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\author{Arsh.b}
\pdfinfo{
  /Title (19-1-electrochemical-cells.pdf)
  /Creator (Cheatography)
  /Author (Arsh.b)
  /Subject (19.1 Electrochemical Cells Cheat Sheet)
}

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\noindent
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    {\parbox{\dimexpr\textwidth-2\fboxsep\relax}{\noindent
        \hspace*{-6pt}\includegraphics[width=5.8cm]{/web/www.cheatography.com/public/images/cheatography_logo.pdf}}
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\begin{tabulary}{11cm}{L}
    \vspace{-2pt}\large{\bf{\textcolor{DarkBackground}{\textrm{19.1 Electrochemical Cells Cheat Sheet}}}} \\
    \normalsize{by \textcolor{DarkBackground}{Arsh.b} via \textcolor{DarkBackground}{\uline{cheatography.com/179523/cs/37904/}}}
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\noindent
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  \mymulticolumn{2}{p{5.377cm}}{\bf\textcolor{white}{Cheatographer}}  \\
  \vspace{-2pt}Arsh.b \\
  \uline{cheatography.com/arsh-b} \\
  \end{tabulary}
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  \mymulticolumn{1}{p{5.377cm}}{\bf\textcolor{white}{Cheat Sheet}}  \\
   \vspace{-2pt}Published 25th March, 2023.\\
   Updated 25th March, 2023.\\
   Page {\thepage} of \pageref{LastPage}.
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  %\includegraphics[width=48px,height=48px]{dave.jpeg}
  Measure your website readability!\\
  www.readability-score.com
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\begin{multicols*}{2}

\begin{tabularx}{8.4cm}{X}
\SetRowColor{DarkBackground}
\mymulticolumn{1}{x{8.4cm}}{\bf\textcolor{white}{Definitions}}  \tn
% Row 0
\SetRowColor{LightBackground}
\mymulticolumn{1}{x{8.4cm}}{{\bf{Electromotive force}} is the voltage generated by any source of electrical energy.} \tn 
% Row Count 2 (+ 2)
% Row 1
\SetRowColor{white}
\mymulticolumn{1}{x{8.4cm}}{{\bf{Potential difference}} is the difference in voltage between the anode and the cathode in a cell.} \tn 
% Row Count 4 (+ 2)
% Row 2
\SetRowColor{LightBackground}
\mymulticolumn{1}{x{8.4cm}}{{\bf{Standard hydrogen electrode}} is a reference half-cell that is used to measure the electrode potentials of other half-cell.} \tn 
% Row Count 7 (+ 3)
% Row 3
\SetRowColor{white}
\mymulticolumn{1}{x{8.4cm}}{{\bf{Electrode potential}} is the EMF that is generated by a half-cell when it is connected to the standard hydrogen electrode.} \tn 
% Row Count 10 (+ 3)
% Row 4
\SetRowColor{LightBackground}
\mymulticolumn{1}{x{8.4cm}}{{\bf{Standard conditions}} is a set of consistent reaction conditions that is sued when measuring cell potentials.} \tn 
% Row Count 13 (+ 3)
% Row 5
\SetRowColor{white}
\mymulticolumn{1}{x{8.4cm}}{{\bf{Standard electrode potential}} is the EMF that is generated by a half-cell when it is connected to the standard hydrogen electrode under standard conditions.} \tn 
% Row Count 17 (+ 4)
% Row 6
\SetRowColor{LightBackground}
\mymulticolumn{1}{x{8.4cm}}{{\bf{Electrolysed}}occurs when a compound is converted into simpler compounds using electrolysis.} \tn 
% Row Count 19 (+ 2)
% Row 7
\SetRowColor{white}
\mymulticolumn{1}{x{8.4cm}}{{\bf{Electrolysis}} is a process in which electrical current is used to make non-spontaneous redox reactions occur.} \tn 
% Row Count 22 (+ 3)
% Row 8
\SetRowColor{LightBackground}
\mymulticolumn{1}{x{8.4cm}}{{\bf{Faraday constant}} is the total charge carried by 1 mol of electrons. It has the symbol {\emph{F}} and has a value 96 500 C mol\textasciicircum{}-1\textasciicircum{}.} \tn 
% Row Count 25 (+ 3)
% Row 9
\SetRowColor{white}
\mymulticolumn{1}{x{8.4cm}}{{\bf{Electroplating}} is a process that uses electrolysis to deposit a layer of metal on another conducting object.} \tn 
% Row Count 28 (+ 3)
\hhline{>{\arrayrulecolor{DarkBackground}}-}
\end{tabularx}
\par\addvspace{1.3em}

