Physical Properties of p-Block Elements – Physical Properties of Group 13, 14, 15, 16, 17 and 18 Elements, Practice Problems and FAQ

Jack went shoe shopping. He had two pairs of shoes left to choose from, but he wasn't sure which pair.

What steps should Jack take to choose his footwear?

Typically, before making a purchase, we examine the product's quality, durability, and features. We can determine the aforementioned qualities as well as other ones by looking at the product.

Jack followed suit. He carefully weighed the two pairs of shoes and chose the best option for him.

Just like how Jack zeroed on the shoe of his preference, we can perform certain tests to differentiate between elements or groups in our chemistry lab. These properties that we check with performing any chemical reactions are called the physical properties. These include colour, malleability, ductility, solubility, melting point, boiling point, hardness, etc,.

On this concept page, we will get to know about the physical properties of the p-block elements.

TABLE OF CONTENTS

• p- Block Elements
• p- Block Elements – Physical Properties
• Practice Problems
• Frequently Asked Questions – FAQ

p- Block Elements

• The elements in which the last electron moves into the outermost p orbital are called p-block elements.

• Consequently, the periodic table has six groups of p-block elements. Groups with numbers 13 to 18 belong to p- block elements. The groups are led by boron, carbon, nitrogen, oxygen, fluorine and helium.

p- Block Elements – Physical Properties

Group 13: The elements in Group 13 of the periodic table are boron (B), aluminium (Al), gallium (Ga), indium (In), and thorium (Tl). The following are a few important physical properties of group 13 elements.

1. Electronic Configuration: The general valence shell electronic configuration of the group 13 elements is ns2np1, where ‘n’ represents the valence shell number. For example, Boron has [He] 2s2 2p1 as its electronic configuration.

2. Atomic Radius: From boron to thallium, the atomic radii of the group 13 elements all increase, with the exception of gallium, which has an abnormally small size.

3. Oxidation States: This group of elements exhibits two distinct oxidation states, +1 and +3. Moving down the group, the stability of +1, the oxidation state decreases due to the inert pair effect. In general, the stability of the lower oxidation state rises down the group.

4. Ionisation Enthalpy: The values of the ionisation energy do not fall down the group smoothly. The decrease in ionisation enthalpy decreases from B to Al is the general trend of descending a group associated with increased size. The values for the later elements are impacted by the ineffective d-electron shielding and the subsequent d-block shrinkage. Each element's first three ionisation energies added together are extraordinarily high.

5. Electronegativity: From boron to aluminium, group 13 elements become less electronegative down the group, while from gallium to thallium in this group, their electronegativity actually somewhat rises.

6. Melting and Boiling Points: There is no consistent pattern in the melting points of group 13 elements. From boron to gallium, the melting point first drops, then rises for indium and thallium. The boiling point of the elements in group 13 decreases from boron (B) to thallium (Tl).

7. Density: From boron to thallium, the density of the elements steadily rises because the atomic weight of an element grows more than its atomic volume. The correct order of the density of group-13 can be represented by B<Al<Ga<In<Tl.

Group 14: The elements in group 14 are carbon (C), silicon (Si), germanium (Ge), tin (Sn), and lead (Pb). The following are a few important physical properties of group 14 elements.

1. Electronic Configuration: The general valence shell electronic configuration of group 14 elements is ns2np2. For example, carbon has [He] 2s2 2p2 as its electronic configuration.

2. Atomic Radius: From C to Si, the covalent radius increases significantly; from Si to Pb, only a little rise is seen. This is brought on by the fact that heavier members have fully filled d and f orbitals which exhibit poor screening effects.

3. Oxidation states: The common oxidation states shown by the group 14 elements are +4 and +2. On descending the group, the (+IV) oxidation state's stability decreases while the (+II) state's stability increases due to the inert pair effect.

4. Ionisation Enthalpy: There is a small decrease in I.E from Si to Ge to Sn and a slight increase in I.E from Sn to Pb is due to the poor shielding effect of inner d and f orbitals and also due to an increase in the size of the atom.

