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For the ones like Cu and Ag, such as Au, they have different electron configurations because it's more stable to have a full d-orbital so they borrow an electron from the s-orbital. So whenever you see one of those you know to make that change.
Those we unfortunately just have to memorize. It happens because of electron-electron repulsion in orbitals when filling up the d orbital. In the case of Copper, it has 9/10 electrons in the d orbital which causes an imbalance of paired vs the one non-paired electrons. A more stable structure is one where an electron is "stolen" from the 4s subshell and placed into the last 3d orbital to completely fill it up.
For the D orbital, a rule of thumb is to assume that the group 6 and group 11 elements take an s electron from the closest filled s state to make the d shell full, which is also a state of lower energy than having a for example.
The exceptions would be in the same group of the periodic table. The most common exceptions are Cr, Mo, Ag, & Cu. Cr & Mo are exceptions because of the greater stability that a half-filled subshell (d10) provides and Ag & Cu are exceptions because of the greater stability that a fully filled subshell (d10) provides.
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