ethyl alcohol and 80 in case of water. The higher the value of the dielectric constant the greater is the dissociation of the electrolyte dissolved in it because the electrostatic forces vary inversely as the dielectric constant of the medium. Water, which has a high value of dielectric constant is, therefore, a strong dissociating solvent. The electrostatic forces of attraction between the ions are considerably weakened when electrolytes are dissolved in it and as a result, the ions begin to move freely and there is an increase in the conductance of the solution. (3) Concentration The extent of dissociation of an electrolyte is inversely proportional to the concentration of its solution. The less concentrated the solution, the greater will be the dissociation of the electrolyte. This is obviously due to the fact that in a dilute solution the ratio of solvent molecules to the solute molecules is large and the greater number of solvent molecules will separate more molecules of the solute into ions. (4) Temperature The dissociation of an electrolyte in solution also depends on temperature. The higher the temperature greater is the dissociation. At high temperature the increased molecular velocities overcome the forces of attraction between the ions and consequently the dissociation is great.
Hydrogen sulphide (H2S) is a weak electrolyte. It is weakly ionized in its aqueous solution. There exists an equilibrium between unionized molecules and the ions in an aqueous medium as follows:
H2S = 2H+ + S2-
To the above solution of H2S , if we add hydrochloric acid, then it ionizes completely as
HCl = H+ + Cl-
This makes H+ a common ion and creates a common ion effect. Due to the increase in concentration of H+ ions, the equilibrium of dissociation of H2S shifts to the left and keeps the value of Ka constant. Thus the ionization of H2S is decreased. The concentration of unionized H2S is increased. As a result, the concentration of sulphide ions is decreased.
At equilibrium the solute continues to dissolve at a rate that exactly matches the reverse process, the return of solute from the solution. Now the solution is said to be saturated.
A Saturated solution is a solution in which the dissolved and undissolved solute are in equilibrium. A saturated solution represents the limit of a solute’s ability to dissolve in a given solvent. This is a measure of the “solubility” of the solute.
The Solubility (S) of a substance in a solvent is the concentration in the saturated solution. Solubility of a solute may be represented in grams per 100 ml of solution. It can also be expressed in moles per litre. Molar Solubility is defined as the number of moles of the substance per one litre (l) of the solution. The value of solubility of a substance depends on the solvent and the temperature
The value of K sp for a particular solubility equilibrium is constant at a given temperature. The product [Ag+] [Cl–] in the K sp expression above is also called the Ionic Product or Ion Product. The K sp expression may be stated as : the product of the concentration of ions (mol/l) in the saturated solution at a given temperature is constant. This is sometimes called the Solubility product principle.
However in a solution that is 0.0200 M in barium nitrate, Ba(NO3)2, the increase in the common ion barium leads to a decrease in iodate ion concentration. The solubility is therefore reduced to 1.40 x 10-4 M, about five times smaller.
.
Adding a common ion prevents the weak acid or weak base from ionizing as much as it would without the added common ion. The common ion effect suppresses the ionization of a weak acid by adding more of an ion that is a product of this equilibrium.e.g.t he common ion effect of H3O+ on the ionization of acetic acid is shown above in figure. The common ion effect suppresses the ionization of a weak base by adding more of an ion that is a product of this equilibrium.
Consider the common ion effect of OH- on the ionization of ammonia as shown in above figure. dding the common ion of hydroxide shifts the reaction towards the left to decrease the stress (in accordance with Le Châtelier's Principle), forming more reactants. This decreases the reaction quotient, because the reaction is being pushed towards the left to reach equilibrium. The equilibrium constant, Kb=1.8*10-5, does not change. The reaction is put out of balance, or equilibrium.
Qa=[NH+4][OH−][NH3](23)(23)Qa=[NH4+][OH−][NH3]
At first, when more hydroxide is added, the quotient is greater than the equilibrium constant. The reaction then shifts right, causing the denominator to increase, decreasing the reaction quotient and pulling towards equilibrium and causing Q to decrease towards K.
An example of such an effect can be observed when acetic acid and sodium acetate are both dissolved in a given solution, generating acetate ions. However, sodium acetate completely dissociates but the acetic acid only partly ionizes. This is because acetic acid is a weak acid whereas sodium acetate is a strong electrolyte.
As per Le Chatelier’s principle, the new acetate ions put forth by sodium acetate facilitate the suppression of the ionization of acetic acid, thereby shifting the equilibrium to the left. Since the dissociation of acetic acid is reduced, the pH of the solution is increased.
