9th Class Chemistry Unit 8 Periodic Table & Periodicity Long Questions and Answers

1.Which information is needed to locate the elements in the periodic table if you do not know its atomic number? Is atomic mass helpful for this purpose?

To locate elements in the periodic table without knowing the atomic number, you need to know the element's properties like its atomic mass and chemical behavior.

Atomic mass
When elements are arranged in increasing order of their atomic mass, they show periodic repetition of properties e.g. Li, Na and K have increasing atomic mass and all in Group - 1.

Chemical behavior Co-relates with groups and periodic trends confirm an elements position.

Example Transition metals in d-block show variable states.

The atomic mass can help in some cases, but it may not be as precise as the atomic number for identifying elements. The atomic number is the most reliable way to locate elements in the periodic table as elements are arranged in increasing atomic number order.

2.How many blocks of elements are present in the periodic table? Are these blocks helpful in studying the properties of elements?

There are four blocks of elements in the periodic table: s-block, p-block, d-block and f-block. These blocks categorize elements based on their electronic configurations.

i. s-block: The elements of Group 1 (alkali metals) and Group 2 (alkaline earth metals) are called s-block elements because in them s-sub shell of outermost shell is being filled.

ii. p-block: The elements of Group 13 to Group 18 (except He) have outermost electrons in the p-subshell.

iii. d-block: The elements of Group 3 to Group 12 have the outermost electrons in the d-subshell. The d block elements lie between the s and p blocks.

iv. f-block: The elements of lanthanides and actinides have outermost electrons in the f-subshell. f-block lies separate at the bottom of the periodic table.

Advantage of Blocks -

These blocks are very helpful in studying the properties of elements. They highlight common characteristics within each block, making it easier to understand trends in reactivity, electronegativity, and other chemical properties. By organizing elements this way, scientists can predict how different elements will behave in chemical reactions and interactions.

3.Explain the variation in the following properties in the periods giving reasons. (a)Atomic radius (b)Ionization energy

(a) Variations of Atomic Radius in Period

When we move from left to right in a period, the size of atoms decreases generally.

Reason -
It is because as we move from lithium (Li) to neon (Ne) in the second period, we are adding electrons to the outermost shell. The charge on the nucleus also increases from +3 to +10. This tends to pull the electrons closer to the nucleus and hence the sizes of atoms decrease from lithium to neon as shown in Table data.

Atomic Radii of Second Period Elements
2nd period elements | Li | Be | B | C | N | O | F | Ne
Atomic radii (pm) | 152 | 113 | 88 | 77 | 75 | 73 | 71 | 69

(b)Variation of Ionization Energy in Period

As you move from left to right across a period, ionization energy generally increases.

Reasons • As you move to the right, the number of protons in the nucleus increases. This increase in nuclear charge, leading to an increase in ionization energy.
• The atomic radius decreases as you move from left to right across a period. A smaller atomic radius means that the outermost electrons are closer to the nucleus, experiencing a stronger attraction. This stronger attraction makes it harder to remove an electron, thus increasing the ionization energy.
• When moving from left to right, new electrons are added to the same energy level without the addition of new shells. This means that the shielding effect does not increase significantly. As a result, the outermost electrons feel a stronger attraction from the nucleus, which contributes to the increase in ionization energy.

2nd period elements | Li | Be | B | C | N | O | F | Ne
Ionization energy (kJ/mol) | 520 | 899 | 801 | 1086 | 1402 | 1314 | 1681 | 2081

4.Which physical properties of elements may lead us to know what type of bond it will form?

The physical properties of elements that can help us to understand the type of bond they will form include:

Electronegativity This is a measure of how strongly an atom attracts electrons in a chemical bond. Two elements with a large difference in electronegativity values, are likely to form an ionic bond. For example, sodium (Na) has low electronegativity, while chlorine (Cl) has high electronegativity, leading to the formation of NaCl.

