Atomic Electronic Structure

• The electronic structure, or electronic configuration, refers to the arrangement of electrons around an atom's nucleus.
• Electrons in atoms exist in energy levels or shells around the nucleus called principal quantum levels.
• The closer an energy level is to the nucleus, the lower its energy.
• The lowest energy levels are occupied first by electrons.
• The electronic configuration represents the number of electrons in each energy level as they get further from the nucleus.
• For example:
  • Na is 2.8.1
  • Na+ is 2.8
  • Cl is 2.8.7
  • Cl is 2.8.8
• Recall: The maximum number of electrons in a principal quantum level n is given by 2n².

Maximum number of electrons for each level

• Level 1 can contain 2 electrons
• Level 2 can contain 8 electrons
• Level 3 can contain 18 electrons
• Level 4 can contain 32 electrons
• At high school level, note the maximum electrons for energy level n=3 is 18, not 8.

Electronic structures for potassium and calcium:

• Potassium is 2.8.8.2
• Calcium is 2.8.8.2
• Two electrons go to n=1, eight to n=2, eight to n=3, followed by two in n=4 and then n=3 takes up to a further ten electrons.
• The electronic configuration for zinc is 2.8.18.2 and n=4 continues to fill with electrons.
• For elements Sc to Zn, the 10 electrons are added ton = 3 if you lose electrons to form ions, Zn²⁺ and Mg²⁺; the two electrons in n=4 are removed first.
• Zinc, with electronic configuration 2.8.18.2, forms Zn²⁺ by losing 2 electrons in n=4.
  • Zinc is not in group IIA of the periodic table.
• Manganese has an electronic structure is 2.8.13.2
• Electrons in an atom occupy specific energy levels, not between them: energy is quantized.
• An atom has a nucleus surrounded by a cloud of electrons.
• Electronic configuration describes how electrons arrange in an atom.
• Each period corresponds to a shell and successive shells can hold 2, 8, 8, 18, 18, 32, 32 electrons.
• This is the simplest electronic configuration description, considering main group elements (s-block and p-block).

Quantum Numbers and Orbitals

• Each electron is described by four quantum numbers:
  • n = principal quantum number
  • l = angular momentum or azimuthal quantum number
  • ml = magnetic quantum number
  • ms = spin quantum number

Principal Quantum Number (n)

• It can have positive values of 1,2,3, etc.
• Relates to orbital size and energy.
• As n increases, the orbital gets larger and electrons spend more time away from the nucleus.
• Orbitals with the same n belong to the same shell.
• Shells are designated with letters:
  • K = 1
  • L = 2
  • M = 3
  • N = 4

Angular Momentum Quantum Number (l)

• It can have any positive integer from 0 to n-1.
• Within each shell of quantum number n, there are n-1 kinds of orbitals, each with a distinctive shape, denoted by l.
  • For n = 3, l is 0,1,2, referred to as s, p, d.
• Within the M shell (n=3), there are three kinds of orbitals, each unique in shape.
• Orbitals with the same n, but different l are said to be in different subshells of a given shell.

Subshells are denoted as follows:

• s has l=0
• p has l=1
• d has l=2
• f has l=3
• g has l=4
• To define a subshell within a particular shell, we write the value of n quantum number for the shell, followed by the letter designating the subshell.
• 2p denotes a subshell with quantum number n=2 and l=1.

Magnetic Quantum Number (ml)

• Allowed values are -l...0...+l.
• Distinguishes orbitals of a given n and l, i.e., orbitals of the same energy and shape but with different orientations.
• Orbitals in a given subshell differ only in their orientation in space, not their shape.
  • For l=0, ml=0 (s subshell), for l=1 (p subshell), ml= -1, 0, +1. The p subshell has three different orbitals.
• These orbitals have the same shape but different orientation in space.
• There are 2l + 1 orbitals in each subshell of quantum number l.

Permissible values of quantum number for atomic orbital's

• N=1, has l=0, ml=0 with subshells 1s, with 1 orbital
• N=2, has l=0, ml=0 with subshells 2s, with 1 orbital
• N=2, has l=1, ml=-1,0,+1 with subshells 2p, with 3 orbitals
• N=3, has l=0, ml=0 with subshells 3s, with 1 orbital
• N=3, has l=1, ml=-1,0,+1 with subshells 3p, with 3 orbitals
• N=3, has l=2, ml=-2,-1,0,+1,+2 with subshells 3d, with 5 orbitals
• N=4, has l=0, ml=0 with subshells 4s, with 1 orbital
• N=4, has l=1, ml=-1,0,+1 with subshells 4p, with 3 orbitals
• N=4, has l=2, ml=-2,-1,0,+1,+2 with subshells 4d, with 5 orbitals
• N=4, has l=3, ml=-3,-2,-1,0,+1,+2,+3 with subshells 4f, with 7 orbitals

Spin Quantum Number

• Possible values are +½ and -½.
• This refers to the two possible orientations of the spin axis of an electron.

