Name: common ion effects on solubility
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K12 Standards

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Ohio’s Learning Standards and Model Curriculum

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Chemistry

Interactions Of Matter

Chemical reactions

Factors Affecting Solubility

İyonik bir bileşiğin çözünürlüğü, çözelti içinde bulunan diğer çözünen maddelere bağlı olarak değişebilir. Çözünen madde, bileşikte ortak bir iyon veya bir asit veya baz olabilir. Bu nedenle çözünürlüğü etkileyen iki ana faktör, ortak iyon etkisi ve bir çözeltinin pH'ıdır.Bir sodyum klorür çözeltisine kurşun klorür eklenirse, her iki tuz da suda iyonlaşarak sodyum ve kurşun katyonları ve ortak anyon olan klorür üretir. Çözeltideki klorür iyonları, hem sodyum klorürün tam iyonlaşmasından hem de kurşun klorürün kısmi iyonlaşmasından geldiğinden, konsantrasyonları her ikisinden de daha yüksek olacaktır. Telafi etmek için, katı kurşun klorür ile çözelti içindeki iyonları arasındaki denge, çözülmemiş tuza doğru kayarak daha fazla kurşun klorürün çözülmeden kalmasına neden olur.Dolayısıyla, ortak bir iyonun varlığı, eser miktarda çözünür bir maddenin çözünürlüğünü azaltır. Örneğin, 0, 100 molar sodyum klorür çözeltisi içindeki kurşun klorürün molar çözünürlüğü x, ICE tablosundan hesaplanabilir. Çözeltideki kurşun iyonlarının başlangıç konsantrasyonu sıfırdır, klorür ise 0, 100 molar'dır.Her bir kurşun klorür molekülü, bir kurşun ve iki klorür iyonu halinde ayrışacaktır. Yani, kurşun iyonlarının konsantrasyonundaki değişim x, klorürünki ise 2x olacaktır. Kurşun iyonlarının denge konsantrasyonu x olacaktır, ancak klorür iyonları için 2x ve 0, 100 toplamı olacaktır.25°C'de, kurşun klorür için Ksp'nin değeri 1, 17×10⁻⁵'dir ve denge ifadesi, kurşun ve klorür iyonlarının denge konsantrasyonlarının çarpımıdır;bu da x çarpı karesi olur. Ksp küçük olduğundan, 2x 0, 100 molardan çok daha küçük olacaktır ve böylece toplamları 0, 100 molara yuvarlanabilir. Değerler yerine koyulunca, 1, 17×10⁻⁵ eşittir x çarpı 0, 100 molar kare eşitliği elde edilir.Denklem x için çözüldüğünde, kurşun klorürün 0, 100 molar sodyum klorür çözeltisindeki çözünürlüğü 1, 17x10⁻³ molar şeklinde hesaplanır. Kurşun klorürün suda 1, 43×10⁻² molar çözünürlüğü ile karşılaştırıldığında, ortak iyon kurşun klorürün çözünürlüğünü 12 kat azaltır. Bir bileşiğin çözünürlüğü, çözeltinin pH'ından da etkilenebilir.Kalsiyum hidroksitin kalsiyum iyonlarına ve hidroksite kısmi ayrışmasını düşünün. Örneğin potasyum hidroksit ilave edilerek pH yükseltilirse, ortak hidroksit iyonu kalsiyum hidroksitin çözünürlüğünü düşürecektir. Tersine, örneğin hidroklorik asit ekleyerek pH düşürülürse, protonlar hidroksit iyonları ile birleşerek çözeltideki hidroksit iyonlarının konsantrasyonunu düşürür.Kalsiyum hidroksitin çözünürlük dengesi daha sonra ürünlere doğru kayacak ve asidik bir çözelti içinde çözünürlüğün artmasına yol açacaktır.

<ol>
	<li>Koubek, E. "Ortak İyon Etkisinin Gösterilmesi." Kimya eğitimi dergisi 70, no. 2 (1993): 155.</li>
	<li>Amaral, L. F., I. R. Oliveira, R. Salomão, E. Frollini ve V. C. Pandolfelli. "Magnezyum oksit (MgO) hidrasyonu üzerindeki sıcaklık ve ortak iyon etkisi." Seramik Uluslararası 36, hayır. 3 (2010): 1047-1054.</li>
	<li>Cassens, Jan, Anke Prudic, Feelly Ruether ve Gabriele Sadowski. "PH'ın bir fonksiyonu olarak farmasötiklerin ve tuzlarının çözünürlüğü." Endüstri ve Mühendislik Kimyası Araştırma 52, no. 7 (2013): 2721-2731.</li>
</ol>

