LEC (6-8)

Question 1

1. How did John Newlands contribute to the development of the periodic table, and what was its limitation?

Answer 1

• Answer: Newlands arranged elements by atomic mass and observed that every 8th element had similar properties, a pattern called the “Law of Octaves.” However, the trend didn’t work past calcium, which limited its accuracy.

Question 2

2. How did Mendeleev’s arrangement of the periodic table differ from Newlands’ model?

Answer 2

• Answer: Mendeleev also arranged elements by atomic mass but left gaps for undiscovered elements. He predicted the properties of these elements, such as germanium, which were later confirmed. This made his table more successful than Newlands’.

Question 3

3. What did Henry Moseley discover, and how did it impact the organization of the periodic table?

Answer 3

• Answer: Moseley discovered that the atomic number, not atomic mass, determines an element’s properties. This led to the modern organization of the periodic table by atomic number, resolving inconsistencies found in Mendeleev’s table.

Question 4

4. What is effective nuclear charge (Zeff), and how is it calculated?

Answer 4

• Answer: Zeff is the net positive charge experienced by an electron in the valence shell. It is calculated as Zeff = Z - S , where Z is the atomic number and S is the shielding constant, which accounts for the electron repulsion from inner shells.

Question 5

5. How does shielding affect the atomic size across a period and down a group?

Answer 5

• Answer: Shielding reduces the effective nuclear charge experienced by outer electrons. Across a period, shielding is less effective, so atomic size decreases as Zeff increases. Down a group, increased shielding from additional electron shells causes atomic size to increase.

Question 6

6. What is Slater’s Rule, and how is it applied to calculate shielding?

Answer 6

• Answer: Slater’s Rule provides a method to calculate the shielding constant, S , for electrons. Electrons in the same shell contribute 0.35 each, electrons in n-1 shells contribute 0.85, and electrons in n-2 shells or lower contribute 1.0.

Question 7

7. How do alkali metals react with water, and what trend can be observed as you move down the group?

Answer 7

• Answer: Alkali metals react with water to produce metal hydroxides and hydrogen gas. As you move down the group, their reactivity increases due to lower ionization energies, making it easier for them to lose electrons.

Question 8

8. Why does the reaction of alkali metals with oxygen produce different types of oxides, and how do these oxides change across the group?

Answer 8

• Answer: Alkali metals form oxides when they react with oxygen, and the specific type of oxide depends on the metal’s position in the group. As the metals become larger and more reactive down the group, the oxides become more basic.

Question 9

9. Explain the periodic trend in ionization energy across a period and down a group.

Answer 9

• Answer: Ionization energy increases across a period due to increasing Zeff, which holds electrons more tightly. It decreases down a group as the increased atomic radius and shielding make it easier to remove an electron.

Question 10

10. What is electron affinity, and how does it change across the periodic table?

Answer 10

• Answer: Electron affinity is the energy change when an atom gains an electron. It generally becomes more negative across a period due to increasing Zeff and less negative down a group as electrons are added farther from the nucleus.

Question 11

11. How do transition metals differ from main group elements in terms of their oxidation states?

Answer 11

• Answer: Transition metals can lose electrons from both their outermost s-orbitals and their d-orbitals, allowing them to have multiple oxidation states. Unlike main group elements, transition metals rarely achieve noble gas configurations.

Question 12

12. Why does the 4s orbital fill before the 3d orbital, but electrons are lost from the 4s orbital first in transition metals?

Answer 12

• Answer: The 4s orbital fills before the 3d because it is lower in energy when unoccupied. However, once filled, the 3d orbitals become lower in energy than 4s, so the 4s electrons are removed first during ionization.

Question 13

13. What is metallic character, and how does it change across a period and down a group?

Answer 13

• Answer: Metallic character refers to how easily an atom loses electrons. It decreases across a period as Zeff increases, making it harder to lose electrons. It increases down a group as larger atoms lose electrons more easily due to weaker attraction from the nucleus.

Question 14

14. How do metals and non-metals typically form cations and anions?

Answer 14

• Answer: Metals tend to lose electrons, forming cations to achieve a noble gas configuration. Non-metals tend to gain electrons, forming anions, also aiming to achieve a noble gas configuration.

