Alkalinity measures the ability of a solution to neutralize acids to the equivalence point of carbonate or bicarbonate, defined as pH 4.5 for many oceanographic/limnological studies. The alkalinity is equal to the stoichiometric sum of the bases in solution.  In most Earth surface waters carbonate alkalinity tends to make up most of the total alkalinity due to the common occurrence and dissolution of carbonate rocks and other geological weathering processes that produce carbonate anions. Other common natural components that can contribute to alkalinity include borate, hydroxide, phosphate, silicate, dissolved ammonia, and the conjugate bases of organic acids (e.g., acetate). Solutions produced in a laboratory may contain a virtually limitless number of species that contribute to alkalinity. Alkalinity is frequently given as molar equivalents per liter of solution or per kilogram of solvent. In commercial (e.g. the swimming pool industry) and regulatory contexts, alkalinity might also be given in parts per million of equivalent calcium carbonate (ppm CaCO). Alkalinity is sometimes incorrectly used interchangeably with basicity.   For example, the addition of CO lowers the pH of a solution, thus reducing  basicity while alkalinity remains unchanged (see example below). 
A variety of titrants, endpoints, and indicators are specified for various alkalinity measurement methods.  Hydrochloric and sulfuric acids are common acid titrants, while phenolpthalein, methyl red, and bromocresol green are common indicators.