Name: solubility limits and selective precipitation
Uploaded: 2020-09-24T15:06:51.000Z
Duration: 3 min 11 s
Channel: JoVE

K12 Standards

ohios-learning-standards-and-model-curriculum

Ohio’s Learning Standards and Model Curriculum

9-12

Chemistry

Interactions Of Matter

Chemical reactions

Precipitation of Ions

إذا مُزج محلولان من مركبان أيونيان،مثل يوديد الصوديوم ونترات الرصاص الثنائي،تستطيع الأيونات الموجبة من أحد المحلولين أن تتحد مع الأيونات السالبة من المحلول الآخر. كواحد من النواتج العابرة نترات الصوديوم-قابل للذوبان في الماء،تبقى أيونات الصوديوم والنترات على شكل محلول،بينما يُشكل الرصاص وأيونات اليوديد راسب يوديد الرصاص. حاصل التفاعل،Q،للتفاعل الانحلالي،يساوي منتج تركيزات أيونات الرصاص وأيونات اليوديد.وخلافًا لمنتج الذوبان Ksp،فإن حاصل التفاعل يشمل تركيز الأيونات في أي مرحلة،ولا يقتصر على مرحلة التوازن. يمكن مقارنة قيم Q و Ksp للتنبؤ فيما إذا كان سيحدث تفاعل ترسيب. لنتأمل إضافة قطرات من محلول يوديد الصوديوم إلى محلول نترات الرصاص.في البداية،تكون Q أقل من Ksp،وتكون أيونات الرصاص واليوديد عل شكل محلول مع أيونات الصوديوم وأيونات النترات. هذا محلول غير مشبع. بإضافة مزيد من يوديد الصوديوم،يستمر ارتفاع تركيز أيونات اليوديد.يكون التفاعل قد وصل التوازن عندما تتساوى Q مع Ksp. في هذه المرحلة،تكون كمية قليلة من يوديد الرصاص في حالة توازن ديناميكي مع الأيونات،مما يشكل محلولًا مشبعًا. إضافة المزيد من يوديد الصوديوم تجعل Q أكبر من Ksp،ويتحرك التفاعل باتجاه الراسب.هذا محلول فائق التشبع،يستمر فيه الترسيب إلى أن يُخفّض تركيز الأيونات إلى قيمتها عند حالة التوازن. على سبيل المثال،لنفترض أن خلط محلولي يوديد الصوديوم و نترات الرصاص الثنائي ينتج عنه محلول يحتوي على أيونات رصاص بتركيز 1.6 10⁻⁴ مولار،وأيونات يوديد بتركيز 4.0 10⁻⁴ مولار. هنا Q تساوي 2.6 10⁻¹¹،بينما تكون قيمة Ksp ليوديد الرصاص 1.4 10⁻⁸.لأن قيمة Q من قيمة Ksp،فإن يوديد الرصاص لن يترسب. التنبؤ بتفاعلات الترسيب،قد يكون مفيدًا للغاية أثناء فصل المركبات الأيونية. لنتأمل محلولًا يتكون من أيونين فلزيين الرصاص الثنائي،والنحاس الثنائي.إذا أضيف حمض الهيدروكلوريد إلى هذا المحلول،سيترسب كلوريد الرصاص الثنائي لأن قيمة Ksp فيه صغيرة،بينما يبقى النحاس على شكل محلول لأن كلوريد النحاس الثنائي يمتاز بدرجة ذائبية عالية في الماء. تسمى هذه التقنية الترسيب الانتقائي.

