Enthalpy (H) is the transfer of energy in a reaction (for chemical reactions, it is in the form of heat) and ΔH is the change in enthalpy. ΔH is a state function, meaning that ΔH is independent of processes occurring between initial and final states. In other words, it does not matter which steps are taken to get from initial reactants to final products, as ΔH will always be the same. ΔHrxn, or the change in enthalpy of a reaction, has the same value of ΔH as in a thermochemical equation; however, ΔHrxn is measured in units of kJ/mol, meaning that it is the enthalpy change per moles of any particular substance in an equation. Values of ΔH are determined experimentally under standard conditions of 1 atm and 25 °C (298.15K).
As discussed earlier, ΔH can have a positive or negative sign. If ΔH has a positive sign, the system uses heat and is endothermic; if ΔH is negative, then heat is produced and the system is exothermic.
Endothermic: A + B + Heat → C, ΔH > 0
Exothermic: A + B → C + Heat, ΔH < 0
Since enthalpy is a state function, the ΔH given for a particular reaction is only true for that exact reaction.  Physical states of reactants and products matter, as do molar concentrations.
Since ΔH is dependent on the physical state and molar concentrations in reactions, thermochemical equations must be stoichiometrically correct. If one agent of an equation is changed through multiplication, then all agents must be proportionally changed, including ΔH.
The multiplicative property of thermochemical equations is mainly due to the first law of thermodynamics, which says that energy can neither be created nor destroyed; this concept is commonly known as the conservation of energy. It holds true on a physical or molecular scale.