Çözünürlüğü etkileyen faktörler

Saf su ile karşılaştırıldığında, bir iyonik bileşiğin çözünürlüğü, ortak bir iyon içeren sulu çözeltilerde daha azdır (biri ayrıca iyonik bileşiğin çözünmesiyle üretilir). Bu, Le Chȃtelier ilkesi kullanılarak açıklanabilecek kütle eylem yasasının bir sonucu olan ve ortak iyon etkisi olarak bilinen bir fenomen örneğidir. Gümüş iyodürün çözünmesini düşünün:
 
 <img alt=""Eq1"" src=""/files/ftp_upload/11418/11418_Equation_1.png"" style=""text-align: center; height: 50px;"" />
 
 Bu çözünürlük dengesi, gümüş veya iyodür iyonlarının eklenmesiyle sola kayabilir, bu da AgI'nin çökelmesine ve çözünmüş Ag+ ve I– konsantrasyonlarının düşmesine neden olur. Bu iyonlardan herhangi birini zaten içeren çözeltilerde, bu iyonların bulunmadığı çözeltilere göre daha az AgI çözünebilir.
 
 Bu etki, çözünürlük ürünü ifadesinde gösterildiği gibi kütle etkisi açısından da açıklanabilir:
 
 <img alt=""Eq2"" src=""/files/ftp_upload/11418/11418_Equation_2.png"" style=""text-align: center; height: 50px;"" />
 
 Gümüş ve iyodür iyon molaritelerinin matematiksel ürünü, iyonların kaynağına bakılmaksızın bir denge karışımında sabittir ve bu nedenle, bir iyon konsantrasyonundaki artış, diğerinde orantılı bir azalma ile dengelenmelidir.
 
 Atık Su Arıtımında Çökeltmeleri Rolü
 
 Çözünürlük dengesi, bir şehir veya kasabadaki belediye suyunu arıtabilen tesislerde gerçekleştirilen atık suyun arıtılmasında yararlı araçlardır. Özellikle, seçici çökeltme, kirletici maddeleri atık sudan doğal su kütlelerine geri salınmadan önce çıkarmak için kullanılır. Örneğin, üretim tesislerinden boşaltılan suda genellikle fosfat iyonları (PO43−) bulunur. Bol miktarda fosfat, fazla alg büyümesine neden olur, bu da deniz yaşamı için mevcut oksijen miktarını etkiler ve suyu insan tüketimi için uygunsuz hale getirir.
 
 Sudan fosfatları gidermenin yaygın bir yolu, kalsiyum hidroksit veya kireç (Ca(OH)2) eklemektir. Su daha bazik hale getirildikçe, kalsiyum iyonları fosfat iyonlarıyla reaksiyona girerek hidroksilapatit (Ca5(PO4)3·OH) üretir ve bu daha sonra çözeltiden çöker:
 
 <img alt=""Eq3"" src=""/files/ftp_upload/11418/11418_Equation_3.png"" style=""text-align: center; height: 50px;"" />
 
 Eklenen kalsiyum iyonu miktarı diğer kalsiyum tuzları için çözünürlük ürünlerinin aşılmasına neden olmadığından, bu tuzların anyonları atık suda geride kalır. Çökelti daha sonra süzülerek çıkarılır ve su, yeniden karbonlama işleminde CO2 ilave edilerek nötr bir pH'ya geri getirilir. Demir (III) klorür ve alüminyum sülfat dahil olmak üzere diğer kimyasallar da çökeltme yoluyla fosfatların giderilmesi için kullanılabilir.
 
  Bu metin bu kaynaktan uyarlanmıştır: <a href=""https://openstax.org/books/chemistry-2e/pages/15-1-precipitation-and-dissolution"">Openstax, Chemistry 2e, Section 15.1: Precipitation and Dissolution.</a>

Çözünürlüğü Etkileyen Faktörler

Saf su ile karşılaştırıldığında, iyonik bir bileşiğin çözünürlüğü, ortak bir iyon içeren sulu çözeltilerde daha azdır (biri aynı zamanda iyonun çözünmesiyle de üretilir ...

