Name: acid-base color changes
Uploaded: 2020-09-24T15:06:51.000Z
Duration: 2 min 39 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

Indicators

Gli indicatori cambiano colore a seconda del pH e possono essere usati per monitorare il pH di una soluzione. Un indicatore di pH contiene un acido debole, rappresentato da HIn, che cambia colore quando viene convertito nella sua base coniugata, lo ione In.Il colore cambia nell'intervallo di pH dell'indicatore, che dipende dal suo pKa. Se una soluzione ha un pH inferiore al pKa dell'indicatore, la concentrazione di idronio è alta.

L'equilibrio della reazione si sposta verso sinistra, incrementando la concentrazione di HIn. Quando viene aggiunta una base, la concentrazione di idroni diminuisce. Questa variazione di pH converte HIn nello ione In, e il colore della soluzione cambia.

Il colore HIn domina quando il pH è inferiore al pKa di almeno uno;il colore degli ioni domina quando il pH è maggiore di almeno uno, e il colore è una miscela a pH intermedi. Esistono numerosi indicatori di pH che possono essere utilizzati per monitorare il cambiamento di colore del pH in intervalli di pH distinti. Il punto in cui un indicatore cambia colore è detto endpoint.

L'indicatore ideale per una titolazione ha un punto finale vicino al punto di equivalenza. La curva di titolazione di un acido e di una base forti è ripida, per via della rapida variazione del pH vicino al punto di equivalenza. Queste titolazioni possono usare una gamma relativamente vasta di indicatori.

Durante la titolazione di acido cloridrico 0, 1 molare con idrossido di sodio 0, 1 molare, il pH aumenta rapidamente da 3 a 11. O la fenolftaleina, un indicatore con un intervallo di pH compreso fra 8, 3 e 10, o il rosso di metilene, un indicatore con un intervallo fra 4, 2 e 6, potrebbe essere utilizzato qui poiché i loro endpoint si sovrappongono alla parte ripida della curva. Al contrario, una titolazione che coinvolge un acido o una base debole ha una curva meno ripida di una con solo un acido forte e una base forte;pertanto, è possibile utilizzare una gamma di indicatori relativamente ristretta.

Durante la titolazione di acido acetico 0, 1 molare con idrossido di sodio 0, 1 molare, il pH aumenta rapidamente da 7 a 11. La fenolftaleina può essere utilizzata qui poiché il suo endpoint si sovrappone al punto di equivalenza, e il rosso di metilene sarebbe un indicatore inefficace.

Coleman, William F. "Modelli molecolari di indicatori". Journal of Chemical Education 87 no.1, (2008):1152 <a href="https://pubs.acs.org/doi/pdf/10.1021/ed800038w">https://pubs.acs.org/doi/pdf/10.1021/ed800038w.</a>

indicatori

Alcune sostanze organiche cambiano colore nella soluzione diluita quando la concentrazione dello ione idronio raggiunge un valore particolare. Ad esempio, la fenolftaleina è una sostanza incolore in qualsiasi soluzione acquosa con una concentrazione di ioni idronio maggiore di 5,0 × 10^−9 M (pH <8,3). Nelle soluzioni più basiche in cui la concentrazione di ioni idronio è inferiore a 5,0 × 10^−9 M (pH > 8,3), è rosso o rosa. Sostanze come la fenolftaleina, che possono essere utilizzate per determinare il pH di una soluzione, sono chiamate indicatori acido-base. Gli indicatori acido-base sono acidi organici deboli o basi organiche deboli.

L'equilibrio in una soluzione dell'indicatore acido-base metilarancio, un acido debole, può essere rappresentato da un'equazione in cui utilizziamo HIn come semplice rappresentazione della complessa molecola di metilarancio:

<img alt="Eq1" src="/files/ftp_upload/11415/11415_Equation_1.png" style="text-align: center; height: 115px;" />

L'anione dell'arancio metilico, In^−, è giallo e la forma non ionizzata, HIn, è rossa. Quando aggiungiamo un acido a una soluzione di arancio metilico, l’aumento della concentrazione di ioni idronio sposta l’equilibrio verso la forma rossa non ionizzata, secondo il principio di Le Châtelier. Se aggiungiamo la base, spostiamo l'equilibrio verso la forma gialla. Questo comportamento è del tutto analogo all'azione dei buffer.

