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Summary
Elements are pure substances, and each one is made of a unique kind of atom. Think of atoms as the tiny building blocks of everything. Scientists wanted to organize these elements to study them better. Early on, some tried grouping elements by threes based on similar properties, or by eights, noticing patterns when arranged by their atomic weights. A scientist named Mendeleev created a more detailed table using atomic masses, which was very helpful and even predicted elements that hadn’t been found yet. However, this table had a few puzzles.
Later, Henry Moseley discovered that arranging elements by their atomic number, which is the number of protons in an atom’s center, worked even better. This led to the modern periodic table we use today. This table has vertical columns called groups and horizontal rows called periods. Elements in the same group behave similarly in chemical reactions because they have the same number of electrons in their outermost layer or shell. Elements in the same period have the same number of electron shells. The table is organized into 18 groups and 7 periods. Some groups have special family names: Group 1 elements are alkali metals, Group 2 are alkaline earth metals, Group 17 are halogens, and Group 18 are noble gases.
The term periodicity describes how the properties of elements repeat in a predictable pattern as you move through the table. This repetition happens because the way electrons are arranged in the outermost shell repeats. The number of shells an atom has determines which period it belongs to. Valency tells us about an element’s combining power with other elements. For elements in groups 1, 2, 13, and 14, valency is often the number of electrons in the outermost shell. For groups 15, 16, and 17, it is often calculated as 8 minus the number of outermost electrons. Generally, elements in the same group have the same valency, while across a period, valency changes.
Several properties of elements show clear trends. Atomic size refers to how big an atom is. Going down a group, atoms get bigger because new electron shells are added, like adding layers to an onion. Going across a period from left to right, atoms usually get smaller. This is because the center of the atom, the nucleus, gets more positive charge and pulls the electron shells closer. When an atom loses electrons, forming a positive ion (cation), it becomes smaller than its original atom. When an atom gains electrons, forming a negative ion (anion), it becomes larger.
Metallic character describes how easily an element can lose electrons. This property increases as you go down a group and decreases as you go across a period. Non-metallic character is the tendency of an element to gain electrons. This decreases down a group and increases across a period.
Ionization potential is the amount of energy needed to remove an electron from an atom. This energy generally decreases down a group and increases across a period. Electron affinity is the energy change when an atom gains an electron. It tends to decrease down a group and increase across a period, though there are some exceptions. Electronegativity measures how strongly an atom attracts electrons when it is bonded to another atom in a molecule. This property decreases down a group and increases across a period. For example, alkali metals are known for being very reactive metals, while halogens are very reactive non-metals.
The atomic number (often shown as Z) of an element is equal to the number of protons in its nucleus. The mass number (A) is the total count of protons and neutrons in an atom’s nucleus.
Workbook Solutions (Concise/Selina)
Intext Questions and Answers I
1. (i) What is the modern periodic law?
(ii) Who stated this law?
(iii) How many groups and periods are there in the modern periodic table?
Answer:
(i) The modern periodic law states that “the physical and chemical properties of elements are the periodic functions of their atomic number”.
(ii) Henry Moseley stated the modern periodic law.
(iii) The modern periodic table has eighteen vertical columns, known as groups, and seven horizontal rows, known as periods.
2. What are horizontal rows and vertical columns in a periodic table known as?
Answer: In a periodic table, the horizontal rows are called periods, and the vertical columns of elements with similar properties are called groups.
3. Periodicity is observed due to similar ______ (number of valence electrons/atomic number/electronic configuration).
Answer: Periodicity is observed due to similar number of valence electrons. The cause of periodicity is the recurrence of similar electronic configuration, i.e. having the same number of electrons in the outermost orbit.
4. How does the electronic configuration in atoms change
(i) in a period from left to right?
(ii) in a group from top to bottom?
Answer: (i) In a period from left to right, the number of electrons in the valence (outermost) shell increases from left to right, while the number of shells remains the same.
(ii) In a group from top to bottom, the number of electrons in the outermost orbit remains the same, meaning the electronic configuration is similar in terms of valence electrons. However, the number of shells increases successively by one.
5. Name two elements in each case:
(i) Alkali metals
(ii) Alkaline earth metals
(iii) Halogens
(iv) Inert gas
Answer: (i) Two Alkali metals are Lithium (Li) and Sodium (Na).
(ii) Two Alkaline earth metals are Beryllium (Be) and Magnesium (Mg).
(iii) Two Halogens are Fluorine (F) and Chlorine (Cl).
(iv) Two Inert gases are Helium (He) and Neon (Ne).
6. Elements of group 1 and elements of group 17 both have valency 1. Explain.
Answer: Valency depends on the number of electrons in the outermost shell. Elements of group 1 have 1 electron in their outermost shell. If the number of electrons present in the outermost shell is 1, then their valency is 1. Elements of group 17 have 7 electrons in their outermost shell. If the number of electrons present in the outermost shell is 7, then their valency is 8 – 7 = 1. Therefore, both group 1 and group 17 elements have a valency of 1.
7. Correct the statements.
(i) Elements in the same period have the same valency.
(ii) Valency depends upon the number of shells in an atom.
(iii) Copper and zinc are representative elements.
(iv) Transition elements are placed at extreme right of the periodic table.
Answer: (i) Correct statement: Elements in the same period do not have the same valency; valency generally increases up to Group 14 and then decreases.
(ii) Correct statement: Valency depends upon the number of electrons in the outermost shell (i.e. valence shell) of an atom.
(iii) Correct statement: Copper and zinc are transition elements.
(iv) Correct statement: Transition elements are placed in the middle of the periodic table, between s-block and p-block elements.
8. What do you understand by?
(i) Periodicity
(ii) Typical elements
(iii) Orbits
Answer: (i) Periodicity is the phenomenon where properties reappear at regular intervals, or in which there is gradual variation (i.e. increase or decrease) at regular intervals. These properties are called ‘periodic properties’.
(ii) Typical elements are the third period elements, Na, Mg, Al, Si, P, S and Cl, which summarise the properties of their respective groups.
(iii) Orbits, or shells, are certain definite circular paths in which electrons revolve around the nucleus.
9. Name two elements that you would expect to show chemical reactions similar to calcium. What is the basis of your choice?
Answer: Two elements that would be expected to show chemical reactions similar to calcium are Magnesium (Mg) and Strontium (Sr).
The basis of this choice is that elements of the same group have similar chemical properties. Calcium, Magnesium, and Strontium all belong to Group 2 (Alkaline earth metals). Their similarity in chemical properties is due to the fact that they have the same number of electrons (two electrons) in their outermost shell, leading to a similar electronic configuration in the valence shell.
10. Name the
(i) metals
(ii) metalloids and
(iii) non-metals in the first twenty elements.
Answer: In the first twenty elements:
(i) Metals are: Lithium (Li), Beryllium (Be), Sodium (Na), Magnesium (Mg), Aluminium (Al), Potassium (K), Calcium (Ca).
(ii) Metalloids are: Boron (B), Silicon (Si).
(iii) Non-metals are: Hydrogen (H), Helium (He), Carbon (C), Nitrogen (N), Oxygen (O), Fluorine (F), Neon (Ne), Phosphorus (P), Sulphur (S), Chlorine (Cl), Argon (Ar).
11. Fluorine, Chlorine and Bromine are put in one group on the basis of their similar properties.
(i) What are those similar properties?
(ii) What is the common name of this group or family?
Answer: (i) Those similar properties of Fluorine, Chlorine, and Bromine include:
1 They each possess seven valence electrons and therefore show similar properties.
2 They are highly reactive, electronegative non-metals.
3 They form negative ions carrying a single charge (e.g., F⁻, Cl⁻, Br⁻).
4 They are diatomic in their molecular form (e.g., F₂, Cl₂, Br₂).
5 They form salts.
(ii) The common name of this group or family is Halogens.
12. What is the main characteristic of the last element in each period of the Periodic Table? What is the general name of such elements?
Answer: The main characteristic of the last element in each period of the Periodic Table is that these elements have their outermost orbit (valence shell) completely filled.
