Periodic Table & Periodicity — Short Questions

The elements in the periodic table are arranged in order of their increasing atomic number because the atomic number represents the number of protons in an atom. This arrangement allows elements with similar chemical properties to be grouped together in the same column, known as a group or family. The atomic number increases across each period from left to right, reflecting the increase in the number of protons.
This arrangement helps in understanding the trends in physical and chemical properties of elements and provides a systematic way to organize and study the elements.

The word "periodic" in the periodic table refers to the repeating pattern of properties that elements exhibit when they are arranged in order of increasing atomic number. This repeating pattern occurs in groups or periods, that is why it is called the periodic table.

As you go down, the number of electronic shells increases. Addition of d-sub shell elements or transition metals, presence of f-sub-shell lanthanide and actinide cause increase in period size.

Elements in the same group have similar chemical properties because they have the same number of electrons in their outer shell. The number of electrons in the valence shell determines how an element will interact chemically with other elements.

No, calcium less reactive than sodium because the atomic size of sodium is larger than calcium. Moreover, sodium has one electron in its outer shell while calcium has two electrons in its outer shell.

In the 3rd period, the element with the maximum atomic radius is sodium (186 pm) and the element with the minimum atomic radius is chlorine (79 pm).

The most electronegative elements are found in the sixth (16th) and seventh (17th) groups, specifically oxygen and fluorine respectively because they have high effective nuclear charge due to smallest atomic sizes among rest of the elements in the respective groups.

Once the first electron is removed, the magnesium atom becomes a positively charged ion (Mg+). This ion has a stronger positive charge which pulls the remaining electrons closer to the nucleus. As a result, the second electron is held more tightly, requiring more energy to remove it. So, the increase in effective nuclear charge after the first ionization energy makes the second ionization energy significantly higher.
Mg(g)→ Mg+(aq) + e- ΔH= + 737 KJ/mol
Mg+(aq)→ Mg2+(aq) + e- ΔH= + 1450 KJ/mol

No, two metals cannot form ionic bond and two non-metals also cannot form ionic bond. Two metals form metallic bond.

The element with the least value of ionization energy is caesium (377 kJ/mol), and the element with the highest value of electronegativity is fluorine (4.0).

Period Number The number of the period corresponds to the energy level of the outermost electrons. In this case, since the p orbital is being filled, the element is in the 2nd energy level. Therefore, the period number is 2.
Group Number: The group number is determined by the number of valence electrons in the outermost shell. In this case, there are total 5 valence electrons (2 from the s orbital and 3 from the p orbital). Elements with 5 valence electrons are typically found in Group 15 (also known as Group VA) of the periodic table.
Block: The element with the electronic configuration ns2 m3 will be placed in the p-block of the periodic table. Elements in the p-block have their outermost electrons in the p orbital. Therefore, the element with the electron configuration ns2 np3 is in Period-2, Group-15 (Group VA) and belongs to the p-block of the periodic table.

Having eight electrons in the outermost shell indicates that the element is in Group 18 of the periodic table, also known as the Noble Gases group. Elements in Group 18 have a completely filled outermost shell, making them stable and unreactive. As for the physical state in which this element exists, elements in Group 18 are known as Noble Gases, and they exist as gases at room temperature and pressure.

The element belonging to the 6th group will be in Period 2 of the periodic table.

i. Is it a metal or a non-metal?
Ans. Barium is a metal.
ii. Will it be electropositive or electronegative?
Ans. Barium is electropositive in nature.
iii. What is the nature of its oxide?
Ans. Barium oxide is basic oxide.
iv. In which physical state you expect this element to exist?
Ans. Barium is expected to exist in a solid state at room temperature and pressure.

You would expect to find an element with the largest atomic radius in Group 1 and Period 7 of the periodic table.

The most electropositive element is found in Group 1 and the most electronegative element is found in Group 17 of the periodic table.

The element with the lowest ionization energy is potassium (K). The element with the highest electron affinity is fluorine (F).
Potassium (K) is in Group 1 and Period 4.
Fluorine (F) is in Group 17 and Period 2.

i. Sodium oxide (Na2O) is formed when sodium reacts with oxygen.
4 Na(s) O2(g) → 2 Na2O
ii. Sulphur dioxide (SO2) is formed when sulphur burns in the presence of oxygen.
S8(s) + O2(g) → SO2(g)

So, sodium produces a basic oxide, while sulphur produces an acidic oxide.