Ionization Energy Elements with low ionization energy tend to lose electrons easily and can form cations, which is common in ionic bonds. On the other hand, elements with high ionization energy are less likely to lose electrons and may form covalent bonds by sharing electrons.

Electron Affinity Elements with high electron affinity tend to gain electrons easily, which is a characteristic of non-metals that often form anions in ionic bonds.

Metallic Character Elements that are metals typically have lower electronegativity and low ionization energy, making them more likely to lose electrons and form metallic bonds with other metals.

By examining these properties, you can predict whether elements will form ionic, covalent, or metallic bonds based on their tendency to lose, gain, or share electrons.

5.Write down the names of four non-metals which exist in a solid state at normal temperature.

Four non-metals that exist in a solid state at normal temperature are:
(i) Phosphorus (P)
(ii) Iodine (I)
(iii) Sulphur (S)
(iv) Carbon (C)

6.Why do second and third periods have equal number of elements while all other periods contain different number of elements?

Each period also represents the completion of a shell. Since the number of electrons to be accommodated in a particular shell is fixed, the number of elements in a period is also fixed. 2nd and 3rd periods have an equal number of elements, which is eight, is due to the way electrons fill the available energy levels or shells around the nucleus of an atom.

2nd Period

In the 2nd period, elements fill the 2s and 2p orbitals. The 2s orbital can hold 2 electrons and the 2p orbitals can hold 6 electrons, making a total of 8 electrons for the 2nd period.

3rd Period

Similarly, in the 3rd period, elements fill the 3s and 3p orbitals. Again, the 3s can hold 2 electrons and the 3p can hold 6 electrons, total making of 8 electrons for the 3rd period as well.

For periods beyond the 3rd, the number of elements varies because of the presence of d and f orbitals, which can hold more electrons. As a result, those periods can accommodate more elements, leading to differences in the number of elements in each period.

7.Arrangement of the elements in the form of a periodic table is a remarkable achievement of chemists. Comment on this statement citing the benefit of this table.

The periodic table is indeed a remarkable achievement of chemists. This table organizes all the known elements based on their atomic number, electronic configuration, and chemical properties. The benefits of this table are numerous:

Element Organization The periodic table organizes elements in a systematic way, making it easier to understand and predict their properties based on their position.

Identification of Trends The table allows chemists to identify trends in properties such as atomic size, electronegativity and reactivity as you move across a period or down a group.

Predictive Capabilities By using the periodic table, scientists can predict the properties of undiscovered elements and how they might react with other elements.

Chemical Bonding It helps in understanding how elements bond with each other to form compounds based on their electronic configurations.

8.Both lithium and beryllium show behavior different from rest of the alkali and alkaline earth metals respectively. Can you think of the possible reasons for this difference?

Lithium and beryllium exhibit different behaviors compared to the rest of the alkali and alkaline earth metals respectively due to their unique characteristics.

Lithium (Li)

i. Lithium is the lightest metal in the alkali metal group.

ii. Lithium has a smaller atomic size and higher ionization energy compared to other alkali metals.

iii. Due to its small size and high ionization energy, lithium tends to form covalent compounds rather than purely ionic compounds like other alkali metals.

Beryllium (Be)

i. Beryllium is the lightest alkaline earth metal.

ii. Beryllium has a small atomic size and high ionization energy, similar to lithium.

iii. The small size and high ionization energy of beryllium make it less reactive and less likely to form ionic compounds compared to other alkaline earth metals.

iv. Beryllium forms covalent compounds and exhibits amphoteric behavior, meaning it can act as both an acid and a base.

9.Modern periodic table is the amended form of the earlier table developed by Mendeleev. Elaborate how these two tables are different from each other.

The earlier table developed by Mendeleev and the modern periodic table have some differences:

Mendeleev's Table

i. Mendeleev's table was based on atomic mass and chemical properties of elements.

ii. Mendeleev arranged only 63 known elements. Gaps were left in Mendeleev's table for undiscovered elements and he could predict the properties of these missing elements.

iii. Mendeleev's table had less organized groups and periods than in the modern periodic table.

