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\newpage  \section{ES2} \subsection{Oxidation and Reduction}  Oxidation is:  \begin{itemize}  \item Loss of electrons  \item Gaining oxygen  \item An increase in oxidation state  \end{itemize}  Reduction is the opposite of oxidation. A redox reaction involves both oxidation and reduction.  \subsection{Oxidation States}  The oxidation state assigned to an element represents the number of electrons lost or gained in comparison to the unreacted element. \subsubsection{Oxidation states in elements}  Atoms in an element will always have an oxidation state of 0.  \subsubsection{Oxidation states for ions}  The oxidation state is the same as the charge on an individual ion, and the sum of the oxidation states is equal to the overall charge on the ion.  \subsubsection{Oxidation states in compounds}  If a compound has no overall charge, the sum of the oxidation states is 0. Some elements have oxidation states that rarely/never change;  \begin{table}[!htb]  \centering  \caption{Common Oxidation States}  \label{Common Oxidation States}  \begin{tabular}{ll}  \textit{Element} & \textit{Oxidation State} \\ \hline  Fluorine & -1 \\  Oxygen & -2, unless with $F$ or in peroxide ion ($O_2^{2-}$) \\  Chlorine & -1, unless with $O$ or $F$ \\  Bromine & -1, unless with $O$, $F$, or $Cl$ \\  Iodine & -1, unless with $O$, $F$, $Cl$, or $Br$ \\  Hydrogen & +1, except in a metal hydride \\  Group 1 & +1 \\  Group 2 & +2 \\  Aluminium & +3 \\  \end{tabular}  \end{table}  \subsubsection{Systematic names}  \begin{itemize}  \item Oxidation states used for elements with variable oxidation states  \item Shows oxidation state of preceding element  \item No space between element and number  \end{itemize}  \subsubsection{Naming Oxyanions}  An oxyanion is a negative ion containing oxygen, thus ending in \textit{-ate}. They follow the same convention as systematic naming.  \subsubsection{Balancing equations}  In a redox reaction, the number of electrons gained must be equal to the number of electrons lost.