Activities and molar concentrations
Most thermodynamic equations are derived assuming a perfect gas or ideal solution. However, for real a solution the chemical potential $\mu_i=\mu^0_i +RT\ln(x_i)$ may not be accurate for the solvent where $\mu^0_i$ is the chemical potential of the liquid solvent at 1 atm and $x_i$ is the mole fraction. The solute may, similarly, not follow ideal behaviour.  The activity is introduced to allow us to use the familiar equations but now as $\mu=\mu^0_{liq} +RT\ln(\gamma x_i)$ where $\gamma$ is the activity coefficient thus $\mu_i=\mu^0_{liq} +RT\ln(a_i)$ where $a=\gamma x$. Thus we can think of activity as effective concentration or effective pressure (now called fugasity) relative to its standard state.  The activity coefficient is unity in the ideal case. The activity has to be measured experimentally because $\gamma$ depends on exactly what intermolecular interactions exits as any temperature and pressure and cannot be easily calculated.  As an example, naphthalene dissolved in benzene behaves almost ideally and the activity coefficient $\gamma = 1$ but when dissolved in hexane this is not the case and $\gamma=2.5$. A positive deviation from Raoult's law indicates that the vapour pressure is larger than that expected from perfect behaviour and so the activity is larger than the mole fraction and $\gamma &gt;1$.