Orbital Shapes

• An s orbital has a spherical shape, the size depends on the value of n which is the specific distance of the probability of finding the electron distribution.
• There are three p orbitals in each p subshell and all have the same basic shape which constitute of two lobes arranged along a straight line with the nucleus.
  • They are denoted as 2px, 2py, and 2pz.
  • 2px has its greatest electron probability along the x-axis, 2py has the same along the y-axis, and 2pz along the z-axis.
• Other p-orbitals, such as 3p, have the same shape with the value of n dictating difference.
• There are five d-orbitals, which have more complicated shapes than s and p.
• The f-orbitals are seven in number.

Electron Configuration

• This is the particular distribution of electrons among available subshells and can be specified in two ways:
  • Orbital diagrams: showing how the orbitals of a subshell are occupied by electrons. Electrons are shown by arrows, pointing up when ms = +½ and down when ms = -½.
  • Spectroscopic notation: lists subshell symbols, one after the other, with right superscripts of giving number of electrons in each subshell.
• H atom spectroscopic notation is 1s¹.
• Noble gas notation H atom orbital notation or 1s¹ spectroscopic notation. ex 1s22s22p¹

Pauli Exclusion Principle

• No two electrons in an atom can have all four quantum numbers the same (i.e., identical sets of quantum numbers).
• No atomic orbital contains more than two electrons.
• For example, an electron in a 1s orbital has the label 1s¹ where n=, l=0, ml=0, ms=-½ or +½ where the electron spin arrow may pint either up or down.
• Two electrons in the 1s orbital can be named as: 1s² with n=1, l=0, ml=0, ms=+½ (for electron facing up), and 1s¹ n=1, l=0, ml=0, and ms=-½ (for electron facing down).

Order of Subshell Energies and Assignments

• The energies of the subshell for multi-electrons depend on both n and l
• the arrangement of different energies for a given atom where n=3, are in the order; 3s < 3p < 3d.
• Electrons are assigned to subshells in order of increasing (n and l ) values.
• Where two subshells share the same (n & l) value , the electrons are first assigned to the subshell of lower n.
• Building up a configuration follows the order; write the subshells in rows, each row having subshells of a given n.
• Within a row, arrange the increasing subshells by l (starting at nf subshells) because elements containing g or higher subshells are not known. Starting then with the 1s subshell, draw a series of diagnal.

Aufbau Principle

• A scheme used to reproduce the electron configuration of ground states atoms by filling subshells with electrons in a specific manner and order .

Applying this principle produces; the electron configuration for an atom by filling subshells successively in this following order;

• 1s2s2p3s3p4s3d4p5s4d5p6s4f5d6p7s5f6d7p
• Ground state is the configuration associated with lower-energy level in an atom is its ground state configuration, ex for Sodium atom is 1s22s22p63s¹m
• The electron configuration of the main group elements with Z = 1 to 10.
• For H, with Z=1, electron configuration is 1s1.
• For He, with Z=2, electron configuration is 1s2 or [He].
• For Li, with Z=3, electron configuration is 1s2 2s1 or [He]2s1.
• For Be, with Z=4, electron configuration is 1s2 2s2 or [He]2s2.
• For B, with Z=5, electron configuration is 1s2 2s2 2p1 or [He] 2s2 2p1
• For C, with Z=6, electron configuration is 1s2 2s2 2p2 or [He] 2s2 2p2.
• For N, with Z=7, electron configuration is 1s2 2s2 2p3 or [He] 2s2 2p3.
• For O, with Z=8, electron configuration is 1s2 2s2 2p4 or [He] 2s2 2p4.
• For F, with Z=9, electron configuration is 1s2 2s2 2p5 or [He] 2s2 2p5.
• For Ne, with Z=10, electron configuration is 1s2 2s2 2p6 or [He] 2s2 2p6.
• All group 1A elements have a single electron assigned to an s orbital in the nth shell, where n is the period for the element.

For Example

• K is the first element in period 4 (n=4) where its electron configuration is 1s2 2s2 2p6 3s2 3p6 4s1.
• The noble gas preceding K is argon (Z=18), where its electron configuration is 1s2 2s2 2p6 3s2 3p6.
• K electron configuration can be written as [Ar]4s1 as a noble gas notation.
• All group 2A elements configuration of [electrons preceding noble gas ] ns² where n is the period in which the element is found.
• Elements of group IA have configuration ns1 and those of group IIA have configuration of ns2, these elements are called S-block elements.
• All elements from group IIIA to VIIIA have electron configuration of [electrons preceding noble gas ] ns²npx.
• As they have electrons in the p-orbitals, they are called the p-block elements.

