DOCSLIB.ORG

Equilibrium Constant Equilibrium Constant

Total Page：16

File Type：pdf, Size：1020Kb

Download full-text PDF Read full-text
Equilibrium Constant Equilibrium Constant Equilibrium Constant Most chemical systems are governed by equilibria Chemical Equilibrium such that if: aA + bB cC + dD, then & Activity c d (ac) (ad ) K = a b where ax = γx[X], γx is the activity coefficient (aa) (ab) In dilute solutions, γx →1, if we assume γx =1, [C]c[D]d K = [A]a[B]b Equilibrium constants may be written for dissociations, associations, reactions, or distributions. Equilibrium Constant aA + bB cC + dD [C]c[D]d K = [A]a[B]b • If K is very large, that the equilibrium lies far to the right (or towards products). If K is small, the reaction lies towards reactants. • Knowledge of reaction stoichiometry and the equilibrium ©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley) constant allows us to make some predictions about the system. Important items to regarding K expressions Equilibrium and Thermodynamics 1. All solute concentrations should be in mol/L (M). • Gibbs free energy: 2. All gas concentrations should be in atmospheres. 3. By convention, all K’s are calculated relative to 1 M G o -RT ln K solutions or 1 atm gas, so the resulting constants are = o dimensionless. K = e−∆G / RT 4. Concentrations of pure solids, pure liquids and solvents are omitted from the equilibrium constant expression. • If K >1 Go<0 Spontaneous Equilibrium constant expressions are • If K<1 Go NOT Spontaneous thermodynamic relations. 1 Using equilibria to characterize Reaction Quotient (Q) systems The equilibrium constant expression with non- equilibrium concentrations plugged in. Solubility: solubility product (Ksp) Complexation: formation constants (Kn), • If Q>K, the reaction must proceed to the left cumulative formation constants (n) • If Q<K, the reaction must proceed to the right Acids and Bases: acid dissociation constant • If Q=K the reaction is at equilibrium (Ka), base hydrolysis constant (Kb) Solubility Solubility - Separation by The solubility product (Ksp) describes the concentrations of precipitation species present when ions are in equilibrium with undissolved salt. It is possible to quantitatively separate two or more species based on their solubility. 2+ 2- Ability to do so is related to the magnitudes of the Ksp for each Example: CaCO3 = Ca + CO3 ion. 2+ 2- 2+ 2- -9 K= [Ca ][CO3 ]/[CaCO3]= [Ca ][CO3 ] = Ksp = 6.0 x 10 Example: What is the concentration of calcium in a saturated solution Is it possible to precipitate 99% of 0.010M Ce3+ by adding 2- 2+ of calcium carbonate? oxalate (C2O4 ) without precipitate 0.010M Ca -8 CaC2O4 Ksp = 1.3 x 10 -29 Ce2(C2O4)3 Ksp = 3.0 x 10 Acids and Bases pH - H O OH +H + • Definitions: Lewis – Electrons (acid: electron pair acceptor); 2 BrØnsted-Lowry (acid: proton donor) P-function: pX = - log10[X] • Conjugate Acid-Base Pairs: related by the gain or loss of one proton (ex. Acetic acid & acetate ion). pH = -log [H+] • Neutralization Reactions: reactions of acid and base to form salts and water pH > 7 is basic, pH<7 is acidic • Solvent Autoprotolysis or self-ionization: water is the most common, in which it acts both an acid and a base: pOH=-log[OH-] 2H O H O ++O H - 2 3 pH + pOH = pKw = 14.00 at 25oC - OH +H + H 2O + - o Autoprotolysis constant (equilibrium constants) Ex. Concentration of H and OH in pure water at 25 C -7 + − + − −14 ANS: 1.0x10 M Kw = [H 3O ][OH ] = [H ][OH ] =10 2 Acid and Base Strength Solubility – Common Ion Effects Based on percent dissociation in solution A salt will be less soluble if one of its constituent ions is