Question 15

15. What factors contribute to the trend in atomic size within a group?

Answer 15

• Answer: Atomic size increases within a group because as you move down, new electron shells are added, which increase the distance between the nucleus and the outermost electrons, reducing the effect of Zeff.

Question 16

16. Why do the alkali metals show increasing reactivity with water as you move down the group?

Answer 16

• Answer: The increasing reactivity is due to the decreasing ionization energy as you move down the group, which makes it easier for alkali metals to lose their outer electron and react with water.

Question 17

17. What is the significance of half-filled and fully-filled d-subshells in transition metals?

Answer 17

• Answer: Half-filled and fully-filled d-subshells provide additional stability to transition metals due to electron exchange and symmetry. This stability influences the oxidation states and chemical behavior of transition metals.

Question 18

18. How does bonding type correlate with electronegativity differences between atoms?

Answer 18

• Answer: Ionic bonds form between atoms with a large electronegativity difference, where electrons are transferred. Covalent bonds form between atoms with similar electronegativity, where electrons are shared. Metallic bonds occur between metal atoms where electrons are delocalized.

Question 19

1. How is the periodic table organized, and what trends does it reflect?

Answer 19

The periodic table is organized by increasing atomic number, which corresponds to the number of protons in the nucleus. The table reflects several trends:

*	Blocks: Elements are grouped into blocks based on their valence orbitals (s, p, d, f).
*	Categories: Elements are classified into three main categories:
*	Metals: Good conductors of heat and electricity, malleable and ductile.
*	Semi-metals: Have properties between metals and non-metals, often semiconductors.
*	Non-metals: Poor conductors of heat and electricity, brittle if solid, often gases or liquids.

Question 20

2. What are the key groups in the periodic table, and what general properties do they exhibit?

Answer 20

• Group 1 (Alkali Metals): Highly reactive, especially with water, soft, have low melting points.
  
  • Group 2 (Alkaline Earth Metals): Less reactive than alkali metals, but still reactive; harder than alkali metals.
  • Group 17 (Halogens): Very reactive non-metals, exist as diatomic molecules (e.g., F₂, Cl₂), form salts with metals.
  • Group 18 (Noble Gases): Inert gases with very low reactivity due to their full valence electron shells.

Each group shares similar chemical properties because they have the same number of valence electrons.

Question 21

3. What is the effective nuclear charge (Zeff), and how is it determined?

Answer 21

Effective nuclear charge (Zeff) is the net positive charge felt by an electron in an atom, after accounting for shielding by core electrons. It is given by the formula:

Z_{eff} = Z - S

Where:

*	Z is the atomic number (total number of protons),
*	S is the shielding constant, representing the repulsion from core electrons.

Shielding occurs because inner electrons reduce the attraction between the nucleus and outer electrons. However, outer electrons do not shield each other effectively. This means that Zeff increases moving across a period because core electrons stay the same, while Z increases.

Question 22

4. How does shielding affect the behavior of outer electrons in a multi-electron atom?

Answer 22

In multi-electron atoms, outer electrons experience shielding from the nucleus due to the core electrons. This causes outer electrons to experience a reduced effective nuclear charge (Zeff), as they feel less pull from the nucleus. Shielding makes it easier for outer electrons to be farther from the nucleus, contributing to larger atomic sizes for elements lower on the periodic table. However, shielding is not perfect; electrons in the same orbital do not shield each other effectively, so the outermost electrons are still somewhat attracted to the nucleus.

Question 23

5. What are Slater’s rules, and how are they used to calculate shielding constants (S)?

Answer 23

Slater’s rules provide a method to estimate the shielding constant (S) for an electron by considering the contributions of other electrons in an atom. The steps are:

1.	Write the electron configuration of the atom.
2.	Group the electrons into shells and subshells (e.g., (1s), (2s, 2p), (3s, 3p), (3d)).
3.	Apply the following values to calculate the shielding constant (S):
*	Electrons in the same group contribute 0.35 to S (except for 1s, which contributes 0.30).
*	Electrons in the n-1 shell contribute 0.85 to S.
*	Electrons in the n-2 shell and below contribute 1.00 to S.

For example, in potassium (K, Z=19), the electron configuration is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹. To calculate Zeff for the 4s electron:

*	Electrons in the same 4s group contribute 0.35 each (0 in this case since there’s only one electron in 4s).
*	Electrons in the n-1 (3s, 3p) group contribute 0.85 each.
*	Electrons in the n-2 (2s, 2p) and below contribute 1.00 each.

