1) Would copper be considered a hard or soft Lewis acid? What about zinc? Describe what makes these metals hard or soft

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1) Would copper be considered a hard or soft Lewis acid? What about zinc? Describe what makes these metals hard or soft.

2.What trends would you predict for EDTA-metal complex Kf and ΔG values when:

a) going down a row of transition metals of the same oxidation state?

b) of the same transition metal with increasing oxidation state? (For example, Iron can exist as Fe(II), Fe(III), Fe(IV), and higher oxidation states)

3. Discussion (formal report style): Compare and contrast the concepts of Lewis and Brønsted-Lowry acids and bases. Support these concepts with your data over the last two weeks and with research. Cite your sources.

4. Select and summarize a peer review article detailing a method for calculating equilibrium concentrations in a particular metal-chelator reaction.

Lewis acids and bases Intro/Background At the core of general acid-base theories, there is he concep of ran fer of one entity to another. In the case of Brønsted-Lowry acid base theory, the transferred entity is the proton, and their transfer processes are widely exploited in chemical methods. In addition to proton transfer, electron pairs can also be donated/shared between donors and acceptors. In this case, an atom/molecule, which can accept a pair of electrons, is referred to as a Lewis acid while the entity donating the pair of electrons is referred to as a Lewis base. The most prevalent Lewis acid-base systems involve metal complexes with vacant nd-orbitals, as well as electron deficient systems including B, Si, Ge, etc. The Lewis acid base theory is also known as the Hard-Sof acid ba e heor (HSABT) i h hardne defined as small highly charged species that are weakly polarizable. In con ra , of ne i de cribed b pecie ha are larger and highly polarizable species with low charge. In the case of metal chemistry, Lewis acid-base theory provides a means to describe he abili of ligand o coordina e o me al ion a ell a he abili of metal-ligand complexes. Recall that coordination bonds are the donation/acceptance of lone pairs of electrons to empty metal orbitals to another atom; the definition of Lewis acid-base chemistry! The stability of Lewis acid-base complexes can be described through the stability constant which is analogous to the Ka that describes Brønsted-Lowry acidities. Many metal ions can coordinate (bind) many Lewis base ligands in a step-wise fashion according to: M + L = ML ML + L = ML2 ML2 + L = ML3 MLn-1 + L = MLn K1 = [ML]/[M][L] K2 = [ML2]/[ML][L] K3 = [ML3]/[ML2][L] Kn = [ML4]/[ML3][L] and for which the free energy for this dissociation reaction is given by: ?G0 = -RTLnKn Eq. 5. where R is the gas constant and T is the temperature in Kelvin. Figure 1: Left- EDTA. Right- Typical EDTA-transition metal complex. Eq. 1 Eq. 2 Eq. 3 Eq. 4 The ethylene diamine tetraacetic acid ligand (EDTA, a Lewis base) provides a mechanism to probe the Lewis acid strength of various metal complexes (Figure 1). This multi-dentate ligand (multiple Lewis base sites on a single ligand) forms complexes with metals of the form: Mn+ + Y4- = MYn-4 Eq. 6 K = [MYn-4]/[Mn+][Yn-4] Eq. 7 where Y4- is the EDTA ligand. In this experiment, the Lewis acid strength for two different metals will be examined by measuring the stability constants for the metal-EDTA complex. Of specific interest is the how the Lewis acid strength of the metal ion correlates with: 1) atomic/ionic radii, 2) oxidation state, 3) electron configuration, and 4) polarizability. You will determine the stability constants of both a Cu2+- and Zn2+-EDTA complex as well as the corresponding free energy. Using th data, you will compare the relative Lewis acidities of both metals and describe differences in terms of the metal ion properties described above. Chemicals CuSO4 ZnSO4 EDTA Murexide Instruments Buret Experimental You and your group will determine the concentration and ultimately the free energy of Cu(II) and Zn(II) using the hexadentate chelating ligand, EDTA. You will be given a ~100 mM standardized solution of EDTA (check the label for exact concentration prior to beginning). You will need to weigh ~0.1 g per metal salt sample. Dissolve the metal salt into 50 mL of DI water. Add 2-3 drops of murexide to your metal. After properly preparing your buret with EDTA, titrate your metal with EDTA. Repeat for a total of 3 titrations per metal. 