already present in the solution. • Strong acids/bases dissociate essentially completely in water • Weak acids/bases only partially dissociate, results in Example: What is the solubility of ferric hydroxide equilibrium concentrations of both the acid and its -38 conjugate base in pH 8.0 buffered aqueous solution (Ksp=4x10 ) Ka - + Ka - HA +H2O A +H3O H A A + H + Kb + - B + H2O BH +OH Weak Acid/Base Equilibria Diprotic acids/Dibasic bases Polyprotic acids/Polybasic bases • Weak Acid: acid dissociation constant (Ka) + − Ka [H ][HA ] A- +H O+ − + − + HA +H2O 3 [A ][H O ] [A ][H ] + - Ka1 = 3 H A H + HA Ka = = 2 [H 2 A] Ka - + H A A + H [HA] [HA] - + 2- + 2− HA H + A [H ][A ] Ka = • Weak Base, hydrolysis reaction: base hydrolysis 2 [HA− ] constant (Kb) Ka1, Ka2 … [BH + ][OH − ] Kb + - + - -3 B + H2O BH +OH Kb = H PO H + H PO Ka1 = 7.11 x 10 [B] 3 4 2 4 - + 2- -8 H2PO4 H + HPO4 Ka2 = 6.32 x 10 • Relationship between Ka and Kb for conjugate HPO 2- H+ + PO 3- Ka3 = 7.1 x 10-13 acid/base pairs: 4 4 + − + − −14 Kw = Ka × Kb = [H 3O ][OH ] = [H ][OH ] =10 Complex Formation Activity Formation of coordinate bonds between Lewis Acids/Bases + K1 + If aA + bB cC + dD, then Ag (aq ) + NH 3 (aq) Ag ( NH 3 ) (aq ) + K2 + Ag ( NH 3 ) (aq ) + NH 3 (aq) Ag ( NH 3 ) 2 (aq ) c d c d (ac ) (ad ) (γ c[C]) (γ d [D]) + + K = = [Ag(NH 3 ) ] [Ag(NH 3 )2 ] a b a b (aa ) (ab ) (γ a[A]) (γ b[B]) K1 = + K 2 = + [Ag ][NH 3 ] [Ag(NH 3 ) ][NH 3 ] where ax is the activity of x Formation constants (Kf) are the equilibrium constants for complex ion a = γ [X], γ is the activity coefficient formation. The overall, or cumulative, formation constants are x x x denoted i + Kf + If γ x →1, ax → [X ] Ag (aq) + 2NH3 (aq) Ag(NH3 )2 (aq) + c d c d c d (ac ) (ad ) (γ c[C]) (γ d [D]) [C] [D] [Ag(NH 3 )2 ] K = = = K f = + 2 = β 2 = K1 ⋅ K 2 a b a b a b (aa ) (ab ) (γ a[A]) (γ b[B]) [A] [B] [Ag ][NH 3 ] 3 Activity Coefficient Ionic Strength • Related to the size of the hydrated species • Calculate using the extended Debye-Hückel equation, 1 2 2 1 2 which relates activity coefficients to the ability of ions in µ = (c1 z1 +c2 z2 +...) = ci zi i solution to interact with one another. 2 2 where c is concentration and z is charge of each ion 0.51z 2 µ log r = − α µ Ex. What is the ionic strength of a 0.010 M Na2SO4? 1+ ½*{(0.020*1)+(0.010*(-2)2)}=0.030 M 305 If add 0.020 M KBr? z = charge of the ion ½*{(0.020*1) + (0.020*1) + (0.020*1) + (0.010*4)}=0.050 M = effective diameter "hydrated" of the ion in nanometers = ionic strength of the solution Ionic Strength and Ionic Atmospheres Hydrated Radius • Ions with small ionic radii and large charge tend to more strongly bind to solvent molecules (Ion-dipole interactions). • The result of this binding is a larger hydrated radius, causing diminished interaction The greater the ionic strength of a solution, the higher the charge in the with other ions. ionic atmosphere. Each ion-plus-atmosphere contains less net charge and there is less attraction between any particular cation and anion. Ionic Strength, Ion Charge, and Ion Size effect pH Revisited 0.51z 2 µ • Increased ionic conc. log r = − α µ • Concentration is replaced with activity decreased activity coefficient 1+ 305 + • Increased ion charge (±) pH = -log aH+ = -log [H ]γH+ increased departure of activity coefficient from unity (Multiply charged ions are generally more likely to interact with other ions than Examples : singly charged.) • Calculate the pH of pure water using activity coefficients • Smaller hydrated radius correctly. increased importance of activity effects • Calculate the pH of water containing 0.10 M KCl at 25oC. 4.
