Effect of basicity of ammine ligands on thermodynamic stability of Ni (II)-ammine complexes

How does the thermodynamic stability, measured in terms of log k (where k is the thermodynamic stability constant) of a Nickel (II) ammine complex produced by reacting Ni(II) with excess of ligand depends on the basic strength (expressed in terms of pk_b) of the ammine ligand used, determined using spectrophotometry?

The moment, I went through the HL part of Topic-3 (Periodic Table), I was fascinated to know about the co- ordination complexes as they find immense use whether it is cosmetics or pharmaceuticals or even industries. Though the ligands EDTA, ethylene diamine was introduced as a part of the course, I always wanted to delve more to understand the way these ammine ligands interact with the central metal ion and form the bonds. The immediate question that bothered me was – do ligands have dative bonds with the metal ion? or Is it just a strong electrostatic interaction between the transition metal ion and the ligands that donates the lone pair? This question became more interesting when I understood how the stability of the complex is an issue while synthesizing them. Several drugs used are metal ion complexes and they need to be designed in such a manner that they must decompose only at a certain part of the body and not anywhere else. Thus, it is imperative to understand the stability of the complex made and choose the ligands accordingly. What factors of the ligands determines the stability of the complex they form? Further research led me to know that among various factors the ability of the ligand to donate lone pairs or in other wordsthe basicity is an important factor. As the ammine ligands of various types with different basic strengths was easily procurable, I chose to use multiple ammine ligands of different basicity and elucidate how they could impact the stability of the metal ion complex they produce.

Ni is a transition metal from the first row of d block. At an oxidation state of +2, the Ni(II) ion shows an electronic configuration of [Ar]3d⁸ . It can react with ammine ligands (ligands that contains the ammine – NH₂ group) to form metal-ammine complexes. The ammine ligands used in this investigation are of the type X- NH₂ where X = H for ammonia (NH₃), X = OH for hydroxyl amine (OH-NH₂), X = H₂N-CH₂-CH₂ for ethan- 1,2-diamine (H₂N-CH₂-CH₂-NH₂), X = C₆H₅ for amino benzene (C₆H₅-NH₂) and X = NH₂ for hydrazine (H₂N- NH₂). All of these ligand act as Lewis base due to the presence of lone pair on N atom. According to Valence bond theory, the ligands form a dative bond with the central metal ion by donating lone pairs to the empty d subshells of the central metal ion while according to the Crystal Field theory, the interaction between a metal ion and a ligand is considered as a strong electrostatic force of attraction between point charges of opposite nature and is thus simply an ionic interaction. Ni (II) can show both co-ordination number 4 in square planar complexes and 6 in octahedral complexes. This investigation deals with all octahedral complexes and thus the metal ion displays a co-ordination number of 6 showing six metal ligand bonds. In octahedral complexes, the Ni(II) shows an electronic state of t⁶₂e²_g. All the metal-ammine complexes are low spin complexes and does not involve any pairing of ligand compromising the CFSE (Crystal Field Stabilization energy).

Ni²⁺ (aq) + 6 NH₃ (aq) ------→ [Ni(NH₃)₆] ²⁺ (aq)

Hexaammine Nickel(II) ion

Ni²⁺ (aq) + 6 H₂N-OH (aq) --------→ [Ni(NH₂OH)₆]²⁺ (aq)

Hexahydroxylaminonickel(II) ion

Ni²⁺ (aq) + 3 H₂N-CH₂-CH₂-NH₂ (aq) --------→ [Ni(H₂N-CH₂-CH₂-NH₂)₃])²⁺ (aq)

Trisethan-1,2-diammine nickel(II)

Ni²⁺ (aq) + 6 C₆H₅NH₂ (aq) -------→ [Ni(C₆H₅NH₂)₆] ²⁺ (aq)

Hexaaminobenzenenickel(II) ion

Ni²⁺ (aq) + 3 H₂N-NH₂ (aq) -------→[Ni(H₂N-NH₂)₃] ²⁺

Hexahydrazinonickel(II) ion

All the complexes formed are octahedral and water soluble.

The thermodynamic stability constant is a measure of the ability of the complex to disassociate to separate out the ligands and the metal ion. Higher the thermodynamic stability, more is the un-reactivity of the complex towards any disassociation or ligand exchange reaction. The discussion below aims to deduce a mathematical formula to calculate the magnitude of thermodynamic stability constant using the value of molar concentration of Ni(II) ion at equilibrium.