The general name of such elements is noble gases or inert gases.
13. According to atomic structure, what determines which element will be the first and which will be the last in a period?
Answer: According to atomic structure, the arrangement of electrons in shells determines which element will be the first and which will be the last in a period. Each period begins with an element having one electron in its valence shell, which will be the first element. Each period ends with an element having a completely filled outermost orbit (valence shell), which will be the last element in that period.
14. How do the following vary on moving from left to right in the third period of the periodic table?
(i) valence electrons and
(ii) valency
Answer: On moving from left to right in the third period of the periodic table:
(i) The number of valence electrons increases from 1 (for Sodium) to 8 (for Argon). For example, Na has 1, Mg has 2, Al has 3, Si has 4, P has 5, S has 6, Cl has 7, and Ar has 8 valence electrons.
(ii) The valency first increases from 1 to 4, and then decreases to 1, and is 0 for the last element. For example, Na (1), Mg (2), Al (3), Si (4), P (3), S (2), Cl (1), Ar (0).
15. Name the type of elements which have their
(i) outermost shell complete
(ii) outermost shell incomplete
(iii) two outermost shell incomplete
(iv) one electron short of octet
(v) two electrons in the outermost orbit.
Answer: (i) Elements which have their outermost shell complete are Noble gases or Inert gases (Group 18).
(ii) Elements which have their outermost shell incomplete are Main group elements or Representative elements (except Group 18).
(iii) Elements which have their two outermost shells incomplete are Transition elements (Groups 3 to 12).
(iv) Elements which have one electron short of octet (i.e., 7 valence electrons) are Halogens (Group 17).
(v) Elements which have two electrons in the outermost orbit are Alkaline earth metals (Group 2).
16. An element has 2 electrons in its N shell.
(i) What is its atomic number?
(ii) State its position in the periodic table.
(iii) Is it a metal or a non-metal?
(iv) State the name assigned to this group.
(v) What is the valency of this element?
Answer: If an element has 2 electrons in its N shell (the fourth shell), and assuming the K, L, and M shells are filled as 2, 8, 8 respectively, its electronic configuration is 2, 8, 8, 2.
(i) Its atomic number is 2 + 8 + 8 + 2 = 20.
(ii) It has 4 shells, so it is in Period 4. It has 2 valence electrons, so it is in Group 2.
(iii) It is a metal. (Elements of Group 2 are metals).
(iv) The name assigned to this group (Group 2) is Alkaline earth metals.
(v) Its valency is 2, as it has 2 valence electrons.
17. Answer the following in respect of element ³²₁₆S. (i) Give its electronic configuration. (ii) To which group and period does it belong? (iii) What is its valency?
Answer: For element ³²₁₆S (Sulphur), the atomic number is 16.
(i) Its electronic configuration is 2, 8, 6.
(ii) It has 3 shells (K, L, M), so it belongs to Period 3. It has 6 valence electrons, so it belongs to Group 16 (or VIA).
(iii) Its valency is 8 – 6 = 2.
18. Name
(a) An alkali metal in period 3 and halogen in period 2.
(b) The noble gas with 3 shells.
(c) The non-metals present in period 2 and metals in period 3.
(d) The element of period 3 with valency 4.
(e) The element in period 3 which does not form oxide.
(f) The element of lower nuclear charge out of Be and Mg.
Answer: (a) An alkali metal in period 3 is Sodium (Na). A halogen in period 2 is Fluorine (F).
(b) The noble gas with 3 shells (Period 3) is Argon (Ar).
(c) Non-metals present in period 2 are Carbon (C), Nitrogen (N), Oxygen (O), and Fluorine (F). Metals present in period 3 are Sodium (Na), Magnesium (Mg), and Aluminium (Al).
(d) The element of period 3 with valency 4 is Silicon (Si).
(e) The element in period 3 which does not form oxide is Argon (Ar), as noble gases are generally unreactive.
(f) Out of Be (atomic number 4) and Mg (atomic number 12), the element with lower nuclear charge is Beryllium (Be).
19. The electronic configuration of an element T is 2, 8, 8, 1.
(i) What is the group number of T?
(ii) What is the period number of T?
(iii) How many valence electrons are there in an atom of T?
(iv) What is the valency of T?
(v) Is it a metal or a non-metal?
Answer: The electronic configuration of element T is 2, 8, 8, 1.
(i) The group number of T is 1, as it has 1 valence electron.
(ii) The period number of T is 4, as it has 4 shells.
(iii) There is 1 valence electron in an atom of T.
(iv) The valency of T is 1.
(v) T is a metal (it belongs to Group 1, alkali metals).
20. Match the atomic number 19, 15, 8, 4 and 2 with each of the following:
(i) A metal of valency one.
(ii) A solid non-metal of period 3.
(iii) A rare gas.
(iv) A gaseous element with valency 2.
(v) An element of group 2.
Answer: (i) A metal of valency one: Atomic number 19 (Potassium).
(ii) A solid non-metal of period 3: Atomic number 15 (Phosphorus).
(iii) A rare gas: Atomic number 2 (Helium).
(iv) A gaseous element with valency 2: Atomic number 8 (Oxygen).
(v) An element of group 2: Atomic number 4 (Beryllium).
Intext Questions and Answers II
1. What do you understand by atomic size? State its unit.
Answer: Atomic size (atomic radius) is the distance between the centre of the nucleus of an atom and its outermost shell.
Its units are Angstrom (Å) where 1Å = 10⁻¹⁰ m, and Picometre (pm) where 1 pm = 10⁻¹² m.
2. Give the trends in atomic size on moving:
(i) down the group,
(ii) across the period left to right.
Answer: (i) Down a group, the size of an atom increases as one proceeds from top to bottom.
(ii) Across a period, from left to right, the size of an atom decreases.
3. Arrange the elements of second and third period in increasing order of their atomic size (excluding noble gases).
Answer: The elements of the second period in increasing order of their atomic size are:
Fluorine (64 pm) < Oxygen (66 pm) < Nitrogen (70 pm) < Carbon (77 pm) < Boron (88 pm) < Beryllium (112 pm) < Lithium (152 pm).
The elements of the third period in increasing order of their atomic size are:
Chlorine (99 pm) < Sulphur (104 pm) < Phosphorus (110 pm) < Silicon (117 pm) < Aluminium (143 pm) < Magnesium (160 pm) < Sodium (186 pm).
4. Why is the size of (i) neon greater than fluorine? (ii) sodium greater than magnesium?
Answer: (i) The size of neon is greater than fluorine because, as an exception, the size of the atoms of inert gases are bigger than halogens of the same period. This is because the outermost shell of inert gases is complete. They have the maximum number of electrons in their outermost orbit thus the electronic repulsions are maximum. The effect of nuclear pull over the valence electrons is not seen. Hence, the size of the atom of an inert gas is bigger.
(ii) The size of sodium is greater than magnesium because, in a period, the size of an atom decreases from left to right. This is because the nuclear charge, i.e., the atomic number increases from left to right in the same period, thereby bringing the outermost shell closer to the nucleus. Sodium is to the left of magnesium in the third period.
5. (i) Which is greater in size?
(a) an atom or a cation
(b) an atom or an anion
(c) Fe²⁺ or Fe³⁺
(d) Fluorine or oxygen
(ii) Which has the maximum metallic character – Na, Li or K?
Answer: (i) (a) An atom is greater in size than a cation. A cation is always smaller than the parent atom from which it is formed.
(b) An anion is greater in size than an atom. An anion is larger than the parent atom.
(c) Fe²⁺ is greater in size than Fe³⁺. A cation is formed by the loss of electron(s); the more electrons lost, the smaller the cation becomes as the remaining electrons are more strongly attracted to the nucleus.
(d) Oxygen is greater in size than fluorine. In the second period, atomic size decreases from left to right; oxygen (66 pm) is to the left of fluorine (64 pm).
(ii) Among Na, Li, or K, potassium (K) has the maximum metallic character. Metallic nature increases as one moves down a group, and K is below Li and Na in Group 1.
