9th Class Chemistry Chapter No. 3 Periodic Table and Periodicity of Properties Notes
Notes on Periodic Table and Periodicity of Properties
Historical Background
- In the 19th century, chemists attempted to arrange elements systematically.
- Discovery of the Periodic Law led to the creation of the Periodic Table.
- The table predicted properties of undiscovered elements.
Dobereiner’s Triads
- Dobereiner observed triads of elements with similar properties and atomic mass.
- Triads consisted of three elements, and the central element had an average atomic mass of the other two.
- This classification had limited acceptance.
Long Answer Questions Notes
Newlands Octaves
- Newlands proposed the “law of octaves,” where every eighth element showed similar properties.
- His work lacked recognition and did not account for undiscovered elements and noble gases.
Mendeleev’s Periodic Table
- Mendeleev arranged 63 known elements in increasing atomic mass.
- Elements with similar properties were placed in the same vertical columns (groups).
- This table was based on the “periodic law” stating that properties of elements are periodic functions of their atomic masses.
Limitations of Mendeleev’s Table
- Couldn’t explain isotopes or the wrong order of atomic masses for some elements.
- Atomic mass alone couldn’t serve as a basis for element arrangement.
Short Answer Questions Notes
Modern Periodic Table
- Moseley’s discovery of atomic number in 1913 led to a change in the periodic law.
- Elements were arranged by increasing atomic number instead of atomic mass.
- This arrangement led to periodicity in properties based on electronic configurations.
Structure of the Modern Periodic Table
- Elements are arranged in periods (horizontal rows) and groups (vertical columns).
- Periods have continuously increasing atomic numbers and changing electronic configurations.
- Elements in a group have similar electronic configurations and chemical properties.
MCQ’s Notes
Long Form of Periodic Table
- Consists of seven periods, with the 1st having two elements, the 2nd and 3rd having 8 elements each, and so on.
- Groups are numbered 1 to 18 from left to right, and elements in a group have similar properties.
- Elements are classified into four blocks based on the subshell that receives the last electron.
Periodicity of Properties
Atomic Size and Atomic Radius
- Atomic size is difficult to measure due to the small size of atoms.
- The common method is to assume atoms as spheres and measure half the distance between their nuclei (atomic radius).
- Atomic size decreases from left to right in a period due to increased effective nuclear charge and constant valence shell size.
- Atomic size increases from top to bottom in a group as new shells are added, reducing the effective nuclear charge.
Shielding Effect
- Electrons in inner shells shield the outermost electrons from the full nuclear charge, reducing the effective nuclear charge (Zeff).
- Shielding effect increases down a group and decreases across a period.
Ionization Energy
- Ionization energy is the energy required to remove the most loosely bound electron from an isolated gaseous atom.
- Ionization energy increases from left to right in a period due to smaller atomic size and stronger nuclear attraction.
- Ionization energy decreases from top to bottom in a group due to increased shielding effect and larger atomic size.
Electron Affinity
- Electron affinity is the energy released when an electron is added to the outermost shell of an isolated gaseous atom.
- Electron affinity trends are not as consistent as other properties in the periodic table.
Electronegativity
Electronegativity is a measure of an atom’s ability to attract the shared pair of electrons towards itself in a chemical bond, especially in a covalent bond. It is an important property that helps to understand the nature of chemical bonding and the distribution of electrons in molecules.
Trend of Electronegativity
- Across a Period (Left to Right)
Electronegativity generally increases from left to right in a period of the periodic table. This is because, as we move from left to right, the atomic number (Z) increases, leading to a greater effective nuclear charge (Zeff). The higher effective nuclear charge pulls the shared pair of electrons closer to the nucleus, making the atom more electronegative. Fluorine (F) has the highest electronegativity among all elements. - Down a Group
Electronegativity generally decreases from top to bottom in a group of the periodic table. This is because, as we move down a group, the atomic size increases due to the addition of more electron shells. The increased atomic size reduces the effective nuclear charge experienced by the valence electrons, resulting in weaker attraction for the shared pair of electrons and lower electronegativity.
Example:
Electronegativity is an essential concept in understanding chemical bonding and the polarity of molecules. It helps to predict the nature of chemical bonds, whether they are covalent, polar covalent, or ionic. The Pauling scale is commonly used to express electronegativity values, and the higher the value, the more electronegative an element is.