Electrons

• Electrons included in the noble gas configuration [ ] are called core electrons.
• Electrons beyond the core are called valence electrons.
• Hunds rule: the most stable arrangements of electrons is where the maximum number of unpaired electrons present, all with same spin
• For Example: C configuration can be stated as 1s22s22p² or [He]2s²2p² in respect to Hunds rule
• Transition elements are elements whose atoms are filling the d-subshell.
• Lanthanides are also referred to as inner transition elements because those elements whose atoms are filling the f-subshell, are sometimes called inner transition elements and fill the 4f orbitals. Actinides =those elements which have filling of 5f-orbitals.
• Transition elements are always preceded by s-block elements.
• The first element in the transition series, scandium (Sc), and had a electron configuration follows as[Ar]3d14s2

Transition elements

• The electron distribution of the Transition elements are in accordance with; Ar3d14s² such as Sc, Ti and V

For Example

• Cr is [Ar]3d44s²
• Mn is [Ar]3d54s2
• Fe is [Ar]3d64s2
• Co is [Ar]3d74s²
• Ni is [Ar]3d84s2
• Cu is [Ar]3d104s1
• Zn is [Ar]3d104s2
• The expected configuration of the chromium atom is [Ar]3d44s² but the actual configuration is [Ar]3d54s1.
• This is explained by assuming that the 4s and 3d orbitals have approximately the same. Occasionally, there are minor configuration differences.

Electron Configuration of Ions

• To produce cations from a neatral atom. the general rule is that one more electrons are be ejected from the highest shell number and so if choice of subshell exists, electron/s of max L are ejected.
  • For example Na = [1s²2s²2p⁶3s¹] will remove highest shell, thus producing 1+ as [1s²2s²2p⁶]

Atoms

• Atoms and ions with are unpaired and become, attracted to the magnetic field, known as Paramagnetic
• While if have electrons paired, are referred to as diamagnetic and or repelled by the magnetic freld.

Atomic properties and periodic trends.

• Similarities in properties of elements result from a very similar shell configuration during the valence state.

Atomic radius

• Sizes; for main group elements, atomic radii increases top to bottom but decreases along the period.
• The radius will by experimental result is 2* atoms by dividing by diameter (centre of molecules).
• So atom in 2x Chlorine by Pico meter/ 10-12, value the covalent by 100pm.
• So diameter distance for 2x Carbon will C-C determined at covalent is at 77pm by 154pm.
• Predicted with Carbon and Chloride between CCl4 is value and experimental for carbon is 176pm. And the estimate will provide correct.

Atomic Radius Trend

• It is shown that these trends depend on the orbital in question
• The principal quantum number n of the outer orbital shell must take in to consideration
• A lower number corresponds to smaller orbital

The other important component is the effective nuclei charge. Effective nuclei charge

it electrons, increasing the size of the orbitals. For any particular electron subshell, an increasing effective nuclei charge will reduce orbital numbers. Beryllium, being the effective has be taken in that outer are large, with electron charge.

• The periodic trends of atoms transition metals.
• The sizes middle, smaller than the largest.
• But effect realization variation electron most is small atom determine of is is most electron orbital.

For example

• outer increases with (Z number) by electron.

Periodic trends

• Recall the orbital sizes the electrons The charge must equal net balance of effective nucleus Its the shielding and inter electron . (1s²2s¹). charge nucleus is .

The effective nucleus

• Shielded by one electron is not higher than positive charge, that outer electrons require much more effective nuclear charge.

• Its principal quantum number (n), the value of n correspond to the shells around the nucleus. Thus as n increases the orbital becomes larger and electrons spends more time further away from the nucleus

The Effective nuclear charge Zeff acting on the orbital electrons this is a number

• It is equal to the positive the number of neutrons (nucleus), but minus the average the elections numbers

for example

Zeff or shield the inner effects positive , shield the inner effects (screening)

Atomic orbitals and quantum numbers:

• Increasing instant increase value of atomic number
• The shell means atomic weight be stable value

Remember

• Outer in lower is is for is core
• And value
• Effective steady quantum will atomic size

More orbital size and effects:

Ionization Energy and Shell Propreties

Orbitals of Shell number, or, or, En

Explore atomic electronic structure, energy levels, and electronic configurations. Understand quantum numbers, subshells, and electron behavior in atoms. Learn about electronic configuration of elements like Cl, Na+, and K.