DOCSLIB.ORG

Chapter 8: Ionic Compounds I

Total Page：16

File Type：pdf, Size：1020Kb

Download full-text PDF Read full-text
Chapter 8: Ionic Compounds I Chapter 8: Ionic Compounds I. Forming Chemical Bonds A.Chemical bond: force that holds two atoms together 1. Ionic bonds form through the attraction of a positive ion for a negative ion 2. Covalent bonds form through the attraction of a positive nucleus for negative electrons of another atom 3. Bonding behavior is determined by valence electrons i. Valence electrons are generally in s & p orbitals B. Ion Formation 1. Positive ions a. Formed when an atom loses electrons to reach a noble gas configuration b. Cation: positively charged ion c. Example: sodium ● 1s22s22p63s1 ● Loses 3s1 electron to get to noble gas configuration that it’s closest to (neon) ● Sodium ion: 1s22s22p6 ● Notation: Na+ d. Example 2: aluminum ● 1s22s22p63s23p1 ● Loses 3s2 and 3p1 electrons to get to noble gas configuration that it’s closest to (neon) ● Aluminum ion: 1s22s22p6 ● Notation: Al3+ 2. Negative ions a. Formed when an atom gains electrons to reach a noble gas configuration b. Anion: negatively charged ion c. Example: chlorine ● 1s22s22p63s23p5 ● Gains an electron in the 3p orbital to get to noble gas configuration that it’s closest to (argon) ● Ion: 1s22s22p63s23p6 ● Notation: Cl- d. Example 2: phosphorus ● 1s22s22p63s23p3 ● Gains 3 electrons in the 3p orbital to get to noble gas configuration that it’s closest to (argon) ● Ion: 1s22s22p63s23p6 ● Notation: P3- 3. Ions of Transition Metals a. Usually have an outer energy level of s2, then fill d orbital of second highest energy level b. Most transition metals lose the 2 electrons in their outer s orbital, forming ions with a 2+ charge c. Sometimes, the d orbital gets involved, forming ions of 3+ or greater d. All commonly formed ions of each element are listed to the upper right of the element symbol on your periodic table Homework: p. 214 #1-5 II.Formation & Nature of Ionic Bonds A.Ionic bond: electrostatic force that holds oppositely charged particles (ions) together 1. Oxides: ionic compounds composed of a metal ion & oxygen 2. Salts: ionic compounds not containing oxygen 3. Binary compounds: contain only 2 elements 4. Example: sodium & chlorine ● Sodium (Na+) gives up its valence electron to chlorine (Cl-) ● Positive sodium is attracted to the negative chlorine ● Forms sodium chloride (NaCl) 5. Example 2: calcium & fluorine ● Calcium (Ca2+) gives up 2 valence electrons ● Fluorine (F-) only needs 1 electron, but the number of electrons lost must equal the number of electrons gained ● So… calcium gives one electron each to 2 fluorine atoms ● Forms calcium fluoride (CaF2) B. Properties of Ionic Compounds 1. Depend upon the bonds between the atoms involved 2. General characteristics: a. Form a crystal lattice: three-dimensional arrangement of particles; positive ions surrounded by negative ions & negative ions surrounded by positive ions i. Stronger attraction, harder & stronger crystal b. High melting & boiling points due to strong attraction between particles i. Stronger attraction means more energy is required to break the particles apart -- therefore, more heat required c. Soluble: dissolve in water i. When dissolved, these compounds become electrolytes, meaning they conduct electricity C. Energy of Ionic Bonds 1. Chemical reactions always release or absorb energy 2. The formation of ionic compounds always releases energy, making it exothermic 3. Lattice energy: energy required to separate one mole of ions of an ionic compound a. More negative lattice energy = stronger attraction b. Smaller ions have a more negative lattice energy than larger ions c. Bonds between ions with larger numerical charges have a more negative lattice energy Homework: ➔p. 217 #7-11 ➔p. 220 #12-17 III.Names & Formulas of Ionic Compounds A.Formulas for Ionic Compounds 1. Formula unit: simplest ratio of ions represented in an ionic compound a. Examples: NaCl, KBr, MgCl2 2. Ion: charged atom or a group of atoms with a charge a. Monatomic ion: one-atom ion with a definite charge i. Examples: Na+, Mg2+, O2-, F- ii. Charge is equal to an atom’s oxidation number: the number of electrons transferred from an element to form an ion (ex: Na+ has an oxidation number of +1) b. Polyatomic ion: group of atoms bonded