6. Arrange:
(i) Be, Li, C, B, N, O, F (in increasing metallic character).
(ii) Si, Na, Al, Mg, Cl, P, S (in decreasing non-metallic character).
Answer: (i) The elements Be, Li, C, B, N, O, F in increasing metallic character are:
F < O < N < C < B < Be < Li.
(Metallic character decreases across a period from left to right, so it increases from right to left).
(ii) The elements Si, Na, Al, Mg, Cl, P, S in decreasing non-metallic character are:
Cl > S > P > Si > Al > Mg > Na.
(Non-metallic character increases across a period from left to right, so it decreases from right to left).
7. State the trend in chemical reactivity:
(i) across the third period from left to right,
(ii) down the group.
a. in group IA (1)
b. in group VIIA (17)
Answer: (i) Across the third period from left to right, the chemical reactivity of elements first decreases and then increases. For example, in the third period, reactivity decreases from Na to Si (least reactive), and then increases from P to Cl.
(ii) a. In group IA (1), down the group, the tendency of losing electrons increases. Since chemical reactivity in metals depends upon the tendency to lose electrons, thus reactivity increases on going down the group.
b. In group VIIA (17), down the group, the chemical reactivity of non-metals decreases as it depends upon the tendency to gain electrons, which decreases down the group.
8. A metal M forms an oxide having the formula M₂O₃. It belongs to third period. Write the atomic number and valency of the metal.
Answer: The formula M₂O₃ indicates that the metal M has a valency of 3. Since the metal M belongs to the third period and has a valency of 3, it is Aluminium (Al).
The atomic number of Aluminium is 13.
The valency of Aluminium is 3.
9. An element X belongs to 3rd period and 17th group, state
(i) no. of valence electrons in it.
(ii) name of the element.
Answer: (i) An element in the 17th group has 7 valence electrons.
(ii) The element X, belonging to the 3rd period and 17th group, is Chlorine (Cl).
10. The given table shows elements with same number of electrons in its valence shell.
State:
(i) Whether these elements belong to the same group or period.
(ii) Arrange them in order of increasing metallic character.
Answer: (i) Since the elements have the same number of electrons in their valence shell, they belong to the same group.
(ii) The melting points of metals generally decrease on going down a group, and metallic character increases down a group.
The melting points are A (63.0 °C), C (97.0 °C), B (180.0 °C).
Comparing with Group 1 metals: K (63.5°C), Na (94.5°C), Li (180.5°C).
So, A is like K, C is like Na, and B is like Li.
The order of increasing metallic character (Li < Na < K) is therefore B < C < A.
11. Which one of the following has the largest atomic radius?
(i) Sodium (ii) Potassium (iii) Magnesium (iv) Aluminium
Answer: Among Sodium, Potassium, Magnesium, and Aluminium, Potassium (K) has the largest atomic radius. Atomic radius increases down a group (K is below Na) and decreases across a period (Na > Mg > Al).
12. Which one has the largest size?
(i) Br (ii) I (iii) I⁻ (iv) Cl
Answer: Among Br, I, I⁻, and Cl, the iodide ion (I⁻) has the largest size. Atomic size increases down the halogen group (Cl < Br < I), and an anion (I⁻) is larger than its parent atom (I).
13. The metals of group 2 from top to bottom are Be, Mg, Ca, Sr and Ba
(i) Which one of these elements will form ions most readily and why?
(ii) State the common feature in their electronic configuration.
Answer: (i) Barium (Ba) will form ions most readily. This is because, on moving down a group, the atomic size increases, and the tendency to lose electrons increases. Elements at the bottom of a group are most metallic, their atomic size is large, hence electrons are loosely held, and ions are readily formed, making them more reactive.
(ii) The common feature in their electronic configuration is that they all have two electrons in their outermost shell.
14. Write the number of protons, neutrons and electronic configuration of ³⁹₁₉K, ³¹₁₅P. Also state their position in the periodic table.
Answer: For ³⁹₁₉K (Potassium):
- Number of protons = 19
- Number of neutrons = 39 – 19 = 20
- Electronic configuration = 2, 8, 8, 1
- Position in periodic table = Period 4, Group 1
For ³¹₁₅P (Phosphorus):
- Number of protons = 15
- Number of neutrons = 31 – 15 = 16
- Electronic configuration = 2, 8, 5
- Position in periodic table = Period 3, Group 15
15. Name the element which has:
(i) two shells, both of which are completely filled with electrons?
(ii) the electronic configuration 2, 8, 3?
(iii) a total of three shells with five electrons in its valence shell?
(iv) a total of four shells with two electrons in its valence shell?
(v) twice as many electrons in its second shell as in its first shell?
Answer: (i) Neon (Ne), with electronic configuration 2, 8.
(ii) Aluminium (Al).
(iii) Phosphorus (P), with electronic configuration 2, 8, 5.
(iv) Calcium (Ca), with electronic configuration 2, 8, 8, 2.
(v) Carbon (C), with electronic configuration 2, 4.
16. An element Barium has atomic number 56. Look up its position in the Periodic Table and answer the following questions.
(i) Is it a metal or a non-metal?
(ii) Is it more or less reactive than calcium?
(iii) What is its valency?
(iv) What will be the formula of its phosphate?
(v) Is it larger or smaller than caesium (Cs) in size?
Answer: Barium (Ba), atomic number 56, is in Period 6, Group 2.
(i) It is a metal (an alkaline earth metal).
(ii) It is more reactive than calcium (Ca) because reactivity of metals in Group 2 increases down the group.
(iii) Its valency is 2.
(iv) The formula of its phosphate is Ba₃(PO₄)₂.
(v) It is smaller than caesium (Cs) in size. Caesium (Group 1) is to the left of Barium (Group 2) in the same period (Period 6), and atomic size decreases across a period.
17. In group I of the Periodic Table, three elements X, Y and Z have ionic radii 1.33 Å, 0.95 Å and 0.60 Å respectively. Giving a reason, arrange them in the order of increasing atomic numbers in the group.
Answer: The ionic radii are Z (0.60 Å), Y (0.95 Å), and X (1.33 Å).
In Group I, ionic radius increases down the group as the atomic number increases, due to the addition of electron shells.
Therefore, the element with the smallest ionic radius (Z) will have the smallest atomic number, and the element with the largest ionic radius (X) will have the largest atomic number.
The order of increasing atomic numbers is Z < Y < X.
Reason: In a group, as the atomic number increases, the number of electron shells increases, leading to an increase in ionic size. Thus, a smaller ionic radius corresponds to a lower position (smaller atomic number) and a larger ionic radius corresponds to a higher position (larger atomic number) when moving from top to bottom in the group.
18. Explain why are the following statements not correct:
(i) All groups contain metals and non metals.
(ii) Atoms of elements in the same group have the same number of electron(s).
(iii) Non-metallic character decreases across a period with increase in atomic number.
(iv) Reactivity increases with atomic number in a group as well as in a period.
Answer: (i) This statement is not correct because some groups consist entirely of metals (e.g., Group 1 Alkali metals, Group 2 Alkaline earth metals) or entirely of non-metals (e.g., Group 18 Noble gases).
(ii) This statement is not correct. Atoms of elements in the same group have the same number of electrons in their outermost shell (valence electrons), not the same total number of electrons.
(iii) This statement is not correct. Non-metallic character increases across a period from left to right with an increase in atomic number.
(iv) This statement is not correct. Across a period, chemical reactivity of elements first decreases and then increases. In a group, for metals like alkali metals, reactivity increases with atomic number, but for non-metals like halogens, reactivity decreases with atomic number down the group.
19. (i) State the number of elements in Period 1, Period 2, and Period 3 of the periodic table. Name them.
(ii) What is the common feature of the electronic configuration of the elements at the end of Period 2 and Period 3?
(iii) If an element is in Group 17, it is likely to be [metallic/non-metallic] in character, while with one electron in its outermost energy level (shell), then it is likely to be [metallic/non-metallic].
(iv) In Period 3, the most metallic element is (sodium/magnesium/aluminium).