\begin{tabularx}{8.4cm}{X}
\SetRowColor{DarkBackground}
\mymulticolumn{1}{x{8.4cm}}{\bf\textcolor{white}{Standard hydrogen electrode}}  \tn
\SetRowColor{LightBackground}
\mymulticolumn{1}{p{8.4cm}}{\vspace{1px}\centerline{\includegraphics[width=5.1cm]{/web/www.cheatography.com/public/uploads/arsh-b_1679749213_Screenshot 2023-03-25 at 12.59.52 PM.png}}} \tn 
\hhline{>{\arrayrulecolor{DarkBackground}}-}
\end{tabularx}
\par\addvspace{1.3em}

\begin{tabularx}{8.4cm}{X}
\SetRowColor{DarkBackground}
\mymulticolumn{1}{x{8.4cm}}{\bf\textcolor{white}{Sample setup for standard electron potential}}  \tn
\SetRowColor{LightBackground}
\mymulticolumn{1}{p{8.4cm}}{\vspace{1px}\centerline{\includegraphics[width=5.1cm]{/web/www.cheatography.com/public/uploads/arsh-b_1679749271_Screenshot 2023-03-25 at 1.01.00 PM.png}}} \tn 
\hhline{>{\arrayrulecolor{DarkBackground}}-}
\end{tabularx}
\par\addvspace{1.3em}

