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3 Atomic Structure

An atom is the smallest representative particle of an element. Each element has a unique kind of

atoms that differs from the atoms of all other elements. The chemical behavior of an atom is directly

related to its atomic structure. Atoms are composed of three major kinds of subatomic particles:

Characteristics of the basic particles of an atom

Particle Symbol Absolute mass (kg) Relative mass Charge (C)

Electron e 9.109 × 10

−31

5.486 × 10

−4

−1.602 × 10

−19

Proton p 1.673 × 10−27 1.007 +1.602 × 10−19

Neutron n 1.675 × 10−27 1.009 0.000

An atom consists of a very small, extremely dense, and positively charged nucleus around which neg-

atively charged electrons of relatively very low masses are moving (Rutherford 1911). Nuclei are

clusters of protons and neutrons. For a neutral atom, the number of protons always equals the number

of electrons.

Atomic Nucleus

The number of protons in the nucleus is represented by the proton (atomic) number Z. This number

defines the position of an element in the periodic table. The number of all nucleons is given by the

nucleon (mass) number A. According to the convention, each nuclide can be characterized by the ex-

pression

. Atoms with the same atomic and different nucleonic number are called isotopes. The

majority of natural elements exist as a mixture of several isotopes, one usually predominating over the

others. For example hydrogen is in the form of three isotopes

(protium),

(D, deuterium), H

(T, tritium), and the participation of the isotope

is 99.985 %. Tabulated relative atomic masses are

stated as the mean values of the given mixture of isotopes. Isotopes do not differ in the number of

valence electrons; therefore, they do not differ in their chemical properties. However, their physical

properties can be different.

The existence of isotopes is reflected in the average mass of an atom of an element: The relative atom-

ic masses (weights) listed in tables are the average relative masses of an atom of elements determined

by considering the contribution of each natural isotope.

Most naturally occurring nuclides are stable and retain their structure indefinitely.

However, some nuclides spontaneously decay over time - they are radioactive. Radioactivity depends

on the instability of nuclei which decay by giving off nuclear particles (alpha or beta accompanied, in

some cases by gamma radiation) and forming new nuclei with different atomic numbers.

For example: 1

3H → 2

3He + β-

(β− particle is a high-speed electron emitted from the original nucleus when a neutron decays to form a proton)

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3 Atomic Structure

An atom is the smallest representative particle of an element. Each element has a unique kind of

atoms that differs from the atoms of all other elements. The chemical behavior of an atom is directly

related to its atomic structure. Atoms are composed of three major kinds of subatomic particles:

Characteristics of the basic particles of an atom

Particle Symbol Absolute mass (kg) Relative mass Charge (C)

Electron e 9.109 × 10 −31^ 5.486 × 10 −4^ −1.602 × 10 −

Proton p 1.673 × 10 −27^ 1.007 +1.602 × 10 −

Neutron n 1.675 × 10 −27^ 1.009 0.

An atom consists of a very small, extremely dense, and positively charged nucleus around which neg-

atively charged electrons of relatively very low masses are moving (Rutherford 1911). Nuclei are

clusters of protons and neutrons. For a neutral atom, the number of protons always equals the number

of electrons.

Atomic Nucleus

The number of protons in the nucleus is represented by the proton (atomic) number Z. This number

defines the position of an element in the periodic table. The number of all nucleons is given by the

nucleon (mass) number A. According to the convention, each nuclide can be characterized by the ex-

pression X

A

Z. Atoms with the same atomic and different nucleonic number are called isotopes. The

majority of natural elements exist as a mixture of several isotopes, one usually predominating over the

others. For example hydrogen is in the form of three isotopes H

1 (protium),^ H

1 (D, deuterium),^ H

3 1

(T, tritium), and the participation of the isotope 11 His 99.985 %. Tabulated relative atomic masses are

stated as the mean values of the given mixture of isotopes. Isotopes do not differ in the number of

valence electrons; therefore, they do not differ in their chemical properties. However, their physical

properties can be different.

The existence of isotopes is reflected in the average mass of an atom of an element: The relative atom-

ic masses (weights) listed in tables are the average relative masses of an atom of elements determined

by considering the contribution of each natural isotope.

Most naturally occurring nuclides are stable and retain their structure indefinitely.

However, some nuclides spontaneously decay over time - they are radioactive. Radioactivity depends

on the instability of nuclei which decay by giving off nuclear particles (alpha or beta accompanied, in

some cases by gamma radiation) and forming new nuclei with different atomic numbers.