The generalized equation is:

Ni²⁺ (aq) + n L --------→ [MLn]²⁺ (aq)

Stability constant (K) = \(\frac{[[ML_n]^{2+}]}{[Ni^{2+}][L]^n}\)

The table below narrates the molar concentration of the reactants and products in the reversible reaction of the complex formation at various stages:

k = \(\frac{(0.01-x)}{x(0.10-0.01n+nx)^n}\)

Taking logarithm on both sides,

Using the formula log \((\frac{a}{b})\) = log (a) - log (b)

log(k) = log(0.01 − x) − log[x (0.10 − 0.01 n + nx)ⁿ]

using the formula log(ab) = log(a) + log (b) and the formula : log a^b  = b log a

log(k) = log(0.01 -x) - log (x) + n log(0.10 -0.01n + nx) ........(equation -1)

log(k) = stability constant

n = moles of ligand that reacts with 1 mole of the metal ion

x = molar concentration of Ni(II) ions at equilibrium

As the value of pk_b decreases, the ammine becomes more basic and thus it has higher tendency to donate the lone pair from N atom and can form a stronger bond with the metal ion. Thus, it is predicted that with the decrease in pk_b, the thermodynamic stability of the complex increases. There is a negative correlation between the pk_b and the thermodynamic stability of the complex.

The purpose of the investigation is to use a variety of amine bases (that contains the group-NH₂) which varies in basicity and check how the basicity impacts the stability of the complex they form with Ni(II). The ligands used are – ammonia (NH₃), hydroxyl amine (HO-NH₂), ethan-1,2-diamine (NH₂-CH₂-CH₂-NH₂), amino benzene (C₆H₅-NH₂) and hydrazine (NH₂-NH₂). pKb is a measure of the basic strength. Lower the value of pKb, higher the pH of the aqueous solution of the base and stronger the basicity. For a more reliable and generalized conclusion, the bases have been chosen in such a way that they cover various electronic and structural effects like the electron withdrawing effect of the benzene ring in aminobenzene that reduces the basicity, negative electron withdrawing effect (-I) of the OH group in OH-NH₂ that reduces the basicity. To quantify the basicity, the pK_b values have been used.

The values of pK_b has been taken from three different sources and a mean value has been used. The sources used are: Source-1-pubchem.com³ , Source-2-chem.libretexts.org⁴ and Source-3:chemguide.co.uk⁵ . All of these sources are academic databases and thus are reliable.

The thermodynamic stability of the complex measured as log k, where k is the stability constant will be computed using equation (1). The value of x (molar concentration of Ni(II) at equilibrium) will be determined using the values of absorbance of the solution at equilibrium using a spectrophotometer. The stability of a complex can also be expressed as kinetic stability but that is more of a qualitative indicator of how fast the complex can be made while thermodynamic stability indicates how difficult or easy it is to disassociate the complex. As, the investigation aims to deal with the stability of the complex when they are used in a particular chemical reaction, thermodynamic stability has been measured instead of kinetic stability.

The ammine compounds used are potentially harmful and corrosive in nature. Exposure to these chemicals
may cause allergic reactions, respiratory disorders and even nausea.

• A laboratory coat was worn always.
• To restrict inhaling the fumes and smokes, a safety masks was used and gloves as well.
• All solutions were prepared carefully under strict supervision of a laboratory technician to avoid
  spillage.

An attempt has been made to minimize the use of consumable resources. For example, dilute solutions have
been used in the investigation to use least possible amount of chemicals.

All waste chemicals were diluted and thrown into a safety bin for disposal.

• A 100 cm³ glass beaker was taken.
• A watch glass was taken and placed on a top pan digital mass balance.
• The balance was tared to a reading of 0.00 ± 0.01 g.
• Nickel(II) chloride, green solid was transferred to the watch glass using a spatula until it reads 0.13 ± 0.01 g.
• The weighed solid was transferred to a 100 cm³ glass beaker.
• Distilled water was added till the mark.
• A glass rod was used to dissolve the solid.
• Two cuvettes were taken and one of them was filled with distilled water while the other one was filled with the 0.10 moldm^-3 Ni(II) solution.
• The distilled water was used as a sample and the Ni(II) solution was used as blank.
• The absorbance of the Ni(II) sample solution was recorded in the range of 400 nm to 700 nm at an
  interval of 5 nm.

Note: Nickel(II) chloride is a green solid. When dissolved in water, the chloride ions are replaced by water
and thus hexaaquanickel (II) complex ion is formed.