\begin{tabularx}{8.4cm}{p{0.912 cm} x{5.472 cm} x{1.216 cm} }
\SetRowColor{DarkBackground}
\mymulticolumn{3}{x{8.4cm}}{\bf\textcolor{white}{Electrolysis of aqueous solutions}}  \tn
% Row 0
\SetRowColor{LightBackground}
\mymulticolumn{3}{x{8.4cm}}{{\emph{When aqueous solutions are electrolysed, water can be oxidized to oxygen at the anode and reduced to hydrogen at the cathode.}}} \tn 
% Row Count 3 (+ 3)
% Row 1
\SetRowColor{white}
\mymulticolumn{3}{x{8.4cm}}{{\bf{Cathode}}} \tn 
% Row Count 4 (+ 1)
% Row 2
\SetRowColor{LightBackground}
\seqsplit{Cation} & M\textasciicircum{}+\textasciicircum{}(aq) + e\textasciicircum{}-\textasciicircum{} -\textgreater{} M & \seqsplit{Reduced} \tn 
% Row Count 6 (+ 2)
% Row 3
\SetRowColor{white}
Water & 2H2O(l) + 2e\textasciicircum{}-\textasciicircum{} -\textgreater{} H2(g) + 2OH\textasciicircum{}-\textasciicircum{}(aq) & \seqsplit{Reduced} \tn 
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% Row 4
\SetRowColor{LightBackground}
\mymulticolumn{3}{x{8.4cm}}{{\bf{Anode}}} \tn 
% Row Count 9 (+ 1)
% Row 5
\SetRowColor{white}
Anion & A\textasciicircum{}-\textasciicircum{} -\textgreater{} A + e\textasciicircum{}-\textasciicircum{} & \seqsplit{Oxidised} \tn 
% Row Count 11 (+ 2)
% Row 6
\SetRowColor{LightBackground}
Water & 2H2O(l) -\textgreater{} 4H\textasciicircum{}+\textasciicircum{} + O2(g) + 4e\textasciicircum{}-\textasciicircum{}(aq) & \seqsplit{Oxidised} \tn 
% Row Count 13 (+ 2)
\hhline{>{\arrayrulecolor{DarkBackground}}---}
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\begin{tabularx}{8.4cm}{x{1.292 cm} x{3.116 cm} x{3.192 cm} }
\SetRowColor{DarkBackground}
\mymulticolumn{3}{x{8.4cm}}{\bf\textcolor{white}{Electrolysis of NaCl(aq)}}  \tn
% Row 0
\SetRowColor{LightBackground}
\mymulticolumn{3}{x{8.4cm}}{{\bf{Possible reactions at electrodes}}} \tn 
% Row Count 1 (+ 1)
% Row 1
\SetRowColor{white}
\seqsplit{Cathode} & 2Na\textasciicircum{}+\textasciicircum{}(aq) + 2e\textasciicircum{}-\textasciicircum{} -\textgreater{} Na(s) & {\emph{E\textasciicircum{}⍬\textasciicircum{}}} = -2.71V \tn 
% Row Count 3 (+ 2)
% Row 2
\SetRowColor{LightBackground}
 & 2H2O(l) + 2e\textasciicircum{}-\textasciicircum{} -\textgreater{} H2(g) + 2OH\textasciicircum{}-\textasciicircum{}(aq) & {\emph{E\textasciicircum{}⍬\textasciicircum{}}} = -0.83V \tn 
% Row Count 6 (+ 3)
% Row 3
\SetRowColor{white}
Anode & 2Cl\textasciicircum{}-\textasciicircum{} -\textgreater{} Cl2(g) + 2e\textasciicircum{}-\textasciicircum{} & -{\emph{E\textasciicircum{}⍬\textasciicircum{}}} = -1.36V \tn 
% Row Count 8 (+ 2)
% Row 4
\SetRowColor{LightBackground}
 & 2H2O(l) -\textgreater{} 4H\textasciicircum{}+\textasciicircum{} + O2(g) + 4e\textasciicircum{}-\textasciicircum{}(aq) & -{\emph{E\textasciicircum{}⍬\textasciicircum{}}} = -1.23V \tn 
% Row Count 11 (+ 3)
% Row 5
\SetRowColor{white}
\mymulticolumn{3}{x{8.4cm}}{{\bf{Low concentration of NaCl(aq)}}} \tn 
% Row Count 12 (+ 1)
% Row 6
\SetRowColor{LightBackground}
\seqsplit{Cathode} & 2H2O(l) + 2e\textasciicircum{}-\textasciicircum{} -\textgreater{} H2(g) + 2OH\textasciicircum{}-\textasciicircum{}(aq) & The reaction with the most positive reduction potential is favoured \tn 
% Row Count 17 (+ 5)
% Row 7
\SetRowColor{white}
Anode & 2H2O(l) -\textgreater{} 4H\textasciicircum{}+\textasciicircum{} + O2(g) + 4e\textasciicircum{}-\textasciicircum{}(aq) & The reaction with the most positive oxidation potential is favoured \tn 
% Row Count 22 (+ 5)
% Row 8
\SetRowColor{LightBackground}
\seqsplit{Overall} \seqsplit{equation} & 2H2)(l) -\textgreater{} 2H2(g) + O2(g) &  \tn 
% Row Count 25 (+ 3)
% Row 9
\SetRowColor{white}
\mymulticolumn{3}{x{8.4cm}}{Colourless H2(g) bubbles at cathode, and colourless O2(g) bubbles at anode} \tn 
% Row Count 27 (+ 2)
% Row 10
\SetRowColor{LightBackground}
\mymulticolumn{3}{x{8.4cm}}{{\bf{High concentration of NaCl(aq)}}} \tn 
% Row Count 28 (+ 1)
% Row 11
\SetRowColor{white}
\seqsplit{Cathode} & 2H2O(l) + 2e\textasciicircum{}-\textasciicircum{} -\textgreater{} H2(g) + 2OH\textasciicircum{}-\textasciicircum{}(aq) & The reaction with the most positive reduction potential is favoured \tn 
% Row Count 33 (+ 5)
\end{tabularx}
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\vfill
\columnbreak
\begin{tabularx}{8.4cm}{x{1.292 cm} x{3.116 cm} x{3.192 cm} }
\SetRowColor{DarkBackground}
\mymulticolumn{3}{x{8.4cm}}{\bf\textcolor{white}{Electrolysis of NaCl(aq) (cont)}}  \tn
% Row 12
\SetRowColor{LightBackground}
Anode & 2Cl\textasciicircum{}-\textasciicircum{} -\textgreater{} Cl2(g) + 2e\textasciicircum{}-\textasciicircum{} & The two possible oxidation reactions have similar potentials, at high {[}Cl\textasciicircum{}-\textasciicircum{}{]} the oxidation of Cl\textasciicircum{}-\textasciicircum{} is favoured \tn 
% Row Count 7 (+ 7)
% Row 13
\SetRowColor{white}
\seqsplit{Overall} \seqsplit{equation} & 2H2O(l) + 2NaCl(aq) ➝ 2H2(g) + Cl2(g) + 2NaOH(aq) &  \tn 
% Row Count 11 (+ 4)
% Row 14
\SetRowColor{LightBackground}
\mymulticolumn{3}{x{8.4cm}}{Colourless H2(g) bubbles at cathode, green Cl2(g) bubbles at anode, and the pH increases as OH\textasciicircum{}-\textasciicircum{} is formed} \tn 
% Row Count 14 (+ 3)
\hhline{>{\arrayrulecolor{DarkBackground}}---}
\SetRowColor{LightBackground}
\mymulticolumn{3}{x{8.4cm}}{As the oxidation reactions take place at the anode, the sign of the electrode potentials has been reversed to convert it from a reduction potential to an oxidation potential.}  \tn 
\hhline{>{\arrayrulecolor{DarkBackground}}---}
\end{tabularx}
\par\addvspace{1.3em}