Answers of Test Questions
i. Atomic Radius:
Atomic radius refers to the size of an atom, which is the distance from the nucleus to the outermost electron in an atom. Since atoms do not have well-defined boundaries, the atomic radius is typically measured as half the distance between the nuclei of two bonded atoms of the same element.
ii. SI Units of Atomic Radius:
The SI unit of atomic radius is picometer (pm), which is equal to 1 × 10^-12 meters. It is commonly used due to the extremely small size of atoms.
iii. Why the Size of Atoms Decreases in a Period:
The size of atoms decreases in a period (from left to right) of the periodic table because the atomic number increases, leading to more protons in the nucleus. The increased positive charge in the nucleus exerts a stronger force of attraction on the electrons in the same energy level (valence shell), pulling them closer to the nucleus. As a result, the atomic size reduces across a period.
iv. Ionization Energy:
Ionization energy is the amount of energy required to remove the most loosely bound electron from an isolated gaseous atom to form a positive ion. It is often measured in kilojoules per mole (kJ/mol).
v. Why the 2nd Ionization Energy of an Element is Higher than the First One:
The second ionization energy of an element is higher than the first one because, after losing the first valence electron, the remaining electrons experience a stronger effective nuclear charge (Zeff). The loss of the first electron leaves a positive charge on the ion, causing the remaining electrons to be more strongly attracted to the positively charged nucleus, making it more difficult to remove additional electrons.
vi. Trend of Ionization Energy in a Group:
The ionization energy generally decreases from top to bottom in a group of the periodic table. This is due to the increase in atomic size as we move down the group. The outermost electrons are farther from the nucleus, and the effective nuclear charge experienced by them decreases, making it easier to remove them.
vii. Why the Ionization Energy of Sodium is Less than that of Magnesium:
The ionization energy of sodium is less than that of magnesium because sodium has one valence electron in its 3s orbital, while magnesium has two valence electrons in its 3s orbital. The extra electron in magnesium experiences a stronger effective nuclear charge compared to the single valence electron in sodium, making it more difficult to remove the second electron in magnesium, resulting in a higher ionization energy.
viii. Why is it Difficult to Remove an Electron from Halogens:
Halogens are located in Group 17 of the periodic table and have seven valence electrons in their outermost energy level (7 valence electrons in the p orbital). These elements require only one additional electron to achieve a stable noble gas electron configuration. As a result, halogens have high electron affinity, meaning they strongly attract and hold onto additional electrons. Therefore, it is difficult to remove an electron from halogens.
ix. Shielding Effect:
The shielding effect is the reduction in the attractive force between the positively charged nucleus and the valence electrons in an atom due to the presence of inner-shell electrons. Inner-shell electrons act as a shield, repelling and partially offsetting the attractive force from the protons in the nucleus. As a result, the outermost electrons experience a reduced effective nuclear charge.
x. How Does Shielding Effect Decrease the Forces of Electrostatic Attractions between Nucleus and Outermost Electrons:
The shielding effect decreases the forces of electrostatic attractions between the nucleus and outermost electrons because the inner-shell electrons create a repulsive force that opposes the attractive force from the protons in the nucleus. This repulsive force weakens the net force experienced by the outermost electrons, making them less tightly held to the nucleus.
xi. Why Do Bigger Size Atoms Have More Shielding Effect:
Bigger size atoms have more shielding effect because they have additional electron shells between the valence electrons and the nucleus. The presence of more inner-shell electrons increases the repulsive force acting on the valence electrons, effectively reducing the net attractive force from the protons in the nucleus. As a result, the larger atoms experience a greater shielding effect.
xii. Why Does the Trend of Electron Affinity and Electronegativity is the Same in a Period:
The trend of electron affinity and electronegativity is the same in a period because both properties are influenced by the atomic size and effective nuclear charge. As we move from left to right in a period, the atomic size decreases, and the effective nuclear charge increases, leading to higher electron affinity and electronegativity values.
xiii. Which Element has the Highest Electronegativity:
Fluorine (F) has the highest electronegativity among all elements. It is the most electronegative element with the highest ability to attract the shared pair of electrons in a chemical bond. The electronegativity of fluorine is 3.98 on the Pauling scale.