together which have a net charge + - 2- i. Examples: NH4 , CN , SO4 B. Writing Formulas of Ionic Compounds 1. Rule: the sum of the charges in a stable compound must be zero a. Examples: i. Na+ + Cl- → NaCl + - 3 ii. Ag + NO3 → AgNO iii.Al3+ + Cl- → __(3) + __(-1) = 0 __(3) + 3(-1) = 0 AlCl3 (aluminum chloride) 2. How to write a formula: a. Step 1: find the chemical symbols or polyatomic ions & the charges they have based on the compound’s name (textbook, PT, common ions list) b. Step 2: write the charges in an algebraic formula, setting the sum of the charges equal to zero c. Step 3: determine how many of each ion are necessary to make the total charge equal to zero 3. Examples!! a. magnesium oxide d. ammonium sulfate 2+ 2- + 2- i. Mg O i. NH4 SO4 ii. __(2) + __(-2) = 0 ii. __(1) + __(-2) = 0 iii.MgO iii.2(1) + __(-2) = 0 b. calcium chloride iv.(NH4)2SO4 i. Ca2+ Cl- e. aluminum sulfate ii. __(2) + __(-1) = 0 i. Al3+ 2- iii.__(2) + 2(-1) = 0 SO4 iv.CaCl2 ii. __(3) + __(-2) = 0 c. aluminum sulfide iii.2(3) + 3(-2) = 0 3+ 2- i. Al S iv.Al2(SO4)3 ii. __(3) + __(-2) = 0 iii.2(3) + 3(-2) = 0 iv.Al2 S3 Homework: ➔p. 224 #19-23 (write out all 3 steps) ➔p. 225 #24-28 (write out all 3 steps) C. Naming Ions & Ionic Compounds 1. Naming Ions a. If two polyatomic ions are identical except for the number of oxygen atoms, the ion with more oxygens takes the suffix “-ate” and the ion with less oxygen takes the suffix “-ite” - - i. Examples: NO3 : nitrate, NO2 : nitrite 2- 3- SO4 : sulfate, SO4 : sulfite b. If four polyatomic ions are identical except for the number of oxygen atoms, use the prefix “per-” for the ion with the most oxygens, suffix “- ate” for the second most, suffix “-ite” for the third, and prefix “hypo-” with suffix “-ite” for the least - i. Examples: ClO4 : perchlorate - ClO3 : chlorate - ClO2 : chlorite ClO-: hypochlorite 2. Naming Binary (Ionic) Compounds a. Written from left to right on the periodic table b. Positive ion first -- use the full element name c. Negative ion last -- if it’s just one element (example: bromine), use the root of the element name and suffix “- ide” (bromide) d. If there is a polyatomic ion involved, just use the name of the polyatomic ion - - do not change the suffix to “-ide” + - i. Example: Na + NO3 → sodium + nitrate → sodium nitrate e. Transition metals & some metals on the right side of the table have more than one possible charge; to distinguish between these, use Roman numerals after the name of the element to indicate the charge i. Example: FeO → since there are no subscripts, they must have the same numerical charge. Oxygen is always 2-, so iron in this case has a charge of 2+. Name would be iron (II) oxide ii. Example: Fe2O3 → the subscript on one element is the same as the charge of the other. Oxygen’s charge is 2-, and iron in this case is 3+. Name would be iron (III) oxide Homework: ➔copy figure 8-8 (p. 227) ➔p. 226 #29-33 ➔p. 227 #34-39 IV.Metallic Bonds & Properties of Metals A.Metals are not ionic compounds, but they do have some similar properties 1. Similarities are due to the fact that bonds are based on attractions between opposite charges B.Metallic Bonds 1. Form lattices in solid state, similar to crystal lattices 2. Metals’ outer electrons are loosely held & overlap when bonded a. Electron sea model: all of the valence electrons of the metals contained in a bond group together to form a “sea”; positive nuclei are “floating” in this b. These valence electrons are free to move from atom to atom -- called delocalized electrons c. Contributes to strong conductivity of metals C. Metallic Properties 1. Properties of metals depend upon # of valence electrons a. Group 1: one valence electron, soft b. Group 2: two valence electrons, harder 2. More valence electrons = greater hardness & strength a. Strongest bonds are in structural metals like nickel, chromium, & iron -- tend to share more electrons b. Transition metals usually form hardest & strongest materials 3. Properties of metallic bonds include high ductility, conductivity, malleability, durability D. Metal Alloys 1. Alloy: different metals combined to form one even stronger material 2. Two types of alloys: a. Interstitial i. Small holes in a metallic crystal are filled with smaller atoms ii. Example: carbon steel (98.5-99.95% Fe, 0.05-1.5% C); carbon fills the holes between atoms b. Substitutional i. Atoms of original metallic solid are replaced by other metal atoms of similar size ii. Example: brass (67-90% Cu, 10-33% Zn) Homework: ➔copy table 8-8 (p. 231) ➔p. 231 #40-44 ➔ionic compounds worksheet (activity Monday, quiz Tuesday) Chapter 8 Review: ➔p. 235: Vocabulary (all - 15 words) ➔p. 236: Concept Mapping 46 (DRAW IT) ➔p. 236: Mastering Concepts 47-59 ➔p .236: Mastering Problems 60-84 ➔p. 237: Mixed Review 85-96 ➔HONORS ONLY p. 238: Thinking Critically #97-104 ➔p. 239: Standardized Test Practice 1-10 .