\begin{tabularx}{8.4cm}{x{2.052 cm} x{1.748 cm} x{3.8 cm} }
\SetRowColor{DarkBackground}
\mymulticolumn{3}{x{8.4cm}}{\bf\textcolor{white}{Quantitative Electrolysis}}  \tn
% Row 0
\SetRowColor{LightBackground}
\mymulticolumn{3}{x{8.4cm}}{{\bf{{\emph{ΔG\textasciicircum{}⍬\textasciicircum{} = –nFE\textasciicircum{}⍬\textasciicircum{}}} . When E is positive, ΔG is negative, indicating of a spontaneous process. When E is negative, ΔG is positive, indicating of a non-spontaneous process. When E is 0, ΔG is 0.}}} \tn 
% Row Count 5 (+ 5)
% Row 1
\SetRowColor{white}
\mymulticolumn{3}{x{8.4cm}}{The Cell potential of a voltaic cell is related to the Gibbs free energy change for the overall redox reaction occurring in the cell: {\emph{ΔG\textasciicircum{}⍬\textasciicircum{} = –nFE\textasciicircum{}⍬\textasciicircum{}}}} \tn 
% Row Count 9 (+ 4)
% Row 2
\SetRowColor{LightBackground}
\mymulticolumn{3}{x{8.4cm}}{Where,\{\{nl\}\}{\emph{ΔG}} is the difference in the standard free energy of reactants and products,\{\{nl\}\}{\emph{n}} is the number of moles of electrons transferred in the reaction,\{\{nl\}\}{\emph{F}} is the Faraday constant,\{\{nl\}\}{\emph{E\textasciicircum{}⍬\textasciicircum{}}} is the standard cell potential.} \tn 
% Row Count 14 (+ 5)
% Row 3
\SetRowColor{white}
\mymulticolumn{3}{x{8.4cm}}{-{\emph{ΔG\textasciicircum{}⍬\textasciicircum{}}} is spontaneous, and +{\emph{ΔG\textasciicircum{}⍬\textasciicircum{}}} is non-spontaneous.} \tn 
% Row Count 16 (+ 2)
% Row 4
\SetRowColor{LightBackground}
\mymulticolumn{3}{x{8.4cm}}{This means that the cell potential {\emph{E\textasciicircum{}⍬\textasciicircum{}}}cell, also indicates if the redox reaction occurring is spontaneous or non-spontaneous.} \tn 
% Row Count 19 (+ 3)
% Row 5
\SetRowColor{white}
{\emph{E\textasciicircum{}⍬\textasciicircum{}}}cell & {\emph{ΔG\textasciicircum{}⍬\textasciicircum{}}} & Spontaneity \tn 
% Row Count 21 (+ 2)
% Row 6
\SetRowColor{LightBackground}
positive & negative & spontaneous \tn 
% Row Count 22 (+ 1)
% Row 7
\SetRowColor{white}
negative & positive & non-spontaneous \tn 
% Row Count 23 (+ 1)
% Row 8
\SetRowColor{LightBackground}
zero & zero & cell is at equilibrium \tn 
% Row Count 25 (+ 2)
\hhline{>{\arrayrulecolor{DarkBackground}}---}
\end{tabularx}
\par\addvspace{1.3em}

\begin{tabularx}{8.4cm}{X}
\SetRowColor{DarkBackground}
\mymulticolumn{1}{x{8.4cm}}{\bf\textcolor{white}{Electroplating}}  \tn
\SetRowColor{LightBackground}
\mymulticolumn{1}{p{8.4cm}}{\vspace{1px}\centerline{\includegraphics[width=5.1cm]{/web/www.cheatography.com/public/uploads/arsh-b_1679769621_Screenshot 2023-03-25 at 6.39.51 PM.png}}} \tn 
\hhline{>{\arrayrulecolor{DarkBackground}}-}
\end{tabularx}
\par\addvspace{1.3em}