Answer: (i)
- Period 1 has 2 elements: Hydrogen (H) and Helium (He).
- Period 2 has 8 elements: Lithium (Li), Beryllium (Be), Boron (B), Carbon (C), Nitrogen (N), Oxygen (O), Fluorine (F), and Neon (Ne).
- Period 3 has 8 elements: Sodium (Na), Magnesium (Mg), Aluminium (Al), Silicon (Si), Phosphorus (P), Sulphur (S), Chlorine (Cl), and Argon (Ar).
(ii) The common feature of the electronic configuration of the elements at the end of Period 2 (Neon: 2,8) and Period 3 (Argon: 2,8,8) is that their outermost orbit (valence shell) is completely filled, typically with 8 electrons (an octet), making them stable.
(iii) If an element is in Group 17, it is likely to be non-metallic in character, while with one electron in its outermost energy level (shell), then it is likely to be metallic.
(iv) In Period 3, the most metallic element is sodium.
20. Complete the following sentences choosing the correct word or words from those given in brackets at the end of each sentence:
(i) The properties of the elements are a periodic function of their ___________(atomic number, mass number, relative atomic mass).
(ii) Moving across a ____________ of the Periodic Table the elements show increasing _ character (group, period, metallic, non-metallic).
(iii) The elements at the bottom of a group would be expected to show _________ metallic character than the element at the top (less, more).
(iv) The similarities in the properties of a group of elements are because they have the same ___________ (electronic configuration, number of outer electrons, atomic numbers).
Answer:
(i) The properties of the elements are a periodic function of their atomic number.
(ii) Moving across a period of the Periodic Table the elements show increasing non-metallic character.
(iii) The elements at the bottom of a group would be expected to show more metallic character than the element at the top.
(iv) The similarities in the properties of a group of elements are because they have the same number of outer electrons.
21. Give reasons for the following:
(i) The size of a Cl⁻ ion is greater than the size of a Cl atom.
(ii) Argon atom is bigger than chlorine atom.
(iii) Chlorine is less reactive than fluorine.
(iv) Inert gases do not form ion.
Answer: (i) The size of a Cl⁻ ion is greater than the size of a Cl atom because an anion is formed by the gain of electron(s). Thus, the number of electrons are more than protons. These electrons repel each other, which reduces the overall attraction between the nucleus and the electrons, thereby reducing the effective nuclear charge, leading to a larger size.
(ii) Argon atom is bigger than chlorine atom, although both are in the third period, because the size of the atoms of inert gases are bigger than halogens of the same period. This is because the outermost shell of inert gases is complete, leading to maximum electronic repulsions, and the effect of nuclear pull over the valence electrons is not as pronounced.
(iii) Chlorine is less reactive than fluorine because both are halogens (Group 17), and the chemical reactivity of non-metals decreases on going down the group. Fluorine is at the top of Group 17 and is the most reactive non-metal.
(iv) Inert gases do not form ions readily because these elements have their outermost orbit complete. Due to this stable electronic configuration, they hardly react with other elements, meaning they do not easily lose or gain electrons to form ions. They have very high ionisation energies and zero or positive electron affinities.
Intext Questions and Answers III
1. (a) Define the term ‘ionisation potential’.
Answer: The energy required to remove an electron from a neutral isolated gaseous atom and convert it into a positively charged gaseous ion is called ionisation potential (I.P.) or ionisation energy (I.E.) or first ionisation energy (IE₁).
(b) Represent it in the form of an equation. In which unit is it measured?
Answer: Ionisation potential can be represented by the equation: M(g) + I.E. → M⁺(g) + e⁻, where M can be any element.
I.E. is measured in electron volts per atom (eV/atom) and its S.I. unit is kilojoule per mole (kJ mol⁻¹).
2. Ionisation Potential values depend on (a) atomic size (b) nuclear pull. Explain.
Answer: Ionisation Potential values depend on:
(a) Atomic size: The greater the atomic size, the lesser the force of attraction. Since the electrons of the outermost shell lie further away from the nucleus, it makes their removal easier, i.e., the ionisation energy required is less.
(b) Nuclear pull (Nuclear charge): The greater the nuclear charge, greater is the attraction for the electrons of the outermost shell. Therefore, the electrons in the outermost shell are more firmly held because of which greater energy is required to remove the electron(s).
3. State the trends in ionisation energy :
(a) across the period, (b) down the group.
Answer: The trends in ionisation energy are as follows:
(a) Across a period: The ionisation energy tends to increase as one moves from left to right across a period (with exceptions), because the atomic size decreases due to an increase in the nuclear charge, and thus, more energy is required to remove the electron(s).
(b) Down a group: There is an increase in atomic number (nuclear charge) and atomic size down the group due to the addition of extra shells. This increase in the atomic size overcomes the effect of an increase in the nuclear charge. Therefore, ionisation energy decreases with an increase in the atomic size, i.e., it decreases as one moves down a group.
4. Name the elements with highest and lowest ionisation energies in first three periods.
Answer: In the first three periods, Helium (He) will have the highest ionisation energy (2372.0 kJ mol⁻¹). The element with the lowest ionisation energy in the first three periods is Sodium (Na) with an ionisation energy of 496 kJ mol⁻¹.
5. (a) Arrange the elements of second and third periods in increasing order of their ionisation energy.
Answer: The elements of the second period in increasing order of their ionisation energy are:
Li (520 kJ mol⁻¹) < B (801 kJ mol⁻¹) < Be (899 kJ mol⁻¹)* < C (1086 kJ mol⁻¹) < O (1314 kJ mol⁻¹) < N (1402 kJ mol⁻¹)* < F (1681 kJ mol⁻¹) < Ne (2080 kJ mol⁻¹).
The elements of the third period in increasing order of their ionisation energy are:
Na (496 kJ mol⁻¹) < Al (577 kJ mol⁻¹) < Mg (737 kJ mol⁻¹)* < Si (786 kJ mol⁻¹) < S (999 kJ mol⁻¹) < P (1011 kJ mol⁻¹)* < Cl (1256 kJ mol⁻¹) < Ar (1520 kJ mol⁻¹).
(b) The element with the highest ionisation potential is:
A. Hydrogen
B. Caesium
C. Radon
D. Helium (2020)
Answer: D. Helium. Helium will have the highest ionisation energy.
6. (a) Define the term ‘electron affinity’. State its unit.
Answer: The amount of energy released while converting a neutral gaseous isolated atom into a negatively charged gaseous ion (anion) by the addition of electron is called Electron Affinity (E.A.).
Its unit is electron volts per atom (eV/atom) or kJ mol⁻¹.
(b) Arrange the elements of second period in increasing order of their electron affinity. Name the elements which do not follow the trend in this period.
Answer: The elements of the second period, based on available electron affinity values (in kJ mol⁻¹ where more negative means higher affinity), arranged in increasing order of electron affinity (from least tendency to accept an electron/most positive or least negative value, to highest tendency/most negative value) are approximately:
Ne (+48) < Be (positive value, exception) < N (positive value, exception) < Li (-60) < B (-88) < C (-122) < O (-141) < F (-328).
The elements which do not follow the general trend (of increasing electron affinity across the period) in this period are Be (Beryllium), N (Nitrogen), and Ne (Neon). Groups 2 (like Be) and 15 (like N) do not show regular trends and are exceptions. Inert gases (like Ne) have zero electron affinity or positive values.
(c) Which has higher E.A., Fluorine or Neon ?
Answer: Fluorine has a higher electron affinity than Neon. Fluorine (F) has an E.A. of -328 kJ mol⁻¹, while Neon (Ne) has an E.A. of +48 kJ mol⁻¹, indicating Fluorine releases significant energy on gaining an electron, whereas Neon requires energy or releases very little.
7. Electron affinity values generally ……….. across the period left to right and ……….. down the group top to bottom.
Answer: Electron affinity values generally increases across the period left to right and decreases down the group top to bottom.
8. (a) Define the term ‘Electronegativity’. State its unit.
