Period 4 Subshell Electronic Configuration Tutorial

Key Concepts

The number of the period of the Periodic Table tells us the highest energy level available to an atom in the ground state that can be filled with electrons:

period 1: 1^st energy level (K shell) being filled
period 2: 2^nd energy level (L shell) being filled
period 3: 3^rd energy level (M shell) being filled
period 4: 4^th energy level (N shell) being filled

The Periodic Table can be divided up into s, p and d blocks showing us which atoms are having their s, p and d subshells filled.

s block: Group 1, Group 2, hydogen (H) and helium (He)
d block: Groups 3 to 12 (transition metals)
p block: Groups 13 to 18

NOTE that the s and p subshells in a period belong to the same energy level, BUT the d subshell belongs to the preceding energy level.

	Group
	1	2	3	4	5	6	7	8	9	10	11	12	13	14	15	16	17	18
	s block 4^th Energy Level		d block 3^rd Energy Level										p block 4^th Energy Level
Period 4	Z=19 K	Z=20 Ca	Z=21 Sc	Z=22 Ti	Z=23 V	Z=24 Cr	Z=25 Mn	Z=26 Fe	Z=27 Co	Z=28 Ni	Z=29 Cu	Z=30 Zn	Z=31 Ga	Z=32 Ge	Z=33 As	Z=34 Se	Z=35 Br	Z=36 Kr

The Aufbau principle (or Aufbau rule) states that electrons orbiting an atom's nucleus fill the lowest available energy levels before filling higher energy levels.

An atom gains extra stability in one of two ways:
- exactly half-filling a subshell composed of more than 1 orbital (especially relevant to d subshells)
  Electronic configuration of chromium is [Ar] 3d⁵ 4s¹ NOT [Ar] 3d⁴ 4s²
- completely filling a subshell composed of more than one orbital (especially relevant to d subshells)
  Electronic configuration of copper is [Ar] 3d¹⁰ 4s¹ NOT [Ar] 3d⁹ 4s²

When written down, the subshell electronic configuration of period 4 elements follows the pattern:
1s²2s²2p⁶3s²3p⁶ 3d 4s 4p
[Ar]3d 4s 4p

Aufbau Principle (Aufbau Rule)

"Aufbau" is a german word meaning construction or "building up". The Aufbau principle (or Aufbau rule), formulated in the 1920s by Neils Bohr and Wolfgang Pauli, is used to help us write the electronic configuration of an atom in terms of energy levels and subshells. It is important to remember that this is a hypothetical "building-up" of electrons around an atom's nucleus, it is a generalised tool that enables us to write the electronic configuration of an atom but it does not necessarily represent the experimentally determined energy values for the subshells of particular atoms.¹

The Aufbau principle tells us to fill the lower energy subshells with electrons first, before filling subshells of higher energy with electrons.

For the first 20 elements we can represent the energy levels and subshells as shown below with the order of filling shown by red arrows:

↑	4^th energy level (N shell)	____	4s____(maximum of 2 electrons)	↑ 6^th
↑	3^rd energy level (M shell)	____	3p____(maximum of 6 electrons) 3s____(maximum of 2 electrons)	↑ 5^th ↑ 4^th
↑
↑	2^nd energy level (L shell)	____	2p____(maximum of 6 electrons) 2s____(maximum of 2 electrons)	↑ 3^rd ↑ 2^nd
↑
↑	1^st energy level (K shell)	____	1s____ (maximum of 2 electrons)	↑ 1^st
↑ energy	levels (shells)		sublevels (subshells)	filling order

But when we come to the first of the transition metals, d subshell in period 4, we hit a snag.

The s subshell of this period belongs to the 4^th energy level, BUT the d subshell belongs to the 3^rd energy level.
The subshells for these higher energy levels "bunch up" so that the 3d subshell, in general, is represented as being of slightly higher energy than the 4s subshell.

