Chemistry Class 9
Chapter 3 Important Question Answers 9th Class Chemistry
Periodic Table and Periodicity of Properties
Q: How do you find the Group number of the periodic table elements?
Trick to find Periodic Table Elements Group number
Group Number = Number of valence electron
Q: What is the electron arrangement of an element with an atomic number of 6?
Write down the electronic configuration of that element.
Ans: Electronic configuration of an element with atomic no. 6 is Carbon (C) .
With a total of six electrons, carbon is the sixth element. The first two electrons will be placed in the 1s orbital when writing the electron configuration for carbon. The following two electrons for C are placed in the 2s orbital because 1s can only accommodate two electrons. The 2p orbital will hold the final two electrons. The carbon (C) electron configuration will therefore be 1s22s22p2.
Q: How to calculate the period number and block of the periodic table elements?
The super easiest trick is to calculate the period number of an element.
You can determine an element’s period number if you are provided its atomic number.
A substance with the atomic number 6
The atomic number 6 substance is a carbon (C) and its electron configuration is 1s22s22p2.
Block of Elements
Period & Block of Sodium (Na)
Electronic Configuration of Na
What is meant by valence electrons?
The electrons present in the outermost shell of an atom are known as valence electrons.
Importance of valence electrons
1. Valence electrons decide the reactivity of an element.
2. Valence electrons decide the type of chemical bond an atom will form with another atom.
What is the purpose of valence electrons?
The valence electrons in an atom determine its reactivity.
Q.What is periodic table?
A periodic table is a table in which the elements are arranged according to their rising atomic number and recurrent chemical characteristics. They are arranged in a tabular format with rows denoting periods and columns denoting groups.
In the periodic tables the horizontal rows are known as periods. The total number of periods in periodic table are 7. There are 2 elements in period 1 and 32 in period 6, respectively. The first three periods are known as short periods, while the rest are known as long periods. Moving from left to right in a period causes the characteristics of its constituents to progressively alter. However, the pattern of attributes within a period repeats as you travel from one to the next. The periodic law is followed in this situation.
The vertical rows in the periodic table are called groups. The total number of groups in the periodic table is 18. Elements that share a similar valence shell electronic configuration are grouped together. A number and the letter A or B are used to identify each group. Normal or representative elements are the names given to Group A elements. They are sometimes referred to as primary group components. Transition elements are Group B elements.
Q.Write a note on the shape of the periodic table.
illustrate the shape of the periodic table using groups and periods.
SHAPE OF PERIODIC TABLE
The order of the elements is based on their increasing atomic number. Only two elements, H and He, are present in the first period. They both have valence electrons in the K shell. K shells are limited to two electrons. Lithium (Li), atomic number 3, has one electron in the L shell, hence the second period begins with Li. As the K shell is finished at He, the period also finishes at He. Since the L shell can hold eight electrons, the second period contains eight elements. Second period terminates at Ne, which possesses eight electrons(2s2,2p6).
The next element is Na [sodium] whose valence electron is present in the M shell.In Na valence electron is present in 3s sub-shell,which has a similar electronic configuration to that of Li (2s1). So it is located under Li.
Mg with 3s2 valence shell electronic configuration is present below Be (2s2).In the same way the upcoming elements i.e Al,Si,P,S,Cl and Ar on the basis of similarity in valence shell electronic configuration are present under B, C, N,O,F and Ne respectively.Ar has 3s2,3p6 valence shell electronic configuration that is similar to the valence shell electronic configuration of Ne (2s2,2p6).
Next element i.e K has 4s1 valence shell electronic configuration, which is comparable to Na (3s1).So K is located under Na and a new period is (4th) started with K. In this manner, elements with comparable valance shell configurations are grouped.
The way the elements are arranged into periods has a significant impact. In the periodic table, elements with related properties are grouped.
Q. Write a note on s and p block elements?
s block elements
s-block elements are those found in Groups 1 and 2 of the modern periodic table.
Two different kinds of s-block elements can exist: those with one electron (s1) and those with two electrons (s2) in their inner s sub-shell.
s-block Is made of 14 elements that are hydrogen (H), lithium (Li), helium (He), sodium (Na), beryllium (Be), potassium (K), magnesium (Mg), rubidium (Rb), calcium (Ca), cesium (Cs), strontium (Sr), francium (Fr), barium (Ba), and radium (Ra).
p block elements
Group IIIA to VIIIA elements (excluding He) are referred to as p-block elements. Due to the location of their valence electrons in the p subshell.
