Topic 2 Electronic Structures of Atoms

2.1 Electronic energy levels of atomic hydrogen

What is the Bohr model of the hydrogen atom?

Niels Bohr introduced the atomic Hydrogen model in 1913. He described it as a positively charged nucleus, comprised of protons and neutrons, surrounded by a negatively charged electron cloud. The atom is held together by electrostatic forces between the positive nucleus and negative surroundings.

What are the spectrum of H-atom?

The spectrum of Hydrogen atom studied by Lyman, Balmer, Paschen, Brackett and Pfund can be explained on the basis of Bohr’s Model. It is clear that when an electron jumps from a higher energy state to a lower energy state, the radiation is emitted in form of photons

How is the emission spectrum for hydrogen created?

Hydrogen molecules are first broken up into hydrogen atoms (hence the atomic hydrogen emission spectrum) and electrons are then promoted into higher energy levels. Suppose a particular electron was excited into the higher energy level. This would tend to lose energy again by falling back down to a lower level.

What is the energy type produced in the hydrogen emission spectrum series of Lyman, Balmer, Paschen, Brackett and Pfund?

Lyman : ultraviolet emission lines

Balmer : visible region lines

Paschen : infrared region lines

Brackett : infrared region lines

Pfund : infrared region lines.

What is the visible light region?

The visible spectrum is the portion of the electromagnetic spectrum that is visible to the human eye. Electromagnetic radiation in this range of wavelengths is called visible light or simply light. A typical human eye will respond to wavelengths from about 380 to 740 nanometers.

How is the Balmer series produced in the hydrogen emission spectrum?

The Balmer series is a hydrogen spectral series of transitions and resulting ultraviolet emission lines of the hydrogen atom as an electron goes from n ≥ 3 to n = 2 (where n is the principal quantum number).

How is the Lyman series produced in the hydrogen emission spectrum?

The Lyman series is a hydrogen spectral series of transitions and resulting ultraviolet emission lines of the hydrogen atom as an electron goes from n ≥ 2 to n = 1 (where n is the principal quantum number), the lowest energy level of the electron.

How many emission lines are possible for hydrogen?

four. The electron energy level diagram for the hydrogen atom. It is found that the four visible spectral lines corresponded to transitions from higher energy levels down to the second energy level (n = 2).

Why are only 4 lines seen in the hydrogen emission spectrum?

Based on the wavelengths of the spectral lines, Bohr was able to calculate the energies that the hydrogen electron would have in each of its allowed energy levels. It is found that the four visible spectral lines corresponded to transitions from higher energy levels down to the second energy level (n = 2, Balmer series).

Who is balmer?

Johann Jakob Balmer (1 May 1825 – 12 March 1898) was a Swiss mathematician best known for his work in physics, the Balmer series of Hydrogen atom.

What is balmer series?

The Balmer series, or Balmer lines in atomic physics, is one of a set of six named series (Lyman, Balmer, Paschen, Brackett, Pfund) describing the spectral line emissions of the hydrogen atom, that result from electron transitions from higher levels down to the energy level with principal quantum number 2.

Which element has the simplest spectral lines?

Hydrogen has the simplest spectrum with 4 spectral lines, some show 5.

Why a single atom of hydrogen cannot produce all four hydrogen spectral lines simultaneously?

A single hydrogen atom only has one electron so it can’t have all four transitions at the same time.

What is the minimum energy required to ionize a hydrogen atom?

For a hydrogen atom, composed of an orbiting electron bound to a nucleus of one proton, an ionization energy of 2.18 × 10−18 joule (13.6 electron volts) is required to force the electron from its lowest energy level entirely out of the atom.

How electrons are arranged in an atom?

Electrons are arranged in shells around an atom’s nucleus. Electrons closest to the nucleus will have the lowest energy. Electrons further away from the nucleus will have higher energy. In a more realistic model, electrons move in atomic orbitals, or subshells.

What is the spectral line emissions of the hydrogen atom?

The emission spectrum of atomic hydrogen has been divided into a number of spectral series, with wavelengths given by the Rydberg formula. These observed spectral lines are due to the electron making transitions between two energy levels in an atom.