1. Approximately how much EDTA will you need for each titration assuming the metalal ma e are acc ra e? 2. What type of indicator is murexide and what color will its endpoint be? Lewis’s acid and bases results ZnSO4 Trail 1 0.1019g znso4 Trail 2 0.1266g znso4 Ph:11.83 Ph:11.85 6.2ml EDTA 6.4ml EDTA CuSO4 Trail 1 0.1985g cuso4 Trail 2 0.1995g cuso4 Ph:11.56 Ph:11.81 5.7ml EDTA 5.5ml EDTA Post-Lab Questions 1. a) Using the equations given in the introduction and your titration results, what are the Kf of the EDTA-Copper complex? Show calculations. b) Using the equations given in the introduction and your titration results, what are the Kf of the EDTA-Zinc complex? Show calculations. 2.What are the corresponding ΔG values for the EDTA-metal complexes? Show your calculations 3.Would copper be considered a hard or soft Lewis acid? What about zinc? Describe what makes these metals hard or soft. 4.What trends would you predict for EDTA-metal complex Kf and ΔG values when: a) going down a row of transition metals of the same oxidation state? b) of the same transition metal with increasing oxidation state? (For example, Iron can exist as Fe(II), Fe(III), Fe(IV), and higher oxidation states) 5. Discussion (formal report style): Compare and contrast the concepts of Lewis and BrønstedLowry acids and bases. Support these concepts with your data over the last two weeks and with research. Cite your sources. 6. Select and summarize a peer review article detailing a method for calculating equilibrium concentrations in a particular metal-chelator reaction. Received: 26 July 2017 | Revised: 16 October 2017 | Accepted: 17 October 2017 DOI: 10.1111/apha.12988 REVIEW A practical guide to the preparation and use of metal ionbuffered systems for physiological research F. Neumaier | S. Alpdogan | J. Hescheler | T. Schneider Institute for Neurophysiology, University of Cologne, Cologne, Germany Correspondence F. Neumaier, Institute of Neurophysiology, University of Cologne, Cologne, Germany. Email: felix@neumaier-net.de Funding information This work was financially supported by the Deutsche Forschungsgemeinschaft (DFG, SCHN 387/21-1). Abstract Recent recognition that mobile pools of Zn2+ and Cu2+ are involved in the regulation of neuronal, endocrine and other cells has stimulated the development of tools to visualize and quantify the level of free trace metal ions. Most of the methods used to measure or control loosely bound metals require reference media that contain exactly defined free concentrations of the target ions. Despite the central importance of proper metal ion buffering, there is still a lack of international standards and beginners in the field may have difficulties finding a coherent description of how to prepare trace metal ion buffers, especially when experiments are to be performed in multimetal systems. To close this gap, we provide a guide for the design, preparation and use of metal ionbuffered systems that facilitate immediate application under physiologically relevant ionic conditions. Thermodynamic and kinetic concepts of chemical speciation as well as general protocols and specific examples are outlined for the accurate preparation of single- and dual-metal ion buffers. In addition, experiments have been performed with FluoZin-3 to illustrate that metal ion-buffered systems are required for reliable preparation of nanomolar Zn2+ solutions and that dual-metal ion buffers can be used to calibrate suitable fluorescent Zn2+ sensors in the presence of millimolar Ca2+ concentrations. Together, the information provided should sensitize readers to the many potential pitfalls and uncertainties that exist when working with physiologically relevant concentrations of trace metal ions and