Load more

Recommended publications
  • Chapter 11: Solutions: Properties and Behavior
    Chapter 11: Solutions: Properties and Behavior Learning Objectives 11.1: Interactions between Ions • Be able to estimate changes in lattice energy from the potential energy form of Coulomb’s Law, Q Q Cation Anion , in which is lattice energy, is a proportionality constant, is the charge on E = k E k QCation r the cation, Q is the charge on the anion, and r is the distance between the centers of the ions. Anion 11.2: Energy Changes during Formation and Dissolution of Ionic Compounds • Understand the concept of lattice energy and how it is a quantitative measure of stability for any ionic solid. As lattice energy increases, the stability of the ionic solid increases. • Be able to estimate changes in lattice energy from the potential energy form of Coulomb’s Law, Q Q Cation Anion , in which is lattice energy, is a proportionality constant, is the charge on E = k E k QCation r the cation, Q is the charge on the anion, and r is the distance between the centers of the ions. Anion • Be able to determine lattice energies using a Born-Haber cycle, which is a Hess’s Law process using ionization energies, electron affinities, and enthalpies for other atomic and molecular properties. • Be able to describe a molecular view of the solution process using Hess’s Law: 1. Solvent → Solvent Particles where ∆Hsolvent separation is endothermic. 2. Solute → Solute Particles where ∆Hsolute separation is endothermic. 3. Solvent Particles + Solute Particles → Solution where ∆Hsolvation can be endo- or exothermic. 4. ∆Hsolution = ∆Hsolvent separation + ∆Hsolute separation + ∆Hsolvation This is a Hess’s Law thinking process, and ∆Hsolvation is called the ∆Hhydration when the solvent is water.
    [Show full text]
  • Assembly of Multi-Flavored Two-Dimensional Colloidal Crystals† Cite This: Soft Matter, 2017, 13,5397 Nathan A
    Soft Matter View Article Online PAPER View Journal | View Issue Assembly of multi-flavored two-dimensional colloidal crystals† Cite this: Soft Matter, 2017, 13,5397 Nathan A. Mahynski, *a Hasan Zerze,b Harold W. Hatch, a Vincent K. Shena and Jeetain Mittal *b We systematically investigate the assembly of binary multi-flavored colloidal mixtures in two dimensions. In these mixtures all pairwise interactions between species may be tuned independently. This introduces an additional degree of freedom over more traditional binary mixtures with fixed mixing rules, which is anticipated to open new avenues for directed self-assembly. At present, colloidal self-assembly into non-trivial lattices tends to require either high pressures for isotropically interacting particles, or the introduction of directionally anisotropic interactions. Here we demonstrate tunable assembly into a plethora of structures which requires neither of these conditions. We develop a minimal model that defines a three-dimensional phase space containing one dimension for each pairwise interaction, then employ various computational techniques to map out regions of this phase space in which the system self-assembles into these different morphologies. We then present a mean-field model that is capable of reproducing these results for size-symmetric mixtures, which reveals how to target different structures by tuning pairwise interactions, solution stoichiometry, or both. Concerning particle size asymmetry, we find Received 19th May 2017, that domains in this model’s phase space, corresponding to different morphologies, tend to undergo a Accepted 30th June 2017 continuous ‘‘rotation’’ whose magnitude is proportional to the size asymmetry. Such continuity enables DOI: 10.1039/c7sm01005b one to estimate the relative stability of different lattices for arbitrary size asymmetries.