Some p-block elements are Carbon, silicon, phosphorus, sulphur, boron, germanium, tin, arsenic etc.
Q. Discuss the placement of different blocks of the elements in the periodic table.
Blocks in Periodic Table
The periodic table is divided into four blocks:
Group IA and IIA on the left side of the periodic table, form the s block. The final electron of the elements of these groups enter into the s-orbital hence they are called s-block elements.
Group IIIA to VIIIA(except helium) on the right side of the periodic table forms the p-block. The final electron of the elements of these groups enter into the p-orbital hence they are called p-block elements.
Group III B to group VIIIB and group I B and II B form d-block of the periodic table.
Lanthanides and Actinides series located at the bottom of the periodic table constitute the f-block.
Q. Enlist the names of different families (groups) in the periodic table.
There are families, which are groups of elements having related properties, on the periodic table.
There are 8 groups in the periodic table and their names are as follows
Group I A Alkali metals
Group II A Alkaline earth metals
Group III A Boron family
Group IV A Carbon family
Group V A Nitrogen family
Group VI A Oxygen family
Group VII A Halogen family
Group VIII A Noble or inert gases
Q.What do you mean by representative elements and transition elements?
The periodic table’s groups 1 and 2, as well as groups 13 to 18, contain the representative elements.
In other words, the terms “representative elements” collectively refer to both s-block and p-block elements.
The term “representative elements” refers to elements with complete inner shells but incomplete outer shells, i.e., substances with fewer than 8 electrons in the outermost shell. Except for inert gas, the s and p – block elements are known as representative elements.
Transition elements also known as transition metals
Transition elements are of two types
1. Inner transition elements
2. Outer transition elements
Inner transition elements
Inner transition elements are those where the final electron enters the f-orbital and are called inner transition elements.
Outer transition elements
Outer transition elements are those where the final electron enters the d-orbital and are called outer transition elements.
Q. What is the difference between transition metals and representative metals?
Representative elements are chemical elements in groups 1, 2, and 13–18, whereas transition elements are chemical elements in groups 3–12, including Lanthanides and Actinides. This is the main difference between representative elements and transition elements.
Q. What is meant by the term periodicity of properties?
Periodicity Of Properties
According to the fundamental law regulating the contemporary periodic table, an element’s properties are periodic functions of its atomic number. These characteristics repeat frequently or follow a specific pattern frequently. The periodicity of properties is the name given to this phenomenon.
Due to the repetition of comparable electronic configurations with the same amount of electrons in the outermost orbit, periodic characteristics of elements develop. The number of valence electrons in a given group stays constant. On the other hand, when we move from left to right throughout a period, the number of valence electrons rises. The quantity of electrons in the valence shell determines an element’s chemical properties.
Periodicity of properties examples
- Shielding effect
- Ionization energy
- Electron affinity
- Atomic size
Q. Explain the phenomena of the shielding effect along with its periodicity.
The shielding effect refers to the decrease in the force of attraction between the nucleus and the valence electrons caused by the electrons in the inner subshells.
When there are more inner electrons, they shield the outermost electron from the nucleus, allowing it to neglect the nuclear pull. This is known as the shielding or screening effect. In a multielectron atom, the electrons in the valence shell are drawn toward the nucleus but are repelled by the electrons in the inner shells. The real force of attraction between the nucleus and the valence electrons is significantly diminished as a result of the repulsive forces acting in opposite directions. Due to the presence of electrons in the inner shells, a phenomenon known as the screening effect or shielding effect causes the nucleus’s force of attraction on valence electrons to weaken.
• Trends of shielding effect concerning groups
As we go from top to bottom in a group the influence of the shielding effect increases because the number of the shell of elements increases, as well as the number of inner shell electrons, also increases.
• Trends of shielding effect concerning periods
As we go from left to right in a periodic table the number of shells remains the same. Therefore the shielding effect remains unchanged or constant.
What is the atomic size? How do atomic radii differ across the periodic table?
The distance between an atom’s nucleus and valence shell is known as its atomic size.It is also called atomic radii.