What are energy levels in an atom?

Energy levels (also called electron shells) are fixed distances from the nucleus of an atom where electrons may be found. Electrons are tiny, negatively charged particles in an atom that move around the positive nucleus at the center. Energy levels are a little like the steps of a staircase.

 

Explain the formation of the emission spectrum of hydrogen atom in the Balmer series?

When the electrons in the atom are excited, for example by being heated, the additional energy pushes the electrons to higher energy orbitals. When the electrons fall back down and leave the excited state, energy is re-emitted in the form of a photon. These emitted photons form the element’s spectrum.

What is the difference between emission and excitation?

The excitation spectrum and absorption spectrum of a molecule probe the excited states, whereas an emission spectrum probes the ground state. … When recording an excitation spectrum, the emission is measured at fixed wavelength while varying the excitation wavelength.

 

What is the purpose of the Bohr model?

The Bohr model shows that the electrons in atoms are in orbits of differing energy around the nucleus (think of planets orbiting around the sun). Bohr used the term energy levels (or shells) to describe these orbits of differing energy.

How many emission lines do you see for hydrogen?

Four. The four visible hydrogen emission spectrum lines in the Balmer series.

Why does hydrogen only emit 4 colors?

Although hydrogen has only one electron, it contains many energy levels. When its electron jumps from higher energy level to a lower one, it releases a photon. Those photons cause different colours of light of different wavelengths due to the different levels.

How many emission lines do you see for helium?

The 12 lines of the visible helium spectrum correspond to wavelengths of 388.8, 447.1, 471.3, 492.1, 501.5, 504.7, 587.5, 667.8, 686.7, 706.5, 728.1 and 781.3 nanometres (nm).

What are the main postulates (claim or assumption) of Bohr model?

Bohr’s model of the hydrogen atom is based on three postulates:

(1) an electron moves around the nucleus in a circular orbit,

(2) an electron’s angular momentum in the orbit is quantized, and

(3) the change in an electron’s energy as it makes a quantum jump from one orbit to another is always accompanied by the emission or absorption of a photon. Bohr’s model is semi-classical because it combines the classical concept of electron orbit (postulate 1) with the new concept of quantization (postulates 2 and 3).

The Bohr model postulates that electrons orbit the nucleus at fixed energy levels. Orbits further from the nucleus exist at higher energy levels. When electrons return to a lower energy level, they emit energy in the form of light.

Why the Bohr model works only for the hydrogen atom?

The Bohr model works only for hydrogen because it considers only the interactions between one electron and the nucleus. The Bohr model is based on the energy levels of one electron orbiting a nucleus at various energy levels. Any other electrons in the atom will repel the one electron and change its energy level.

Why does the Bohr model not work?

The main problem with Bohr’s model is that it works very well for atoms with only one electron, like H or He+, but not at all for multi-electron atoms. Bohr’s model breaks down when applied to multi-electron atoms. It does not account for sublevels (s,p,d,f), orbitals or elecrtron spin.

What are the limitations of Bohr’s theory of hydrogen atom?

Bohr’s model explains the elusive hydrogen spectral lines. But this model has its limitations / drawbacks. Bohr’s atomic model can not explain the Zeeman effect or Stark effect. And it also fails to explain the complicated spectral lines for atoms with more than one electron.

What is Zeeman effect and Stark effect?

The Zeeman and Stark effects are modifications of spectral lines (displacements, splittings, and polarization changes) induced respectively by magnetic and electric fields. Application of an external static field splits the energy levels of a degenerate multiplet of states.

What is the difference between Zeeman effect and Stark effect?

The main difference between Zeeman effect and Stark effect is that Zeeman effect is observed in the presence of an external magnetic field whereas Stark effect is observed in the presence of an external electrical field.

What is Zeeman effect?

Zeeman effect, in physics and astronomy, the splitting of a spectral line into two or more components of slightly different frequency when the light source is placed in a magnetic field.

What is meant by Stark effect?