Modern Periodic Table

i. The modern periodic table is based on atomic number rather than atomic mass.

ii. Elements are arranged in increasing atomic number and grouped based on similar chemical properties.

iii. The modern periodic table has periods (rows) and groups (columns) that help in organizing elements based on their properties.

iv. The modern periodic table includes noble gases as a separate group, which was not part of Mendeleev's original table.

v. The modern periodic table includes transition metals in a distinct block, which was not defined in Mendeleev's table.

10.Discuss in detail the periods in periodic table.

Definition

The horizontal rows in the modern periodic table are called periods. There are seven periods in total.

Explanation

Each period except the first starts with an alkali metal and ends at a noble gas. Each period also represent the completion of a shell. Since the number of electrons to be accommodated in a particular shell is fixed the number of elements in a period is also fixed.

Arrangement of Lanthanides and Actinides:

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 from.

Lanthanides

The first series starts after lanthanum (La = 57) and is called lanthanides.

Actinides

The second series starts after actinium (Ac = 89) and it is called actinides.

Together the elements present in these two series are also called rare earths or f-block elements.

11.Discuss the groups in modern periodic table.

Definition

The vertical columns present in the periodic table are called groups. There are in total eighteen groups.

Explanation

Elements present in the group resemble one another in their chemical properties since they contain the same number of electrons in their outermost shell. Elements present in a group are also called a family and each group has also have been given a family name. The distribution of electrons in the outermost shells (electronic configuration) and other information about the groups are given below in the following Table

Electronic Configuration of Elements in the Outermost Shell

Normal Element The groups 1 to 2 and 13 to 17 contain the normal elements. In the normal elements all the inner shells are completely filled while the outermost shell is incomplete.

Transition Element The groups 3 to 12 are called transition elements and in these elements the inner sub-shells are in the process of completion.

12.Why the elements placed in same group show similar chemical properties?

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. For Example:

Alkali Metal All elements of group 1 have one electron in their outermost shells so they show a strong tendency to use their 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₍ₑ₎ ⟶ NaOH₍ₐₑ₎ + H₂ ₍ₑ₎
Alkali

Oxides Oxides of these metals are also strongly basic in nature. They are readily soluble in water giving alkalies
K₂O ₍ₛ₎ + H₂O₍ₑ₎ ⟶ 2KOH ₍ₐₑ₎

Halides Alkali metals also react with halogens giving halides
Na₍ₛ₎ + Cl₂₍ₑ₎ ⟶ 2NaCl₍ₛ₎

Reactivity The reactivity of alkali metals gradually increases down the group.

Alkaline Earth Metal The second group elements also show a tendency to lose both of their outermost electrons forming dipositive ions

Mg ₍ₛ₎ + 2eˉ ⟶ Mg²⁺₍ₛ₎

The tendency to lose electrons down the group increases due to gradual increase in their atomic sizes.

Oxides The oxides of these metals are also basic in nature and form alkalis in water .

CaO₍ₛ₎ + H₂O₍ₑ₎ ⟶ Ca(OH)₂₍ₐₑ₎

Halogens All elements in group 17 have ns²np⁵ configuration in their outer shells. They show a strong tendency to gain one electron to become an anion. They are thus called electronegative elements. These elements are very reactive non-metals and form salts with metals. Hence they are called halogens (salt forming). Unlike alkali metals, the reactivity of halogens decreases down the group.

Cl₂₍ₑ₎ + 2eˉ ⟶ 2Cl⁻₍ₐₑ₎
Ca ₍ₛ₎ + Cl₂₍ₑ₎ ⟶ CaCl₂ ₍ₛ₎

Group16 Elements present in group 16 have ns²,np⁴ configuration in their outermost shells so they have a tendency to accept two electrons to form a dinegative ion. Oxygen behaves as a strong electronegative element.