Load more

Recommended publications
  • Characteristics of Chemical Equilibrium
    Characteristics of Chemical Equilibrium Chapter 14: Chemical Equilibrium © 2008 Brooks/Cole 1 © 2008 Brooks/Cole 2 Equilibrium is Dynamic Equilibrium is Independent of Direction of Approach Reactants convert to products N2(g) + 3 H2(g) 2 NH3(g) a A + b B c C + d D Species do not stop forming OR being destroyed Rate of formation = rate of removal Concentrations are constant. © 2008 Brooks/Cole 3 © 2008 Brooks/Cole 4 Equilibrium and Catalysts The Equilibrium Constant For the 2-butene isomerization: H3C CH3 H3C H C=C C=C H H H CH3 At equilibrium: rate forward = rate in reverse An elementary reaction, so: kforward[cis] = kreverse[trans] © 2008 Brooks/Cole 5 © 2008 Brooks/Cole 6 1 The Equilibrium Constant The Equilibrium Constant At equilibrium the concentrations become constant. We had: kforward[cis] = kreverse[trans] kforward [trans] or = kreverse [cis] kforward [trans] Kc = = = 1.65 (at 500 K) kreverse [cis] “c” for concentration based © 2008 Brooks/Cole 7 © 2008 Brooks/Cole 8 The Equilibrium Constant The Equilibrium Constant For a general reaction: a A + b B c C + d D [NO]2 N2(g) + O2(g) 2 NO(g) Kc = Products raised to [N2] [O2] stoichiometric powers… k [C]c [D]d forward …divided by reactants Kc = = a b kreverse [A] [B] raised to their stoichiometric [SO ] 1 2 powers 8 S8(s) + O2(g) SO2(g) Kc = [O2] © 2008 Brooks/Cole 9 © 2008 Brooks/Cole 10 Equilibria Involving Pure Liquids and Solids Equilibria in Dilute Solutions [Solid] is constant throughout a reaction. density g / L • pure solid concentration = mol.
    [Show full text]
  • Laboratory 1: Chemical Equilibrium 1
    1 Laboratory 1: Chemical Equilibrium 1 Reading: Olmstead and Williams, Chemistry , Chapter 14 (all sections) Purpose: The shift in equilibrium position of a chemical reaction with applied stress is determined. Introduction Chemical Equilibrium No chemical reaction goes to completion. When a reaction stops, some amount of reactants remain. For example, although we write → ← 2 CO 2 (g) 2 CO (g) + O 2 (g) (1) as though it goes entirely to products, at 2000K only 2% of the CO 2 decomposes. A chemical reaction reaches equilibrium when the concentrations of the reactants and products no longer change with time. The position of equilibrium describes the relative amounts of reactants and products that remain at the end of a chemical reaction. The position of equilibrium for reaction (1) is said to lie with the reactants, or to the left, because at equilibrium very little of the carbon dioxide has reacted. On the other hand, in the reaction → ← H2 (g) + ½ O2 (g) H2O (g) (2) the equilibrium position lies very far to the right since only very small amounts of H 2 and O 2 remain after the reaction reaches equilibrium. Since chemists often wish to maximize the yield of a reaction, it is vital to determine how to control the position of the equilibrium. The equilibrium position of a reaction may shift if an external stress is applied. The stress may be in the form of a change in temperature, pressure, or the concentration of one of the reactants or products. For example, consider a flask with an equilibrium mixture of CO 2, CO, and O 2, as in reaction (1).