enable them to formulate their own metal ion buffers for most in vitro applications. KEYWORDS endogenous mobile zinc and copper pools, fluorescent detection of trace metals, loosely bound or free trace metals, polyaminopolycarboxylate chelator, thermodynamic stability constants, tricine 1 | INTRODUCTION It is increasingly recognized that pools of loosely bound zinc (Zn2+) and copper (Cu2+) ions exist in the brain, which have been implicated in synaptic transmission and linked to several pathophysiological conditions. These and other findings have led to an impressive growth in attention and stimulated the development of new tools for studying Acta Physiologica. 2018;222:e12988. https://doi.org/10.1111/apha.12988 trace metal biology. Physiologically relevant concentrations of loosely bound metal ions span several orders of magnitude. They range from less than one Cu2+ molecule per cell1 and picomolar Zn2+ in the intracellular compartment2,3 to nanomolar extracellular levels, which may transiently rise to (high) micromolar concentrations.4,5 This wide range of very low concentrations provides many sources of potential error and necessitates the use of wileyonlinelibrary.com/journal/apha © 2017 Scandinavian Physiological Society. Published by John Wiley & Sons Ltd | 1 of 18 2 of 18 | suitable metal ion buffers. For example, most methods used to measure or control the level of a given trace metal require reference media that contain exactly defined free concentrations of this ion. The central importance of proper metal ion buffering is well established, but there is still a lack of international standards for physiological Zn2+, Cu2+ and other metal ions,6,7 such as are available for calibration of pH electrodes.8 Moreover, while buffers containing specified free Ca2+ levels based on calculated values are available from several suppliers, the preparation of trace metal ion buffers is still much more empirically. Unfortunately, beginners in the field may have difficulties finding a coherent description of how to prepare trace metal ion buffers themselves, especially when experiments are to be performed in the presence of other divalent cations. As a consequence, many of the solutions reported in the literature are not adequately buffered, and significant differences may exist between assumed and actual free metal ion concentrations.9 Aim of this article is to provide an introduction into the theory and practice of trace metal ion buffering, with a focus on the physiologically relevant range of free Zn2+ and Cu2+ concentrations. Particular emphasis is placed on the preparation and use of dualmetal ion buffers, which allow for the control of free Zn2+ or Cu2+ at a fixed concentration of Ca2+ or Mg2+. To this end, section 2 will briefly familiarize the reader with the concept of metal ion buffering, summarize the general requirements for construction of well-defined buffer systems, and discuss problems associated with the calculation of free metal ion concentrations and with the lack of a simple method for their verification. The third section deals with practical aspects, such as the choice of reagents and labware or general sample handling techniques, and provides a step-by-step guide on the preparation of single- and dual-metal ion buffers with minimal experimental uncertainty. Calibration of the fluorescent probe FluoZin-3 with nanomolar Zn2+ in the presence of millimolar Ca2+ concentrations is used to exemplify both the application of dual-metal ion buffers and the consequences of improper trace metal ion buffering. 