NiCl₂ (s) + 6H₂0 (l) ------→ [Ni(H₂O)₆]²⁺ (aq) + 2Cl^- (aq)

Moles of  NiCl₂ added = Moles of  [Ni(H₂O]₆]²⁺

Mass of NiCl₂ added = 0.13 g

Moles of NiCl2 added = \(\frac{mass}{molar\ mass}\) = \(\frac{0.13}{129}\) ≅ 0.001

Moles of [Ni(H₂O)₆] ²⁺ = 0.01

Molar concentration of [Ni(H₂O)₆] ²⁺ = \(\frac{mass}{Volume}\) = 0.01 mol dm^-3

This will be a smooth line scatter plot and the maxima of the curve will be marked. A perpendicular will be drawn from the maxima to the x axes to determine the wavelength at which the absorbance of [Ni(H20)6]2+ complex is maximum. The complex is green in color. Thus, according to the color wheel, this complex should absorb red color and the wavelength of maximum absorbance is supposed to lie within the range of 640 nm to 700 nm. The value reported in literature is 670 nm⁶.

• A 100 cm³ glass beaker was taken.
• On a watch glass, 0.13 ± 0.01 g (0.001 moles) of NiCl₂ was weighed using a top pan digital mass balance.
• The weighed solid was transferred to the same beaker.
• Distilled water was added till the mark of 100 cm³
• A glass rod was used to stir the solution.
• A cuvette was filled with the solution (to be used as sample) and another one was filled with (to be used as blank).
• The absorbance of the sample was recorded at (insert the wavelength entered from Figure - 6).
• Steps 6 and 7 were repeated for two more times.
• All of the above steps were repeated using 0.26 ± 0.01 g (0.002 moles), 0.39 ± 0.01 g (0.003 moles), 0.52 ± 0.01 g (0.004 moles) and 0.65 ± 0.01 g (0.05 moles) of NiCl₂.

This will be a scatter plot. A linear trend line passing through the origin will be drawn. The equation of the trend line will give a mathematical relationship between absorbance and molar concentration. For example, of the equation is y = mx ; y is absorbance and x is the molar concentration.

Thus, molar concentration (y) = \(\frac{absorbance\ (x)}{gradient\ of\ Graph - 2\ (m)}\) mol dm^-3 ..........(equation - 2)

All the ligands used are in the physical state of liquid at room temperature. The number of moles of the ligand used is 1.00 moles in all cases and it is taken in excess of the metal ion so that the metal ion is the limiting reactant and is completely consumed. The table below shows the volume of ligand to be used.

Formula:

Volume of ligand to be used = \(\frac{mass}{density}\) = \(\frac{moles\ ×\ molar \ mass\ (in\ g)}{density\ at\ room \ temperature\ (in \ g\ cm^{-3} )}\) cm³

*All values of density are taken from the pubchem.ncbi.nlm.nih.gov database and as this is a government website, the data can be relied upon.

• A 100 cm³ glass beaker was taken.
• Following this, 1.29 ± 0.01 g (0.01 moles) of NiCl₂ (green solid) was weighed on a watch glass using a top pan digital mass balance.
• After that, 50.0 ± 0.10 cm³ of distilled water was added to the beaker using a burette.
• A glass rod was used to stir the solution, dissolve the solid and obtain a clear green colored solution.
• A 10.00 cm³ graduated pipette was used to add 5.10 ± 0.05 cm³ of liquid ammonia (NH₄OH) to the same beaker.
• A glass rod was used to stir the reaction mixture and mix the ligand homogenously.
• The stop-watch was started.
• As soon as the stop-watch reads 30.00 ± 0.01 mins, a 1.00 cm³ graduated pipette was used to fill one of the cuvettes and the other cuvette was filled with distilled water. The color of the reaction mixture was noted down.
• The UV-Visible spectrophotometer was set at (insert the value of wavelength you get in Graph-1)
• The absorbance of the sample was noted.
• Steps 8-10 were repeated for two more times.

The same process was repeated for other ligands. Refer to Table - 2 for the volume of the ligand to be used in Step-5.

Formula to be used: Refer to equation-2 to calculate molar concentration from absorbance and equation-1 to calculate log k. The values of n are there in the equations written in the background information.

Graph-3: Thermodynamic stability (log k) of Ni(II)-ammine complexes at room temperature against the basic strength (pk_b) of the ammine compound used as a ligand.