    [Show full text]
  • Answer Key Chapter 16: Standard Review Worksheet – + 1
    Answer Key Chapter 16: Standard Review Worksheet – + 1. H3PO4(aq) + H2O(l) H2PO4 (aq) + H3O (aq) + + NH4 (aq) + H2O(l) NH3(aq) + H3O (aq) 2. When we say that acetic acid is a weak acid, we can take either of two points of view. Usually we say that acetic acid is a weak acid because it doesn’t ionize very much when dissolved in water. We say that not very many acetic acid molecules dissociate. However, we can describe this situation from another point of view. We could say that the reason acetic acid doesn’t dissociate much when we dissolve it in water is because the acetate ion (the conjugate base of acetic acid) is extremely effective at holding onto protons and specifically is better at holding onto protons than water is in attracting them. + – HC2H3O2 + H2O H3O + C2H3O2 Now what would happen if we had a source of free acetate ions (for example, sodium acetate) and placed them into water? Since acetate ion is better at attracting protons than is water, the acetate ions would pull protons out of water molecules, leaving hydroxide ions. That is, – – C2H3O2 + H2O HC2H3O2 + OH Since an increase in hydroxide ion concentration would take place in the solution, the solution would be basic. Because acetic acid is a weak acid, the acetate ion is a base in aqueous solution. – – 3. Since HCl, HNO3, and HClO4 are all strong acids, we know that their anions (Cl , NO3 , – and ClO4 ) must be very weak bases and that solutions of the sodium salts of these anions would not be basic.
    [Show full text]
  • Drugs and Acid Dissociation Constants Ionisation of Drug Molecules Most Drugs Ionise in Aqueous Solution.1 They Are Weak Acids Or Weak Bases
    Drugs and acid dissociation constants Ionisation of drug molecules Most drugs ionise in aqueous solution.1 They are weak acids or weak bases. Those that are weak acids ionise in water to give acidic solutions while those that are weak bases ionise to give basic solutions. Drug molecules that are weak acids Drug molecules that are weak bases where, HA = acid (the drug molecule) where, B = base (the drug molecule) H2O = base H2O = acid A− = conjugate base (the drug anion) OH− = conjugate base (the drug anion) + + H3O = conjugate acid BH = conjugate acid Acid dissociation constant, Ka For a drug molecule that is a weak acid The equilibrium constant for this ionisation is given by the equation + − where [H3O ], [A ], [HA] and [H2O] are the concentrations at equilibrium. In a dilute solution the concentration of water is to all intents and purposes constant. So the equation is simplified to: where Ka is the acid dissociation constant for the weak acid + + Also, H3O is often written simply as H and the equation for Ka is usually written as: Values for Ka are extremely small and, therefore, pKa values are given (similar to the reason pH is used rather than [H+]. The relationship between pKa and pH is given by the Henderson–Hasselbalch equation: or This relationship is important when determining pKa values from pH measurements. Base dissociation constant, Kb For a drug molecule that is a weak base: 1 Ionisation of drug molecules. 1 Following the same logic as for deriving Ka, base dissociation constant, Kb, is given by: and Ionisation of water Water ionises very slightly.
    [Show full text]
  • Electrochemistry –An Oxidizing Agent Is a Species That Oxidizes Another Species; It Is Itself Reduced
    Oxidation-Reduction Reactions Chapter 17 • Describing Oxidation-Reduction Reactions Electrochemistry –An oxidizing agent is a species that oxidizes another species; it is itself reduced. –A reducing agent is a species that reduces another species; it is itself oxidized. Loss of 2 e-1 oxidation reducing agent +2 +2 Fe( s) + Cu (aq) → Fe (aq) + Cu( s) oxidizing agent Gain of 2 e-1 reduction Skeleton Oxidation-Reduction Equations Electrochemistry ! Identify what species is being oxidized (this will be the “reducing agent”) ! Identify what species is being •The study of the interchange of reduced (this will be the “oxidizing agent”) chemical and electrical energy. ! What species result from the oxidation and reduction? ! Does the reaction occur in acidic or basic solution? 2+ - 3+ 2+ Fe (aq) + MnO4 (aq) 6 Fe (aq) + Mn (aq) Steps in Balancing Oxidation-Reduction Review of Terms Equations in Acidic solutions 1. Assign oxidation numbers to • oxidation-reduction (redox) each atom so that you know reaction: involves a transfer of what is oxidized and what is electrons from the reducing agent to reduced 2. Split the skeleton equation into the oxidizing agent. two half-reactions-one for the oxidation reaction (element • oxidation: loss of electrons increases in oxidation number) and one for the reduction (element decreases in oxidation • reduction: gain of electrons number) 2+ 3+ - 2+ Fe (aq) º Fe (aq) MnO4 (aq) º Mn (aq) 1 3. Complete and balance each half reaction Galvanic Cell a. Balance all atoms except O and H 2+ 3+ - 2+ (Voltaic Cell) Fe (aq) º Fe (aq) MnO4 (aq) º Mn (aq) b.