So our energy level diagram looks more like the one shown below with the "order of filling" given by the red arrows:

↑	4^th energy level (N shell)	____	4p____(maximum of 6 electrons) 3d____(maximum of 10 electrons) 4s____(maximum of 2 electrons)	↑ 8^th ↑ 7^th ↑ 6^th
↑	3^rd energy level (M shell)	____	3p____(maximum of 6 electrons) 3s____(maximum of 2 electrons)	↑ 5^th ↑ 4^th
↑
↑	2^nd energy level (L shell)	____	2p____(maximum of 6 electrons) 2s____(maximum of 2 electrons)	↑ 3^rd ↑ 2^nd
↑
↑	1^st energy level (K shell)	____	1s____ (maximum of 2 electrons)	↑ 1^st
↑ energy	levels (shells)		sublevels (subshells)	filling order

The Aufbau principle is usually represented as a "filling diagram" as shown below:

We follow the direction of the arrows to place electrons 1 at a time into their subshells in order to "build up" the electronic configuration of an atom.
We could write this "filling order" out as a series of steps:

Step 1: add electrons to fill the 1s subshell (maximum of 2 electrons), when this is full, go to step 2.
Step 2: add electrons to fill the 2s subshell (maximum of 2 electrons), when this is full, go to step 3.
Step 3: add electrons to fill the 2p subshell (maximum of 6 electrons), when this is full, go to step 4.
Step 4: add electrons to fill the 3s subshell (maximum of 2 electrons), when this is full, go to step 5.
Step 5: add electrons to fill the 3p subshell (maximum of 6 electrons), when this is full, go to step 6.
Step 6: add electrons to the 4s subshell (maximum of 2 electrons), when this is full, go to step 7.
Step 7: add electrons to the 3d subshell (maximum of 10 electrons), when this is full, go to step 8.
Step 8: add electrons to the 4p subshell (maximum of 6 electrons), when this is full, go to step 9.
etc

You can use the Aufbau principle to correctly predict the electronic configuration of the atoms of most elements.
However, if we consider an atom of chromium, Z=24, the Aufbau principle predicts its electronic configuration to be:

1s²2s²2p⁶3s²3p⁶4s²3d⁴
which, when we collect subshells of the same energy level together, is written as:
1s²2s²2p⁶3s²3p⁶3d⁴4s²

Experimental evidence tells us that the electronic configuration of an atom of chromium is actually:

1s²2s²2p⁶3s²3p⁶3d⁵4s¹

that is, the atom's ground state electronic configuration is more stable, of lower energy, if both the 4s and 3d subshells are half-filled.²

Similarly, when we come to copper, Z=29, the Aufbau principle predicts its electronic configuration to be:

1s²2s²2p⁶3s²3p⁶4s²3d⁹
which, when we collect subshells of the same energy level together, is written as:
1s²2s²2p⁶3s²3p⁶3d⁹4s²

Experimental evidence tells us that the electronic configuration of an atom of copper is actually:

1s²2s²2p⁶3s²3p⁶3d¹⁰4s¹

that is, the atom's ground state electronic configuration is more stable, of lower energy, if the 3d subshell is filled leaving the 4s subshell half-filled.

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s Block Electronic Configuration

There are 2 elements in the s block of period 4, potassium (K, Group 1) and calcium (Ca, Group 2).

Group 1

Group 2

Group 3

Group 4

Group 5

Group 6

Group 7

Group 8

Group 9

Group 10

Group 11

Group 12

Group 13

Group 14

Group 15

Group 16

Group 17

Group 18

s block

d block

p block

Period 4

Z=19
K

Z=20
Ca

Z=21
Sc

Z=22
Ti

Z=23
V

Z=24
Cr

Z=25
Mn

Z=26
Fe

Z=27
Co

Z=28
Ni

Z=29
Cu

Z=30
Zn

Z=31
Ga

Z=32
Ge

Z=33
As

Z=34
Se

Z=35
Br

Z=36
Kr

Just like we did for the atoms of elements in Periods 1, 2 and 3, we can draw a diagram representing the energy levels and subshells available to these s block elements of the fourth period:

↑	4^th energy level (N shell)	____	4s____(maximum of 2 electrons)
↑	3^rd energy level (M shell)	____	3p____(maximum of 6 electrons) 3s____(maximum of 2 electrons)
↑
↑	2^nd energy level (L shell)	____	2p____(maximum of 6 electrons) 2s____(maximum of 2 electrons)
↑
↑	1^st energy level (K shell)	____	1s____ (maximum of 2 electrons)
↑ energy	levels (shells)		sublevels (subshells)

The order for filling the subshells with electrons is therefore:

1s < 2s < 2p < 3s < 3p < 4s

Atoms of both potassium and calcium have the same "core" electronic configuration, that of the preceding Noble Gas (Group 18 element) which is argon,

