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An outer electron is mostly more distant from the nucleus than an inner electron. Therefore the mean radial repulsion of the outer one by the inner one aims outwards.

There is said mostly above, as there is significant probability that e.g. 2s or 2p electrons are closer to the nucleus than 1s electrons, even if vice versa is (much) more probable.

The resulting effect is like if the outer electron were attracted by a nucleus of less than actual charge. This effect is called nucleus shielding/screening.

The effect intensity for 2 considered orbitals depends on their mutual probabilistic overlapping (more overlapping = less screening) and inner orbital shape (closer to spherical symmetry = better screening)

Electrons in s orbitals have significant occurance probability near nucleus and in inner regions. Therefore, they are less shielded than electrons in other orbitals with the same $n$. and have the lowest energy of them.

OTOH, due their spherical symmetry, they are the best in shielding, unless overlap level overrules that.

Note that orbitals are theoretical constructs and not physical objects. The term orbital is primarily the wave function describing the particular electron quantum state. But it can be also understood as this quantum state itself or the 3D geometrical representation of the spatial presence probability of such an electron.

An outer electron is mostly more distant from the nucleus than an inner electron. Therefore the mean radial repulsion of the outer one by the inner one aims outwards.

There is said mostly above, as there is significant probability that e.g. 2s or 2p electrons are closer to the nucleus than 1s electrons, even if vice versa is (much) more probable. That is because of orbital geometrical overlapping.

The resulting effect is like if the outer electron were attracted by a nucleus of less than actual charge. This effect is called nucleus shielding/screening.

The effect intensity for 2 considered orbitals depends on their mutual probabilistic overlapping (more overlapping = less screening) and inner orbital shape (closer to spherical symmetry = better screening)

Electrons in s orbitals have significant occurance probability near nucleus and in inner regions. Therefore, they are less shielded than electrons in other orbitals with the same $n$ and have the lowest energy of them.

OTOH, due their spherical symmetry, they are the best in shielding, unless overlap level overrules that.

An outer electron is mostly more distant from the nucleus than an inner electron. Therefore the mean radial repulsion of the outer one by the inner one aims outwards.

There is said mostly above, as there is significant probability that e.g. 2s or 2p electrons are closer to the nucleus than 1s electrons, even if vice versa is (much) more probable.

The resulting effect is like if the outer electron were attracted by a nucleus of less than actual charge. This effect is called nucleus shielding/screening.

The effect intensity for 2 considered orbitals depends on their mutual probabilistic overlapping (more overlapping = less screening) and inner orbital shape (closer to spherical symmetry = better screening)

ns orbitals have significant n-1 local inner radial probability maxima and therefore are generally shielded less than p or d orbitals, therefore have lower energy.

OTOH, due their spherical symmetry, they are the best in shielding, unless overlap level overrules that.

An outer electron is mostly more distant from the nucleus than an inner electron. Therefore the mean radial repulsion of the outer one by the inner one aims outwards.

There is said mostly above, as there is significant probability that e.g. 2s or 2p electrons are closer to the nucleus than 1s electrons, even if vice versa is (much) more probable.

The resulting effect is like if the outer electron were attracted by a nucleus of less than actual charge. This effect is called nucleus shielding/screening.

The effect intensity for 2 considered orbitals depends on their mutual probabilistic overlapping (more overlapping = less screening) and inner orbital shape (closer to spherical symmetry = better screening)

Electrons in s orbitals have significant occurance probability near nucleus and in inner regions. Therefore, they are less shielded than electrons in other orbitals with the same $n$. and have the lowest energy of them.

OTOH, due their spherical symmetry, they are the best in shielding, unless overlap level overrules that.

An outer electron is mostly more distant from the nucleus than an inner electron. Therefore the mean radial repulsion of the outer one by the inner one aims outwards.

There is said mostly above, as there is significant probability that e.g. 2s or 2p electrons are closer to the nucleus than 1s electrons, even if vice versa is (much) more probable.

The resulting effect is like if the outer electron were attracted by a nucleus of less than actual charge. This effect is called nucleus shielding/screening.

The effect intensity for 2 considered orbitals depends on their mutual probabilistic radial overlapping (more overlapping = less screening) and inner orbital shape (closer to spherical symmetry = better screening)

ns orbitals have significant n-1 local inner radial probability maxima and therefore are generally shielded less than p or d orbitals, therefore have lower energy.

OTOH, due their spherical symmetry, they are the best in shielding, unless overlap level overrules that.

An outer electron is mostly more distant from the nucleus than an inner electron. Therefore the mean radial repulsion of the outer one by the inner one aims outwards.

There is said mostly above, as there is significant probability that e.g. 2s or 2p electrons are closer to the nucleus than 1s electrons, even if vice versa is (much) more probable.

The resulting effect is like if the outer electron were attracted by a nucleus of less than actual charge. This effect is called nucleus shielding/screening.

The effect intensity for 2 considered orbitals depends on their mutual probabilistic overlapping (more overlapping = less screening) and inner orbital shape (closer to spherical symmetry = better screening)

ns orbitals have significant n-1 local inner radial probability maxima and therefore are generally shielded less than p or d orbitals, therefore have lower energy.

OTOH, due their spherical symmetry, they are the best in shielding, unless overlap level overrules that.