5. Electronegativity: These elements are slightly more electronegative than group 13 elements because of their smaller size. From Si to Pb, the electronegativity values are almost the same.

6. Melting and Boiling Points

7. Density: The below-mentioned table describes the density values for group 14 elements.

Group 15: The elements in group 15 are nitrogen (N), phosphorus(P), arsenic (As), antimony (Sb), and bismuth (Bi). The following are a few important physical properties of group 15 elements.

1. Electronic Configuration: These elements have ns2np3 as the valence shell electrical configuration. For example, Nitrogen has [He] 2s2 2p3 as its electronic configuration.

2. Atomic Radius: Covalent radius increases down the group. The covalent radius increases significantly from N to P. However, only a little increase in covalent radius is seen from As to Bi. This is caused by heavier members having fully filled d and/or f orbitals.

3. Oxidation State: These elements often exist in the oxidation states of -3, +3, and +5. As size and metallic character increase along the group, the propensity to display the -3 oxidation state diminishes.

4. Ionisation Enthalpy: The ionisation enthalpy of group 15 elements in their relevant periods is significantly higher than that of group 14 elements due to their extra stable half-filled p orbitals electrical configuration and smaller size.

5. Electronegativity:

6. Melting and Boiling Points: The melting point rises till arsenic and then lowers until bismuth, the boiling points in the group generally increase from top to bottom.

7. Density: The below-mentioned table describes the density values for group 15 elements.

Group 16: The elements in group 16 are oxygen (O), sulphur (S), selenium (Se), tellurium (Te), and polonium (Po). The following are a few important physical properties of group 16 elements.

1. Electronic Configuration: These elements have a general valence shell configuration of ns2np4. For example, Oxygen has [He] 2s2 2p4 as its electronic configuration.
2. Atomic Radius: Atomic and ionic radii in the group increases from top to bottom as a result of a rise in the number of shells.

3. Oxidation States: There are many oxidation states present in the Group 16 elements but -2 is the most frequent one. Down the group, the -2 oxidation state becomes less stable. Rarely does polonium exhibit a -2 oxidation state.
4. Ionisation Enthalpy: Down the group, ionisation enthalpy falls. It is because of the growth in size. In contrast to Group 15, this group's elements have lower ionisation enthalpy values in their respective periods. This is because Group 15 elements have very stable electronic structures i.e. half-filled p orbitals.
5. Electron Gain Enthalpy: The oxygen atom is more compact than sulphur, which results in a lower negative electron gain enthalpy. However, the value decreases again from sulphur to polonium.
6. Electronegativity:

7. Melting and Boiling Points: With increasing atomic numbers down the group, the melting and boiling points rise. Due to their different atomicities, there is a large difference in melting and boiling points of oxygen and sulphur.
8. Density: The below-mentioned table describes the density values for group 16 elements.

Group 17: The elements in group 17 are fluorine (F), chlorine (Cl), bromine (Br), iodine (I) and astatine (At). The following are a few important physical properties of group 17 elements.

1. Electronic Configuration: The general valence shell electronic configuration of group 17 elements is ns2np5. For example, fluorine has [He] 2s2 2p5 as its electronic configuration.
2. Atomic Radius: In their respective periods, halogens have the shortest atomic radii because of the highest nuclear charge. Fluorine's atomic radius is really little, much like the other components of the second period. Due to an increase in the nuclear charge, atomic radii from fluorine to iodine increase.

3. Oxidation States: Halogens exhibit –1 as their common oxidation state. However, the oxidation states + 1, + 3, + 5 and + 7 can also be found in compounds of chlorine, bromine, and iodine.
4. Ionisation Enthalpy: They do not typically lose electrons. They, therefore, possess a very high ionisation enthalpy. Ionisation enthalpy falls across the group as the atomic size grows.
5. Electron Gain Enthalpy: Electron gain enthalpy decreases down the group due to the increase in atomic radii, but chlorine has a higher electron gain enthalpy than fluorine. This is because of the smaller size of fluorine and its subsequently interelectronic repulsion. Halogens have the maximum negative electron gain enthalpy in their corresponding periods. This is because of the fact that the atoms of these elements have only one electron less to attain the nearest noble gas configuration.
6. Electronegativity:

7. Melting and Boiling Points: With increasing atomic numbers down the group, the melting and boiling points rise.
8. Density: The below-mentioned table describes the density values for group 17 elements.