1s² 2s² 2p⁶ 3s² 3p⁶
and electrons are being added to the s subshell of the fourth energy level, that is, electrons are being added to fill the 4s subshell which can house a maximum of 2 electrons:

1s² 2s² 2p⁶ 3s² 3p⁶ 4s
OR
[Ar] 4s

potassium (Z=19) has one electron positioned in the 4s subshell:

subshell electronic configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹
condensed electronic configuration: [Ar] 4s¹

calcium (Z=20) has two electrons located in the 4s subshell:

subshell electronic configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s²
condensed electronic configuration: [Ar] 4s²

With an atom of calcium, the 4s subshell is now full of electrons.

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d Block (transition metal) Electronic Configuration

There are 10 transition metal elements in the d block of this period:

Group 1

Group 2

Group 3

Group 4

Group 5

Group 6

Group 7

Group 8

Group 9

Group 10

Group 11

Group 12

Group 13

Group 14

Group 15

Group 16

Group 17

Group 18

s block
4^th Energy Level

d block
3^rd Energy Level

p block
4^th Energy Level

Period 4

Z=19
K

Z=20
Ca

Z=21
Sc

Z=22
Ti

Z=23
V

Z=24
Cr

Z=25
Mn

Z=26
Fe

Z=27
Co

Z=28
Ni

Z=29
Cu

Z=30
Zn

Z=31
Ga

Z=32
Ge

Z=33
As

Z=34
Se

Z=35
Br

Z=36
Kr

Scandium is the first of the transition metal or d block elements.
The d subshell for atoms of period 4 elements belongs to the third energy level, that is, the 3d subshell.
We now need to include this 3d subshell in our generalised "filling diagram":

↑	4^th energy level (N shell)	____	3d____(maximum of 10 electrons) 4s____(maximum of 2 electrons)
↑	3^rd energy level (M shell)	____	3p____(maximum of 6 electrons) 3s____(maximum of 2 electrons)
↑
↑	2^nd energy level (L shell)	____	2p____(maximum of 6 electrons) 2s____(maximum of 2 electrons)
↑
↑	1^st energy level (K shell)	____	1s____ (maximum of 2 electrons)
↑ energy	levels (shells)		sublevels (subshells)

The order for filling the subshells with electrons is therefore:

1s < 2s < 2p < 3s < 3p < 4s < 3d

Note that when we write the electronic configuration of these atoms we keep all the subshells of the same energy level together³: 1s 2s 2p 3s 3p 3d 4s
We can follow this sequence to write the electronic configurations of the first few d block elements:

scandium (Z=21)
Using the filling sequence: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹
Writing the electronic configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹ 4s² or [Ar] 3d¹4s²
titanium (Z=22)
Using the filling sequence: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d²
Writing the electronic configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d² 4s² or [Ar] 3d²4s²
vanadium (Z=23)
Using the filling sequence: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d³
Writing the electronic configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d³ 4s² or [Ar] 3d³4s²

But what happens next?
If we follow the filling sequence for chromium (Z=24) we would end up with a filled 4s subshell (4s²) and nearly half-filled 3d subshell (3d⁴).
So what?
The most stable electronic configurations for an atom occur in one of two ways:

completely filling a subshell composed of more than orbital (eg, d¹⁰)
or
exactly half-filling a subshell composed of more than orbital (eg, d⁵)

So the most stable electronic configuration for an atom of chromium is to have both its 4s and 3d subshells half-filled!

chromium (Z=24)
Using the filling sequence with half-filled 4s and 3d subshells: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d⁵
Writing the electronic configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵ 4s¹ or [Ar] 3d⁵4s¹

With manganese (Z=25) we can pair up that lone 4s electron while keeping the stable half-filled 3d subshell:

manganese (Z=25)
Using the filling sequence: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁵
Writing the electronic configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵ 4s² or [Ar] 3d⁵4s²

And we can continue to add more electrons to the 3d subshell for the next few atoms:

iron (Z=26)
Using the filling sequence: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶
Writing the electronic configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s² or [Ar] 3d⁶4s²
cobalt (Z=27)
Using the filling sequence: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁷
Writing the electronic configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁷ 4s² or [Ar] 3d⁷4s²
Nickel (Z=28)
Using the filling sequence: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁸
Writing the electronic configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁸ 4s² or [Ar] 3d⁸4s²