The periodic table is arranged according to ascending order of their atomic numbers. From left to right in a horizontal row, properties of elements are found repeating after regular intervals. This results in the form of a table in which elements of similar properties are placed in the same vertical columns.

Groups Vertical columns in the periodic table are called groups. They are studied from top to bottom. There are 18 groups in the periodic table.
Periods: Horizontal rows in the periodic table are called periods. They are studied from left to right. There are 7 periods in the periodic table.

In the modern periodic table, the electronic configuration of the elements continues changing when we move from left to right in period. Due to this, the elements in a period show a gradual change in the properties while the elements in a group show similar properties due to having same electronic configuration of last shell electrons.

The periodic table shows repetition of the properties of elements after regular intervals. The study of properties in a sequence is called periodicity in the properties of elements.

In 6th and 7th periods, two series of fourteen elements each have been accommodated. Because of the space problem these two series were placed at the bottom of the periodic table to keep it in a manageable and presentable form. The first series starts after lanthanum (La=57) and it is called lanthanides The second series starts after actinium (Ac=89) and it is called actinides.

The elements which follow the Lanthanum with atomic number 58 to 71 are called lanthanides. They belong to 6th period. The name of the series is based upon element named lanthanum.

The elements which follow the actinium having atomic number 90 to 103 are called actinides. They belong to 7th period. The name of the series is based upon element named actinium.

Elements which belong to 's' and 'p' blocks are called normal or representative elements. In these elements all the inner shells are completely filled with electrons. Only their outermost shells are incomplete.

Those elements which belong to 'd'and 'f' blocks are called transition elements. Their d – subshells or f – subshells are in the process of completion with electrons.

Group 1 and Group 2 elements have the electronic configuration that ends at s-subshell. Therefore, these elements are called s-block elements.

Elements in groups 13 – 18 (except He) are known as p-block elements because their electronic configuration ends at p-subshell.

Lanthanides and actinides are known as f-block elements since their valence electrons lie in f sub-shell.

All elements in group 17 have ns2 np5 configuration in their outer shells. They show a strong tendency to gain one electron to become an anion. They are called electronegative elements. These elements are very reactive and form salt with metals. Hence they are called halogens (salt forming).
Example: F, Cl, Br, I

Elements present in group 16 have s2 p6configuration in their outermost shell, so they have a tendency to accept two electrons to form a dinegative ion. Oxygen behaves as a strong electronegative element.
O2(g) + 2e- → 2O2-(g)

Mendeleev arranged only 63 elements in his periodic table because only these elements were discovered at that time. He left many space vacant for those elements which were yet to be discovered.

EG. Mazurs collected 700 different published versions of the periodic table. Many forms retain the regular rectangular structure. Some forms had spirals circles and triangular shapes.

The chemical properties of elements depend largely upon the number of electrons present in their outermost shells. Since in a group of the periodic table all the elements have the same number of electrons in the outermost shell they are expected to show similar chemical properties.

Group Number It tells about the number of electrons present in the outermost shell of an atom.
Period Number: It tells about the number of electronic shells present in an atom.

Although you might expect atoms in become larger with the increase in their atomic numbers, this does not always occur because the size of atoms is determined by the diameter of its electron shells.

It increases from left to right in a period and decreases from top to bottom in a group. Thus the most electronegative atoms are found at the top right-hand corner of the periodic table. The most electronegative atoms are F, O, N and Cl.

Electronegativity is one of the most well-known property for explaining why chemical reactions take place.

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

The atomic radius increases from top to bottom in a group.
Reason: Because the increase of more electronic shell in atoms of successive period decreases the effective nuclear charge.

Metallic character of a metal generally determines its level of reactivity.

In a period, the atomic radii gradually decreases form left to right. Because of increase of atomic number, the effective nuclear charge increases gradually left to right due to addition of one more proton in the nucleus of next every atom. This increased nuclear force pulls more strongly the outermost shell towards the nucleus.

