Each carbon atom in a diamond is covalently bonded to four other carbons in a tetrahedron. These tetrahedrons together form a 3-dimensional network of six-membered carbon rings (similar to cyclohexane), in the chair conformation, allowing for zero bond angle strain. This stable network of covalent bonds and hexagonal rings is the reason that diamond is so strong. Although graphite is the most stable allotrope of carbon under standard laboratory conditions (273 or 298 K, 1 atm), a recent computational study indicated that under idealized conditions ("T" = 0, "p" = 0), diamond is the most stable allotrope by 1.1 kJ/mol compared to graphite.