Stark effect, the splitting of spectral lines observed when the radiating atoms, ions, or molecules are subjected to a strong electric field. The electric analogue of the Zeeman effect (i.e., the magnetic splitting of spectral lines), it was discovered by a German physicist, Johannes Stark (1913).

What is the limitation of Bohr model?

The Bohr Model is very limited in terms of size. Poor spectral predictions are obtained when larger atoms are in question. It cannot predict the relative intensities of spectral lines. It does not explain the Zeeman Effect, when the spectral line is split into several components in the presence of a magnetic field.

What are the weaknesses of the Bohr model?

The two weaknesses of Bohr’s atomic model are : 1) the path of electron around nucleus is considered to be circular of definite radius but in reality it can be at any distance from the nucleus. 2) the model is based on spectra of atoms and newton’s laws of motion which are not applicable to microscopic particles.

2.2 Atomic orbitals: s, p and d

What is a principal quantum number in chemistry? What does the principal quantum energy number indicate?

The principal quantum number is the quantum number denoted by n and which indirectly describes the size of the electron orbital. It is always assigned an integer value (e.g., n = 1, 2, 3…), but its value may never be 0. The principal quantum number describes on the energy of an electron and the most probable distance of the electron from the nucleus (energy level an electron is placed in).

What is subshells in chemistry?

A subshell, or l, is a subdivision of electron shells separated by electron orbitals. Subshells are labelled s, p, d, and f in an electron configuration.The number of subshells describes the shape of the orbital.

What are subshells and orbitals?

A subshell is a group of orbitals. Orbitals each hold two electrons, and electrons in an orbital with the same principle quantum number, angular momentum quantum number, and magnetic quantum number, ml are part of the same orbital. The ml is what distinguishes different orbitals in a subshell.

Why are orbitals called SPDF?

The orbital names s, p, d, and f stand for names given to groups of lines originally noted in the spectra of the alkali metals. These line groups are called sharp, principal, diffuse, and fundamental.

 

2.3 Electronic configuration

What is the electronic configuration of an element?

An atom’s electron configuration is a numeric representation of its electron orbitals. Electron orbitals are differently-shaped regions around an atom’s nucleus where electrons are mathematically likely to be located.

Correction on how to write electronic configuration?

Cr: 3d5 4s1 not 4s1 3d5
Cu : 3d10 4s1 not 4s1 3d10
Mo: 4d5 5s1 not 5s1 4d5

Why not we write Cr ‘s electronic configuration as [Ar] 4s¹ 3d5 or 4s² 3d⁴?

The correct way is 3d5 4s¹.
As 4s filled first, because at first 4s orbital is a lower energy. Then, 4s become higher energy after 3d get filled. One of the 4s electron is transfer to 3d orbital to make 3d more stable which is half-filled.

How do you write the electron configuration?

Electron configurations are written so as to clearly display the number of electrons in the atom as well as the number of electrons in each orbital. Each orbital is written in sequence, with the number of electrons in each orbital written in superscript to the right of the orbital name.