<ol>
	<li>فيرشينج ، فرديناند هنري. "الترسيب الانتقائي لهاليدات الفضة من محلول متجانس. فصل اليوديد والبروميد والكلوريد باستخدام تطاير الأمونيا." الكيمياء التحليلية 32 ، العدد 13 (1960): 1876-1878.</li>
	<li>رينولدز ، جون ب. "تجربة Ksp: منتج الذوبان لهيدروكسيد الباريوم". مجلة التعليم الكيميائي 52 ، العدد 8 (1975): 521.</li>
	<li>هو ، مياولين ، وجورج إل بوغمان. "التنبؤ بهطول الأمطار الحمضية والأصباغ المباشرة في المياه الطبيعية". الأصباغ والأصباغ 18 ، العدد 1 (1992): 35-46.</li>
</ol>

ترسيب الأيونات

توقع هطول الأمطار المعادلة التي تصف التوازن بين كربونات الكالسيوم الصلبة وأيوناتها المنحلة هي: <img alt="Eq1" src="/files/ftp_upload/11420/11420_Equation_1.png" style="text-align: center; height: 50px;" /> من المهم أن ندرك أن هذا التوازن يتم إنشاؤه في أي محلول مائي يحتوي على أيونات Ca2+ and CO32–، ليس فقط في محلول يتكون من تشبع الماء بكربونات الكالسيوم. ضع في اعتبارك، على سبيل المثال، خلط المحاليل المائية للمركبات القابلة للذوبان كربونات الصوديوم ونترات الكالسيوم. إذا كانت تركيزات أيونات الكالسيوم والكربونات في الخليط لا تنتج حاصل قسمة تفاعل، Q، الذي يتجاوز منتج الذوبان، Ksp، فإذاً لن يحدث ترسيب. إذا أسفرت تركيزات الأيونات عن ناتج تفاعل أكبر من منتج الذوبان، فسيحدث الترسيب، مما يؤدي إلى خفض تلك التركيزات حتى يتحقق التوازن (Q = Ksp). تعد مقارنة Q بـ Ksp للتنبؤ بالترسيب مثالاً على النهج العام لتوقع اتجاه التفاعل. بالنسبة للحالة المحددة لتوازن الذوبان: Q < Ksp: يستمر التفاعل في الاتجاه الأمامي (المحلول ليس مشبعًا ، لم يحدث ترسيب) Q > Ksp: يستمر التفاعل في الاتجاه العكسي (المحلول مفرط التشبع؛ سيحدث الترسيب) في المحاليل التي تحتوي على اثنين أو أكثر من الأيونات التي قد تشكل مركبات غير قابلة للذوبان لها نفس الأيون المضاد، يمكن استخدام استراتيجية تجريبية تسمى الترسيب الانتقائي لإزالة أيونات مفردة من المحلول. عن طريق زيادة تركيز الأيونات المضادة بطريقة مضبوطة، قد تترسب الأيونات في المحلول بشكل فردي، بافتراض أن قابلية ذوبانها المركبة مختلفة بشكل كافٍ. <span style="color: rgb(0, 0, 0); font-family: Roboto, RobotoDraft, Helvetica, Arial, sans-serif; font-size: 13px; font-style: normal; font-variant-ligatures: normal; font-variant-caps: normal; font-weight: 400; letter-spacing: normal; orphans: 2; text-align: start; text-indent: 0px; text-transform: none; white-space: pre-wrap; widows: 2; word-spacing: 0px; -webkit-text-stroke-width: 0px; background-color: rgb(255, 255, 255); text-decoration-thickness: initial; text-decoration-style: initial; text-decoration-color: initial; display: inline !important; float: none;"> في المحاليل ذات التركيزات المتساوية من الأيونات المستهدفة، فإن الأيون الذي يشكل أقل مركب قابل للذوبان سوف يترسب أولاً (عند أدنى تركيز للأيون المضاد) . تترسب الأيونات الأخرى لاحقًا عند الوصول إلى ذوبان’ مركبها. ترسيب هاليدات الفضة يحتوي المحلول على 0.00010 مول من KBr و 0.10 مول من KCl لكل لتر. AgNO3 يضاف تدريجياً إلى هذا المحلول. ما هي الأشكال الأولى، AgBr الصلب أم AgCl الصلب؟ التوازنان المتضمنان هما: <img alt="Eq2" src="/files/ftp_upload/11420/11420_Equation_2.png" style="text-align: center; height: 90px;" /> إذا احتوى المحلول على تركيزات متساوية من Cl– و Br–، فإن الملح الفضي الذي يحتوي على أصغر Ksp (AgBr) سيترسب أولاً. لكن التركيزات غير متساوية، لذلك يجب حساب [Ag+] الذي يبدأ عنده AgCl في الترسب وحساب [Ag+] الذي يبدأ عنده AgBr في الترسي. الملح الذي يتكون في المستوى الأقل لـ [Ag+] يترسب أولاً. يترسب AgBr عندما Q = Ksp لـ AgBr <img alt="Eq3" src="/files/ftp_upload/11420/11420_Equation_3.png" style="text-align: center; height: 112px;" /> يبدأ AgBr في الترسب عندما يكون [Ag+] 5.0 × 10−9 M. بالنسبة إلى AgCl: يترسب AgCl عندما Q يساوي Ksp لـ AgCl (1.6  × 10-10). عندما [Cl–] = 0.10 M:<img alt="Eq4" src="/files/ftp_upload/11420/11420_Equation_4.png" style="text-align: center; height: 112px;" />
 يبدأ AgCl في الترسب عندما يكون [Ag+] هو 1.6  × 10-9 M. يبدأ AgCl في الترسيب عندما يكون[Ag+] أقل من AgBr، لذلك يبدأ AgCl في الترسب أولاً. لاحظ أن تركيز أيون الكلوريد للخليط الأولي كان أكبر بكثير من تركيز أيون البروميد، ولذا ترسب كلوريد الفضة أولاً على الرغم من احتوائه على Ksp أكبر من تلك الموجودة في بروميد الفضة. هذا النص مقتبس من <a href="https://openstax.org/books/chemistry-2e/pages/15-1-precipitation-and-dissolution">Openstax, Chemistry 2e, Chapter 15.1: Precipitation and Dissolution</a>.