Variation in atomic radii in period
In any particular period, the atomic radius decreases as you move through it. This is due to the fact that, within a period, you shift from one element to the next to its right. In the same valence shell, an additional electron is added. The nucleus’s positive charge likewise rises by 1 at the same moment. The nucleus’s attraction to the electron in the valence shell grows. As a result, the atomic radius and shell size both decrease. For instance, the atomic size decreases from lithium to beryllium. This is seen from the electronic arrangement of Li (2s1) and Be valence shell (2s2). The atomic number grows from 3 to 4 while the shell number n remains unchanged while traveling from Li to Be. As a result, the nucleus’s force on the electron in the valence shell increases. So, atomic radius or atomic size decreases.
Variation in atomic radii in group
Moving down the group of elements in any particular major group causes the atomic radius to increase. This is so because an atom’s valence shell size determines how big it is. With each subsequent lower element in the group, the atom gains an extra electron shell. This increases the atomic radius. For example, the atomic radius increases when moving from Li to Na.
Q. Define ionization energy. Draw attention to the variables that affect the ionization energies of elements in groups and periods.
The least amount of energy necessary to remove the outermost electron from a gaseous atom in isolation is known as “ionization energy.”
Ionization Energy General Representation
M(g) + ionization energy⎯⎯→M (g) + e
Ionization Energy Example
Ionization Energy Unit
The unit of ionization energy is Kj/Mol.
The extent to which the nucleus attracts the outermost electron is measured by the ionization energy. The stronger attraction between the nucleus and the outermost electron is indicated by a high value of ionization energy. The nucleus and the outermost electron are more weakly attracted to one another when the ionization energy is low.
Trends in the values of ionization energy
1. In Groups
In a group, the ionization energy value decreases from top to bottom. This is due to the fact that as you descend in the group, the shielding effect in atoms gets stronger. The nucleus’s attraction to the valence electrons is weaker as a result of the greater shielding effect. Valence electrons are therefore simpler to remove. This causes a decrease in ionization energy in a group from top to bottom.
For example, the ionization energy of Li is greater than that of Na.
Ionization energy of Li= 520 KJ/mol
Ionization energy of Na= 496 KJ/mol
2. In Periods
The shielding effect is consistent as you move in a period from left to right. However, the nuclear charge increases. The nucleus’s and the valence electron’s attraction grows stronger. So, Ionization energy increases when moving from left to right in a period.
For example, the ionization energy of Be is less than B.
Ionization energy of C = I086 Kj/mol
Ionization energy of N = 1402 Kj/mol
What are the different factors that affect ionization energy?
Q. Define electron affinity and discuss its variation in groups and periods.
The amount of energy released when an electron adds up in the valence shell of an isolated atom to produce an uninegative gaseous ion is described as electron affinity. The formation of anion is explained by electron affinity.
Electron Affinity General Representation
X(g) + e- ⎯⎯→ X- (g) + electron affinity
The energy change that occurs when an electron is added to a gaseous atom is known as its electron affinity. For example, the associated energy change is -328 kJ/mol when a fluorine atom gains an electron to create F – (g) in the gaseous state.
Electron Affinity Unit
Variation of electron affinity in the periodic table
1) Variation of electron affinity in a period
The electron affinity generally increases across a period as you move from left to right. This is because the nuclear charge increases and the atomic radius decreases, which attaches the extra electron more tightly to the nucleus. However, the shielding effect is constant throughout the period. As a result, in each period, halogens have the highest electron affinities and alkali metals the lowest.
2) Variation of electron affinity in a group
In a group, the electron affinity decreases from top to bottom. This is because the shielding effect has increased. Because of the increased shielding effect, the additional electron binds less tightly to the nucleus. As a result, less energy is released.
Q.What is electronegativity? Briefly explain its variation in groups and periods.
An atom’s ability to attract electrons to itself in a chemical bond is known as electronegativity.
Electronegativity is a chemical property that describes an atom’s ability to attract electrons to itself.
An atom’s electronegativity is influenced by both its atomic number and the separation of its valence electrons from its charged nucleus.