Il colore percepito di una soluzione indicatrice è determinato dal rapporto tra le concentrazioni delle due specie In^− e HIn. Se la maggior parte dell'indicatore (tipicamente circa il 60−90% o più) è presente come In^−, il colore percepito della soluzione è giallo. Se la maggior parte è presente come HIn, il colore della soluzione appare rosso. L'equazione di Henderson-Hasselbalch è utile per comprendere la relazione tra il pH di una soluzione indicatrice e la sua composizione (quindi il colore percepito):

<img alt="Eq1" src="/files/ftp_upload/11415/11415_Equation_2.png" style="text-align: center; height: 75px;" />

Nelle soluzioni in cui pH > pK_a, il termine logaritmico deve essere positivo, indicando un eccesso della forma base coniugata dell'indicatore (soluzione gialla). Quando pH > pK_a, il termine logaritmico deve essere negativo, indicando un eccesso di acido coniugato (soluzione rossa). Quando il pH della soluzione è vicino all'indicatore pK_a, sono presenti quantità apprezzabili di entrambi i partner coniugati e il colore della soluzione è quello di una combinazione additiva di ciascuno (giallo e rosso, che produce l'arancione). L'intervallo di cambiamento di colore (o intervallo di pH) per un indicatore acido-base è definito come l'intervallo di valori di pH entro il quale si osserva un cambiamento di colore e per la maggior parte degli indicatori questo intervallo è di circa pK_a ± 1.

Esistono molti indicatori acido-base diversi che coprono un'ampia gamma di valori di pH e possono essere utilizzati per determinare il pH approssimativo di una soluzione sconosciuta mediante processo di eliminazione. Gli indicatori universali e la cartina pH contengono una miscela di indicatori e mostrano colori diversi a pH diversi.

Questo testo è adattato da<a href="https://openstax.org/books/chemistry-2e/pages/14-7-acid-base-titrations"> Openstax, Chemistry 2e, Section 14.7: Acid-Base Titrations</a>.

Indicatori

Alcune sostanze organiche cambiano colore in soluzione diluita quando la concentrazione di ioni idronio raggiunge un determinato valore. Ad esempio, la fenolftaleina ...

Education

Didattica

JoVE Core

Chimica

Acid-base and Solubility Equilibria

Equilibri acido-base e di solubilità

KEY TERMS AND CONCEPTS

TERMINI E CONCETTI CHIAVE

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 indicator shifts between two forms. In this weak acid, the anion In− is yellow, while the nonionized form HIn is red. Adding acid increases the hydronium ion concentration and shifts the equilibrium toward the red form, following Le Châtelier’s principle. Adding base shifts the equilibrium toward the yellow form. This is similar to how buffers respond to changes in solution conditions.

The color seen in an indicator solution depends on the ratio of In− to HIn. If most of the indicator, usually about 60% to 90% or more, is present as In−, the solution looks yellow. If most is present as HIn, the solution looks red. When the pH is close to the indicator pKa, both forms are present in noticeable amounts, so the color is a mix of the two. That mix can appear orange.

The Henderson-Hasselbalch equation helps connect pH, pKa, and color. When pH is greater than pKa, the log term is positive and the conjugate base form is favored, so the solution appears yellow. When pH is less than pKa, the log term is negative and the conjugate acid is favored, so the solution appears red. The color change interval for most indicators is about pKa ± 1. Different indicators cover different pH ranges, so students can estimate the pH of an unknown solution by process of elimination. Universal indicators and pH paper use mixtures of indicators and show different colors across the pH scale.

Acid-Base Color Changes

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 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

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

Explore how acid-base indicators change color with pH. Learn about phenolphthalein, methyl orange, pKa, and pH paper for high school chemistry.

Acid-Base Color Changes | JoVE K12

chemistry

JoVE Core: Chimica illustra concetti chiave di chimica attraverso video animazioni concise e di immediata comprensione. Inoltre, i video della sezione "scienziati in azione" mostrano procedure metodologiche realizzate da ricercatori provenienti dai migliori laboratori internazionali.