    [Show full text]
  • CHAPTER 20: Lattice Energy
    CHAPTER 20: Lattice Energy 20.1 Introduction to Lattice Energy 20.2 Born-Haber Cycles 20.3 Ion Polarisation 20.4 Enthalpy Changes in Solutions Learning outcomes: (a) explain and use the term lattice energy (∆H negative, i.e. gaseous ions to solid lattice). (b) explain, in qualitative terms, the effect of ionic charge and of ionic radius on the numerical magnitude of a lattice energy. (c) apply Hess’ Law to construct simple energy cycles, and carry out calculations involving such cycles and relevant energy terms, with particular reference to: (i) the formation of a simple ionic solid and of its aqueous solution. (ii) Born-Haber cycles (including ionisation energy and electron affinity). (d) interpret and explain qualitatively the trend in the thermal stability of the nitrates and carbonates in terms of the charge density of the cation and the polarisability of the large anion. (e) interpret and explain qualitatively the variation in solubility of Group II sulfates in terms of relative magnitudes of the enthalpy change of hydration and the corresponding lattice energy. 20.1 Introduction to Lattice Energy What is lattice energy? 1) In a solid ionic crystal lattice, the ions are bonded by strong ionic bonds between them. These forces are only completely broken when the ions are in gaseous state. 2) Lattice energy(or lattice enthalpy) is the enthalpy change when one mole of solid ionic lattice is formed from its scattered gaseous ions. 3) Lattice energy is always negative. This is because energy is always released when bonds are formed. 4) Use sodium chloride, NaCl as an example.
    [Show full text]
  • Calculating Lattice Energies Using the Born-Haber Cycle an Extension Activity for AP Chemistry Students
    Calculating Lattice Energies Using the Born-Haber Cycle An Extension Activity for AP Chemistry Students A particular set of equations known as a Born-Haber cycle demonstrates how chemists are able to use the first law of thermodynamics—that the energy of the universe is conserved in any chemical or physical change—to find an unknown energy value that is difficult or impossible to measure experimentally. Some energy quantities, such as the lattice energy of a mineral or the electron affinity of an atom, can be difficult to measure in the lab. Examining a Born-Haber cycle we see that there is more than one path to the formation of a substance in a particular state, and that if we use consistent definitions, an energy value that we seek can be calculated from energy values that we already know. The following exercise will help us see the way these energy values relate to one another, give us practice with their definitions and symbols, and deepen our insight to their meaning when we see them in other types of problems. Each physical or chemical change represented has: • an equation that represents a clearly defined physical or chemical change; • a definition of the particular type of energy change; • a symbol or abbreviation for the energy change equal to a value for the change in enthalpy (ΔH), the energy that is released or absorbed during the change, expressed in kJ/mol; and • a name by which that change in enthalpy is commonly known. To set up the equations in a Born-Haber cycle, cut out the cards for names, equations, defini- tions, and symbols with energy values.
    [Show full text]
  • A.P. Chemistry Quiz: Ch. 8, Bonding Lattice Energy, Bond Enthalpy, Electronegativity, and Polarity Name___
    A.P. Chemistry Quiz: Ch. 8, Bonding Lattice Energy, Bond Enthalpy, Electronegativity, and Polarity Name___________________________________ MULTIPLE CHOICE. Choose the one alternative that best completes the statement or answers the question. The diagram below is the Born-Haber cycle for the formation of crystalline potassium fluoride. 1) Which energy change corresponds to the negative lattice energy of potassium fluoride? A) 5 B) 1 C) 6 D) 4 E) 2 2) Which energy change corresponds to the electron affinity of fluorine? A) 2 B) 5 C) 6 D) 4 E) 1 3) Which energy change corresponds to the first ionization energy of potassium? A) 2 B) 4 C) 3 D) 5 E) 6 4) Given the electronegativities below, which covalent single bond is most polar? Element: H C N O Electronegativity: 2.1 2.5 3.0 3.5 A) C-H B) O-C C) O-H D) N-H E) O-N 5) Of the molecules below, the bond in __________ is the most polar. A) HBr B) HCl C) HF D) HI E) H2 A-1 6) Using the table of average bond energies below, the DH for the reaction is __________ kJ. Bond: C”C C-C H-I C-I C-H D (kJ/mol): 839 348 299 240 413 A) +160 B) -63 C) +63 D) -160 E) -217 7) Of the bonds C-N, ChN, and C”N, the C-N bond is __________. A) strongest/longest B) weakest/longest C) intermediate in both strength and length D) strongest/shortest E) weakest/shortest A-2 Answer Key Testname: QUIZ_BONDING_CH_08.TST MULTIPLE CHOICE.