Group 18: The elements in group 18 are helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe) and radon (Rn). The following are a few important physical properties of group 18 elements.

1. Electronic Configuration: These elements have ns2np6 as the valence shell electronic configuration. For example, neon has [He] 2s2 2p6 as its electronic configuration.

2. Atomic Radius: With an increase in atomic number down the group, the atomic radius increases with an increase in the number of shells.

3. Oxidation States: Noble gases have no vacancies in their outermost shells because their outermost shells are entirely filled with electrons. As a result, these gases typically exhibit zero oxidation state, which is why they are frequently referred to as zero group elements.

4. Ionisation Enthalpy: Since their electronic configuration is stable, these gases display extremely high enthalpy of ionisation.

5. Electron Gain Enthalpy: Noble gas electronic configurations are stable, so they don't have any propensity to receive the electron and, as a result, possess high positive values enthalpy of electron gain.

6. Melting and Boiling Points: Since weak dispersion forces are the sole interatomic forces that occur in these elements, their melting and boiling temperatures are extremely low. Of all known substances, helium has the lowest boiling point (4.2 K).

7. Density: The below-mentioned table describes the density values for group 18 elements.

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Practice Problems

1. _____________ has the highest and _____________ has the lowest melting point among group 13 elements respectively.

a. B , Al
b. Al, B
c. Tl, Al
d. B , Gl

Answer: D

Solution: There is no consistent pattern in the melting points of group 13 elements. From boron to gallium, the melting point first drops, then rises for indium and thallium. Hence, born (2453 K) has the highest melting point whereas gallium (303 K) has the lowest melting point among group 13 elements.

So, option D is the correct answer.

2. Halogens like fluorine and chlorine are _______ at room temperature.

a. Gaseous
b. Solid
c. Liquid
d. Both solid and liquid

Answer: A

Solution: Halogens like fluorine and chlorine are gases at room temperature, whereas bromine is liquid at room temperature and iodine is solid at room temperature.

So, option A is the correct answer.

3. Which of the following electronegativity orders is correct?

a. B>F>In> C
b. C> He> F>N
c. F>O>Br>Te
d. F>N>O>Cl

Answer: C

Solution: The following elements have the following electronegativity values.

• Boron - 2.04
• Fluorine - 3.98
• Indium - 1.78
• Carbon - 2.55
• Nitrogen - 3.04
• Oxygen - 3.44
• Tellurium - 2.1
• Bromine - 2.96
• Chlorine - 3.16

The correct order of electronegativity among the given p-block elements is F>O>Br>Te. The electronegative values for these elements are F (3.98)>O (3.44)>Br (2.96)>Te (2.1).

So, option C is the correct answer.

4. ______________ is the most common oxidation state for chalcogens.

a. 2
b. 1
c. -2
d. -1

Answer: C

Solution: There are many oxidation states present in the Group 16 elements or chalcogens but -2 is the most frequent one.

So, option C is the correct answer.

Frequently Asked Questions – FAQ

1. How are group 16 elements classified as metals, non-metals and metalloids?
Answer: Among the group 16 elements, oxygen and sulphur are non-metals whereas selenium and tellurium are metalloids. Polonium is a metal which is radioactive in nature.

2. Are all halogen compounds coloured?
Answer: Almost all halogen compounds have colour. This is caused by radiation absorption in the visual region that causes the higher energy excitation of outer electrons level. Because they absorb various amounts of radiation, they exhibit various colours. As an illustration, F2 is yellow in colour, Cl2 is greenish yellow, Br2 is red, and I2 is violet.

3. What unique properties do noble gases possess?
Answer: Noble gases are low chemically reactive, odourless, colourless, nonflammable, and monoatomic gases.

4. What are the physical characteristics of nitrogen?
Answer: Nitrogen is a typical diatomic non-metal gas. The most common physical properties are that they are typically colourless, odourless, and tasteless.