\begin{tabularx}{8.4cm}{X}
\SetRowColor{DarkBackground}
\mymulticolumn{1}{x{8.4cm}}{\bf\textcolor{white}{Voltaic cells}}  \tn
% Row 0
\SetRowColor{LightBackground}
\mymulticolumn{1}{x{8.4cm}}{{\emph{A voltaic cell generates an electromotive force (EMF) resulting in the movement of electrons from the anode (negative electrode) to the cathode (positive electrode) via the external circuit. The EMF is termed the cell potential (E ).}}} \tn 
% Row Count 5 (+ 5)
% Row 1
\SetRowColor{white}
\mymulticolumn{1}{x{8.4cm}}{The electromotive force (EMF) of a voltaic cell:\{\{nl\}\} is also known as the cell potential\{\{nl\}\}is the potential difference between the anode and cathode\{\{nl\}\}has the symbol {\emph{E}}\{\{nl\}\}changes with different combinations of half cells.} \tn 
% Row Count 10 (+ 5)
\hhline{>{\arrayrulecolor{DarkBackground}}-}
\end{tabularx}
\par\addvspace{1.3em}

\begin{tabularx}{8.4cm}{X}
\SetRowColor{DarkBackground}
\mymulticolumn{1}{x{8.4cm}}{\bf\textcolor{white}{Standard hydrogen electrode}}  \tn
% Row 0
\SetRowColor{LightBackground}
\mymulticolumn{1}{x{8.4cm}}{{\emph{The standard hydrogen electode (SHE) consists of an inert platinum electrode in contact with 1 mol dm–3 hydrogen ion and hydrogen gas at 100 kPa and 298 K. The standard electrode potential (E ) is the potential (voltage) of the reduction half-equation under standard conditions measured relative to the SHE. Solute concentration is 1 mol dm–3 or 100 kPa for gases. The E of the SHE is 0 V.}}} \tn 
% Row Count 8 (+ 8)
% Row 1
\SetRowColor{white}
\mymulticolumn{1}{x{8.4cm}}{The standard hydrogen electrode(SHE):\{\{nl\}\}- is a half-cell that can be combined with other half-cells to make a voltaic cells\{\{nl\}\}- is used as a reference half-cell\{\{nl\}\}- has a standard electrode potential of 0V\{\{nl\}\}- has the redox equilibrium between aqueous H\textasciicircum{}+\textasciicircum{} ions and H2(g)\{\{nl\}\}- uses an inert platinum electrode\{\{nl\}\}- uses standard conditions} \tn 
% Row Count 16 (+ 8)
% Row 2
\SetRowColor{LightBackground}
\mymulticolumn{1}{x{8.4cm}}{Electrode potentials are affected by temperature, concentration, and pressure. This means that a standard set of conditions must be used when measuring electrode potentials.} \tn 
% Row Count 20 (+ 4)
% Row 3
\SetRowColor{white}
\mymulticolumn{1}{x{8.4cm}}{The standard conditions for electrochemical cells are:\{\{nl\}\}- a concentration of 1.0mol dm\textasciicircum{}-3\textasciicircum{} for all solutions\{\{nl\}\}- a pressure of 100kPa for all gases\{\{nl\}\}- a temperature of 298K\{\{nl\}\}- a platinum electrode if the half-cell does not include a metal} \tn 
% Row Count 26 (+ 6)
% Row 4
\SetRowColor{LightBackground}
\mymulticolumn{1}{x{8.4cm}}{the standard electrode potential ({\emph{E\textasciicircum{}⍬\textasciicircum{}}}) of a half-cell is the EMF that is measured when it is connected to a SHE to make a voltaic cell under standard conditions.} \tn 
% Row Count 30 (+ 4)
\end{tabularx}
\par\addvspace{1.3em}