Answer: The tendency of an atom in a molecule to attract the shared pair of electrons towards itself is called its electronegativity.
Electronegativity is a dimensionless property; it has no unit.
(b) Among the elements given below, the element with least Electronegativity is :
(i) Lithium (ii) Boron (iii) Carbon (iv) Fluorine
Answer: (i) Lithium. Lithium has an electronegativity value of 1.0, Boron is 2.0, Carbon is 2.5, and Fluorine is 4.0.
9. Explain the following:
(a) Group 17 elements are strong non-metals, while group 1 elements are strong metals.
Answer: Group 17 elements (halogens) are strong non-metals because they have 7 valence electrons and a strong tendency to gain one electron to complete their octet, due to their relatively small atomic size and high nuclear charge for their respective periods. This makes them highly electronegative.
Group 1 elements (alkali metals) are strong metals because they have only one valence electron, which they tend to lose easily to form a positive ion and achieve a stable electronic configuration. This is due to their large atomic size and low nuclear pull on the outermost electron, making them highly electropositive.
(b) Metallic character of elements decreases from left to right in a period while it increases in moving down a group.
Answer: Metallic character of elements decreases from left to right in a period because, on moving across a period, the nuclear pull increases due to the increase in atomic number, and the atomic size decreases. Hence, elements cannot lose electrons easily.
Metallic character increases in moving down a group because, on moving down a group, the atomic size increases due to the addition of new shells. The effect of an increased atomic size is greater as compared to the increased nuclear charge. Therefore, the tendency to lose electrons increases, and elements can lose electrons easily.
(c) Halogens have a high electron affinity.
Answer: Halogens (Group 17 elements) have a high electron affinity. This is because they have a small atomic size and a high nuclear charge for their respective periods, leading to a strong effective attractive force between the nucleus and an incoming electron. They require only one electron to achieve a stable octet configuration, so energy is readily released when they accept an electron.
(d) The reducing power of an element increases down a group while decreases along a period.
Answer: The reducing power of an element (characteristic of metals which are good reducing agents) increases down a group because the tendency to lose electron(s) increases down a group due to increasing atomic size and decreasing effective nuclear charge on the valence electrons. Greater the tendency to lose electron(s), stronger is the reducing agent.
The reducing power decreases along a period from left to right because the tendency to lose electrons decreases across a period due to increasing nuclear charge and decreasing atomic size.
(e) Size of atoms progressively becomes smaller when we move from sodium (Na) to chlorine (Cl) in the third period of the Periodic Table.
Answer: The size of atoms progressively becomes smaller when we move from sodium (Na) to chlorine (Cl) in the third period. This is because the nuclear charge, i.e., the atomic number, increases from left to right in the same period, thereby attracting the outermost shell electrons more strongly and bringing the outermost shell closer to the nucleus. For example, in the third period, the atomic size trend is Na > Mg > Al > Si > P > S > Cl.
10. Name the periodic property which relates to the :
(a) amount of energy required to remove an electron from an isolated gaseous atom,
Answer: Ionisation potential or Ionisation energy.
(b) character of element which loses one or more electrons when supplied with energy,
Answer: Metallic character or electropositivity.
(c) tendency of an atom in a molecule to attract the shared pair of electrons.
Answer: Electronegativity.
11. This question refers to the elements of the Periodic Table with atomic numbers from 3 to 18. Some of the elements are shown by letters, but the letters are not the usual symbols of the elements.
Which of these:
(a) is the most electronegative element ?
Answer: G (which represents Fluorine, Z=9) is the most electronegative element among these.
(b) is a halogen?
Answer: G (Fluorine, Z=9) and O (Chlorine, Z=17) are halogens.
(c) is an alkali metal ?
Answer: A (Lithium, Z=3) and I (Sodium, Z=11) are alkali metals.
(d) is an element with valency 4?
Answer: D (Carbon, Z=6) and L (Silicon, Z=14) are elements with valency 4.
(e) has the least Ionisation Energy ?
Answer: I (Sodium, Z=11) has the least Ionisation Energy among these elements.
(f) has the least atomic size in period 3?
Answer: O (Chlorine, Z=17) has the least atomic size in period 3 (excluding the noble gas P).
12. A group of elements in the Periodic Table are given ahead (boron is the first member of the group and Thallium is the last).
Boron, Aluminium, Gallium, Indium, Thallium. Answer the following questions in relation to this group of elements:
Answer: The statement describes Group 13 of the Periodic Table, also known as the Boron group. Boron is the first member, and Thallium is a later member in this group.
(a) Which element has the most metallic character ?
Answer: In a group, metallic nature increases as one moves down a group. Elements at the bottom of a group are most metallic. Since Thallium is at the bottom of this group (Boron, Aluminium, Gallium, Indium, Thallium), Thallium has the most metallic character.
(b) Which element would be expected to have the highest electronegativity ?
Answer: Electronegativity decreases down a group. Therefore, the element at the top of the group, Boron, would be expected to have the highest electronegativity.
(c) If the electronic configuration of aluminium is 2, 8, 3, how many electrons are there in the outer shell of thallium ?
Answer: Elements in the same group have the same number of electrons in the outermost shell. Since Aluminium (Al) has an electronic configuration of 2, 8, 3, it has 3 electrons in its outermost shell. Thallium (Tl) is in the same group as Aluminium, so Thallium also has 3 electrons in its outer shell.
(d) The atomic number of boron is 5. Write the chemical formula of the compound formed when boron reacts with chlorine.
Answer: Boron has an atomic number of 5, so its electronic configuration is 2, 3. It has 3 valence electrons, and its valency is 3. Chlorine is a halogen and has a valency of 1. Therefore, the chemical formula of the compound formed when boron reacts with chlorine is BCl₃.
(e) Will the elements in the group to the right of this boron group be more metallic or less metallic in character? Justify your answer.
Answer: The elements in the group to the right of this boron group (Group 13) will be less metallic in character.
This is because the metallic nature decreases across a period, moving from left to right. On moving across a period, nuclear pull increases due to the increase in atomic number, and thus the atomic size decreases. Hence, elements cannot lose electrons easily, leading to a decrease in metallic character.
Exercise
MCQs
1. In the periodic table, alkali metals are placed in group:
(a) 1
(b) 11
(c) 17
(d) 18
Answer: (a) 1
2. Which of the following properties does not match with the elements of the halogen family?
(a) They have seven electrons in their valence shell.
(b) They are highly reactive chemically.
(c) They are metallic in nature.
(d) They are diatomic in their molecular form.
Answer: (c) They are metallic in nature.
3. With reference to the variation of properties in the Periodic Table, which of the following is generally true?
(a) Atomic size increases from left to right across a period.
(b) Ionization potential increases from left to right across a period.
(c) Electron affinity increases going down a group.
(d) Electronegativity increases going down a group.
Answer: (b) Ionization potential increases from left to right across a period.
4. An element in period 3 whose electron affinity is zero is:
(a) Neon
(b) Sulphur
(c) Sodium
(d) Argon
Answer: (d) Argon
5. The number of electrons in the valence shell of a halogen is:
(a) 1
(b) 3
(c) 5
(d) 7
Answer: (d) 7
6. Among period 2 elements, the element which has the highest electron affinity is:
(a) Lithium
(b) Carbon
(c) Chlorine
(d) Fluorine
Answer: (d) Fluorine
7. Ionisation potential increases over a period from left to right because:
(a) Atomic radius and nuclear charge increase
(b) Atomic radius and nuclear charge decrease
(c) Atomic radius increases and nuclear charge decreases
(d) Atomic radius decreases and nuclear charge increases
Answer: (d) Atomic radius decreases and nuclear charge increases
8. An element A belonging to period 3 and group II will have:
(a) 3 shells and 2 valence electrons
(b) 2 shells and 3 valence electrons
(c) 3 shells and 3 valence electrons
(d) 2 shells and 2 valence electrons
Answer: (a) 3 shells and 2 valence electrons
9. Among the elements given below, the element with the least electronegativity is:
(a) Lithium
(b) Carbon
(c) Boron
(d) Fluorine
Answer: (a) Lithium
10. An element with atomic number 19 will most likely combine chemically with the elements whose atomic number is:
(a) 17
(b) 11
(c) 28
(d) 20
Answer: (a) 17
11. Parts (i) to (iv) refer to changes in the properties of elements on moving from left to right across a period of the Periodic Table. For each property, choose the correct answer.
