The given electronic configuration is (n−1)d2ns2. We need to determine the group number for the element when n=4. First, substituting n=4 into the electronic configuration, we have: (4−1)d24s2=3d24s2‌. ‌ This configuration indicates that the element has two electrons in the 3d orbital and two electrons in the 4s orbital. The total number of valence electrons that determine the group of the element in the periodic table is the sum of the electrons in these orbitals. To find the group, we sum up: 2(‌ from ‌3d2)+2(‌ from ‌4s2)=4‌ electrons. ‌ Now, looking at the groups of the Periodic Table where the transition metals lie (elements with partially filled d-orbitals), the counting of the groups begins from Group 3 for elements with outer shell d-electrons starting to fill. This configuration corresponds to the 4th group primarily because d-block elements (transition metals) are categorized from Groups 3 to 12 , and the counting for this block starts from d1 configuration in Group 3. As for 3d2, this increments by two from Group 3 to reach Group 4: For d1, the group is 3 . Thus for d2, the group should be 4 (one step higher). Hence, the element with the configuration 3d24s2 would be located in group 4 of the periodic table, matching with the elements of the titanium family, which includes titanium (Ti) with an electron configuration of [Ar]3d24s2 in its ground state. The correct answer is therefore: Option A: 4.