\vfill
\columnbreak
\begin{tabularx}{8.4cm}{X}
\SetRowColor{DarkBackground}
\mymulticolumn{1}{x{8.4cm}}{\bf\textcolor{white}{Standard hydrogen electrode (cont)}}  \tn
% Row 5
\SetRowColor{LightBackground}
\mymulticolumn{1}{x{8.4cm}}{{\emph{E\textasciicircum{}⍬\textasciicircum{}}}cell = {\emph{E\textasciicircum{}⍬\textasciicircum{}}}half-cell - {\emph{E\textasciicircum{}⍬\textasciicircum{}}}SHE\{\{nl\}\}{\emph{E\textasciicircum{}⍬\textasciicircum{}}}SHE=0, {\emph{E\textasciicircum{}⍬\textasciicircum{}}}cell = {\emph{E\textasciicircum{}⍬\textasciicircum{}}}half-cell} \tn 
% Row Count 2 (+ 2)
% Row 6
\SetRowColor{white}
\mymulticolumn{1}{x{8.4cm}}{{\emph{E\textasciicircum{}⍬\textasciicircum{}}}half-cell values:\{\{nl\}\}- are always given for the reduction reaction of the half-cell\{\{nl\}\}- do not depend on the number of electrons involved in the reduction reaction\{\{nl\}\}- are positive when the half-cell is more easily induced than the SHE\{\{nl\}\}- are negative when the half-cell is harder to reduce than the SHE} \tn 
% Row Count 9 (+ 7)
% Row 7
\SetRowColor{LightBackground}
\mymulticolumn{1}{x{8.4cm}}{When two half-cells are combined:\{\{nl\}\}- reduction will occur at the half-cell with the more positive {\emph{E\textasciicircum{}⍬\textasciicircum{}}} value.\{\{nl\}\}- oxidation will occur at the half-cell with the more negative {\emph{E\textasciicircum{}⍬\textasciicircum{}}}value.} \tn 
% Row Count 13 (+ 4)
\hhline{>{\arrayrulecolor{DarkBackground}}-}
\SetRowColor{LightBackground}
\mymulticolumn{1}{x{8.4cm}}{{\emph{E\textasciicircum{}⍬\textasciicircum{}}} values are given for a large number of half-cells in section 24 of the IB data booklet.}  \tn 
\hhline{>{\arrayrulecolor{DarkBackground}}-}
\end{tabularx}
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\begin{tabularx}{8.4cm}{x{1.292 cm} x{3.116 cm} x{3.192 cm} }
\SetRowColor{DarkBackground}
\mymulticolumn{3}{x{8.4cm}}{\bf\textcolor{white}{Electrolysis of CuSo4(aq)}}  \tn
% Row 0
\SetRowColor{LightBackground}
\mymulticolumn{3}{x{8.4cm}}{{\bf{{\emph{Inert electrodes}}}}} \tn 
% Row Count 1 (+ 1)
% Row 1
\SetRowColor{white}
\mymulticolumn{3}{x{8.4cm}}{{\bf{Possible reactions at electrodes}}} \tn 
% Row Count 2 (+ 1)
% Row 2
\SetRowColor{LightBackground}
\seqsplit{Cathode} & Cu\textasciicircum{}2+\textasciicircum{} + 2e\textasciicircum{}-\textasciicircum{} -\textgreater{} Cu(s) & {\emph{E\textasciicircum{}⍬\textasciicircum{}}} = +0.34V \tn 
% Row Count 4 (+ 2)
% Row 3
\SetRowColor{white}
 & 2H2O(l) + 2e\textasciicircum{}-\textasciicircum{} -\textgreater{} H2(g) + 2OH\textasciicircum{}-\textasciicircum{}(aq) & {\emph{E\textasciicircum{}⍬\textasciicircum{}}} = -0.83V \tn 
% Row Count 7 (+ 3)
% Row 4
\SetRowColor{LightBackground}
Anode & 2SO4\textasciicircum{}2-\textasciicircum{}(aq) -\textgreater{} S2O8\textasciicircum{}2-\textasciicircum{}(g)+2e\textasciicircum{}-\textasciicircum{} & -{\emph{E\textasciicircum{}⍬\textasciicircum{}}} = -2.01V \tn 
% Row Count 10 (+ 3)
% Row 5
\SetRowColor{white}
 & 2H2O(l) -\textgreater{} 4H\textasciicircum{}+\textasciicircum{} + O2(g) + 4e\textasciicircum{}-\textasciicircum{}(aq) & -{\emph{E\textasciicircum{}⍬\textasciicircum{}}} = -1.23V \tn 
% Row Count 13 (+ 3)
% Row 6
\SetRowColor{LightBackground}
\mymulticolumn{3}{x{8.4cm}}{{\bf{Actual reactions}}} \tn 
% Row Count 14 (+ 1)
% Row 7
\SetRowColor{white}
\seqsplit{Cathode} & Cu\textasciicircum{}2+\textasciicircum{} + 2e\textasciicircum{}-\textasciicircum{} -\textgreater{} Cu(s) & The reaction with the most positive reduction potential is favoured \tn 
% Row Count 19 (+ 5)
% Row 8
\SetRowColor{LightBackground}
Anode & 2H2O(l) -\textgreater{} 4H\textasciicircum{}+\textasciicircum{} + O2(g) + 4e\textasciicircum{}-\textasciicircum{}(aq) & The reaction with the most positive oxidation potential is favoured \tn 
% Row Count 24 (+ 5)
% Row 9
\SetRowColor{white}
\seqsplit{Overall} \seqsplit{equation} & 2CuSO4(aq) + 2H2O(l) -\textgreater{} 2Cu(s) + O2(g) + 4H\textasciicircum{}+\textasciicircum{}(aq) + 2SO4\textasciicircum{}2-\textasciicircum{}(aq) &  \tn 
% Row Count 29 (+ 5)
% Row 10
\SetRowColor{LightBackground}
\mymulticolumn{3}{x{8.4cm}}{Pink/brown metal layer of Cu(s) forms on cathode, blue colour of solution fades as Cu2+(aq) is depleted, and colourless O2(g) bubbles at anode pH decreases as H\textasciicircum{}+\textasciicircum{} formed} \tn 
% Row Count 33 (+ 4)
\end{tabularx}
\par\addvspace{1.3em}

