Name: common ion effects on solubility
Uploaded: 2020-09-24T15:06:51.000Z
Duration: 4 min 1 s
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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

이온 화합물의 용해도는 용액에 존재하는 다른 용질에 따라 달라질 수 있습니다. 용질은 화합물 또는 산이나 염기에 공통으로 존재하는 이온일 수 있습니다. 따라서 용해도에 영향을 미치는 두 가지 주요 요인은 공통 이온 효과와 용액의 pH입니다.

염화(II)납을 염화 나트륨 용액에 첨가하면 두 염은 모두 물 속에서 이온화되어 나트륨과 납 양이온 그리고 공통 음이온인 염화물이 생성됩니다. 용액 속의 염화 이온은 염화나트륨의 완전한 이온화와 염화납(II)의 부분적인 이온화에서 모두 나오기 때문에 용액의 농도는 나트륨 이온이나 납 이온의 농도보다 높을 것입니다. 이를 보완하기 위해 용액 속의 고체 염화납(II)과 그것의 이온 사이의 평형은 해리되지 않은 염 쪽으로 이동하며 더 많은 염화 납(II)이 분리되지 않은 상태로 남게 합니다.

따라서 공통 이온의 존재는 용해도가 낮은 물질의 용해도를 감소시킵니다. 예를 들어 0.100 몰의 염화나트륨 용액에서 염화 납(II)의 몰 용해도 x는 ICE 표를 통해 계산할 수 있습니다. 용액 내 납 이온의 초기 농도는 0이지만 염화물의 농도는 0.100 몰입니다.

염화 납(II)의 각 분자는 하나의 납과 두 개의 염화 이온으로 분리됩니다. 따라서 납 이온에서 농도의 변화는 x가 될 것이고 염화물에서는 2x가 될 것입니다. 납 이온의 평형 농도는 x이지만 염화 이온의 경우 2x와 0.100의 합이 됩니다.

25°C에서 염화 납(II)에 대한 Ksp의 값은 1.17 10⁻⁵이고 평형식은 납과 염화 이온의 평형 농도의 적으로 의 두제곱에 x를 곱한 것과 같습니다. Ksp가 작기 때문에 2x는 0.100 몰보다 훨씬 작고 그것의 합계는 0.100 몰에 근사합니다. 식에 대입하면 1.17 10⁻⁵는 x에 0.100 몰 두제곱을 곱한 것과 같습니다.

x를 풀면 0.100 몰의 염화나트륨용액에서 염화 납(II)의 몰용해도는 1.17 10⁻³ 몰입니다. 물에 함유된 염화납(II)의 몰용해도가 1.43 10⁻² 몰인 것과 비교하면 공통 이온은 염화납(II)의 용해도를 12배 감소시킵니다. 화합물의 용해도는 용액의 pH에도 영향을 받을 수 있습니다.

수산화칼슘을 칼슘 이온과 수산화물로 부분적으로 해리되는 반응을 고려해봅시다. 예를 들어 수산화 칼륨을 첨가하여 pH가 증가한다면 일반적인 수산화 이온은 수산화칼슘의 용해도를 낮출 것입니다. 반대로 예를 들어 염산을 첨가하여 pH가 감소하면 양성자는 수산화 이온과 결합하여 용액 내 수산화 이온의 농도를 낮춥니다.

그런 다음 수산화칼슘의 용해도 평형은 생성물 방향으로 이동하면서 산성 용액의 용해도를 증가시킬 것입니다.

<ol>
	<li>Koubek, E. "공통 이온 효과의 실증." 화학 교육 저널 70, no. 2 (1993) : 155.</li>
	<li>아마랄, L.F., I.R. 올리베이라, R. 살로망, E. 프롤리니, V. C. 판돌펠리."산화마그네슘(MgO) 수화에 대한 온도 및 공통 이온 효과." 세라믹 인터내셔널 36, no. 3 (2010): 1047-1054.</li>
	<li>Cassens, Jan, Anke Prudic, Feelly Ruether 및 Gabriele Sadowski. "pH의 함수로서의 의약품 및 그 염의 용해도." Industrial and Engineering Chemistry Research 52, no. 7 (2013): 2721-2731.</li>
</ol>

용해도에 영향을 미치는 요인

순수한 물과 비교하여, 이온 화합물의 용해도는 공통 이온(이온 화합물의 용해에 의해 생성되는 이온)을 포함하는 수용액에서 더 적습니다. 이는 르 샤틀리에의 원리를 사용하여 설명할 수 있는 대량 작용 법칙의 결과인 공통 이온 효과로 알려진 현상의 한 예입니다. 요오드화은의 용해에 대해 생각해 보세요.

<img alt="Eq1" src="/files/ftp_upload/11418/11418_Equation_1.png" style="text-align: center; height: 50px;" />

이러한 용해도 평형은 은이나 요오드 이온을 첨가하면 왼쪽으로 이동하여 AgI가 침전되고 용해된 Ag^+ 및 I^-의 농도가 낮아질 수 있습니다. 이러한 이온 중 하나를 이미 포함하고 있는 용액에서는 이러한 이온이 없는 용액보다 더 적은 양의 AgI가 용해될 수 있습니다.

이 효과는 용해도 곱 표현으로 표현되는 대량 작용의 관점에서 설명될 수도 있습니다.

<img alt="Eq2" src="/files/ftp_upload/11418/11418_Equation_2.png" style="text-align: center; height: 50px;" />

은과 요오드 이온 몰농도의 수학적 곱은 이온 소스에 관계없이 평형 혼합물에서 일정하므로 한 이온 농도의 증가는 다른 이온 농도의 비례적인 감소와 균형을 이루어야 합니다.