Until we arrive at copper (Z=29).
If we continue using the filling sequence we would end up with a filled 4s subshell (4s²), and a nearly completely full 3d subshell (3d⁹).
Now remember that the most stable electronic configurations for an atom occur in one of two ways:

completely filling a subshell composed of more than orbital (eg, d¹⁰)
or
exactly half-filling a subshell composed of more than orbital (eg, d⁵)

So the most stable electronic configuration for an atom of copper is to have a completely filled 3d subshell and a half-filled 4s subshell:

copper (Z=29)
Using filling sequence with half-filled 4s and completely filled 3d subshell: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d¹⁰
Writing the electronic configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s¹ or [Ar] 3d¹⁰4s¹

The last element in the d block of period 4 is zinc (Z=30) and it will have completely filled 4s and 3d subshells:

zinc (Z=30)
Using the filling sequence: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰
Writing the electronic configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² or [Ar] 3d¹⁰4s²

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p Block Electronic Configuration

There are 6 elements in the p block of the fourth period:

Group 1

Group 2

Group 3

Group 4

Group 5

Group 6

Group 7

Group 8

Group 9

Group 10

Group 11

Group 12

Group 13

Group 14

Group 15

Group 16

Group 17

Group 18

s block
4^th Energy Level

d block
3^rd Energy Level

p block
4^th Energy Level

Period 4

Z=19
K

Z=20
Ca

Z=21
Sc

Z=22
Ti

Z=23
V

Z=24
Cr

Z=25
Mn

Z=26
Fe

Z=27
Co

Z=28
Ni

Z=29
Cu

Z=30
Zn

Z=31
Ga

Z=32
Ge

Z=33
As

Z=34
Se

Z=35
Br

Z=36
Kr

The 4p subshell is of higher energy than the 4s and 3d subshells, so the representation we use for filling these subshells looks familiar:

↑	4^th energy level (N shell)	____	4p____(maximum of 6 electrons) 3d____(maximum of 10 electrons) 4s____(maximum of 2 electrons)
↑	3^rd energy level (M shell)	____	3p____(maximum of 6 electrons) 3s____(maximum of 2 electrons)
↑
↑	2^nd energy level (L shell)	____	2p____(maximum of 6 electrons) 2s____(maximum of 2 electrons)
↑
↑	1^st energy level (K shell)	____	1s____ (maximum of 2 electrons)
↑ energy	levels (shells)		sublevels (subshells)

The order for filling the subshells with electrons is therefore:

1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p

Gallium (Z=31) has the 4s and 3d subshells full and just 1 electron in a 4p subshell:

gallium (Z=31)
Using the filling sequence: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p¹
Writing the electronic configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p¹ or [Ar] 3d¹⁰4s²4p¹

1 electron is added to the 4p subshell for each subsequent atom in period 4:

germanium (Z=32)
Using the filling sequence: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p²
Writing the electronic configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p² or [Ar] 3d¹⁰4s²4p²
arsenic (Z=33)
Using the filling sequence: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p³
Writing the electronic configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p³ or [Ar] 3d¹⁰4s²4p³
selenium (Z=34)
Using the filling sequence: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁴
Writing the electronic configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁴ or [Ar] 3d¹⁰4s²4p⁴
bromine (Z=35)
Using the filling sequence: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁵
Writing the electronic configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁵ or [Ar] 3d¹⁰4s²4p⁵
krypton (Z=36)
Using the filling sequence: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶
Writing the electronic configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶ or [Ar] 3d¹⁰4s²4p⁶

The Group 18 (Noble Gas) element krypton is the last element in period 4 of the periodic table, and has all its available subshells completely filled with electrons.