# ‘means number’
What is the electron configuration order?
Note that orbital sets are numbered by electron shell, but ordered in terms of energy. The order for filling orbitals is as follows:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, 8s.
How many electrons are in each shell?
Each shell can contain only a fixed number of electrons: The first shell can hold up to two electrons, the second shell can hold up to eight (2 + 6) electrons, the third shell can hold up to 18 (2 + 6 + 10) and so on. The general formula is that the nth shell can in principle hold up to 2(n2) electrons.
How many electrons can fill each orbital?
There can be two electrons in one orbital maximum. The s sublevel has just one orbital, so can contain 2 electrons max. The p sublevel has 3 orbitals, so can contain 6 electrons max. The d sublevel has 5 orbitals, so can contain 10 electrons max. The f sublevel has 7 orbitals, so can contain 14 electrons max.
What are the 3 rules of electron configuration?
When assigning electrons to orbitals, we must follow a set of three rules: the Aufbau Principle, the Pauli-Exclusion Principle, and Hund’s Rule.
Pauli exclusion principle define that all electrons in an atom have to have a unique set of quantum numbers.
Aufbau principle tells us to “build up” from the bottom of the energy well to the top. Pour water in a bucket and it fills from the bottom up – same idea.
Completely fill a lower level of energy before we advance to the next higher level. Here is the order of filling for all the orbitals in the atom.
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p
Filling an Electron Energy Level Diagram? How to fill those orbitals?
Use the periodic table to help you get the correct order of orbitals. The orbitals themselves are shown on an energy diagram as blanks and we will put in up arrows ↿ and down arrows ⇂ to represent the spin quantum numbers +½ and –½. Follow the rules:
1. Aufbau Principle – always take the lowest possible energy level and fill it before going up to the next level.
2. Hund’s Rule – on any degenerate levels (same energies), always fill singly with up arrows ↿ (+½) before you then pair with the down arrows ⇂ (–½). note: electrons with matching spin states are said to have parallel spin states.
3. Pauli Exclusion Principle – no two electrons can have the same set of four quantum numbers – these means no more that 2 electrons per orbital (blank) and when there are two, they have opposite spins ⥮.
How the orbitals “fit” the table by using Aufbau filling order (periodic table)?
Which Subshell fills first?
Rule 1 – Lowest energy orbitals fill first. Thus, the filling pattern is 1s, 2s, 2p, 3s, 3p, 4s, 3d, etc. Since the orbitals within a subshell are degenerate (of equal energy), the entire subshell of a particular orbital type is filled before moving to the next subshell of higher energy.
Rule 2 – Pauli Exclusion Principle – Only two electrons are permitted per orbital and they must be of opposite spin. If one electron within an orbital possesses a clockwise spin, then the second electron within that orbital will possess a counter clockwise spin. Two electrons with opposite spins found in the same orbital are referred to as being paired.
Rule 3- Hund’s Rule – The most stable arrangement of electrons in a subshell occurs when the maximum number of unpaired electrons exist, all possessing the same spin direction. This occurs due to the degeneracy of the orbitals, all orbitals within a subshell are of equal energy. Electrons are repulsive to one another and only pair after all of the orbitals have been singly filled.
Why is 3d higher than 4s?
According to Aufbau principle , electrons first occupy the lowest energy orbital available to them and enter into higher energy orbitals only after the lower energy orbitals are filled . Therefore , 3d orbital is higher in energy than 4s . And hence electrons fill up in 4s before filling up in 3d.

Which electrons are removed first?

The electronic configuration of cations is assigned by removing electrons first in the outermost p orbital, followed by the s orbital and finally the d orbitals (if any more electrons need to beremoved).
Why 4s electrons are removed before the 3d?
Because the 4s orbitals has the lower energy, it gets filled first. When 3d orbitals are filled, 4s is no longer lower in energy. Hence electrons are lost from 4s orbital first, because electrons lost first will come from the highest energy level (furthest away from the nucleus).

Which Subshell loses electrons first between 3d and 4s?

4s.
Since 4s orbital lies in the fourth shell, and 3d in the third shell, electrons are removed first from 4s. A simple rule to follow: Filling up of orbitals is dependent on orbital energy while removal of electrons from orbitals is dependent on orbital location. 4s has a lower energy than 3d. So it gets filled up first.
Write the electronic configuration of Cr?
Write the electronic configuration of Cu?
Write the electronic configuration of Se?
Write the electronic configuration from elements with atomic number of 40 until 47?
What element is 1s2 2s2 2p6 3s2 3p6 4s2 3d6?
iron, Fe. The rules on how to fill the orbitals, the electronic configuration of iron (for example) is 1s2 2s2 2p6 3s2 3p6 4s2 3d6 , and it is abbreviated form [Ar] 4s2 3d6.
What does 1s2 2s2 2p6 mean?
This represents 2 electrons in the s subshell of the first energy level, 2 electrons in the s subshell of the second energy level and 6 electrons in the p subshell of the second energy level.
What is isoelectronic?
having the same numbers of electrons or the same electronic structure. “carbon monoxide is isoelectronic with nitrogen”.
What is the isoelectronic series?
An Isoelectronic Series is a group of atoms/ ions that have the same number of electrons. Since the number of electrons are the same, size is determined by the number of protons.
Which species have the same electron arrangements?
Atoms and ions that have the same electron configuration are said to be isoelectronic. Examples of isoelectronic species are N3–, O2–, F–, Ne, Na+, Mg2+, and Al3+ (1s22s22p6). Another isoelectronic series is P3–, S2–, Cl–, Ar, K+, Ca2+, and Sc3+ ([Ne]3s23p6).