For example: 13 H → 23 He + β-

(β−^ particle is a high-speed electron emitted from the original nucleus when a neutron decays to form a proton)

The nuclear sciences deal with the structure and behavior of atomic nuclei. On the contrary, chemistry

deals with the structure and behavior of atoms and molecules, with the chemical reactions of elements

and compounds supposing atomic nuclei to be mostly indefinitely stable.

Electronic Configuration of Atoms

In neutral atoms, the number of electrons surrounding the nuclei and forming their electron "covering"

or "casing" equals the number of protons in atomic nuclei. The most stable arrangement of the elec-

trons in the atom (i.e. the one having the lowest energy) is called the electron configuration. Each

element has a unique electronic configuration; the differences in it give different elements different

chemical and physical characteristics.

Electronic Structure

Several models have been created to describe the composition of the electron shell, trying to explain

the physical and chemical properties of the elements satisfactorily. Recently, the quantum-mechanical

model (wave mechanical) model of the atom is accepted.

According to this theory, the electron moving in three-dimensional space surrounding the nucleus has

the properties of a mass particle and a wave at the same time and its movement is characterized by the

wave function Ψ (psi). The energy of electrons is quantized, which means, that electrons appear to be

only in certain energetic states. Their way of movement is not exactly defined, but they are found in

some areas of the atom with a certain degree of probability. The border area, which encloses the place

with the highest probability of the presence of an electron with a certain energetic content, is called

orbital and it is defined as the square of the wave function. In order to fully define the electron in a

certain orbital, we use four quantum numbers (see Table).

Quantum numbers and types of orbital

Quantum number Subshell quantum number l

Name Symbol Value Orbital type

Principal

Subshell

Magnetic

Spin

n

l

m

s

p

d

f

The principal quantum number n describes the main energy level, in which the electron is present

in its ground state. Values of n range from 1 to 7 in the unexcited states of the known elements. As n

increases, electrons are generally farther from the nucleus and have higher potential energy. The elec-

trons with the same principal quantum number are contained in the same electron layer (sphere, shell).

The subshell quantum number l designates the different energy subshells within the main level. It

also indicates the general shape of orbitals and its size in connection with the main quantum number.

Only certain values of l are allowed. These depend on n : l = 0, 1, 2, …, ( n − 1) for the given n. The

size of orbital of the same type increases with the increasing value of the main quantum number. In-

Quantum numbers and their mutual relationships

Quantum number

Type of orbital

Maximal number

Main Subshell Magnetic^ of electrons

1 0 0 1 s 2

2 s

2 p

3 s

3 p

3 d

4 s

4 p

4 d

4 f

The spin quantum number characterizes the spin momentum of an electron in the given orbital. It

has only two possible values +1/2 and −1/2.

Rules for Assigning Electrons to Atomic Orbitals

There are three rules according to which can be with few exceptions determined the electron configu-

ration of an element.

The first rule is called the building up principle (aufbau principle). Electrons are added to orbitals

in the order of increasing energy.

Following this principle, the orbitals are filled in this sequence: 1 s , 2 s , 2 p , 3 s , 3 p , 4 s , 3 d , 4 p , 5 s , 4 d ,

5 p , 6 s , 5 d , 4 f , 6 p, …

The second rule, which specifies the building up of the electron shell, is the Pauli exclusion princi-

ple. No two electrons can have the same four quantum numbers. One orbital can contain two electrons

with the opposite spin maximally and they create an electron pair. According to this rule, the smallest

repulsive forces are between two electrons in one orbital only if they have the opposite spin.

The third rule of filling up of the orbitals is called the Hund’s rule. Orbitals of equal energy (degen-

erate orbitals) are each occupied by a single electron before a second electron, which will have the

opposite spin quantum number, enters any of them. Unpaired electrons in degenerate orbitals have the

same spin.

The Electron Configuration of Elements

Following the above-mentioned rules we can predict the electron distribution of any element and

demonstrate the relationship between the configuration of the electron shell of an element and its posi-

tion in the periodic table.

The electron building up of the first 18 elements of the periodic table

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

H ↑

He ↑↓ Li (^) ↑↓ ↑ Be (^) ↑↓ ↑↓ B ↑↓ ↑↓ ↑ C (^) ↑↓ ↑↓ ↑ ↑ N (^) ↑↓ ↑↓ ↑ ↑ ↑ O (^) ↑↓ ↑↓ ↑↓ ↑ ↑ F ↑↓ ↑↓ ↑↓ ↑↓ ↑ Ne (^) ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ Na (^) ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ Mg ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ Al (^) ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ Si (^) ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ P (^) ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ ↑ S (^) ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ Cl (^) ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ Ar (^) ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ K ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ Ca (^) ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ Sc (^) ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑↓ Ti (^) ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ ↑↓ V (^) ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ ↑ ↑↓ Cr (^) ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ ↑ ↑ ↑↓ Mn (^) ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ ↑ ↑ ↑ ↑↓ Fe ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ ↑ ↑ ↑↓ Co (^) ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ ↑ ↑↓ Ni ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ ↑↓ Cu (^) ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑↓ Zn (^) ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ Ga (^) ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ Ge (^) ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ As (^) ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ ↑ Sc (^) ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ Br (^) ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ Kr (^) ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓

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Download atoms orbitals electron and more Study notes Chemistry in PDF only on Docsity!