Therefore, the common ion solution containing acetic acid and sodium acetate will have an increased pH and will, therefore, be less acidic when compared to an acetic acid solution.
Thus, the common ion effect, its effect on the solubility of a salt in a solution, and its effect on the pH of a solution are discussed in this article
https://byjus.com/chemistry/common-ion-effect/
A practical example used very widely in areas drawing drinking water from chalk or limestone aquifers is the addition of sodium carbonate to the raw water to reduce the hardness of the water. In the water treatment process, highly soluble sodium carbonate salt is added to precipitate out sparingly soluble calcium carbonate. The very pure and finely divided precipitate of calcium carbonate that is generated is a valuable by-product used in the manufacture of toothpaste.
The salting-out process used in the manufacture of soaps benefits from the common-ion effect. Soaps are sodium salts of fatty acids. Addition of sodi
The solubility product (K sp) of an insoluble substance is the product of the concentrations of its ions at equilibrium. However, the ion product is the product of actual concentrations of ions which may or may not be in equilibrium with the solid. The increase in concentration of Cl– will shift the equilibrium to the left to form a precipitate of AgCl.
Essential of physical chemistry By BS Bhal
Since soaps are the sodium salts of carboxylic acids containing a long aliphatic chain (fatty acids), the common ion effect can be observed in the salting-out process which is used in the manufacturing of soaps. The soaps are precipitated out by adding sodium chloride to the soap solution in order to reduce its solubility.
The important applications of Solubility product is Salting out of soap. Soap is sodium salt of higher fatty acid . It is precipitated from the solution by adding concentrated solution of NaCl( sodium chloride) . NaCl is strong electrolyte , it ionises completely in solution then concentration of sodium ions (Na+) increases. Due to common ion effect dissociation of soap decreases and soap gets precipitated and easily removed from soap solution.This is processs of getting solid soap from soap solution by adding salt like NaCl and called salting out of soap.
The common ion effect is the phenomenon in which the addition of an ion common to two solutes causes precipitation or reduces ionization. You know Le Chatelier's principle, which states that a reaction stays in equilibrium unless acted on by an outside force; then the reaction will shift to accommodate the force and re-establish equilibrium. This means that if you want a reaction to shift, you just apply an outside force, such as temperature, concentration or pressure. In the case of the common ion effect, a reaction can be shifted by adding an ion that is common to both solutes, thus changing the concentration of the ion in solution and shifting the equilibrium of the reaction. The net effect of the common ion is that it reduces the solubility of the solute in the solution. The common ion effect can make insoluble substances more insoluble. An example of the common ion effect is when sodium chloride (NaCl) is added to a solution of HCl and water. The hydrochloric acid and water are in equilibrium, with the products being H3O+ and Cl- . Then, some sodium chloride is added to the solution. The NaCl dissolves into the solution, forming Na+ and Cl-. As the NaCl dissolves, the concentration of Cl- ions increases. The system accommodates by combining the Na+ and Cl- back into NaCl, which is a solid and precipitates out of the solution. In effect, more of the Cl- ions, which are common to both of the reactions, were added to the solution in equilibrium, so the equilibrium shifted back to the left.
http://www.csun.edu/~ml727939/coursework/695/common%20ion%20effect/recrystallization%20of%20NaCl%20solution%20with%20HCl.htm
As lime in the form of limewater is added to raw water, the pH is raised and the equilibrium of carbonate species in the water is shifted. Dissolved carbon dioxide (CO2) is changed into bicarbonate (HCO−3) and then carbonate (CO2-3). This action causes calcium carbonate to precipitate due to exceeding the solubility product. Additionally, magnesium can be precipitated as magnesium hydroxide in a double displacement reaction.[4]
In the process both the calcium (and to an extent magnesium) in the raw water as well as the calcium added with the lime are precipitated. This is in contrast to ion exchange softening where sodium is exchanged for calcium and magnesium ions. In lime softening, there is a substantial reduction in total dissolved solids (TDS) whereas in ion exchange softening (sometimes referred to as zeolite softening), there is no significant change in the level of TDS.
Lime softening can also be used to remove iron, manganese, radium and arsenic from water.
https://en.wikipedia.org/wiki/Lime_softening
https://www.suezwatertechnologies.com/handbook/chapter-07-precipitation-softening