The new element would be placed in the periodic table based on its atomic number and electron configuration. 
If it has similar properties to an existing group, it would be placed in that group, and its position would reflected its periodic trends such as atomic size, electronegativity, and ionization energy.

The first element of periodic table is hydrogen (H). 
It will likely lose its single electron to form a cation (H⁺), but it can also gain an electron to form a hydride ion (H⁻), depending on the chemical environment.

Aluminum (Al) is expected to lose electrons more easily than boron (B) because aluminum has a larger atomic radius, which makes it easier for the outermost electron to be removed due to weaker attraction from the nucleus.

Atomic radius:

Atomic radius is defined as half the distance between the nuclei of the two identical bonded atoms. It is expressed in pm (1pm = 10⁻¹²m).

Example:

The distance between the nuclei of two bonded carbon atoms is 154 pm. Half of this distance i.e. 77pm is therefore the radius of carbon atom. This is also called covalent radius of carbon atom.

Oxygen cannot accept these electrons to form O³⁻ because its electron configuration would become unstable. 
Oxygen already has a relatively high electronegativity and only needs two electrons to complete its octet. Adding three electrons would result in significant electron-electron repulsion, making the ion highly unstable. 
In contrast, nitrogen can form N³⁻ more easily due to its smaller atomic size and lower electronegativity.

a. Element Symbol or Name:

If we know the symbol or name of the element, we can look it up in the periodic table.

b. Physical or Chemical properties:

If we know specific physical or chemical properties of the element, such as its group (family) or period, electronegativity, or reactivity, we can use these clues to locate it in the periodic table.

c. Electron configuration:

If we know the electron configuration of the element we can use this information to determine its position in the periodic table.

d. Summary:

In summary, while atomic mass can provide some clues, it's not sufficient on its own to locate an element in the periodic table. Combining atomic mass with other information, such as element symbol, physical or chemical properties, or electron configuration, can help you identify the element.

1. Blocks:

Elements present in the periodic table are also classified into blocks. It depends upon the type of the subshell which is being filled: s, p, d and f.

a. s-block elements:

Elements of group 1 and 2 are called s-block elements, because in them s-subshell of the outermost shell is being filled.

b. p-block elements:

Elements present in groups 13 to 18 are called p-block elements because p-subshell is filled in these elements.

c. d-block elements:

The d-block elements lie between s and p blocks.

f-block elements:

f-block elements in the form of two rows lie at the bottom of the periodic table.

Variation of atomic radius in Periods:

When we move from left to right in a period, the size of atoms decreases generally. It is because as we go from lithium to neon (Ne) in the second period, we are adding elements to the outermost shell. The charge on the nucleus also increases from +3 to +10. This tends to pull the electrons closer to the closer nucleus.

Variation of ionization energy in periods:

Ionization energy value is related to the atomic size. The smaller the radius of an atom, the stronger the attraction between the nucleus and the outer electrons and higher the value for ionization energy. The ionization energy values thus increase from left to right in a period.

1. Electronegativity:

Measures an element's ability to attract electrons. Elements with high EN (e.g., F, O, N) tend to form ionic or covalent bonds with elements of low EN.

2. Ionization energy (IE):

The energy required to remove an electron from an atom. Elements with low IE (e.g., alkali metals) tend to lose electrons and form ionic bonds, while those with high IE (e.g., noble gases) tend to share electrons and form covalent bonds.

3. Electron Affinity (EA):

The energy released when an electron is added to an atom. Elements with high EA (e.g., halogens) tend to gain electrons and form ionic bonds.

4. Atomic Radius:

The size of an atom. 
Elements with large atomic radii (e.g., alkali metals) tend to lose electrons and form ionic bonds, while those with small radii (e.g., noble gases) tend to share electrons and form covalent bonds.

5. Valence Electron configuration:

The arrangement of electrons in an atom's outermost energy level. 
Elements with a full outer energy level (e.g., noble gases) tend to be unreactive, while those with partially filled outer energy levels (e.g., transition metals) tend to form covalent or ionic bonds.

6. Metallic Character:

The tendency of an element to exhibit metallic properties (e.g., conductivity, malleability). 
Metals tend to form ionic bonds, while nonmetals tend to form covalent bonds.

Names of Non-Metals:

i- Sulphur
ii- Phosphorus
iii- Carbon
iv- Iodine

Periods and electron shells:

In the periodic table, each period corresponds to a new electron shell being filled. 
The first period (n=1) has only one orbital (1s), which can hold up to 2 electrons. 
The second and third periods (n=2 and n=3) have four orbitals (2s, 2p, 3s and 3p), which can hold up to 8 electrons each.

Electron configuration and period lengths:

The length of the period is determined by the number of electrons that can fill the available orbitals in that shell.For the second and third periods, the electron configuration is similar:

• Second period (n=2):  2s² 2p⁶ (8 electrons)
• Third period (n=3):  3s² 3p⁶ (8 electrons)

Since both periods have the same number of electrons filling the same type of orbitals, they have the same length, which is 8 elements.

Other Periods:

In contrast, the other periods have different numbers of elements because the electron configuration changes:

• First period (n=1):  1s² (2 electrons)
• Fourth period (n=4):  4s² 3d¹⁰ 4p⁶ (18 electrons)
• Fifth period (n=5):  5s²  4d¹⁰ 5p⁶ (18 electrons)

These changes in electron configuration result in different period lengths. 
This explanation helps to clarify why the second and third periods have the same number of elements!

The periodic table organizes elements in a systematic way, based on their atomic number and properties. This arrangement helps predict the behaviour of elements, their compounds, and reactions.  It also allows scientists to easily identify trends, such as atomic size, electronegativity, and ionization energy, which aids in understanding chemical properties and predicting the behaviour of unknown elements.

Lithium and beryllium behave differently due to their smaller atomic sizes and higher ionization energies compared to other elements in their respective groups. Lithium's small size causes it to have a stronger attraction for its electrons, making it less reactive than other alkali metals. 
Beryllium with its high ionization energy, forms covalent bonds rather than ionic bonds, unlike the typical alkaline earth metals that form ionic bonds.

Mendeleev's periodic table was based on atomic mass and periodic properties and he left gaps for undiscovered elements. He arranged elements with similar properties in columns. However, it had some inconsistencies when elements were ordered by solving these inconsistencies. The modern table also includes elements like the noble gases and accounts for isotopes, providing a more accurate and comprehensive classification of elements.