2 | THEORY OF TRACE METAL ION BUFFERING A thorough understanding of chemical speciation is crucial for all experiments involving physiologically relevant free trace metal concentrations, but the complexity of the topic prevents an exhaustive treatment in this article. Instead, we will provide a brief introduction into the topic and then focus on aspects relevant for the design of metal ion-buffered systems and for estimation of free metal ion concentrations in complex solutions. NEUMAIER ET AL. 2.1 | Chemical speciation and metal ion buffering In complex media such as cytoplasm or extracellular fluid, trace metals become bound by a range of ligands, which include inorganic anions, small organic molecules like amino acids and large biomolecules like proteins. As some of them bind metal ions very tightly, the amount that remains thermodynamically and kinetically accessible (ie the free or loosely bound pool) can be orders of magnitude below the total concentration. The biological effects of metal ions are thus greatly affected by their chemical speciation, which is defined to be the distribution among all possible (free and ligand-bound) species in solution. Even under more defined conditions, like in artificial saline solutions, free metal ion concentrations can markedly deviate from the nominal concentration. This is because at physiologically relevant (ie pico- to micromolar) concentrations, processes like metal ion hydrolysis, surface adsorption or binding by buffer constituents may significantly influence the free metal ion concentration.10 To circumvent this problem and to reduce the impact of contamination with unwanted metal ions, free Zn2+ and Cu2+ concentrations should be buffered with appropriate trace metal chelators. By providing a reservoir of metal complex to replenish ions that have been removed from the system, these ligands can effectively maintain free metal ion concentrations at a constant level. Numerous naturally occurring or synthetic molecules can act as ligands, and there is extensive literature on metal ion-ligand interactions. To be useful for metal buffering, a chelator should be well-characterized in terms of its binding properties, exhibit sufficient specificity for the cation to be buffered and sufficient buffering capacity for the desired concentration range. Additional considerations that are of importance when working in biological systems are chelator mobility across the cell membrane and possible effects not related to metal- ion buffering. Lipophilic chelators (like TPEN, cupral or dithizone) and ionophores (like pyrithione) should generally be avoided, as they can cross the membrane and interfere with intracellular trace metal homeostasis. The same applies to certain naturally occurring ligands like citrate, which may exert metal ion buffering-independent effects.11 Most of the chelators useful for biological research belong to the group of glycine-derived polyaminopolycarboxylates (PAPCs). Thanks to their relevance for a wide variety of environmental, chemical and medical applications, numerous PAPC ligands have been described and there is a large literature on the topic. Here, we will focus on ethylenediaminetetraacetic acid (EDTA), ethyleneglycol-bis(aminoethylether)N,N,N0 ,N0 -tetraacetic acid (EGTA) and N-(2-hydroxyethyl) ethylenediamine-N,N0 ,N0 -triacetic acid (HEDTA), three of the most widely used PAPCs which have been extensively NEUMAIER | ET AL. (monoaminopolycarboxylic) nitrilotriacetic acid (NTA), can interact with metal ions in multiple stoichiometries (ie multiple ligand molecules can bind to a single metal ion). The same applies to N-[tris(hydroxymethyl)methyl] glycine (tricine), a glycine-derived zwitterionic proton buffer and tridentate ligand for Zn2+ and Cu2+ (Figure 1B), which exhibits very low affinity for Ca2+ and Mg2+. As described in more detail in sections 2.2 through 2.4, binding in multiple stoichiometries can complicate speciation modelling and the construction of well-defined buffer systems.12 Depending on the cations present, use of tricine-based buffers can be further complicated by the lack of comprehensive thermodynamic and/or kinetic data on its metal ion binding properties (but see section 2.4). As it remains one of the few compounds that can effectively maintain nano- to micromolar free Zn2+ and Cu2+ concentrations in the presence of physiological Ca2+ or Mg2+ levels however, tricine has been widely used for Zn2+ and Cu2+ buffering in biological systems and will be discussed in this article as well. T A B L E 1 Protonation constants for selected trace metal chelators Protonation constants Hx (at 20°C & I = 0.1 mol/L) and enthalpy