O₂₍ₑ₎ + 2eˉ ⟶ 2O²⁻₍ₑ₎

13.Describe how group number relate with charge of ions?

When we move from left to right in the periodic table the main group elements tend to form cations having a change equal to the group number For example group 1 elements form 1 + ions, group 2 elements forms 2 + ions and group 3 elements forms 3+ ions. The number of charges on the cations also corresponds to the number of electrons present in their outermost shells.

When we move from right to left in the periodic table elements often from atoms with a negative charge equal to the number of group towards the left side of the noble gases.

For example, group 17 elements (which are located one group towards left to the noble gases) form 1-ions. The negative charges presents on these ions correspond to number of electrons which these groups need to complete their octets.

14.Define Atomic Radius. What are variations of atomic radius in the periodic table?

Definition

Atomic radius is defined as half the distance between the nuclei of the two identical bonded atoms.

Units

It is expressed in pm (1pm = 10⁻¹² m) as well as nm (1nm = 10⁻⁹ m)

For example, the distance between the nuclei of two bonded carbon atom 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.

Variation of Atomic Radius in Periods:

When we move from left to right in a period the size of atoms decreases generally

Reasons

It is because as we go from lithium (Li) to neon (Ne) in the second period, we are adding electrons to the outermost shell. The charge on the nucleus also increases from +3 to +10. This tends to pull the electrons closer to the nucleus and hence the sizes of atoms decreases from lithium to neon

Atomic Radii of Second Period Elements
2nd period elements | Li | Be | B | C | N | O | F | Ne
Atomic Radii (pm) | 152 | 113 | 88 | 77 | 75 | 73 | 71 | 69

Variation of Atomic Radius in Group:

The atomic radii of atoms increase from top to bottom in a group.

Reasons

It is because a new shell is being added in the successive period down the group which increases the shielding effect.

15.Define electron affinity. Why it increases in a period and decrease in a group?

Definition

Electron affinity is the amount of energy released when an electron is added up in the outermost shell of an isolated gaseous atom.

For example, 328 kJmol⁻¹ energy is released when an electron enters in the fluorine atom.

F₍ₑ₎ + e⁻ ⟶ F⁻₍ₑ₎ ΔH = -328kJmol⁻¹

Variation of electron affinity in Period:

Electron affinity values are also related to the sizes of the atoms. The smaller the size of an atom, the higher the force of attraction with which the nucleus will attract the entering electron and hence higher is the value of electron affinity

Electron Affinities of Second Period Elements
2nd period elements | Li | Be | B | C | N | O | F
Electron Affinity (kJ mol⁻¹) | -60 | 0 | -29 | -122 | 7 | -141 | -328

Variation of electron affinity in Group:

In a group the electron affinity values decrease from top to bottom because the sizes of atoms increase down the group. In a bigger atom, the nucleus will attract the incoming electron with a weaker force and hence the electron affinity will also be low.

Note While first electron affinities can be negative, positive or zero, second electron affinity are always positive.

16.Explain metallic character and reactivity of elements in periodic table.

Definition

The metallic character is the tendency of an element to lose electrons and form positive ions or cations

Variation in group

Since the ionization energy decreases down the group, the elements have increased ability to lose electrons. For this reason both the metallic character and reactivity increase down the group.

Variation in period

As we know from left to right in a period, the nuclear charge increases due to a gradual increase in the number of protons in the nucleus. Owing to this the valence electrons are pulled strongly by the nucleus making it difficult for the atoms to lose electrons. Hence the metallic character decreases in a period from left to right. The chemical reactivity gradually decreases as we move from left to right in a period.

Example Aluminium and silicon are less reactive than sodium and magnesium. This is because the number of valence electrons increases, making it difficult to lose-electrons. Moving further right in a period towards non-metals, the chemical reactivity gradually increases.

Unit 8: Periodic Table & Periodicity — Long Questions