    [Show full text]
  • Solutions Thermodynamics Dcians
    Solutions Thermodynamics DCI Name _____________________________ Section ______ 1. The three attractive interactions which are important in solution formation are; solute-solute interactions, solvent-solvent interactions, and solute-solvent interactions. Define each of these interactions and describe their importance in determining whether a particular solute-solvent pair will form a homogeneous mixture or a heterogeneous mixture. Solute-solute interactions are the intermolecular attractions between solute particles. Solvent-solvent interactions are the intermolecular attractions between solvent particles. Solute-solvent interactions are the intermolecular attractions between a solute particle and a solvent particle. If the intermolecular attractions between solute particles are different compared to the intermolecular attractions between solvent particles it is unlikely dissolution will occur. Similar intermolecular attractions between solute particles and between solvent particles make it more likely a solution will form when the two are mixed. Shown below are two beakers. One contains liquid The two liquids form a water and the other contains liquid carbon heterogeneous mixture. The tetrachloride. For simplicity a circle is used to drawing should show the carbon represent the molecule. Complete the beaker on the tetrachloride on the bottom and right to illustrate the result of mixing the two liquids. the water on top, due to the difference in densities. Shown below are two beakers. One contains liquid The two liquids form a solution. hexane and the other contains liquid carbon The two components are tetrachloride. For simplicity a circle is used to uniformly distributed represent the molecule. Complete the beaker on the throughout the solution. right to illustrate the result of mixing the two liquids.
    [Show full text]
  • Stoichiometric Control of DNA-Grafted Colloid Self-Assembly
    Stoichiometric control of DNA-grafted colloid self-assembly Thi Voa, Venkat Venkatasubramaniana, Sanat Kumara,1, Babji Srinivasanb, Suchetan Pala,c, Yugang Zhangc, and Oleg Gangc aDepartment of Chemical Engineering, Columbia University, New York, NY 10027; bDepartment of Chemical Engineering, Indian Institute of Technology, Gandhinagar 382424, India; and cCenter for Functional Nanomaterials, Brookhaven National Laboratory, Upton, NY 11973 Edited by Monica Olvera de la Cruz, Northwestern University, Evanston, IL, and approved March 3, 2015 (received for review October 31, 2014) There has been considerable interest in understanding the self- reinforces the commonly held view that the desired crystal will only assembly of DNA-grafted nanoparticles into different crystal struc- form if the right stoichiometry is used. tures, e.g., CsCl, AlB2, and Cr3Si. Although there are important Similarly, several coarse-grained models also allow for the exceptions, a generally accepted view is that the right stoichiometry rapid determination of the crystal structure formed by binary of the two building block colloids needs to be mixed to form the mixtures of DNA-grafted colloids (23–25). The complementary desired crystal structure. To incisively probe this issue, we combine contact model (CCM) is a canonical example of such a coarse- experiments and theory on a series of DNA-grafted nanoparticles at grained description (23). Inspired by experiments, this theory varying stoichiometries, including noninteger values. We show that assumes that only attractive energetics (derived from DNA stoichiometry can couple with the geometries of the building blocks base pairing) determine the standard-state free energy. Thus, to tune the resulting equilibrium crystal morphology. As a concrete the CCM ignores any repulsive interactions.