3 Atomic Structure

An atom is the smallest representative particle of an element. Each element has a unique kind of

atoms that differs from the atoms of all other elements. The chemical behavior of an atom is directly

related to its atomic structure. Atoms are composed of three major kinds of subatomic particles:

Characteristics of the basic particles of an atom

Particle Symbol Absolute mass (kg) Relative mass Charge (C)

Electron e 9.109 × 10 −31^ 5.486 × 10 −4^ −1.602 × 10 −

Proton p 1.673 × 10 −27^ 1.007 +1.602 × 10 −

Neutron n 1.675 × 10 −27^ 1.009 0.

An atom consists of a very small, extremely dense, and positively charged nucleus around which neg-

atively charged electrons of relatively very low masses are moving (Rutherford 1911). Nuclei are

clusters of protons and neutrons. For a neutral atom, the number of protons always equals the number

of electrons.

Atomic Nucleus

The number of protons in the nucleus is represented by the proton (atomic) number Z. This number

defines the position of an element in the periodic table. The number of all nucleons is given by the

nucleon (mass) number A. According to the convention, each nuclide can be characterized by the ex-

pression X

A

Z. Atoms with the same atomic and different nucleonic number are called isotopes. The

majority of natural elements exist as a mixture of several isotopes, one usually predominating over the

others. For example hydrogen is in the form of three isotopes H

1 (protium),^ H

1 (D, deuterium),^ H

(T, tritium), and the participation of the isotope 11 His 99.985 %. Tabulated relative atomic masses are

stated as the mean values of the given mixture of isotopes. Isotopes do not differ in the number of

valence electrons; therefore, they do not differ in their chemical properties. However, their physical

properties can be different.

The existence of isotopes is reflected in the average mass of an atom of an element: The relative atom-

ic masses (weights) listed in tables are the average relative masses of an atom of elements determined

by considering the contribution of each natural isotope.

Most naturally occurring nuclides are stable and retain their structure indefinitely.

However, some nuclides spontaneously decay over time - they are radioactive. Radioactivity depends

on the instability of nuclei which decay by giving off nuclear particles (alpha or beta accompanied, in

some cases by gamma radiation) and forming new nuclei with different atomic numbers.

For example: 13 H → 23 He + β-

(β−^ particle is a high-speed electron emitted from the original nucleus when a neutron decays to form a proton)

The nuclear sciences deal with the structure and behavior of atomic nuclei. On the contrary, chemistry

deals with the structure and behavior of atoms and molecules, with the chemical reactions of elements

and compounds supposing atomic nuclei to be mostly indefinitely stable.

In neutral atoms, the number of electrons surrounding the nuclei and forming their electron "covering"

or "casing" equals the number of protons in atomic nuclei. The most stable arrangement of the elec-

trons in the atom (i.e. the one having the lowest energy) is called the electron configuration. Each

element has a unique electronic configuration; the differences in it give different elements different

chemical and physical characteristics.

Electronic Structure

Several models have been created to describe the composition of the electron shell, trying to explain

the physical and chemical properties of the elements satisfactorily. Recently, the quantum-mechanical

model (wave mechanical) model of the atom is accepted.

According to this theory, the electron moving in three-dimensional space surrounding the nucleus has

the properties of a mass particle and a wave at the same time and its movement is characterized by the

only in certain energetic states. Their way of movement is not exactly defined, but they are found in

some areas of the atom with a certain degree of probability. The border area, which encloses the place

with the highest probability of the presence of an electron with a certain energetic content, is called

orbital and it is defined as the square of the wave function. In order to fully define the electron in a

certain orbital, we use four quantum numbers (see Table).

Quantum numbers and types of orbital

Quantum number Subshell quantum number l

Name Symbol Value Orbital type

Principal

Subshell

Magnetic

Spin

n

l

m

s

s

p

d

f

The principal quantum number n describes the main energy level, in which the electron is present

in its ground state. Values of n range from 1 to 7 in the unexcited states of the known elements. As n

increases, electrons are generally farther from the nucleus and have higher potential energy. The elec-

trons with the same principal quantum number are contained in the same electron layer (sphere, shell).

The subshell quantum number l designates the different energy subshells within the main level. It