The ionization energy is the amount of energy required to remove the most loosely bounded electron from the valence shell of an isolated gaseous atom. Its unit is k/mol^-1. Ionization energy of sodium is 496 kJ/mol^-1
e.g.: Na→ Na^+ + 1e^-
ΔH= +496 KJ/mol

When an electron is removed from an atom, mono-positive ion is formed. Mono-positive ion has one extra proton than the number of electrons. Its nuclear charge attracts the remaining electrons more strongly. Hence more energy is needed to remove second electron. That is why 2nd I.E. has greater value.

Along Group Ionization energy decreases while moving from top to bottom in a group.
Reason:
Because in group more and more shells lie between the valence shell and nucleus reduce the electrostatic force of attraction on valence electrons so they are taken away easily.
Along Period:
Ionization energy increases while moving from left to right in a period.
Reason: Because in period the size of atoms reduce and valence shell electrons are held strongly by the electrostatic force of attraction of nucleus.

The amount of energy released when an electron is added ion in the outermost shell of an isolated gaseous atom. Its unit is kj/mol.
Example: Electron affinity of fluorine is – 328kj/mol.
F + e^- → F^- ΔH= -328 kj/mol

Electron affinity decreases in a group because of increasing atomic size and shielding effect in a group.

Electron affinity increases in a period because of decrease in atomic size from left to right in a period.

First electron affinity can be negative, positive or zero while second electron affinity is always positive.

The ability of an atom to attract the shared pair of electrons towards itself in a molecule is called electronegativity. Electronegativity of Nitrogen is 3.0 while that of Oxygen is 3.4.

All elements of group 1 have one electron in their outermost shell, so they show a strong tendency to lose their valence electron forming cations. They are thus known as electropositive metals. These metals react vigorously with water producing hydrogen and giving alkali in the solution.
Na(s) + H2O(g) → NaOH(aq) + H2(g)
Alkali metals also react with halogens giving halides.
Na(s) + Cl2(g) → 2NaCl(s)
The reactivity of alkali metals gradually increases down the group.

The second group elements also show a tendency to lose both of their outermost electrons forming dispositive ions called alkaline earth metals.
Mg(s) → Mg²+(aq) + 2e^-

The reactivity of metals increase down the group because of increasing atomic size and decreasing ionization energy. As the atomic size increases down the group due to the addition of shells, it becomes easy to remove electron from the outermost shell of a metal which causes increased reactivity of metals.

Electropositivity is the property of an element to readily lose its valence electrons and gain a positive charge. Metals are highly electropositive elements.
e.g. Sodium atom can lose 1 electron to from a positive ion. Na(s) → Na(g) + 1e^-

In a period Atomic radius of elements decreases in a period so removal of electron becomes more difficult which results in decrease of metallic character.
In a group: Atomic radius and shielding increases down the group so removal of electron becomes easier which results in increase of metallic character.

Densities of alkali metals increase down the group in the periodic table due to increase in atomic masses.

Aluminum and silicon are less reactive than sodium and magnesium. This is because comparatively more number of valence electrons making it difficult to lose electrons.

It is mass of a substance in a unit volume. Density of elements generally increases from top to bottom in a group but varies less significantly from left to right in a period. It is expressed in g/cm³.

Densities of First Group Elements

First group Elements Li 0.53, Na 0.97, K 0.89, Rb 1.63, Cs 1.879 density g/cm³

It would be placed in the periodic table according to its atomic number, which is determined by the number of protons in its nucleus.
If the new element has similar properties to existing elements, it would go in the same group.
For example, if it behaves like a metal, it might be placed in one of the groups on the left side. If it has properties similar to non-metals, it could be placed on the right side.
The specific location would also depend on its electron configuration and how it interacts with other elements.

The first element of the periodic table is hydrogen. It can either lose an electron to achieve a stable configuration like a metal or gain an electron to achieve a stable configuration like a non-metal.

Aluminum is expected to lose electrons more easily than boron since it has a larger atomic radius and high shielding affect.

The atomic radius is typically measured as half of the distance between the nuclei of two bonded atoms of the same element. This distance can be determined experimentally using techniques like X-ray crystallography or calculated based on the element's properties.

Oxygen atom cannot accept three electrons to form O³- like nitrogen forming N³-because the electronic configuration of oxygen (O) is 1s² 2s² 2p⁴. Oxygen typically gains 2 electrons to achieve a stable octet configuration, forming O²-. Nitrogen (N) has an electronic configuration of 1s² 2s² 2p³and can accept 3 electrons to complete its octet and achieve a stable configuration, forming N³-.