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Summary

The electronic configurations of the first 36 elements is given below in subshell notation and as a condensed electronic configuration:

Period	name (symbol)	subshell electron configuration	condensed electron configuration
Period 1 (tutorial)	hydrogen (H)	1s¹
Period 1 (tutorial)	helium (He)	1s²
Period 2 (tutorial)	lithium (Li)	1s² 2s¹	[He] 2s¹
	beryllium (Be)	1s² 2s²	[He] 2s²
	boron (B)	1s² 2s² 2p¹	[He] 2s² 2p¹
	carbon (C)	1s² 2s² 2p²	[He] 2s² 2p²
	nitrogen (N)	1s² 2s² 2p³	[He] 2s² 2p³
	oxygen (O)	1s² 2s² 2p⁴	[He] 2s² 2p⁴
	fluorine (F)	1s² 2s² 2p⁵	[He] 2s² 2p⁵
	neon (Ne)	1s² 2s² 2p⁶	[He] 2s² 2p⁶
Period 3 (tutorial)	sodium (Na)	1s² 2s² 2p⁶ 3s¹	[Ne] 3s¹
	magnesium (Mg)	1s² 2s² 2p⁶ 3s²	[Ne] 3s²
	aluminium (Al)	1s² 2s² 2p⁶ 3s² 3p¹	[Ne] 3s² 3p¹
	silicon (Si)	1s² 2s² 2p⁶ 3s² 3p²	[Ne] 3s² 3p²
	phosphorus (P)	1s² 2s² 2p⁶ 3s² 3p³	[Ne] 3s² 3p³
	sulfur (S)	1s² 2s² 2p⁶ 3s² 3p⁴	[Ne] 3s² 3p⁴
	chlorine (Cl)	1s² 2s² 2p⁶ 3s² 3p⁵	[Ne] 3s² 3p⁵
	argon (Ar)	1s² 2s² 2p⁶ 3s² 3p⁶	[Ne] 3s² 3p⁶
Period 4	potassium (K)	1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹	[Ar] 4s¹
	calcium (Ca)	1s² 2s² 2p⁶ 3s² 3p⁶ 4s²	[Ar] 4s²
	scandium (Sc)	1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹ 4s²	[Ar] 3d¹ 4s²
	titanium (Ti)	1s² 2s² 2p⁶ 3s² 3p⁶ 3d² 4s²	[Ar] 3d² 4s²
	vanadium (V)	1s² 2s² 2p⁶ 3s² 3p⁶ 3d³ 4s²	[Ar] 3d³ 4s²
	chromium (Cr)	1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵ 4s¹	[Ar] 3d⁵ 4s¹
	manganese (Mn)	1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵ 4s²	[Ar] 3d⁵ 4s²
	iron (Fe)	1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s²	[Ar] 3d⁶ 4s²
	cobalt (Co)	1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁷ 4s²	[Ar] 3d⁷ 4s²
	nickel (Ni)	1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁸ 4s²	[Ar] 3d⁸ 4s²
	copper (Cu)	1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s¹	[Ar] 3d¹⁰ 4s¹
	zinc (Zn)	1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s²	[Ar] 3d¹⁰ 4s²
	gallium (Ga)	1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p¹	[Ar] 3d¹⁰ 4s² 4p¹
	germanium (Ge)	1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p²	[Ar] 3d¹⁰ 4s² 4p²
	arsenic (As)	1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p³	[Ar] 3d¹⁰ 4s² 4p³
	selenium (Se)	1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁴	[Ar] 3d¹⁰ 4s² 4p⁴
	bromine (Br)	1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁵	[Ar] 3d¹⁰ 4s² 4p⁵
	krypton (Kr)	1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶	[Ar] 3d¹⁰ 4s² 4p⁶

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¹Experimental evidence reveals that for the transition metals, the 3d subshells are actually of lower energy than the 4s subshell. Hence, period 4 transition metals ionise by losing the electrons in the highest energy levels first, that is, from the 4s subshell. These 4s electrons are shielded by the lower energy 3d electrons, making them easier to remove. For the transition metals it is energetically favourable for electrons to occupy both the 4s and 3d subshells rather than just filling the lower energy 3d subshell first.
The graphical representations of "filling order" are therefore only a general guide to writing an electronic configuration, in reality the actual energy values for subshells are different for each atom based on considerations of electrostaic-type attractions and repulsions between the species making up each atom, so that these general representations should be "tweaked" to represent the energy values for each particular atom.

²This is arrived at by the application of "Hund's Rule": if multiple orbitals of the same energy level are available, electrons fill unoccupied orbitals singly first before pairing up with an electron in the same orbital.

³Notice that writing the electronic configuration of the transition metals in this way, ie, 3d before 4s, actually gives a better representation of the relative energies of the 3d and 4s electrons, that is, 4s electrons are of higher energy than 3d electrons and hence are written last in the electronic conifguration.