\vfill
\columnbreak
\begin{tabularx}{8.4cm}{x{1.292 cm} x{3.116 cm} x{3.192 cm} }
\SetRowColor{DarkBackground}
\mymulticolumn{3}{x{8.4cm}}{\bf\textcolor{white}{Electrolysis of CuSo4(aq) (cont)}}  \tn
% Row 11
\SetRowColor{LightBackground}
\mymulticolumn{3}{x{8.4cm}}{{\bf{{\emph{Copper Electrodes}}}}} \tn 
% Row Count 1 (+ 1)
% Row 12
\SetRowColor{white}
\mymulticolumn{3}{x{8.4cm}}{{\bf{Possible reactions at electrodes}}} \tn 
% Row Count 2 (+ 1)
% Row 13
\SetRowColor{LightBackground}
\seqsplit{Cathode} & Cu\textasciicircum{}2+\textasciicircum{} + 2e\textasciicircum{}-\textasciicircum{} -\textgreater{} Cu(s) & {\emph{E\textasciicircum{}⍬\textasciicircum{}}} = +0.34V \tn 
% Row Count 4 (+ 2)
% Row 14
\SetRowColor{white}
 & 2H2O(l) + 2e\textasciicircum{}-\textasciicircum{} -\textgreater{} H2(g) + 2OH\textasciicircum{}-\textasciicircum{}(aq) & {\emph{E\textasciicircum{}⍬\textasciicircum{}}} = -0.83V \tn 
% Row Count 7 (+ 3)
% Row 15
\SetRowColor{LightBackground}
Anode & 2SO4\textasciicircum{}2-\textasciicircum{}(aq) -\textgreater{} S2O8\textasciicircum{}2-\textasciicircum{}(g)+2e\textasciicircum{}-\textasciicircum{} & -{\emph{E\textasciicircum{}⍬\textasciicircum{}}} = -2.01V \tn 
% Row Count 10 (+ 3)
% Row 16
\SetRowColor{white}
 & 2H2O(l) -\textgreater{} 4H\textasciicircum{}+\textasciicircum{} + O2(g) + 4e\textasciicircum{}-\textasciicircum{}(aq) & -{\emph{E\textasciicircum{}⍬\textasciicircum{}}} = -1.23V \tn 
% Row Count 13 (+ 3)
% Row 17
\SetRowColor{LightBackground}
 & Cu(s) -\textgreater{} Cu\textasciicircum{}2+\textasciicircum{}(aq) + 2e\textasciicircum{}-\textasciicircum{} & -{\emph{E\textasciicircum{}⍬\textasciicircum{}}} = -0.34V \tn 
% Row Count 15 (+ 2)
% Row 18
\SetRowColor{white}
\mymulticolumn{3}{x{8.4cm}}{{\bf{Actual reactions}}} \tn 
% Row Count 16 (+ 1)
% Row 19
\SetRowColor{LightBackground}
\seqsplit{Cathode} & Cu\textasciicircum{}2+\textasciicircum{} + 2e\textasciicircum{}-\textasciicircum{} -\textgreater{} Cu(s) & The reaction with the most positive reduction potential is favoured \tn 
% Row Count 21 (+ 5)
% Row 20
\SetRowColor{white}
Anode & Cu(s) -\textgreater{} Cu\textasciicircum{}2+\textasciicircum{}(aq) + 2e\textasciicircum{}-\textasciicircum{} & The reaction with the most positive oxidation potential is favoured \tn 
% Row Count 26 (+ 5)
% Row 21
\SetRowColor{LightBackground}
\mymulticolumn{3}{x{8.4cm}}{No overall change is observed} \tn 
% Row Count 27 (+ 1)
% Row 22
\SetRowColor{white}
\mymulticolumn{3}{x{8.4cm}}{Pink/brown metal layer of Cu(s) forms on cathode, Cu anode reduces in mass as it oxidizes to Cu\textasciicircum{}2+\textasciicircum{}, blue colour of solution is constant, and no change in pH.} \tn 
% Row Count 31 (+ 4)
\hhline{>{\arrayrulecolor{DarkBackground}}---}
\end{tabularx}
\par\addvspace{1.3em}