    [Show full text]
  • Chapter 15 Chemical Equilibrium
    Chapter 15 Chemical Equilibrium Learning goals and key skills: Explain what is meant by chemical equilibrium and how it relates to reaction rates Write the equilibrium-constant expression for any reaction Convert Kc to Kp and vice versa Relate the magnitude of an equilibrium constant to the relative amounts of reactants and products present in an equilibrium mixture. Manipulate the equilibrium constant to reflect changes in the chemical equation Write the equilibrium-constant expression for a heterogeneous reaction Calculate an equilibrium constant from concentration measurements Predict the direction of a reaction given the equilibrium constant and the concentrations of reactants and products Calculate equilibrium concentrations given the equilibrium constant and all but one equilibrium concentration Calculate equilibrium concentrations given the equilibrium constant and the starting concentrations Use Le Chatelier’s principle to predict how changing the concentrations, volume, or temperature of a system at equilibrium affects the equilibrium position. The Concept of Equilibrium Chemical equilibrium occurs when a reaction and its reverse reaction proceed at the same rate. 1 Concept of Equilibrium • As a system approaches equilibrium, both the forward and reverse reactions are occurring. • At equilibrium, the forward and reverse reactions are proceeding at the same rate. • Once equilibrium is achieved, the amount of each reactant and product remains constant. The same equilibrium is reached whether we start with only reactants (N2 and H2) or with only product (NH3). Equilibrium is reached from either direction. 2 The Equilibrium Constant • Consider the generalized reaction aA + bB cC + dD The equilibrium expression for this reaction would be [C]c[D]d K = c [A]a[B]b Since pressure is proportional to concentration for gases in a closed system, the equilibrium expression can also be written c d (PC) (PD) Kp = a b (PA) (PB) Chemical equilibrium occurs when opposing reactions are proceeding at equal rates.
    [Show full text]
  • Acid Dissociation Constant - Wikipedia, the Free Encyclopedia Page 1
    Acid dissociation constant - Wikipedia, the free encyclopedia Page 1 Help us provide free content to the world by donating today ! Acid dissociation constant From Wikipedia, the free encyclopedia An acid dissociation constant (aka acidity constant, acid-ionization constant) is an equilibrium constant for the dissociation of an acid. It is denoted by Ka. For an equilibrium between a generic acid, HA, and − its conjugate base, A , The weak acid acetic acid donates a proton to water in an equilibrium reaction to give the acetate ion and − + HA A + H the hydronium ion. Key: Hydrogen is white, oxygen is red, carbon is gray. Lines are chemical bonds. K is defined, subject to certain conditions, as a where [HA], [A−] and [H+] are equilibrium concentrations of the reactants. The term acid dissociation constant is also used for pKa, which is equal to −log 10 Ka. The term pKb is used in relation to bases, though pKb has faded from modern use due to the easy relationship available between the strength of an acid and the strength of its conjugate base. Though discussions of this topic typically assume water as the solvent, particularly at introductory levels, the Brønsted–Lowry acid-base theory is versatile enough that acidic behavior can now be characterized even in non-aqueous solutions. The value of pK indicates the strength of an acid: the larger the value the weaker the acid. In aqueous a solution, simple acids are partially dissociated to an appreciable extent in in the pH range pK ± 2. The a actual extent of the dissociation can be calculated if the acid concentration and pH are known.