Education

JoVE Core

Kimya

Acid-base and Solubility Equilibria

Asit-Baz ve Çözünürlük Dengeleri

KEY TERMS AND CONCEPTS

TEMEL TERİMLER VE KAVRAMLAR

common ion effects on solubility

The common ion effect lowers the solubility of an ionic compound in water when the solution already contains one of the ions made by that compound. It follows the law of mass action and can be explained with Le Châtelier’s principle. When a common ion is added, the dissolution equilibrium shifts toward the solid, so less of the salt stays dissolved.

Silver iodide shows this effect clearly. If silver ions or iodide ions are added to a solution of AgI, the equilibrium shifts left and more AgI precipitates out. The result is lower concentrations of dissolved Ag+ and I− than in pure water.

The same idea appears in the solubility product expression, or Ksp. In an equilibrium mixture, the product of the silver ion and iodide ion molarities stays constant no matter where the ions came from. If one ion’s concentration rises, the other must fall to keep the equilibrium balance.

Solubility equilibria also help in wastewater treatment. Facilities often use selective precipitation to remove contaminants before water is released into rivers or lakes. One important target is phosphate ion, PO43−, because too much phosphate can cause excess algae growth, reduce oxygen for marine life, and make water unsafe for people.

A common treatment step is adding calcium hydroxide, or lime, Ca(OH)2. This makes the water more basic and allows calcium ions to react with phosphate ions to form hydroxylapatite, Ca5(PO4)3·OH, which precipitates out. The calcium added is kept low enough that other calcium salts do not exceed their solubility products, so their anions stay in the wastewater.

After filtration removes the solid, the water is brought back to a neutral pH by adding CO2 in a recarbonation step. Other chemicals can also remove phosphate by precipitation. Iron(III) chloride and aluminum sulfate are two examples.

Common Ion Effects on Solubility

common ion effect

Common Ion Effect on Solubility

The common ion effect reduces the solubility of an ionic compound in a solution that already contains one of its ions. In aqueous solutions, this means less of the solid dissolves than it would in pure water. The effect follows the law of mass action and can be explained with Le Châtelier’s principle. Silver iodide shows this equilibrium clearly. When silver(I) ions or iodide ions are added, the dissolution equilibrium shifts left. More AgI precipitates, and the concentrations of dissolved Ag+...

buffers

Buffer Solutions and pH Stability

Buffer solutions are mixtures that help keep pH steady. A buffer contains an appreciable amount of a weak conjugate acid-base pair. It resists a change in pH when small amounts of a strong acid or a strong base are added. One common buffer uses a weak acid and its salt. Acetic acid and sodium acetate form this type of system, written as CH3COOH(aq) + CH3COONa(aq). Another buffer uses a weak base and its salt. A solution of ammonia and ammonium chloride, NH3(aq) + NH4Cl(aq), is an example. The...

calculating ph changes in a buffer solution

calculating-ph-changes-in-a-buffer-solution

Buffer pH After Acid or Base Addition

A buffer solution can resist sudden pH changes when a strong acid or strong base is added. It does not stop the pH from changing completely. The shift is usually small, as long as the addition stays within the buffer’s capacity. To find the new pH, first determine how the added acid or base changes the amounts of the buffer components. A buffer contains a weak acid and its conjugate base, or the reverse pair. These amounts are found with stoichiometric calculations, which track how much of...

buffer effectiveness

Buffer Capacity and pH Stability

Buffer solutions help keep pH fairly constant, but they do not work forever. Their resistance to pH change depends on having enough of a weak acid and its conjugate base, or a weak base and its conjugate acid, in the mixture. When too much strong acid or strong base is added, one part of the buffer pair is used up and the buffer action becomes weaker. Buffer capacity describes how much acid or base can be added to a fixed volume of buffer before the pH changes noticeably, usually by about one...

titration calculations: strong acid - strong base

titration-calculations-strong-acid-strong-base

pH at Key Points in Acid-Base Titration

pH at key points in a strong acid and strong base titration can be calculated from the amounts of acid and base present. The example uses 25.00 mL of 0.100 M HCl, a strong acid, and 0.100 M NaOH, a strong base. The pH changes as base is added, and each stage follows a different calculation. At 0 mL of titrant, the pH comes from the ionization of HCl. Because HCl is a strong acid, its ionization is complete. The hydronium ion molarity is 0.100 M, and the pH is found from that value. At 12.50...

