In the methane CH4 molecule, the type of hybrid orbitals present are sp3 hybrid orbitals. To understand how these hybrid orbitals contribute to the bonding in the molecule, let's first look at the electronic configuration of carbon C and hydrogen H atoms. Carbon has an electronic configuration of 1s 2s 2p, while hydrogen has an electronic configuration of 1s.In the ground state, carbon has two unpaired electrons in its 2p orbitals. To form four bonds with four hydrogen atoms, carbon needs to have four unpaired electrons. This is achieved through the process of hybridization, where one 2s orbital and three 2p orbitals of carbon combine to form four new equivalent orbitals called sp3 hybrid orbitals. These sp3 hybrid orbitals are arranged tetrahedrally around the carbon atom, with bond angles of 109.5.Each of the four hydrogen atoms has one unpaired electron in its 1s orbital. When the methane molecule forms, the four sp3 hybrid orbitals of the carbon atom overlap with the 1s orbitals of the four hydrogen atoms, resulting in the formation of four sigma bonds. These bonds are strong covalent bonds that hold the CH4 molecule together.In summary, the sp3 hybrid orbitals present in the methane molecule contribute to the bonding by allowing carbon to form four equivalent and strong sigma bonds with the hydrogen atoms, resulting in a stable tetrahedral molecular geometry.