Variation of electronegativity in the periodic table
1) Variation in electronegativity in a group
The electronegativity of an atom decreases when we move from top to bottom in a group this is because the atomic size increases while moving down the group and as a result shielding effect increases. So, the tendency of an atom to attract electrons decreases. E . g Electronegativity of Oxygen(3.5) is greater than that of Sulphur (2.5)
2) Variation in electronegativity in a period
The electronegativity of an atom increases when we move from left to right in a period because the nuclear charges increase from left to right while the electrons are added in the same shell and these electrons can’t shield each other. So, the electrons are attracted toward the nucleus more effectively. E.g the electronegativity of Oxygen(3.5) is less than that of fluorine(4.0).
Inner shell electrons cut off the force of attraction between the nucleus and the valence electrons.
The greater the number of inner shell electrons,the greater the shielding or screening effect.
Q. Which atom has a greater shielding effect, Li or Na ?
The valence shell electron of Na11 experiences less attraction from the nucleus due to the presence of 10 inner-shell electrons as compared to Li3 which has 2 inner-shell electrons. So,Na will have a greater shielding effect due to a greater number of inner shell electrons as compared to Li 3.
Q. Which atom has higher ionization energy? Li or Be.
Ionization energy increases across a period. The element that has smaller ionization energy will be further to the left in the periodic table. Therefore, Be has higher ionization energy as compared to Li .
Q. What do you mean by nuclear charge?
The nuclear charge (Z) is defined as the total charge in the nucleus due to the presence of positive charge protons.
As we move through the periodic table, the atomic number continuously increases due to the increase in the number of protons in the nucleus, thus the nuclear charge also increases.
Q. How to determine the effective nuclear charge of an electron?
Effective Nuclear Charge ( Zeff):
Effective Nuclear Charge is defined as” the net positive charge experienced by an electron in a poly electronic atom”.
Effective Nuclear Charge Calculation:
The effective nuclear charge calculated for such an electron is given by the following equation:
Zeff = Z- S
Z = Atomic number
S= shielding constant
Q. Differentiate between nuclear charge and effective nuclear charge?
Period 3 Elements
As we move across the period from left to right, sodium (Na) has only one electron while argon(Ar) has eight valence electrons.
Across the period, the atomic number increases due to the increase in the number of protons in the nucleus. Therefore, nuclear charge also increases.
Also on the addition of proton numbers to the nucleus, the electrons feel greater nuclear attraction because the inner shell electron number remains constant. Therefore, experience the same inner electron repulsion and increase nuclear attraction across the period.
Which atom has greater nuclear charge and effective nuclear charge, sodium or chlorine?
As the sodium atom has 11 protons and the chlorine atom has 17 protons so due to the greater number of protons chlorine has a greater nuclear charge.
Also, an effective nuclear charge of chlorine is greater than sodium because both atoms having the same number of inner electrons experience the same inner electron repulsion.
Greater the number of inner shell electrons, the greater the shielding or screening effect results in the decrease of effective nuclear charge.
Activities & Examples
Look at the periodic table and write the number of elements present in the relevant period in the table.
Table 3.1: Number of elements in the periods of the periodic table
Period No. No. of elements
Seventh 32 ( Seventh period is still incomplete)
Example 3.1: Identifying the group and period of an element
Identify the group and period of 27Al13, 9B5 , 24Mg12 based on electronic configuration.
Write the electronic configuration of the element. Identify its valence shell. Remember that the n value of the valence shell indicates a period. The total number of electrons in the valence shells represents the group number.
a) Al13= 1s2,2s2,2p6,3s2,3p1 K=1, L=2, M=3
Valence shell is M
As n = 3, Al is present in the 3rd period. Since total number of electrons in the
valence sub-shells are 2+1=3, it must be present in Group IIIA.
b) B 5 = 1s2,2s2,2p1
Valence shell is L
So n = 2, B is present in the 2nd period. Since total number of electrons in the
valence shell are 2+1=3, it must be present in Group IIIA.
c) Mg 12 = 1s2,2s2,2p6,3s2
K =1 , L=2, M= 3
Valence shell is M
So n = 3, Mg is present in the 3rd period. Since the total number of electrons in the
valence shell are 2, it must be present in Group IIA.
Example 3.2: Classifying or dividing elements into groups and periods
The electronic configuration of atoms of some elements are given below. Classify them in
groups and periods.
1. The elements whose atoms have similar valence shell electronic configuration belong
to the same group.