titration calculations: weak acid - strong base

titration-calculations-weak-acid-strong-base

pH at Key Points in a Weak Acid Titration

pH at key points in a weak acid titration can be calculated for an acetic acid solution as sodium hydroxide is added. The example uses 25.00 mL of 0.100 M CH3CO2H titrated with 0.100 M NaOH. The reaction is shown as the weak acid reacts with the strong base. The initial pH is found from the acetic acid solution using the usual ICE approach. ICE means initial, change, and equilibrium, and it is a common way to set up equilibrium problems. The solution pH is calculated before any titrant is...

indicators

Acid-Base Color Changes

Acid-base indicators show color changes that help identify the pH of a solution. These are often weak organic acids or weak organic bases, and they respond to changes in hydronium ion concentration in dilute aqueous solution. Phenolphthalein is one example. It is colorless when the hydronium ion concentration is greater than 5.0 × 10^-9 M, which means the pH is below 8.3. In more basic solutions, where the hydronium ion concentration is lower, it turns red or pink. Methyl orange shows how an...

titration of a polyprotic acid

Polyprotic Acid Equivalence Points

Polyprotic acids have more than one ionizable hydrogen, so they lose protons in steps during titration. When the acid dissociation constants for those protons are far enough apart, the titration curve shows a separate equivalence point for each hydrogen that can ionize. Diprotic acids have two ionizable hydrogens and produce two equivalence points. Triprotic acids have three ionizable hydrogens and produce three equivalence points. Carbonic acid, H2CO3, is a weak diprotic acid. Carbonic acid...

solubility equilibria

Predicting Precipitation with Ksp

Solubility equilibria explain how a dissolved ionic compound and its solid form can exist in balance. In a saturated solution, the solute dissolves and the ions precipitate at equal rates. These ideas matter in real settings such as tooth decay and water purification, where compound solubility affects the outcome. A substance with finite solubility can form a saturated solution when more solute is added than the solvent can dissolve. The result is a heterogeneous mixture of the saturated...

factors affecting solubility

The common ion effect lowers the solubility of an ionic compound in water when the solution already contains one of the ions made by that compound. It follows the law of mass action and can be explained with Le Châtelier’s principle. When a common ion is added, the dissolution equilibrium shifts toward the solid, so less of the salt stays dissolved. Silver iodide shows this effect clearly. If silver ions or iodide ions are added to a solution of AgI, the equilibrium shifts left and more AgI...

formation of complex ions

Ligands and Complex Ion Stability

Complex ions form when a central atom, often a transition metal cation, binds to ligands. Ligands are ions or molecules around the metal ion. They can be neutral molecules such as H2O or NH3, or ions such as CN− or OH−. In many cases, the ligands act as Lewis bases by donating a pair of electrons to the central atom. These reactions are part of coordination chemistry, which studies coordination complexes. When a metal ion reacts with one or more ligands to form a complex, the equilibrium...

precipitation of ions

Solubility Limits and Selective Precipitation

Solubility limits and selective precipitation explain when ions form a solid in water. The key idea is the balance between an insoluble compound and its dissolved ions. In aqueous solutions that contain calcium ions and carbonate ions, this equilibrium can exist even if the solution was not made by saturating water with calcium carbonate. When two soluble salts are mixed, such as sodium carbonate and calcium nitrate, the ion concentrations determine whether a precipitate forms. If the reaction...

qualitative analysis

Cation Identification by Precipitation Tests

Cation identification by precipitation tests uses selective reactions to find metal ions in a mixture. In qualitative analysis, different reagents are added step by step. Each reaction can produce a characteristic precipitate for a specific group of cations. Other steps, such as changing pH, heating the mixture to dissolve a precipitate again, or adding reagents that form complex ions, help separate ions within the same group. Group IV cations such as Ba2+, Ca2+, and Mg2+ are identified...

henderson-hasselbalch equation

Buffer pH with the Henderson-Hasselbalch Equation

The Henderson-Hasselbalch equation helps relate buffer pH to the weak acid in the solution. It comes from the ionization-constant expression for a weak acid. By rearranging that expression to solve for hydronium ion concentration, [H3O+], and then taking the negative logarithm of both sides, the equation can be written in terms of pH. In this form, pKa is defined as the negative logarithm of the weak acid ionization constant, or pKa = −log Ka. The equation shows the connection between pH, pKa,...

acid-base titration curves

Tracking pH During Acid-Base Titration

Acid-base titration uses pH changes to track how a solution responds as titrant is added. A titration curve is a graph of a solution property versus the amount of added titrant. In acid-base titrations, pH is the main property to follow because it changes in a predictable way with the solution makeup and helps show the endpoint. This pattern can be seen with a strong acid and strong base, a strong acid and weak base, or a strong base and weak acid. For example, 25.00 mL of 0.100 M HCl titrated...

key terms and concepts

Learn how the common ion effect changes solubility, why AgI precipitates, and how selective precipitation removes phosphate in wastewater treatment.