التنبؤ بهطول الأمطار
المعادلة التي تصف التوازن بين كربونات الكالسيوم الصلبة وأيوناتها المذابة هي:

من المهم أن ندرك ...

Education

JoVE Core

الكيمياء

واحد

Acid-base and Solubility Equilibria

اتزان الحموض والقواعد والذوبانية

KEY TERMS AND CONCEPTS

المصطلحات والمفاهيم الأساسية

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 quotient, Q, does not exceed the solubility product, Ksp, no solid appears. If Q is greater than Ksp, precipitation occurs until the solution returns to equilibrium, where Q = Ksp.

For solubility equilibria, the comparison of Q and Ksp predicts the reaction direction. When Q Ksp, the solution is supersaturated and a precipitate will form.

Selective precipitation uses this same idea to remove one ion from a mixture that contains more than one ion that can form insoluble compounds with the same counter ion. By adding the counter ion in a controlled way, ions can be precipitated one at a time if their solubilities are different enough. If the target ions start at equal concentrations, the ion that forms the least soluble compound precipitates first.

A silver halide example shows how concentration affects the order of precipitation. A solution contains 0.00010 mol of KBr and 0.10 mol of KCl per liter, and AgNO3 is added gradually. Although AgBr has the smaller Ksp, AgCl begins to precipitate first because it reaches its precipitation point at a lower Ag+ concentration. AgBr begins to precipitate when [Ag+] is 5.0 × 10−9 M, while AgCl begins when [Ag+] is 1.6 × 10−9 M. The chloride ion concentration is much higher than the bromide ion concentration, so silver chloride forms first.

Solubility Limits and Selective Precipitation

solubility-limits-and-selective-precipitation

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

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...

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 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 Q and Ksp predict precipitation and how selective precipitation separates ions. Includes a silver chloride and silver bromide example.

Solubility Limits and Selective Precipitation | JoVE K12

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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 ...

le-chatelier-s-principle-changing-volume-pressure

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 ...

le-chatelier-s-principle-changing-temperature

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 ...

acidbase-strengths-and-dissociation-constants

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 ...

Lewis Acids and Bases

In 1923, G. N. Lewis proposed a generalized definition of acid-base behavior in which acids and bases are identified by their ability to accept or to ...

Common Ion Effect

Compared with pure water, the solubility of an ionic compound is less in aqueous solutions containing a common ion (one also produced by dissolution of ...

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.

Solubility Limits and Selective Precipitation

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