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Why does the carbon-carbon double bond prevent cis-2-butene and trans-2-butene from easily changing into each other?

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Two diagrams show the three-dimensional structure of a methane molecule: panel A uses a ball-and-stick model with a central carbon atom connected to four hydrogen atoms in a tetrahedral arrangement, while panel B illustrates the bond angle of 109.5 degrees and bond length of 1.09 angstroms between carbon and hydrogen atoms. — isomerism Tetrahedral geometry of methane: (A) stick-and-ball model and (B) diagram showing bond angles and distances. (Plain bonds represent bonds in the plane of the image; wedge and dashed bonds represent those directed toward and away from the viewer, respectively.)

The carbon-carbon double bond prevents cis-2-butene and trans-2-butene from easily interconverting because it imparts rigidity to the molecule by restricting free rotation around the bond axis.1

In the bonding model for alkenes, the double bond consists of a sigma ( $σ$ ) bond and a pi ( $π$ ) bond. The $π$ bond is formed by the side-by-side overlap of half-filled p orbitals, which are oriented perpendicular to the plane of the $σ$ bonds. For rotation to occur around the double bond, this $π$ bond would effectively need to be broken, which requires a significant amount of energy—approximately 66 kcal/mol.23

Because this energy is not available under normal conditions, the two isomers remain distinct and cannot easily change into one another. Consequently, these cis-trans isomers are stable at room temperature and possess their own unique physical properties.34