(i) The non-metallic character of the elements :
(a) decreases,
(b) increases,
(c) remains the same,
(d) depends on the period
Answer: (b) increases,
(ii) The electronegativity :
(a) depends on the number of valence electrons,
(b) remains the same,
(c) decreases,
(d) increases
Answer: (d) increases
(iii) The ionization potential:
(a) goes up and down
(b) decreases
(c) increases
(d) remains the same
Answer: (c) increases
(iv) The atomic size :
(a) decreases,
(b) increases,
(c) remains the same,
(d) sometimes increases and sometimes decreases
Answer: (a) decreases,
12. In the periodic table while going down in the halogen group:
(a) reactivity will increase
(b) electronegativity will increase
(c) ionic radius will increase
(d) ionisation potential will increase
Answer: (c) ionic radius will increase
13. Electron affinity is the :
(a) Power of an atom to attract an electron to itself.
(b) Energy released when an electron is added to an isolated atom in the gaseous state.
(c) Energy absorbed when an electron is added to an isolated atom in the gaseous state.
(d) Energy required to remove an electron from an isolated gaseous atom.
Answer: (b) Energy released when an electron is added to an isolated atom in the gaseous state.
14. Which of these statements gives the correct picture regarding halogens and alkali metals with respect to an increase in the atomic number?
(a) Reactivity decreases in alkali metals but increases in halogens.
(b) Reactivity increases in both.
(c) Reactivity decreases in both.
(d) Reactivity increases in alkali metals but decreases in halogens.
Answer: (d) Reactivity increases in alkali metals but decreases in halogens.
15. The correct order of increasing ionisation energy of Be, Mg, Ca, Sr is:
(a) Be, Mg, Ca, Sr
(b) Ca, Mg, Be, Sr
(c) Sr, Ca, Mg, Be
(d) Mg, Ca, Sr, Be
Answer: (c) Sr, Ca, Mg, Be
16. An element ₁₃X combines with ₁₇Y to form a compound. Which of the following is true ?
P X is a metal, Y is a metal.
Q X is a metal, Y is a non metal.
R X looses electron(s), Y gains electron(s).
(a) Only P
(b) Only Q
(c) Both P and Q
(d) Both Q and R
Answer: (d) Both Q and R
17. Assertion (A) : Second period consists of 8 elements.
Reason (R) : Number of elements in each period is four times the number of atomic orbitals available in the energy level that is being filled.
(1) Both A and R are true and R is the correct explanation of A.
(2) Both A and R are true but R is not the correct explanation of A.
(3) A is true but R is false.
(4) A is false but R is true.
(a) (1)
(b) (2)
(c) (3)
(d) (4)
Answer: (c) (3)
18. Assertion (A): In a Dobereiner’s triad, the three elements present have the same difference of atomic masses.
Reason (R): Elements in a triad have similar properties.
(1) Both A and R are true and R is the correct explanation of A.
(2) Both A and R are true but R is not the correct explanation of A.
(3) A is true but R is false.
(4) A is false but R is true.
(a) (1)
(b) (2)
(c) (3)
(d) (4)
Answer: (d) (4)
19. Assertion (A) : Smaller the size of an atom greater is its electronegativity.
Reason (R) : Electronegativity is the tendency of an atom to attract shared pair of electrons towards itself in a molecule.
(1) Both A and R are true and R is the correct explanation of A.
(2) Both A and R are true but R is not the correct explanation of A.
(3) A is true but R is false.
(4) A is false but R is true.
(a) (1)
(b) (2)
(c) (3)
(d) (4)
Answer: (b) (2)
20. Assertion (A) : Hydrogen is placed in group I.
Reason (R) : Hydrogen can gain an electron to achieve noble gas configuration.
(1) Both A and R are true and R is the correct explanation of A.
(2) Both A and R are true but R is not the correct explanation of A.
(3) A is true but R is false.
(4) A is false but R is true.
(a) (1)
(b) (2)
(c) (3)
(d) (4)
Answer: (b) (2)
21. Assertion (A) : Atomic size increases along a period.
Reason (R) : Effective nuclear charge increases with atomic number.
(1) Both A and R are true and R is the correct explanation of A.
(2) Both A and R are true but R is not the correct explanation of A.
(3) A is true but R is false.
(4) A is false but R is true.
(a) (1)
(b) (2)
(c) (3)
(d) (4)
Answer: (d) (4)
22. Assertion (A): Elements in the same vertical column have similar properties.
Reason (R): Properties depend upon the number of electrons in the valence shell.
(1) Both A and R are true and R is the correct explanation of A.
(2) Both A and R are true but R is not the correct explanation of A.
(3) A is true but R is false.
(4) A is false but R is true.
(a) (1)
(b) (2)
(c) (3)
(d) (4)
Answer: (a) (1)
Very Short Answer Type
1. Fill in the blanks by selecting the correct word from the brackets:
(a) The element below sodium in the same group would be expected to have a …………. electro-negativity than sodium and the element above chlorine would be expected to have a …………… ionization potential than chlorine.
(b) On moving down a group, the number of valence electrons ……………….
(c) Metals are good ……………. because they are electron …………….
(d) Down a group, electron affinity ……………….
(e) Electronegativity across the period …………….
(f) Non-metallic character down the group ……………
(g) In a period, increase in electron affinity increases …………….(oxidising power).
(h) On descending a group, …………….. in ionisation potential as well as electron affinity …………. oxidising capacity.
(i) If an element has a low ionization energy then it is likely to be ……………..
(j) If an element has seven electrons in its outermost shell then it is likely to have the …………….. atomic size among all the elements in the same period.
Answer: (a) The element below sodium in the same group would be expected to have a lower electro-negativity than sodium and the element above chlorine would be expected to have a higher ionization potential than chlorine.
(b) On moving down a group, the number of valence electrons remains the same.
(c) Metals are good reducing agents because they are electron donors.
(d) Down a group, electron affinity decreases.
(e) Electronegativity across the period increases.
(f) Non-metallic character down the group decreases.
(g) In a period, increase in electron affinity increases oxidation (oxidising power).
(h) On descending a group, decrease in ionisation potential as well as electron affinity decreases oxidising capacity.
(i) If an element has a low ionization energy then it is likely to be metallic.
(j) If an element has seven electrons in its outermost shell then it is likely to have the smallest atomic size among all the elements in the same period.
2. Rewrite the following sentences by using the correct symbol > (greater than) or < (less than) in the blanks given :
Answer: (a) The ionization potential of Potassium is < that of Sodium.
(b) The electronegativity of Iodine is < that of Chlorine.
3. In Period 3 of the Periodic Table, element B is placed to the left of element A. On the basis of this information, choose the correct word from the brackets to complete the following statements :
Answer: (a) The element B would have higher metallic character than A.
(b) The element A would probably have higher electron affinity than B.
(c) The element A would have smaller atomic size than B.
4. There are three elements E, F, G with atomic numbers 19, 8 and 17 respectively. Classify the above elements as metals and non-metals.
Answer: Element E, with atomic number 19 (Potassium), has an electronic configuration of 2, 8, 8, 1. It tends to lose its valence electron and is an alkali metal. Therefore, E is a metal.
Element F, with atomic number 8 (Oxygen), has an electronic configuration of 2, 6. It tends to gain electrons to complete its octet. Therefore, F is a non-metal.
Element G, with atomic number 17 (Chlorine), has an electronic configuration of 2, 8, 7. It tends to gain an electron to complete its octet and is a halogen. Therefore, G is a non-metal.
