Why must the total increase in oxidation numbers equal the total decrease of them in redox reactions? In my book, under the topic of redox reactions, it is given,


  In all redox reactions, the total increase in oxidation number must be equal to the total decrease in oxidation number.


Is this a consequence of "Charges can never be created. The total charge of a system remains constant", or something else?

For the reaction,
$$\ce{
10 \color{red}{FeSO4} + 2 \color{red}{KMnO4} + 8 H2SO4 -&gt; 5\color{red}{Fe2(SO4)3} + 2\color{red}{MnSO4} + K2SO4 + 8 H2O
}$$

(Reactants and products in red undergo oxidation or reduction)

For the above reaction, I tried whether the property stated above works or not. 

The oxidation state of Fe and Mn in the reactants side are +2 and +7 respectively. The oxidation states of these elements in the product side are +3 and +2 respectively. The increase in OS is 1 and the decrease is 5.

Finally, the increase and decrease in oxidation states are not equal.

Kindly give an explanation for this fact.
The reaction is with an advantage written in the simplified ion form:

$$\ce{5 Fe^2+ + MnO4- + 8 H+ -&gt; 5 Fe^3+ + Mn^2+ + 4 H2O}$$

Total oxidation numbers $+17=+17$

The oxidation number is formally the net charge an atom within a molecule would have held, if we had broken all bonds and bond electrons  had sticked to atoms according the atom electronegativity.

Therefore, the constant total sum of oxidation numbers across all involved atoms directly relates to the constant total sum of valence electrons and to the law of charge conservation.