2.4 Classification of elements into s, p, d and f blocks in the Periodic Table

2 Electronic Structures of Atoms

2.1 Electronic energy levels of atomic hydrogen

Candidates should be able to:

(a) explain the formation of the emission line spectrum of atomic hydrogen in the Lyman and Balmer series using Bohr’s Atomic Model.

2.2 Atomic orbitals: s, p and d

Candidates should be able to:

(a) deduce the number and relative energies of the s, p and d orbitals for the principal quantum numbers 1, 2 and 3, including the 4s orbitals;
(b) describe the shape of the s and p orbitals.

2.3 Electronic configuration

Candidates should be able to:

(a) predict the electronic configuration of atoms and ions given the proton number (and charge);
(b) define and apply Aufbau principle, Hund’s rule and Pauli Exclusion Principle.

2.4 Classification of elements into s, p, d and f blocks in the Periodic Table

Candidates should be able to:

(a) identify the position of the elements in the Periodic Table as
(i) block s, with valence shell configurations s1 and s2,
(ii) block p, with valence shell configurations from s2p1 to s2p6,
(iii) block d, with valence shell configurations from d1s2 to d10s2;
(b) identify the position of elements in block f of the Periodic Table.

 

2 Electronic Structures of Atoms

2.1 Electronic energy levels of atomic hydrogen

Candidates should be able to:

(a) explain the formation of the emission line spectrum of atomic hydrogen in the Lyman and Balmer series using Bohr’s Atomic Model.

2.2 Atomic orbitals: s, p and d

Candidates should be able to:

(a) deduce the number and relative energies of the s, p and d orbitals for the principal quantum numbers 1, 2 and 3, including the 4s orbitals;

(b) describe the shape of the s and p orbitals.

2.3 Electronic configuration

Candidates should be able to:

(a) predict the electronic configuration of atoms and ions given the proton number (and charge);

(b) define and apply Aufbau principle, Hund’s rule and Pauli Exclusion Principle.

Aufbau principle

The aufbau principle (also called the building-upprinciple or the aufbau rule) states that in the ground state of an atom or ion, electrons fill atomic orbitals of the lowest available energy levels before occupying higher levels. For example, the 1s shell is filled before the 2s subshell is occupied.

Hund’s rule

The German Physicist, Fredrich Hund, formulated it around 1927, which allows to easily draw electron energy diagrams for a ground state multi- electron atom. Hund’s rule states: – Every orbital in a sublevel is occupied before they start doubling up -All electron in a single occupied orbital have the same spin (in order to maximize the total spin).
Hund´s rule is sometimes used to describe people getting seated on a bus. Everyone will first get their own seat, sitting alone. Not until there are no other seats people will start sitting next to a stranger.

Pauli Exclusion Principle

in an atom or molecule, no two electrons can have the same four electronic quantum numbers. As an orbital can contain a maximum of only two electrons, the two electrons must have opposing spins. This means if one is assigned an up-spin ( +1/2), the other must be down-spin (-1/2).

Electron filling in atomic orbitals following Aufbau principle [or building-up principle], Pauli exclusion principle [or Antisymmetry principle] and Hund’s rule of maximum multiplicity.
For copper

2.4 Classification of elements into s, p, d and f blocks in the Periodic Table

Candidates should be able to:

(a) identify the position of the elements in the Periodic Table as
(i) block s, with valence shell configurations s1 and s2,
(ii) block p, with valence shell configurations from s2p1 to s2p6,
(iii) block d, with valence shell configurations from d1s2 to d10s2;

(b) identify the position of elements in block f of the Periodic Table.







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