5. Arrange the following as per instructions given in the brackets.
(a) Mg, Cl, Na, S, Si (decreasing order of atomic size)
Answer: Na > Mg > Si > S > Cl
(b) Cs, Na, Li, K, Rb (increasing metallic character)
Answer: Li < Na < K < Rb < Cs
(c) Na, Mg, Cl, S, Si (increasing ionisation potential)
Answer: Na < Mg < Si < S < Cl
(d) Cl, F, Br, I (increasing electron affinity)
Answer: I < Br < F < Cl
(e) Cs, Na, Li, K, Rb (decreasing electronegativity)
Answer: Li > Na > K > Rb > Cs
(f) K, Pb, Ca, Zn (increasing reactivity)
Answer: Pb < Zn < Ca < K
(g) Li, K, Na, H (decreasing order of their potential ionisation) (2019)
Answer: H > Li > Na > K
6. The electronegativities (according to Pauling) of the elements in Period 3 of the Periodic Table are as follows with the elements arranged in alphabetical order:
Arrange the elements in the order in which they occur in the Periodic Table from left to right.
(The group 1 element first, followed by the group 2 element and so on, up to group 71.
Answer: The elements in Period 3, arranged in the order in which they occur in the Periodic Table from left to right, based on their increasing electronegativity values provided (Na 0.9, Mg 1.2, Al 1.5, Si 1.8, P 2.1, S 2.5, Cl 3 .0) are: Na, Mg, Al, Si, P, S, Cl
7. Name the metal present in period 3 and group 1 of the periodic table.
Answer: The metal present in period 3 and group 1 of the periodic table is Sodium (Na). An element in period 3 has three shells, and an element in group 1 has one electron in its outermost shell. Sodium has atomic number 11, and its electronic configuration is 2, 8, 1.
8. Give one word or phrase for the amount of energy released when an atom in the gaseous state accepts an electron to form an anion.
Answer: The amount of energy released while converting a neutral gaseous isolated atom into a negatively charged gaseous ion (anion) by the addition of an electron is called Electron Affinity (E.A.).
9. The formula of an ion of an element A is A²⁺. Element A probably belongs to which group?
Answer: If the formula of an ion of an element A is A²⁺, it means that element A has lost 2 electrons. Elements that lose 2 electrons to form a positive ion typically have 2 electrons in their outermost shell. Elements with 2 electrons in their outermost shell belong to Group 2 (Alkaline earth metals).
10. State the group and period of the element having three shells with three electrons in its valence shell.
Answer: An element having three shells belongs to Period 3. If it has three electrons in its valence shell, it belongs to Group 13 (or IIIA according to the old notation).
Short Answer Type Questions
1.What is the significance of atomic number in the modern periodic table?
Answer: The physical and chemical properties of elements are the periodic functions of their atomic number. Elements in the modern periodic table are arranged in order of increasing atomic number (proton number). Atomic number is the unique property of an element, because no two elements have the same atomic number. It gives the electronic configuration of an element and helps in finding the position of an element in the periodic table. Hence, atomic number is the fundamental property of an element.
2. Arrange the following in order of :
(i) increasing radii.
(a) Cl, Cl⁻
Answer: The order of increasing radii is Cl < Cl⁻. Anion is formed by the gain of electron(s). Thus, the number of electron(s) are more than proton(s). These electrons repel each other, which reduces the overall attraction between the nucleus and the electrons, thereby reducing the effective nuclear charge, leading to a larger size. Thus chlorine ion is bigger than chlorine atom.
(b) Mg²⁺, Mg, Mg⁺
Answer: The order of increasing radii is Mg²⁺ < Mg⁺ < Mg. A cation is always smaller than the parent atom from which it is formed. A cation is formed by the loss of electron(s), hence proton(s) are more than electron(s) in a cation. So the remaining electrons are strongly attracted by the nucleus and are pulled inward, hence the size decreases. Mg loses one electron to form Mg⁺, and two electrons to form Mg²⁺, so Mg²⁺ will be smaller than Mg⁺, which in turn is smaller than the parent Mg atom.
(c) N, O, P
Answer: The order of increasing radii is O < N < P. In a period, the size of an atom decreases from left to right; nitrogen (N) and oxygen (O) are in the second period, with O to the right of N, so O is smaller than N. In a group, the size of an atom increases as one proceeds from top to bottom; nitrogen (N) and phosphorus (P) are in the same group (Group 15), with P below N, so N is smaller than P.
(ii) increasing ionisation energy.
(a) P, Na, Cl
Answer: The order of increasing ionisation energy is Na < P < Cl. The ionisation energy tends to increase as one moves from left to right across a period. Sodium (Na), Phosphorus (P), and Chlorine (Cl) are in the third period. Their ionisation energies in kJ mol⁻¹ are Na (496), P (1011*), and Cl (1256).
(b) F, O, Ne
Answer: The order of increasing ionisation energy is O < F < Ne. The ionisation energy tends to increase as one moves from left to right across a period. Oxygen (O), Fluorine (F), and Neon (Ne) are in the second period. Their ionisation energies in kJ mol⁻¹ are O (1314), F (1681), and Ne (2080).
(c) Ne, He, Ar
Answer: The order of increasing ionisation energy is Ar < Ne < He. Helium (He), Neon (Ne), and Argon (Ar) are noble gases. Ionisation energy decreases as one moves down a group. Helium is in Period 1, Neon in Period 2, and Argon in Period 3. Helium will have the highest ionisation energy (2372.0 kJ mol⁻¹).
3. Atomic numbers of elements A, B, C, D, E, F are 8, 7, 11, 12, 13 and 9 respectively. State the type of ions they form.
Answer:: Element A (atomic number 8, Oxygen, electronic configuration 2,6) will gain 2 electrons to form an anion, A²⁻.
Element B (atomic number 7, Nitrogen, electronic configuration 2,5) will gain 3 electrons to form an anion, B³⁻.
Element C (atomic number 11, Sodium, electronic configuration 2,8,1) will lose 1 electron to form a cation, C⁺.
Element D (atomic number 12, Magnesium, electronic configuration 2,8,2) will lose 2 electrons to form a cation, D²⁺.
Element E (atomic number 13, Aluminium, electronic configuration 2,8,3) will lose 3 electrons to form a cation, E³⁺.
Element F (atomic number 9, Fluorine, electronic configuration 2,7) will gain 1 electron to form an anion, F⁻.
4. Give reasons:
(a) The oxidising power of elements increases from left to right along a period.
Answer: Greater the value of electron affinity, easier it is to gain electron(s) and more non-metallic or more electronegative or more oxidising is the element. Thus oxidising power increases left to right in a period. Non-metals are oxidising agents, and non-metallic character increases across a period. The tendency to gain electron(s) increases due to an increase in nuclear pull and a decrease in the atomic size across a period.
(b) Ionisation potential of elements increases across a period from left to right.
Answer: The ionisation energy tends to increase as one moves from left to right across a period (with exceptions), because the atomic size decreases due to an increase in the nuclear charge, and thus, more energy is required to remove the electron(s).
(c) Alkali metals are good reducing agents.
Answer: Elements which lose electron(s) to complete their octet are reducing agents. Metals are good reducing agents. Greater the tendency to lose electron(s), stronger is the reducing agent. Alkali metals are metals that usually have low ionisation energy and readily lose electrons. Specifically, alkali metals are strong reducing agents as they lose electrons to complete their octet.
Long Answer Type Questions
1. Chlorine in the Periodic Table is surrounded by the elements with atomic number 9, 16, 18 and 35.
(a) Which of these have Physical and Chemical properties resembling chlorine ?
(b) Which is more electronegative than chlorine ?
Answer: (a) Elements in the same group have similar chemical properties because they have the same number of electrons in their outermost shell. Chlorine (atomic number 17) is a halogen, belonging to Group 17.
The element with atomic number 9 is Fluorine (F), which is also in Group 17.
The element with atomic number 35 is Bromine (Br), which is also in Group 17.
Therefore, the elements with atomic numbers 9 (Fluorine) and 35 (Bromine) will have physical and chemical properties resembling chlorine. Fluorine, Chlorine and Bromine are put in one group on the basis of their similar properties.