    [Show full text]
  • Chapter 8:Chemical Reaction Equilibria 8.1 Introduction Reaction Chemistry Forms the Essence of Chemical Processes
    Chapter 8:Chemical Reaction Equilibria 8.1 Introduction Reaction chemistry forms the essence of chemical processes. The very distinctiveness of the chemical industry lies in its quest for transforming less useful substances to those which are useful to modern life. The perception of old art of ‘alchemy’ bordered on the magical; perhaps in today’s world its role in the form of modern chemistry is in no sense any less. Almost everything that is of use to humans is manufactured through the route of chemical synthesis. Such reactive processes need to be characterized in terms of the maximum possible yield of the desired product at any given conditions, starting from the raw materials (i.e., reactants). The theory of chemical reactions indicates that rates of reactions are generally enhanced by increase of temperature. However, experience shows that the maximum quantum of conversion of reactants to products does not increase monotonically. Indeed for a vast majority the maximum conversion reaches a maximum with respect to reaction temperature and subsequently diminishes. This is shown schematically in fig. 8.1. Fig. 8.1 Schematic of Equilibrium Reaction vs. Temperature The reason behind this phenomenon lies in the molecular processes that occur during a reaction. Consider a typical reaction of the following form occurring in gas phase: AgBgCgDg()+→+ () () (). The reaction typically begins with the reactants being brought together in a reactor. In the initial phases, molecules of A and B collide and form reactive complexes, which are eventually converted to the products C and D by means of molecular rearrangement. Clearly then the early phase of the reaction process is dominated by the presence and depletion of A and B.
    [Show full text]
  • Lab.10. Equilibrium. Determination of Dissociation Constant for a Weak Acid
    Lab.8. Equilibrium. Determination of dissociation constant for a weak acid Key words: Equilibrium, equilibrium constant, dissociation constant, electrolytes, non electrolyte, weak and strong acids (bases), dissociation, titration, buffer solutions, calculation the pH of buffer solution, the Henderson-Hasselbalch equation, the unique properties of buffer solution (effect of dilution, effect of added acids and bases, buffer capacity). Literature J. A. Beran; Laboratory Manual for Principles of General Chemistry, pp. 281-346. M. Hein and S. Arena: Introduction to Chemistry, 13th ed. Wiley 2011; pp. J. Crowe. T. Bradshaw, P. Monk, Chemistry for the Biosciences. The essential concepts., Oxford University Press, 2006; Chapter 16, 482 - 512. J. Brady, N. Jespersen, A. Hysop, Chemistry, International Student Version, 7thed. Wiley, 2015, Chapter 14, 686 – 709. P. Monk, Physical chemistry, Wiley 2004, Chapters, 5, 6, 177 – 229 and 233- 276, Theoretical background The Swedish chemist Arrhenius suggested that solutions that conduct electricity, so called electrolytes, do so because they dissociate into charged species called ions. Compounds of this type: acids, bases and salts may be classified as strong electrolytes dissociate in solution almost completely into ions, weak electrolytes dissociate only to a small extent in solution (only 10% or less ions are present in solution). Acids are compounds that ionize to release hydrogen ions, H+, to their surroundings. Bases are compounds that accept hydrogen ions, The equation for the ionization of a weak acid may be expressed as: HA ↔ H+ + A- (1) However, this suggests that protons exist free in solution. The reality is that protons are solvated in solution, that is they go around attached to a solvent (water) molecules.