values DHx [kcal/mol] H1 |DH1 H2 |DH2 H3 13 10.3 |5.6 6.2 |4.2 EGTA13 9.5 |5.9 8.9 |5.8 HEDTA13 9.9 |6.6 5.4 |3.1 8.0 |7.6 2.4 |1.4 EDTA Tricine 98 |DH3 H4 2.7 |1.5 2.0 |0.3 2.7 |2.6 1.9 |0.4 2.6 |1.1 |DH4 characterized in terms of their protonation (Table 1) and metal-binding constants (Table 2). They are representative for most other members of the group in that their anion species act as strong hexadentate ligands for Zn2+ or Cu2+, giving rise to 1:1 molar complexes where the metal ion is coordinated by six donor groups of the chelator (Figure 1A). Ligands of lower denticity (ie with less donor groups), such as the closely related T A B L E 2 Stability constants for selected 1:1 metal ion-ligand complexes Stability constants logKML (at 20°C at I = 0.1 mol/L) and enthalpy values DHx [kcal/mol] Ca2+ |DH Mg2+ |DH Zn2+ |DH 2.2 | Estimation of free metal ion concentrations Cu2+ |DH EDTA 13 10.7 |6.1 8.7 |3.4 16.6 |4.7 18.9 |8.2 EGTA 13 10.9 |8.4 5.2 |5.2 12.7 |4.3 17.8 |11.0 HEDTA 8.2 |6.5 6.9 |3.4 14.7 |8.4 17.5 |9.4 Tricine99,100 2.4 |n.a. |n.a. |7.9 |9.8 13 1.2 5.2 – (A) – CO2 N CO2– N N The chemical equilibria between ligands (L), metal ions (M) and the corresponding metal complexes (ML) can be described in terms of thermodynamic stability constants (Figure 2). They provide a measure for the strength of metal ion-ligand interactions and can be used to estimate OH CO2 N 7.8 – CO2 O CO2– O CO2 2 CO2 CO2 CO2– HN – N HEDTA Glycine: – N – CO2 – CO2 EDTA 3 of 18 – (B) CO2 – HO HO EGTA CO2– NH Tricine HO O ML N M O– O– O ML O OH 2+ O– N O– HN M O– OH O O O HO OH OH OH O– M NH O 2+ OH HO ML2 O– HN 2+ HO F I G U R E 1 Structure of glycine-derived trace metal chelators and some of their metal complexes. A, The deprotonated species of polyaminopolycarboxylic acids like EDTA are strong hexadentate ligands for Zn2+, Cu2+ and most other divalent cations. Because all amino and carboxylate donor groups (indicated in orange) are coordinated to the metal ion, these ligands form very stable 1:1 complexes, in which the metal is surrounded by a total of five chelate rings. B, The zwitterionic proton buffer tricine acts as a lower affinity tridentate ligand for Zn2+ and Cu2+ that can form both 1:1 and 1:2 metal ion-ligand complexes containing two or four chelate rings respectively 4 of 18 | NEUMAIER the composition of a system at equilibrium. Published constants have been determined at specific conditions for temperature and ionic strength (usually 20-25°C and I = 0.1 mol/L) and are typically expressed on a log scale, where one unit change corresponds to a 10-fold change in the respective constant. A number of critically evaluated compilations (like the NIST,13 IUPAC14 or JESS15 stability constant databases, reviewed in detail in16,17) are available, which provide access to internally consistent sets of thermodynamic constants for thousands of metal ion-ligand reactions. To account for differences in the experimental conditions, they can be converted to apparent or conditional constants, provided that the molar enthalpy changes for the reactions are known (for temperature correction) and non-thermodynamic assumptions are used to calculate the relevant single ion activity coefficients (for ionic strength correction). In principle, conditional constants can be used to solve, by manual (algebraic) techniques, the complexation equilibria in simple systems.12 As illustrated in Figure 2 however, additional reactions must usually be taken into account. When protonated forms of the ligand (HxL) do not significantly contribute to complex formation (ie if there is competition for binding between metal ions KH1 (A) +H L KML +M KH2 HL +H Stability constant: KML = [ML] / ([M][L]) ML (B1) M+L+H M + HL ML + H M + H2L (B3) ML + L M + L2 M + ML +H KH4 H3L +H MHL MHL MHL MH2L KMHL = [MHL] / ([M][L][H]) KMHL = [MHL] / ([M][HL]) KMHL = [MHL] / ([ML][H]) KMH2L = [MH2L] / ([M][H2L]) Formation of hydroxo species MOHL MOHL KMOHL = [MOHL] / ([M][L][OH]) KMOHL = [MOHL] / ([MOH][L]) Deviation from 1:1 stoichiometry ML2 ML2 M2L H4L Protonation constants: KH1 = [HL] / ([L][H]) KH2 = [H2L] / ([HL][H])