Alabama Course of Study Science

Matter and Its Interactions

9-12.C.1.6.e

Georgia Standards of Excellence

Physical Science

SPS6.

Indicators and Color Changes

SC3.

Indicators of Chemical Reactions

Chemical Equations

Chemical equations represent the identities and relative quantities of substances involved in a chemical reaction. The substances undergoing reaction are ...

Reaction Stoichiometry

A balanced chemical equation provides a great deal of information in a very succinct format. Chemical formulas provide the identities of the reactants and ...

Limiting Reactant

The relative amounts of reactants and products represented in a balanced chemical equation are often referred to as stoichiometric amounts. However, in ...

Reaction Yield

The theoretical yield of a reaction is the amount of product estimated to form based on the stoichiometry of the balanced chemical equation. The ...

Chemical Reactions in Aqueous Solutions

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Precipitation Reactions

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Oxidation-Reduction Reactions

Oxidation&ndash;Reduction Reactions

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Oxidation Numbers

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Acids, Bases and Neutralization Reactions

An acid-base reaction is one in which a hydrogen ion, H+, is transferred from one chemical species to another. Such reactions are of central importance to ...

Synthesis and Decomposition Reactions

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Reaction Rate

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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

The Integrated Rate Law: The Dependence of Concentration on Time

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Half-life of a Reaction

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Temperature Dependence on Reaction Rate

The Collision Theory
Atoms, molecules, or ions must collide before they can react with each other. Atoms must be close together to form chemical bonds. ...

Arrhenius Plots

The Arrhenius equation relates the activation energy and the rate constant, k, for chemical reactions. In the Arrhenius equation, k = Ae&#8722;Ea/RT, R is ...

Reaction Mechanisms

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Thermochemical Equations

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Rate-Determining Steps

Relating Reaction Mechanisms
In a multistep reaction mechanism, one of the elementary steps progresses significantly slower than the others. This slowest ...

Constant Pressure Calorimetry

Calorimetry is a technique used to measure the amount of heat involved in a chemical or physical process or to measure the heat transferred to or from a ...

Catalysis

The presence of a catalyst affects the rate of a chemical reaction. A catalyst is a substance that can increase the reaction rate without being consumed ...

Constant Volume Calorimetry

Calorimeters are useful to determine the heat released or absorbed by a chemical reaction. Coffee cup calorimeters are designed to operate at constant ...

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

Enthalpy changes are typically tabulated for reactions in which both the reactants and products are at the same conditions. A standard state is a commonly ...

First Law of Thermodynamics

Energy Conservation

Energy can be converted from one form into another, but all of the energy present before a change occurs always exists in some form ...

Enthalpies of Reaction

Hess&rsquo;s law can be used to determine the enthalpy change of any reaction if the corresponding enthalpies of formation of the reactants and products ...

Internal Energy

The total of all possible kinds of energy present in a substance is called the internal energy (U), sometimes symbolized as E. Suppose a system with ...

Dynamic Equilibrium

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Quantifying Heat

Thermal Energy&#160;

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The Equilibrium Constant

Consider the oxidation of sulfur dioxide:

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

Quantifying Work

As a system undergoes a change, its internal energy can change, and energy can be transferred from the system to the surroundings, or from the ...

homogeneous-equilibria-for-gaseous-reactions

Homogeneous Equilibria for Gaseous Reactions

Homogeneous Equilibria for Gaseous Reactions
For gas-phase reactions, the equilibrium constant may be expressed in terms of either the molar ...

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 ...

relative-strengths-of-conjugate-acid-base-pairs

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

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 ...

le-chatelier-s-principle-changing-concentration

Le Chatelier's Principle: Changing Concentration

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Weak Base Solutions

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

Le Chatelier's Principle: Changing Volume (Pressure)

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Mixtures of Acids

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le-chatelier-s-principle-changing-temperature

Le Chatelier's Principle: Changing Temperature

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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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acidbase-strengths-and-dissociation-constants

Acid/Base Strengths and Dissociation Constants

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Polyprotic Acids

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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.

Acid-Base Color Changes

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