    [Show full text]
  • Solubility Equilibria
    Chapter 16 – Cont. AQUEOUS IONIC EQUILIBRIA 1 SOLUBILITY EQUILIBRIA Carlsbad Caverns, NM. Limestone Big Gypsum, Thornhill Cave, Breckinridge County, KY (CaCO 3) formations inside the cave results from subtle shifts in carbonate equilibria acting over millions of years. http://www.darklightimagery.net/ gypsum.html http://www.traveleze.com/travel_planning/ carlsbad.html 2 1 The Black Widow (Lynn Turner Case) Image available at Matthew E. Johll, “ Investigating Chemistry: A Forensic Science Perspective.” 1st ed. W.H. Freeman, New York: 2007. 3 Kidney stone types : calcium oxalate, calcium phosphate, struvite (magnesium ammonium phosphate), uric acid, and cystine. Image available http://knol.google.com/k/kidney-stones 4 2 Solubility Product Solubility product is the equilibrium constant for a reaction in which a solid dissolves into its constituent ions in solution The corresponding equilibrium constant is called solubility product constant, Ksp Exercise: Write the Ksp expression for the reaction: 2+ - Hg 2Cl 2(s) Hg 2 (aq) + 2Cl (aq) Is this what you got? 2+ - 2 Ksp = [Hg 2 ][Cl ] Table 16.2 (p. 747) lists K sp values for various solids 5 Solubility Calculations Significance of solubility product: 2+ - 2 The equation Ksp = [Hg 2 ][Cl ] indicates that if solid Hg 2Cl 2 is placed in water, it will dissolve until the product of the ion concentrations (raised to a certain power) equals the solubility product for Hg 2Cl 2 When a solution contains excess undissolved solid, it is said to be saturated . 6 3 Solubility Calculations – Cont. Exercises: (Adapted from Silberberg’s “Chemistry: The Molecular Nature of Matter and Change.” 4th ed.) (1) Calculate the solubility of calcium hydroxide (also known as slaked lime , a major component of plaster and cement) -6 in water if the Ksp is 6.5 x 10 .
    [Show full text]
  • Chemical Engineering Thermodynamics
    CHEMICAL ENGINEERING THERMODYNAMICS Andrew S. Rosen SYMBOL DICTIONARY | 1 TABLE OF CONTENTS Symbol Dictionary ........................................................................................................................ 3 1. Measured Thermodynamic Properties and Other Basic Concepts .................................. 5 1.1 Preliminary Concepts – The Language of Thermodynamics ........................................................ 5 1.2 Measured Thermodynamic Properties .......................................................................................... 5 1.2.1 Volume .................................................................................................................................................... 5 1.2.2 Temperature ............................................................................................................................................. 5 1.2.3 Pressure .................................................................................................................................................... 6 1.3 Equilibrium ................................................................................................................................... 7 1.3.1 Fundamental Definitions .......................................................................................................................... 7 1.3.2 Independent and Dependent Thermodynamic Properties ........................................................................ 7 1.3.3 Phases .....................................................................................................................................................
    [Show full text]
  • The Equilibrium Constant
    Chapter 15: Chemical Equilibrium Key topics: Equilibrium Constant Calculating Equilibrium Concentrations The Concept of Equilibrium Consider the reaction k1 A ⌦ B k 1 − At equilibrium there is no net change in [A] or [B], namely d[A] d[B] =0= dt dt But A molecules can react to form B, and B molecules can react to form A, if the forwards and backwards rates are equal. k1[A] = k 1[B] − This is because d[A] = k1[A] + k 1[B] dt − − At the molecular level, typically the reactions do not stop. Initially there is net evaporation of liquid, but after dynamic equilibrium is established, the rate of evaporation = rate of condensation. (from chemtutorial.net) Consider the reaction N2O4(g) 2NO2(g) Begin with only N2O4 (colorless). Begin with only NO2 (brown). 2 The forward rate = kf [N2O4]; the reverse rate = kr [NO2] (the elementary reactions are the same as the overall reaction) In both cases equilibrium is established after some time. The Equilibrium Constant At equilibrium the forwards and reverse rates are equal 2 2 kf [NO2]eq kf [N2O4]eq = kr[NO2]eq = = Kc ) kr [N2O4]eq Kc is the equilibrium constant (the “c” here stands for concentration). In general, to write an equilibrium constant o products over reactants o exponents are the coefficients from the balanced reaction Magnitude of the Equilibrium Constant o If Kc >> 1, products are favored (rxn nearly complete). o If Kc << 1, reactants are favored (rxn hardly proceeds). o If Kc is close to 1, the system contains comparable amounts of products and reactants.
    [Show full text]