\begin{tabularx}{8.4cm}{X}
\SetRowColor{DarkBackground}
\mymulticolumn{1}{x{8.4cm}}{\bf\textcolor{white}{Product of electrolysis}}  \tn
% Row 0
\SetRowColor{LightBackground}
\mymulticolumn{1}{x{8.4cm}}{{\emph{Current, duration of electrolysis, and charge on the ion affect the amount of product formed at the electrodes during electrolysis.}}} \tn 
% Row Count 3 (+ 3)
% Row 1
\SetRowColor{white}
\mymulticolumn{1}{x{8.4cm}}{The amount of product that is produced in an electrolysis reaction depends on:\{\{nl\}\}- the amount of current\{\{nl\}\}- the amount of time that the current is flowing\{\{nl\}\}- the charge on the ion being oxidised or reduced.} \tn 
% Row Count 8 (+ 5)
% Row 2
\SetRowColor{LightBackground}
\mymulticolumn{1}{x{8.4cm}}{{\emph{Determination of the relative amounts of products formed during electrolytic processes.}}} \tn 
% Row Count 10 (+ 2)
% Row 3
\SetRowColor{white}
\mymulticolumn{1}{x{8.4cm}}{Determine the total charge (Q) produced by the current flow: {\emph{Q=It}}.} \tn 
% Row Count 12 (+ 2)
% Row 4
\SetRowColor{LightBackground}
\mymulticolumn{1}{x{8.4cm}}{Determine the number of moles of electrons contained in the total charge: {\emph{n(e\textasciicircum{}-\textasciicircum{}) = Q/F}}.} \tn 
% Row Count 14 (+ 2)
% Row 5
\SetRowColor{white}
\mymulticolumn{1}{x{8.4cm}}{Use the balanced oxidation (or reduction) reaction to determine the moles of product formed.} \tn 
% Row Count 16 (+ 2)
% Row 6
\SetRowColor{LightBackground}
\mymulticolumn{1}{x{8.4cm}}{Determine the mass of product formed: {\emph{m=nM}}} \tn 
% Row Count 17 (+ 1)
\hhline{>{\arrayrulecolor{DarkBackground}}-}
\end{tabularx}
\par\addvspace{1.3em}

\begin{tabularx}{8.4cm}{X}
\SetRowColor{DarkBackground}
\mymulticolumn{1}{x{8.4cm}}{\bf\textcolor{white}{Electroplating}}  \tn
% Row 0
\SetRowColor{LightBackground}
\mymulticolumn{1}{x{8.4cm}}{{\emph{Electroplating involves the electrolytic coating of an object with a metallic thin layer.}}} \tn 
% Row Count 2 (+ 2)
% Row 1
\SetRowColor{white}
\mymulticolumn{1}{x{8.4cm}}{Electroplating is when a layer of metal is deposited on an object made of a different metal or another conductive material such as graphite.} \tn 
% Row Count 5 (+ 3)
% Row 2
\SetRowColor{LightBackground}
\mymulticolumn{1}{x{8.4cm}}{Electroplating uses an electrolytic cell that contains:\{\{nl\}\}- a salt solution containing the cations of the metal to be deposited onto the object\{\{nl\}\}- a cathode made of the conducting object that will be electroplated\{\{nl\}\}- an anode made of the metal being electroplated.} \tn 
% Row Count 11 (+ 6)
\hhline{>{\arrayrulecolor{DarkBackground}}-}
\end{tabularx}
\par\addvspace{1.3em}


% That's all folks
\end{multicols*}

\end{document}