Common Ion Effects on Solubility | JoVE K12

chemistry

JoVE Core Kimya, özlü ve kolay anlaşılır animasyon video derslerini kullanarak genel kimya kavramlarına hayat veriyor. Ek olarak, bilim insanı uygulama videolarımız ile klasik ve orijinal araştırma deneylerini dünya çapında günümüz laboratuvarlarında görselleştiriyoruz.

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Physical Science

SPS6.

actors Affecting Dissolution Rates

Chemical Equations

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Reaction Stoichiometry

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Limiting Reactant

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Chemical Reactions in Aqueous Solutions

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Precipitation Reactions

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Oxidation&ndash;Reduction Reactions

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Oxidation Numbers

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Acids, Bases and Neutralization Reactions

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Measuring Reaction Rates

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Concentration and Rate Law

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Determining Order of Reaction

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The Integrated Rate Law: The Dependence of Concentration on Time

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Half-life of a Reaction

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The Collision Theory
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Arrhenius Plots

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Reaction Mechanisms

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Thermochemical Equations

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Rate-Determining Steps

Relating Reaction Mechanisms
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Constant Pressure Calorimetry

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Catalysis

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Constant Volume Calorimetry

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Enzymes

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Hess's Law

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Energy Basics

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Chemical changes and their accompanying ...

Standard Enthalpy of Formation

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First Law of Thermodynamics

Energy Conservation

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Enthalpies of Reaction

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Internal Energy

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Dynamic Equilibrium

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Quantifying Heat

Thermal Energy&#160;

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The Equilibrium Constant

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For a reaction that begins with a mixture of reactants only, the product concentration is initially equal to ...

Quantifying Work

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Homogeneous Equilibria for Gaseous Reactions

Homogeneous Equilibria for Gaseous Reactions
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Enthalpy

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Calculating the Equilibrium Constant

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Reaction Quotient

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Relative Strengths of Conjugate Acid-Base Pairs

Br&#248;nsted-Lowry acid-base chemistry is the transfer of protons; thus, logic suggests a relation between the relative strengths of conjugate acid-base ...

Calculating Equilibrium Concentrations

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Strong Acid and Base Solutions

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The Small x Assumption

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Weak Acid Solutions

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Le Chatelier's Principle: Changing Concentration

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Ions as Acids and Bases

Salts with Acidic Ions
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Bronsted-Lowry Acids and Bases

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Determining the pH of Salt Solutions

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Acid/Base Strengths and Dissociation Constants

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Polyprotic Acids

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Water: A Bronsted-Lowry Acid and Base

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Acid Strength and Molecular Structure

Binary Acids and Bases
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pH Scale

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Chemical reactions
• Types of reactions
• Kinetics
• Energy
• Equilibrium
• Acids/bases

