Probing Electrochemical Potential Differences over the Solid/Liquid Interface in Li-Ion Battery Model Systems

The electrochemical potential difference (Δμ̅) is the driving force for the transfer of a charged species from one phase to another in a redox reaction. In Li-ion batteries (LIBs), Δμ̅ values for both electrons and Li-ions play an important role in the charge-transfer kinetics at the electrode/electrolyte interfaces. Because of the lack of suitable measurement techniques, little is known about how Δμ̅ affects the redox reactions occurring at the solid/liquid interfaces during LIB operation. Herein, we outline the relations between different potentials and show how ambient pressure photoelectron spectroscopy (APPES) can be used to follow changes in Δμ̅e over the solid/liquid interfaces operando by measuring the kinetic energy (KE) shifts of the electrolyte core levels. The KE shift versus applied voltage shows a linear dependence of ∼1 eV/V during charging of the electrical double layer and during solid electrolyte interphase formation. This agrees with the expected results for an ideally polarizable interface. During lithiation, the slope changes drastically. We propose a model to explain this based on charge transfer over the solid/liquid interface.

S-2 Figure S1 shows the measured Cu 2p, Cu LMM and O 1s spectra of the pristine Cu electrode. By comparing the spectra to reference spectra of Cu and Cu oxide, 1 it can be concluded that the Cu metal is covered with a Cu oxide layer dominated by Cu 2 O. S-3

Electrochemical measurements
Figures S2 and S3 show the first and second cycle of the Cu electrode, performed with cyclic voltammetry. The CVs show Cu reduction (starting at ~3 V with a max around 1.5 V) and SEI formation (starting around 0.8 V). Figure S4 and S5 show the electrochemical measurements performed during APPES for the Au and Cu WE, respectively. APPES measurements are performed under a constant potential step when the current has decayed to a stable value. Currents during APPES measurements are presented in Table S1 and S2.   S-5

Overpotentials and iR-drop in the electrolyte
If a voltage is measured during a faradaic reaction, the measured voltage vs the RE will in general not be equal to the standard/equilibrium potential of the redox reaction predicted from thermodynamics. Whenever a current is running, the measured voltage in addition contain contributions from an overpotential η, stemming from the kinetics of the chemical reaction, and from an iR s -drop, stemming from the ion transport in the bulk solution. Both these contributions can be modelled as resistances in the system, but they are usually treated separately since the overpotential(s) depend on the electrochemical reaction and is attributed to the potential at the electrode (interface), while the iR s -drop is independent of the interfacial chemistry and is a property of the bulk electrolyte. If there is an iR s -drop in the solution, the voltage measured vs. the RE will differ from the voltage drop over the WE/electrolyte interface by the uncompensated iR s -drop. This is illustrated in Figure S6. By choosing electrolytes with good ion conductivity and optimizing the design of the electrochemical cell, the iR s -drop can be minimized.
In our setup the electrolyte conductivity (κ) for 1 M LiClO 4 in PC is about 10 mS cm -1 , 2 the distance between WE and CE (d) is approximately 1 cm and the WE area (A WE ) is approximately 2 cm 2 . This means that the resistance (R s ) in the bulk electrolyte can be estimated. The current (i) is typically a few µA, but at most 20 µA during APPES measurements (see table S1 and S2), so the voltage drop over the electrolyte (U) can be estimated to: U = R x i = 50 Ω x 20 µA = 0.001 V Thus, the iR s -drop in the electrolyte would be ~1 mV. Since the deviation from the expected dependence of 1 eV/V during lithiation is ~0.5 eV/V and ~0.3 eV/V for the Au and Cu WE, respectively, and the voltage change in Region B is ~0.4 V and ~0.9 V, the iR s -drop is too small to account for this deviation.
It can still be noted that we certainly have mass transport limitations in the thin liquid meniscus, that slow down the rate of the reaction and give lower current densities at the same voltage due to the (additional) overpotential needed to drive the reaction. However, this will not affect as ∆ measured by APPES on the bulk electrolyte since this potential drop is located at the electrode/electrolyte interface.