For a reaction at constant temperature and pressure, the Gibbs free energy is:
a negative ΔG would depend on the sign of the changes in enthalpy (ΔH), entropy (ΔS), and the magnitude of the absolute temperature (in kelvins).  Changes in the sign of ΔG cannot be changed directly by temperature, because it can never be less than zero.
When ΔS is positive and ΔH is negative, a<unk>  is<unk> 
When ΔS is positive and ΔH is positive, a<unk>  is<unk>  at high temperatures, where exothermicity plays a small role in the balance.
When ΔS is negative and ΔH is negative, a<unk>  is<unk>  at low temperatures, where exothermicity is important.
When ΔS is negative and ΔH is positive, a<unk>  is not<unk>  at any temperature, but the reverse<unk>  is<unk> .
The second law of thermodynamics states that for any<unk>  the overall change in entropy of the system must be greater than or equal to zero, yet a<unk>  chemical reaction can result in a negative change in entropy.  This does not contradict the second law however, since such a reaction must have a sufficiently large negative change in enthalpy (heat energy) that the increase in temperature of the reaction surroundings (considered to be part of the system in thermodynamic terms) results in a sufficiently large increase in entropy that overall the change in entropy is positive. That is, the ΔS of the surroundings increases enough because of the exothermicity of the reaction that it overcompensates for the negative ΔS of the system, and since the overall ΔS = ΔSsurroundings + ΔSsystem, the overall change in entropy is still positive.  
Another way to view the fact that some<unk>  chemical reactions can lead to products with lower entropy is to realize that the second law states that entropy of a closed system must increase (or remain constant). Since a positive enthalpy means that energy is being released to the surroundings, then the 'closed' system includes the chemical reaction plus its surroundings. This means that the heat release of the chemical reaction sufficiently increases the entropy of the surroundings such that the overall entropy of the closed system increases in accordance with the second law of thermodynamics.
In simplified terms, a<unk>  reaction is one in which a meaningful amount of products are formed from the reactants without the addition of a catalyst.  However, a "nonspontaneous" reaction may still proceed, but will not convert the reactants into an appreciable amount of products. The reaction may proceed to its equilibrium point, but its equilibrium is very small (possibly on the order of 10-23 or smaller).  For example, table salt (NaCl) will not<unk> ly separate into individual ions (Na+ and Cl-), unless it is greatly heated or forcibly separated by electrolysis, yet small, immeasurable amounts of the ions may form.  Additionally, just because a chemist may call a reaction<unk> ” does not mean the reaction happens with great speed. For example, the decay of diamonds into graphite is a<unk>  but this decay is extremely slow and takes millions of years. Thus the rate of a reaction is independent of its spontaneity, and instead depends on the chemical kinetics of the reaction.
Spontaneous process
Sexually transmitted disease
Benign tumor
Opportunistic infection
Congenital abnormality
Malignant transformation
Post operative complications
Idiopathic