Content elaboration: In the Physical Science course, coefficients were used to balance simple equations. Other representations, including Lewis structures and three-dimensional models, were also used and manipulated to demonstrate the conservation of matter in chemical reactions. In this course, more complex reactions will be studied, classified and represented with balanced chemical equations and three-dimensional models. Classifying reactions into types can be a helpful organizational tool for recognizing patterns of what may happen when two substances are mixed. Teachers should be aware that the common reaction classifications that are often used in high school chemistry courses may lead to misconceptions because they are not based on the actual chemistry, but on surface features that can be similar from one system to another (e.g., exchanging partners), even though the underlying chemistry is not the same. However, these classifications may be useful in making predictions about what happens when two substances are mixed. Some general types of chemical reactions are oxidation/reduction, synthesis, decomposition, single replacement, double replacement (including precipitation reactions and some acid-base neutralizations) and combustion reactions. Some reactions can fit into more than one category. For example, a single replacement reaction can also be classified as an oxidation/reduction reaction. Identification of reactions involving oxidation and reduction as well as indicating what substance is being oxidized and what is being reduced are appropriate in this course. However, balancing complex oxidation/reduction reactions is reserved for more advanced study. Organic molecules release energy when undergoing combustion reactions and are used to meet the energy needs of society (e.g., oil, gasoline, natural gas) and to provide the energy needs of biological organisms (e.g., cellular respiration). When a reaction between two ionic compounds in aqueous solution results in the formation of a precipitate or molecular compound, the reaction often occurs because the new ionic or covalent bonds are stronger than the original ion-dipole interactions of the ions in solution. Laboratory experiences (3-D or virtual) with different types of chemical reactions should be provided. Reactions occur when reacting particles collide in an appropriate orientation and with sufficient energy. The rate of a chemical reaction is the change in the amount of the reactants or products in a specific period of time. Increasing the probability or effectiveness of the collisions between the particles increases the rate of the reaction. Therefore, changing the concentration of the reactants, changing the temperature or the pressure of gaseous reactants, or using a catalyst, can change the reaction rate. Likewise, the collision theory can be applied to dissolving solids in a liquid solvent and can be used to explain why reactions are more likely to occur between reactants in the aqueous or gaseous state than between solids. The rate at which a substance dissolves should not be confused with the amount of solute that can dissolve in a given amount of solvent (solubility). Mathematical treatment of reaction rates is reserved for more advanced study. Computer simulations can help visualize reactions from the perspective of the kinetic-molecular theory. In middle school, the differences between potential and kinetic energy and the particle nature of thermal energy were introduced. For chemical systems, potential energy is in the form of chemical energy and kinetic energy is in the form of thermal energy. The total amount of chemical energy and/or thermal energy in a system is impossible to measure. However, the energy change of a system can be calculated from measurements (mass and change in temperature) from calorimetry experiments in the laboratory. Conservation of energy is an important component of calorimetry equations. Thermal energy is the energy of a system due to the movement of its particles. The thermal energy of an object depends upon the temperature and chemical composition. Some materials require little energy to change their temperature and other materials require a great deal to change their temperature by the same amount. Specific heat is a measure of how much energy is needed to change the temperature of a specific mass of material a specific amount. Specific heat values can be used to calculate the thermal energy change, the temperature (initial, final or change in) or mass of a material in calorimetry. Water has a particularly high specific heat capacity, which is important in regulating Earth's temperature. As studied in middle school, chemical energy is the potential energy associated with chemical systems. Chemical reactions involve valence electrons forming bonds to yield more stable products with lower energies. Energy is required to break interactions and bonds between the reactant atoms and energy is released when an interaction or bond is formed between the atoms in the products. Molecules with weak bonds (e.g., ATP) are less stable and tend to react to produce more stable products, releasing energy in the process. Generally, energy is transferred out of the system (exothermic) when the products have stronger bonds than the reactants and is transferred into the system (endothermic) when the reactants have stronger bonds than the products. Predictions of the energy requirements (endothermic or exothermic) of a reaction can be made given a table of bond energies. Graphic representations can be drawn and interpreted to represent the energy changes during a reaction. The role of energy in determining the spontaneity of chemical reactions is dealt with conceptually in this course. Entropy and its influence on the spontaneity of reactions are reserved for more advanced study. All reactions are reversible to a degree and many reactions do not proceed completely toward products but appear to stop progressing before the reactants are all used up. At this point, the amounts of the reactants and the products appear to be constant and the reaction can be said to have reached dynamic equilibrium. Dynamic equilibrium means the rate of the reverse reaction is equal to the rate of the forward reaction so there is no apparent change in the reaction. If a chemical system at equilibrium is disturbed by a change in the conditions of the system (e.g., increase or decrease in the temperature, pressure on gaseous equilibrium systems, concentration of a reactant or product), then the equilibrium system will respond by shifting to a new equilibrium state, reducing the effect of the change (Le Chatelier's Principle). If products are removed as they are formed during a reaction, then the equilibrium position of the system is forced to shift to favor the products. In this way, an otherwise unfavorable reaction can be made to occur. Mathematical treatment of equilibrium reactions is reserved for advanced study. Computer simulations can help visualize the progression of a reaction to dynamic equilibrium and the continuation of both the forward and reverse reactions after equilibrium has been attained. Properties of acids and bases and the ranges of the pH scale were introduced in Physical Science. In this course, the structural features of molecules are explored to further understand acids and bases. Acids often result when hydrogen is covalently bonded to an electronegative element and is easily dissociated from the rest of the molecule to bind with water to form a hydronium ion (H3O+). The acidity of an aqueous solution can be expressed as pH, where pH can be calculated from the concentration of the hydronium ion. Bases are likely to dissociate in water to form a hydroxide ion. Acids can react with bases to form a salt and water. Such neutralization reactions can be studied quantitatively by performing titration experiments.

Common Ion Effects on Solubility

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