What are Dobereiner’s triads? Name two such triads.
Dobereiner arranged certain elements with similar properties in groups of three in such a way that the atomic mass of the middle element was nearly the same as the average atomic masses of the first and third elements.
For example:
Triad: lithium sodium potassium
Atomic mass: 7 23 39
Atomic mass of Na =(39+7)/2= 23
Triad: Chlorine Bromine Iodine
Atomic mass: 35.5 80 127
Atomic mass of Br =127 + 35.5/2 = 81.25
Give reasons:
(i) IE1 of sodium is lower than that of magnesium whereas IE2 of sodium is higher than that of magnesium.
(ii) Noble gases have positive value of electron gain enthalpy.
All transition elements are d-block elements, but all d-block elements are not transition elements. Explain.
What do you understand by ‘Representative elements’? Name the groups whose elements are called representative elements.
Consider the following species:
N3-, O2-, F–, Na+, Mg2+, Al3+
(a) What is common in them?
(b) Arrange them in order of increasing ionic radii?
What are the various factors due to which the ionization enthalpy of the main group elements tends to decrease down the group?
Explain the deviation in ionization enthalpy of some elements from the general trend by using the given figure.
Which of the above elements is likely to be:
(a) the least reactive element (b) the most reactive metal
(c) the most reactive non-metal (d) the least reactive non-metal
(e) the metal which can form a stable binary halide of the formula MX2(X = halogen)
(f) the metal which can form a predominantly stable covalent halide of the formula MX (X = halogen)?
An element belongs to 3rd period and group-13 of the periodic table. Which of the following properties will be shown by the element?
(a) Good conductor of electricity
(b) Liquid, metallic
(c) Solid, metallic
(d) Solid, non metallic
Use periodic table to answer the following questions:
(a) Identify the element with five electrons in the outer subshell.
(b) Identify the element that would tend to lose two electrons.
(c) Identify the element that would tend to gain two electrons.
Discuss and compare the trend in ionization enthalpy of the elements of group 1 with those of group 17 elements.
Discuss the main features of long form of the periodic table. What are the advantages of long . form of periodic table?
Which of the following elements can show covalency greater than 4?
(a) Be (b) P (c) S (d) B
First member of each group of representative elements (i.e., s and p-block elements) shows anomalous behaviour. Illustrate with two examples.
Which important property did Mendeleev use to classify the elements in this periodic table and did he stick to that?
Which of the following pairs of elements would have a move negative electron gain enthalpy? (i) O or F (ii) F or Cl.
In which of the following options order of arrangement does not agree with the variation of property indicated against it?
(a) Al3+ < Mg2+ < Na+ < F– (increasing ionic size)
(b) B < C < N < O (increasing first ionization enthalpy)
(c) I < Br < Cl < F (increasing electron gain enthalpy)
(d) Li < Na < K < Rb (increasing metallic radius)
Nitrogen has positive electron gain enthalpy whereas oxygen has negative. However, oxygen has lower ionization enthalpy than nitrogen explain.
The first ionization enthalpy values (in kJ mol -1) of group 13 elements are:
B Al Ga In Tl
801 577 579 558 589
How would you explain this deviation from the general trend?
What is basic difference between the terms electron gain enthalpy and electro negativity?
Discuss the factors that influence the magnitude of ionization enthalpy. What are the general trends of variation of ionization enthalpy in the periodic table? Explain.
Illustrate by taking examples of transition elements and non-transition elements that oxidation states of elements are largely based on electronic configuration.
What do you understand by exothermic reaction and endothermic reaction? Give one example of each type.
Explain the following:
(a) Electronegativity of elements increases on moving from left to right in the periodic table.
(b) Ionisation enthalpy decreases in a group from top to bottom.
Match the correct atomic radius with the element.
| Column I (Element) | Column II (Atomic radius (pm) |
| Be | 74 |
| C | 88 |
| 0 | 111 |
| B | 77 |
| N | 66 |
Write down the outermost electronic configuration of alkali metals. How will you justify their placement in group 1 of the periodic table?
Would you expect the second electron gain enthalpy of O as positive, more negative or less negative than the first? Justify your answer.
In the modem periodic table, the period indicates the value of
(a)atomic number (b) mass number (c) principal quantum number (d) azimuthal quantum number?
Discuss the main characteristics of four blocks of elements in the periodic table? Give their general electronic configuration.
The first ionisation enthalpy of magnesium is higher than that of sodium. On the other hand, the second ionisation enthalpy of sodium is very much higher than that of magnesium. Explain.
The formation of the oxide ion, 02-(g), from oxygen atom requires first an exothermic and then an endothermic step as shown below:
O(g) + e–→0– (g), ∆H= -141 kJ mol-1
0–(g) + e–→O2 (g), ∆H = +780 kJ mol-1
Thus process of formation of O2- ion in gas phase is unfavourable even though O2- is isoelectronic with neon. It is due to the fact that
(a) Oxygen is more electronegative.
(b) Addition of electron in oxygen results in larger size of the ion.
(c) Electron repulsion outweighs the stability gained by achieving noble gas configuration.
(d) 0– ion has comparatively smaller size than oxygen atom.
Ionisation enthalpies of elements of second period are given below:
Ionisation enthalpy/kJ mol-1: 520, 899, 801, 1086, 1402, 1314, 1681, 2080. Match the correct enthalpy with the elements and complete the graph given in figure. Also write symbols of elements with their atomic number.
p-Block elements form acidic, basic and amphoteric oxides. Explain each property by giving two examples and also write the reactions of these oxides with water.
Write the atomic number of the element present in the third period and seventeenth group of the periodic table.
Energy of an electron in the ground state of the hydrogen atom is- 2.18 x 10-18 J.Calculate the ionization enthalpy of atomic hydrogen in terms of JMol-1.[Hint: Apply the idea of mole concept to derive the answer],
How would you explain the fact that the first ionization enthalpy of sodium is lower than that of magnesium but its second ionization enthalpy is higher than that of magnesium?
How would you react to the statement that the electronegativity ofN on Pauling scale is 3.0 in all the nitrogen compounds?
Would you expect the first ionization enthalpies of two isotopes of the same element to be the same or different? Justify your answer.