2. The n value of the valence shell indicates the period.
3. The elements whose atoms have the same value of n for the valence shell lie in the same
|Period||II A||V A||VII A|
Example 3.3: Obtaining the valence shell configuration
Write the valence shell electronic configuration of the following elements from their position in the periodic table.
(a) Phosphorus (b) Neon
Period number = n value of valence shell
Group number = number of valence electrons
Distribute the electron in the sub-shells of valence shell.
a) Period number of phosphorus is 3,
As n = 3 therefore, valence shell is M
So valence electrons will be present in 3s and 3p sub-shells
The group number is 5, so there are 5 electrons in the valence shell
Two electrons will fill 3s sub-shell and remaining 3p sub-shell. Thus, the valence shell
electronic configuration is 3s2, 3p3
b) Period number of Ne is 2. So, n = 2 and valence shell is L. Valence electrons will be
present in 2s and 2p sub-shells.
Group number for Ne is 8,
This means there are 8 electrons in the valence shell. Two electrons will fill 2s sub-shell
and the remaining six 2p sub-shell. Thus the valence shell electronic configuration for Ne is 2s2,2p6
Example 3.4: Obtaining the position of element in the periodic table from electronic Configuration
Find out the position of the following elements in the periodic table from the electronic configuration
(a) Nitrogen (atomic number: 7) (b) Oxygen (atomic number: 8)
Problem Solving Strategy:
Write electronic configuration of the element. Identify the valence shell configuration, coefficient
of s or p sub-shell represents period number and total number of electrons in valence
shell is equal to the group number.
a) Electronic configuration of N =1s2,2s2,2p3
Valence shell has configuration= 2s2,2p3
Period number = 2
Group number = 2 + 3=5
So,Nitrogen is present in the 2nd period of Group V-A
b) Electronic configuration of oxygen =1s2,2s2,2p4
Valence shell has configuration =2s2,2p4
So, Period number = 2
Group number = 2 + 4 = 6
So,Oxygen is present in the 2nd period of Group VI A
Identifying the element whose atoms have greater shielding effect, using periodic table.
Choose the elements whose atoms you expect to have greater shielding effect.
(a) Be or Mg (b) C or Si
Problem Solving Strategy:
Look at the periodic table and find the relative position of given elements in the periodic
table. Apply the trend of increasing shielding effect in a group.
(a) Mg atoms will have greater shielding effect.
(b) Si atoms will have greater shielding effect.
Example 3.6: Identifying the element that has greater atomic radius
Choose the elements whose atom you expect to have larger atomic radius in each of the following pairs.
(a) Mg, Al (b) C, Si
Problem Solving Strategy:
Remember that the larger atom in any:
(a) Period lies further to the left in the periodic table.
(b) Group lies closer to the bottom in the periodic table.
(c) Check the periodic table and choose the element
(a) The larger atom is Mg
(b) The larger atom is Si
Example 3.7: Identifying the element that has smaller ionization energy
Choose the element whose atom you expect to have smaller ionization energy in each of
the following pairs.
(a) B,C (b) N, P
Problem Solving Strategy
Remember that ionization energy:
(a) Increases across a period. The element that has smaller ionization energy will be further
to the left in the periodic table.
(b) Decreases from top to bottom in a group. The element that has smaller ionization energy
will correspond to the element closer to the bottom.
(c) Check the periodic table to choose the element.
(a) The atom with the smaller ionization energy is B
(b) The atom with the smaller ionization energy is P.
Society, Technology, Science
In 1864, John Newland, an English chemist arranged 24 elements in order of increasing atomic masses. He noticed that every eighth element, starting from any point, has similar properties. A few rows of his arrangement are shown below
H Li Be B C N O
F Na Mg Al Si P S
Cl K Ca Cr Ti Mn Fe
His scheme however, failed because many elements were found out of place in his arrangement.
For instance, Ti does not resemble C and Si, Mn does not resemble N and P and Fe does not resemble O and S. However his arrangement of elements in order of increasing atomic masses formed the basis for the later classification of elements.
In 1869, Mendeleev, a Russian chemist developed a classification scheme of elements. He recognized that if elements were placed in order of increasing atomic masses, the properties of elements repeated at regular intervals. He arranged 65 elements in periods and groups. The development of the periodic table nicely explains how a scientist can build on one another’s work.
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