(b) Electronegativity decreases down a group. Fluorine (atomic number 9) is above chlorine in Group 17. Therefore, Fluorine is more electronegative than chlorine. The electronegativity value for fluorine is 4.0, while for chlorine it is 3.0.
2. First ionisation enthalpy of two elements X and Y are 500 kJ mol⁻¹ and 375 kJ mol⁻¹ respectively. Comment about their relative position in a group as well as in a period.
Answer: Ionisation energy decreases as one moves down a group. Ionisation energy tends to increase as one moves from left to right across a period.
Element Y has a lower ionisation enthalpy (375 kJ mol⁻¹) compared to element X (500 kJ mol⁻¹).
Relative position in a group:
If elements X and Y are in the same group, the element with the lower ionisation energy will be placed below the element with higher ionisation energy. Since Y (375 kJ mol⁻¹) has a lower ionisation energy than X (500 kJ mol⁻¹), Y would be below X in the same group. This is because ionisation energy decreases down a group due to an increase in atomic size which overcomes the effect of an increase in nuclear charge.
Relative position in a period:
If elements X and Y are in the same period, the element with the lower ionisation energy will be placed to the left of the element with higher ionisation energy. Since Y (375 kJ mol⁻¹) has a lower ionisation energy than X (500 kJ mol⁻¹), Y would be to the left of X in the same period. This is because ionisation energy tends to increase across a period from left to right due to a decrease in atomic size and an increase in nuclear charge.
3. The elements of one short period of the Periodic Table are given below in order from left to right: Li Be B C O F Ne
(a) To which period do these elements belong ?
(b) One element of this period is missing. Which is the missing element and where should it be placed ?
(c) Place the three elements fluorine, beryllium and nitrogen in the order of increasing electronegativity.
(d) Which one of the above elements belongs to the halogen series ?
Answer: (a) The elements Li, Be, B, C, O, F, Ne belong to the 2nd period of the periodic table.
(b) The elements of the 2nd period are Lithium (Li), Beryllium (Be), Boron (B), Carbon (C), Nitrogen (N), Oxygen (O), Fluorine (F), and Neon (Ne). Comparing this with the given sequence Li Be B C O F Ne, the missing element is Nitrogen (N). It should be placed between Carbon (C) and Oxygen (O).
(c) Electronegativity increases from left to right in a period. Beryllium (Be), Nitrogen (N), and Fluorine (F) are all in the 2nd period. Their order from left to right is Be, N, F. Therefore, the order of increasing electronegativity is Beryllium < Nitrogen < Fluorine. The electronegativity values are Be (1.5), N (3.0), and F (4.0).
(d) Group 17 elements are known as Halogens. From the given list of elements, Fluorine (F) belongs to Group 17 and thus belongs to the halogen series.
4. The atomic number of an element Z is 16. State
(a) the period to which it belongs.
(b) the number of valence electron(s) in the element.
(c) whether the element is a metal or a non-metal.
(d) State the formula of the compound between Z and hydrogen.
Answer: (a) An element with atomic number 16 has the electronic configuration 2, 8, 6. It has three orbits (shells). A period is determined by the number of shells. Therefore, element Z belongs to the third period.
(b) The electronic configuration of element Z (atomic number 16) is 2, 8, 6. The number of electrons present in the outermost shell are the valence electrons. Thus, element Z has 6 valence electrons.
(c) Non-metals usually have 5, 6 or 7 electrons in their outermost orbits. Since element Z has 6 electrons in its outermost orbit, it is a non-metal. Element 16 is Sulphur (S), which is a non-metal.
(d) Element Z has 6 valence electrons. Its valency is 8 – 6 = 2. Hydrogen has a valency of 1. The formula of the compound between Z and hydrogen will be H₂Z.
5. M is a metal and its oxide has the formula M₂O. This oxide when dissolved in water forms the corresponding hydroxide which is a good conductor of electricity. In the above context answer the following:
(a) State the number of electrons in the outermost shell of M.
(b) Name the group to which M belongs.
(c) Which element has more electronegativity, M or O?
Answer: (a) The formula of the oxide is M₂O. Oxygen has a valency of 2. For the compound to be M₂O, the metal M must have a valency of +1. Metals that form a positive ion carrying a single charge, like Na⁺, possess one valence electron. Therefore, M has 1 electron in its outermost shell.
(b) Elements that have one electron in their outermost shell and form strong alkalis with water belong to Group 1 (Alkali metals). Since the hydroxide of M is a good conductor of electricity, it is a strong alkali. Thus, M belongs to Group 1.
(c) Generally, metals show lower electronegativity as compared to non-metals. M is a metal, and O (Oxygen) is a non-metal. Therefore, Oxygen (O) has more electronegativity than M.
6. The metals of Group 2 from top to bottom are Be, Mg, Ca, Sr and Ba.
(a) Which of these elements will form ions most readily and why ?
(b) State the common feature in the electronic configuration of all these elements.
Answer: (a) Barium (Ba) will form ions most readily. Elements at the bottom of a group are most metallic. The atomic size is large, hence electrons are loosely held, and ions are readily formed. Ionisation energy decreases down the group, meaning less energy is required to remove an electron. Therefore, Barium, being at the bottom of Group 2 among the given elements, will form ions most readily.
(b) The common feature in the electronic configuration of all these Group 2 elements is that they have 2 electrons in their valence shell (outermost shell).
7. In the table below, H does not represent hydrogen. Some elements are given in their own symbol and position in the periodic table while others are shown with a letter.
Answer the following questions.
(i) Identify the most electronegative element.
(ii) Identify the most reactive element of group IA or 1.
(iii) Identify the element from period 3 with the smallest atomic size.
(iv) Identify the element with the highest ionization potential.
(v) How many valence electrons are present in O ?
(vi) Which element from group 2 will have the least ionization energy ?
(vii) Identify the noble gas (L) of the fourth period.
(viii) In the compound between A and H, what type of bond is formed? Give its molecular formula.
Answer: (i) Electronegativity increases from left to right across a period and decreases down a group. Element J is in Period 2, Group 17. Halogens (Group 17) have high electronegativity, and electronegativity is highest at the top of the group. Therefore, J is the most electronegative element among the choices shown.
(ii) In Group 1 (IA), reactivity increases on going down the group because the tendency to lose electrons increases. The elements in Group 1 shown are Li (Period 2), A (Period 3), and B (Period 4). Therefore, B is the most reactive element of Group 1.
(iii) In a period, the size of an atom decreases from left to right. The elements in Period 3 are A, Mg, E, Si, H, K. Element K is in Group 17 (furthest to the right, excluding any noble gas if it were present in that period in the table before group 18). Therefore, K has the smallest atomic size in Period 3 among the elements shown.
(iv) Ionisation potential tends to increase as one moves from left to right across a period and decreases down a group. Noble gases (Group 18) have the highest ionisation potentials in their respective periods. In the table, Ne is a noble gas in Period 2 and L is a noble gas in Period 4. Since ionisation potential decreases down a group, Ne will have a higher ionisation potential than L. Therefore, Ne has the highest ionization potential among the elements shown.
(v) Element O is in Period 2, Group 16 (VIA). Elements of Group 16 have 6 electrons in their valence shell. Therefore, O has 6 valence electrons.
(vi) Ionisation energy decreases down a group. The elements from Group 2 shown are D (Period 2), Mg (Period 3), and C (Period 4). Therefore, C, being at the bottom of these Group 2 elements, will have the least ionization energy.
(vii) L is in Period 4 and Group 18 (0). Group 18 elements are called noble gases or inert gases. Therefore, L is the noble gas of the fourth period.
(viii) Element A is in Period 3, Group 1 (IA), so it is an alkali metal and will form a cation A⁺. Element H is in Period 3, Group 16 (VIA), so it is a non-metal and will form an anion H²⁻ (by gaining 2 electrons to complete its octet, as it has 6 valence electrons). The bond formed between a metal (A) and a non-metal (H) will be an electrovalent (ionic) bond. To balance the charges, the molecular formula of the compound will be A₂H.
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