폐수 처리에서 침전강수의 역할

용해도 평형은 도시나 마을의 생활용수를 처리할 수 있는 시설에서 수행되는 폐수 처리에 유용한 도구입니다. 특히, 선택적 침전은 폐수가 자연 수역으로 다시 방출되기 전에 폐수에서 오염 물질을 제거하는 데 사용됩니다. 예를 들어, 제조 시설에서 배출되는 물에는 인산염 이온(PO_4^3−)이 종종 존재합니다. 풍부한 인산염으로 인해 과도한 조류가 자라게 되고, 이는 해양 생물이 이용할 수 있는 산소의 양에 영향을 미칠 뿐만 아니라 물을 인간이 섭취하기에 부적합하게 만듭니다.

물에서 인산염을 제거하는 일반적인 방법 중 하나는 수산화칼슘 또는 석회인 Ca(OH)_2를 첨가하는 것입니다. 물이 더욱 염기성으로 변하면서 칼슘 이온은 인산염 이온과 반응하여 수산화인회석(Ca_5(PO_4)_3·OH)을 생성하고 용액에서 침전됩니다.

<img alt="Eq3" src="/files/ftp_upload/11418/11418_Equation_3.png" style="text-align: center; height: 50px;" />

첨가된 칼슘 이온의 양은 다른 칼슘 염의 용해도 곱을 초과하지 않기 때문에 해당 염의 음이온은 폐수에 남아 있습니다. 그런 다음 침전물은 여과를 통해 제거되고, 재탄산화 과정에서 CO_2를 첨가하여 물의 pH를 중성으로 되돌립니다. 침전을 통해 인산염을 제거하는 데 염화철(III)과 황산알루미늄 등 다른 화학물질도 사용할 수 있습니다.

 이 문서는 에서 발췌되었습니다  <a href="https://openstax.org/books/chemistry-2e/pages/15-1-precipitation-and-dissolution">Openstax, Chemistry 2e, Section 15.1: Precipitation and Dissolution.</a>

순수한 물과 비교할 때, 이온 성 화합물의 용해도는 공통 이온을 포함하는 수용액에서 더 적습니다 (또한 의 용해에 의해 생성 된 것 ...

Education

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화학

하나

Acid-base and Solubility Equilibria

산-염기 및 용해도 평형

KEY TERMS AND CONCEPTS

주요 용어 및 개념

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

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actors Affecting Dissolution Rates

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Arrhenius Plots

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Thermochemical Equations

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Rate-Determining Steps

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Constant Pressure Calorimetry

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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

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The Equilibrium Constant

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Enthalpy

Chemists ordinarily use a property known as enthalpy (H) to describe the thermodynamics of chemical and physical processes. Enthalpy is defined as the sum ...

Calculating the Equilibrium Constant

The equilibrium constant for a reaction is calculated from the equilibrium concentrations (or pressures) of its reactants and products. If these ...

Reaction Quotient

The status of a reversible reaction is conveniently assessed by evaluating its reaction quotient (Q). For a reversible reaction described by m A + n B ...

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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

Being able to calculate equilibrium concentrations is essential to many areas of science and technology&#8212;for example, in the formulation and dosing ...

Strong Acid and Base Solutions

A strong acid is a compound that dissociates completely in an aqueous solution and produces a concentration of hydronium ions equal to the initial ...

The Small x Assumption

If a reaction has a small equilibrium constant, the equilibrium position favors the reactants. In such reactions, a negligible change in concentration may ...

Weak Acid Solutions

Few compounds act as strong acids. A far greater number of compounds behave as weak acids and only partially react with water, leaving a large majority of ...

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Le Chatelier's Principle: Changing Concentration

A system at equilibrium is in a state of dynamic balance, with forward and reverse reactions taking place at equal rates. If an equilibrium system is ...

Weak Base Solutions

Some compounds produce hydroxide ions when dissolved by chemically reacting with water molecules. In all cases, these compounds react only partially and ...

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Le Chatelier's Principle: Changing Volume (Pressure)

For gas-phase equilibria, changes in the concentrations of reactants and products can occur with altered volume and pressure. The partial pressure, P, of ...

Mixtures of Acids

The pH of a solution containing an acid can be determined using its acid dissociation constant and its initial concentration. If a solution contains two ...

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Le Chatelier's Principle: Changing Temperature

Consistent with the law of mass action, an equilibrium stressed by a change in concentration will shift to re-establish equilibrium without any change in ...

Ions as Acids and Bases

Salts with Acidic Ions
Salts are ionic compounds composed of cations and anions, either of which may be capable of undergoing an acid or base ionization ...

Bronsted-Lowry Acids and Bases

The acid-base reaction class has been studied for quite some time. In 1680, Robert Boyle reported traits of acid solutions that included their ability to ...

Determining the pH of Salt Solutions

The pH of a salt solution is determined by its component anions and cations. Salts that contain pH-neutral anions and the hydronium ion-producing cations ...

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Acid/Base Strengths and Dissociation Constants

The relative strength of an acid or base is the extent to which it ionizes when dissolved in water. If the ionization reaction is essentially complete, ...

Polyprotic Acids

Acids are classified by the number of protons per molecule that they can give up in a reaction. Acids such as HCl, HNO3, and HCN that contain one ...

Water: A Bronsted-Lowry Acid and Base

The reaction between a Br&#248;nsted-Lowry acid and water is called acid ionization. For example, when hydrogen fluoride dissolves in water and ionizes, ...

Acid Strength and Molecular Structure

Binary Acids and Bases
In the absence of any leveling effect, the acid strength of binary compounds of hydrogen with nonmetals (A) increases as the H-A ...

pH Scale

Hydronium and hydroxide ions are present both in pure water and in all aqueous solutions, and their concentrations are inversely proportional as ...

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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