    [Show full text]
  • Arxiv:Cond-Mat/0602428V1
    Ground-State Shapes and Structures of Colloidal Domains Jianlan Wu and Jianshu Cao∗ Department of Chemistry, Massachusetts Institute of Technology Cambridge, MA 02139 (Dated: First submitted August 21, 2004) Abstract In charged colloidal suspensions, the competition between square-well attraction and long-range Yukawa repulsion leads to various stable domains and Wigner supercrystals. Using a continuum model and symmetry arguments, a phase diagram of spheres, cylinders, and lamellae is obtained as a function of two control parameters, the volume fraction and the ratio between the surface tension and repulsion. Above a critical value of the ratio, the microphase cannot be supported by the Yukawa repulsion and macroscopic phase separation occurs. This finding quantitatively explains the lack of pattern formation in simple liquids because of the small hard sphere diameter in comparison with the size of macromolecules. The phase diagram also predicts microphase separation at zero value of the ratio, suggesting the possibility of self-assembly in repulsive systems. PACS numbers: 82.70.Dd, 05.65.+b, 64.70.-p arXiv:cond-mat/0602428v1 [cond-mat.soft] 17 Feb 2006 ∗ Email: [email protected] 1 Microphase separation is ubiquitous in soft matter systems[1, 2, 3, 4, 5, 6, 7, 8]. For example, microphase separation in block copolymers results from mixing of two or more chemically different polymer segments[3]. Competition of hydrophobic and hydrophilic head- groups of amphiphiles leads to micelles in the water solution[4]. The self-assembly processes in charged colloidal suspensions and protein solutions can lead to the formation of stable domains and Wigner supercrystals (see Fig.
    [Show full text]
  • Computational Modelling of Solvent Effects in a Prolific
    Faraday Discussions Cite this: Faraday Discuss.,2018,211,383 View Article Online PAPER View Journal | View Issue Computational modelling of solvent effects in a prolific solvatomorphic porous organic cage† David P. McMahon,a Andrew Stephenson,b Samantha Y. Chong, b Marc A. Little, b James T. A. Jones,‡b Andrew I. Cooper *b and Graeme M. Day *a Received 16th February 2018, Accepted 22nd March 2018 DOI: 10.1039/c8fd00031j Crystal structure prediction methods can enable the in silico design of functional Creative Commons Attribution 3.0 Unported Licence. molecular crystals, but solvent effects can have a major influence on relative lattice energies, sometimes thwarting predictions. This is particularly true for porous solids, where solvent included in the pores can have an important energetic contribution. We present a Monte Carlo solvent insertion procedure for predicting the solvent filling of porous structures from crystal structure prediction landscapes, tested using a highly solvatomorphic porous organic cage molecule, CC1. Using this method, we can understand why the predicted global energy minimum structure for CC1 is never observed from solvent crystallisation. We also explain the formation of three different This article is licensed under a solvatomorphs of CC1 from three structurally-similar chlorinated solvents. Calculated solvent stabilisation energies are found to correlate with experimental results from thermogravimetric analysis, suggesting a future computational framework for a priori Open Access Article. Published on 16 April 2018. Downloaded 10/2/2021 4:27:46 PM. materials design that factors in solvation effects. Introduction A fundamental goal for computational chemistry is the in silico design of func- tional materials.
    [Show full text]
  • IONIC BONDING Electrovalent Or Ionic Bond
    -By S . M . SIR IONIC BONDING Electrovalent or Ionic bond When a bond is formed by complete transfer of electrons from one atom to another so as to complete their octet or duplet and hence acquire the stable nearest noble gas like electronic configuration, then the bond formed is called ionic bond or electrovalent bond. Formation of ionic bond Atoms are electrically neutral as they possess equal number of proton and electrons. If an atom loses electron, number of protons exceeds the number of electrons and hence it becomes positively charged cation. A → A+ + e On the other hand, if an atom gains electron, number of electrons exceeds the number of protons and hence it becomes negatively charged anion. B + e → B- These oppositely charged species attract each other by electrostatic force of attraction and such electrostatic attraction is considered as ionic bond. Q.1. Explain the formation of the following ionic compounds: 1. Sodium chloride (NaCl) 2. Calcium fluoride(CaF2) 3. Magnesium oxide(MgO) 4. Magnesium Bromide(MgBr2) 5. Magnesium fluoride (MgF2) 6. Calcium Sulphide(CaS) 7. Calcium Oxide(CaO) 8. Sodium Sulphide(Na2S) Electrovalency The number of electrons lost or gained during the formation of an electrovalent linkage (ionic bond) is known as the electrovalency of the corresponding element. Example : In the formation of NaCl, the electrovalency of sodium is 1 and that of chlorine is 1. In general, electrovalency can be considered as the magnitude of charge present on the ion. -By S . M . SIR Ionic- Crystal The electrostatic field of an ion is non directional. Each positive ion attracts several negative ions around it depending upon its size (radius).
    [Show full text]
  • SOLUTIONS of EXERCISES 301 Central Role of Gibbs Energy 301:1
    SOLUTIONS OF EXERCISES 301 central role of Gibbs energy 301:1 characteristic Numerical values in SI units: m.p = 2131.3; ΔS = 7.95; ΔH = 16943; ΔCP = −11.47. 301:2 a coexistence problem 1) f = M[T,P] – N[2 conditions] = 0 ; this excludes the first option; 2) by eliminating lnP from the two equilibrium equations in the ideal-gas approximation, the invariant point should follow from [60000 + 16 (T/K) = 0]; this excludes the second option. 301:3 trigonometric excess Gibbs energy – spinodal and binodal . -1 2 RTspin(X) = (4000 J mol )π X(1−X)sin2πX; limits at zero K are X = 0 (obviously) and X = 0.75; at 0.75 Tc the binodal mole fractions are about 0.11 and 0.52 301:4 from experimental equal-G curve to phase diagram ΔΩ = 1342 J.mol-1; Xsol = 0.635; Xliq = 0.435; see also van der Linde et al. 2002. 301:5 an elliptical stability field * The property ΔG A , which invariably is negative, passes through a maximum, when plotted against temperature. The maximum is reached for a temperature which is about ‘in the middle of the ellipse’. The property ΔΩ is positive. 301:6 a metastable form stabilized by mixing The stable diagram has three three-phase equilibria: αI + liquid + β at left; eutectic; liquid + β + αII at right; peritectic; αI + β + αII (below the metastable eutectic αI + liquid + αII); eutectoid. 301:7 the equal-G curve and the region of demixing For the left-hand case, where the EGC has a minimum, the stable phase diagram has a eutectic three-phase equilibrium.
    [Show full text]
  • LATTICE ENERGI (Uo)
    pranjoto utomo Definition and Symbol • The energy given off when oppositely charged ions in the gas phase come together to form a solid. • Symbol = Uo • Signed by (-) • The greater lattice energy: – Cation-anion distance <<< – Ions charge >>> Determinations • Born-Lande equation – Crystal structure are known • Kapustinskii equation – Crystal structure are not known • Born-Haber Cycles – Based on Hess Law Born-Lande Equation N.Z.Z -.M.e 2 1 Uo x 1 - 4πεoro n N.e2 M.Z.Z - 1 Uo x x 1 - 4πεo ro n • N = Avogadro number = 6.02x1023 • Z+ = positive charge • Z- = negative charge • M = Madelung constant • e- = electron charge = 1.6021x10-19C • = 3.14159 • o = dielectric contants of a vacuum = 8.854x10-12C2m-1J-1 • n = factor for compressibility • ro = cation-anion equilibrium distance = r+ + r- Madelung Constant Lattice M Rock salt 1.7476 CsCl 1.7627 Zinc blende 1.6381 Fluorite 2.5194 Compressibility factor Element n He 5 Ne 7 Ar, Cu 9 Kr, Ag 10 Xe, Au 12 Na [ Ne] n 7 7 9 11 10 n 8 - 2 17 Cl [18 Ar] n 9 • Substituting the constants, yields: - -4 M.Z .Z 1 Uo 1.390x10 x x 1 - ro n Kapustinskii Equation 120200..Z .Z - 34.5 Uo x 1 - ro ro = number of ions per molecule NaCl = 2 Z+ = positive charge Z- = negative charge ro = cation-anion equilibrium distance = r+ + r- Comparison • Lattice energy of NaCl – Experiment = -755 kJ.mol-1 – Born-Lande = -770 kJ.mol-1 – Kapustinskii = -753 kJ.mol-1 Born Haber Cycle Unbonded gaseous ions •Breaking bonds •Forming gases H H Lattice energy •Forming (+) & (-) 1 2 ions Metal + non-metal
    [Show full text]