Question 24

6. How does atomic radius change down a group and across a period, and why?

Answer 24

• Down a group: The atomic radius increases because electrons are added to higher principal quantum numbers (n). This means that the valence electrons are farther from the nucleus, despite the increasing nuclear charge, because the inner electrons shield the outer ones effectively.
  
  • Across a period (left to right): The atomic radius decreases because the effective nuclear charge (Zeff) increases. As more protons are added to the nucleus, the added electrons do not fully shield each other. This pulls the outer electrons closer to the nucleus, shrinking the atomic radius.

Question 25

7. What is the difference between atomic radius and ionic radius, and how do they compare for cations and anions?

Answer 25

* Atomic radius is the distance from the nucleus to the outermost electron in a neutral atom. * Ionic radius is the distance from the nucleus to the outermost electron in an ion. For cations (positive ions): * The ionic radius is smaller than the atomic radius because losing electrons reduces electron-electron repulsion and allows the remaining electrons to be pulled closer to the nucleus. For anions (negative ions): * The ionic radius is larger than the atomic radius because gaining electrons increases electron-electron repulsion, causing the electro

Question 26

8. How do ionization energies vary across a period and down a group, and why?

Answer 26

Ionization energy is the energy required to remove an electron from a gaseous atom or ion. * Across a period: Ionization energy increases because the effective nuclear charge (Zeff) increases, making it more difficult to remove an electron. * Down a group: Ionization energy decreases because the valence electrons are farther from the nucleus due to higher principal quantum numbers, and the increased shielding makes it easier to remove an electron.

Question 27

9. What is the first ionization energy (I₁), and how does it relate to subsequent ionization energies?

Answer 27

* First ionization energy (I₁) is the energy required to remove the first electron from a neutral atom. * Second ionization energy (I₂) is the energy required to remove a second electron after the first has been removed, and so on. Ionization energies increase as more electrons are removed because: * The remaining electrons experience a higher effective nuclear charge (Zeff) due to less electron-electron repulsion. * Successive removals are from an increasingly positive ion, requiring more energy. For example, for Mg: * I₁ = 738 kJ/mol for Mg → Mg⁺ + e⁻. * I₂ = 1451 kJ/mol for Mg⁺ → Mg²⁺ + e⁻.

Question 28

10. Why are cations smaller than their neutral atoms, and anions larger?

Answer 28

* Cations: When an atom loses an electron to form a cation, the loss of electrons reduces electron-electron repulsion, allowing the remaining electrons to be pulled closer to the nucleus, shrinking the size of the ion. * Anions: When an atom gains an electron to form an anion, the gain of electrons increases electron-electron repulsion, which causes the electron cloud to expand, making the anion larger than the neutral atom.

Question 29

1. What is the relationship between cation size and the process of ionization?

Answer 29

* Answer: When an atom forms a cation, its valence electrons are removed. This results in the cation being smaller than the neutral atom because the remaining electrons experience a higher effective nuclear charge. The removal of electrons reduces electron-electron repulsions, allowing the nucleus to pull the remaining electrons closer, which further shrinks the ion.

Question 30

2. How does moving down a group affect cation size?

Answer 30

* Answer: As you move down a group, the cation size increases because the principal quantum number (n) of the outermost electrons increases. This means the outer electrons are in orbitals further from the nucleus, making the ion larger despite the increasing nuclear charge.

Question 31

3. Why are anions larger than their corresponding neutral atoms?

Answer 31

* Answer: Anions are larger than neutral atoms because when an atom gains electrons to form an anion, the added electrons experience less effective nuclear charge, which means they are held less tightly by the nucleus. Additionally, the extra electrons increase electron-electron repulsion, causing the electron cloud to expand.

Question 32

4. What is ionization energy (IE), and why is it always positive?

Answer 32

* Answer: Ionization energy is the minimum energy required to remove an electron from an atom or ion in the gas state. It is always positive because removing an electron from an atom requires energy input to overcome the attraction between the negatively charged electron and the positively charged nucleus, making this an endothermic process.

Question 33

5. How does effective nuclear charge (Zeff) affect ionization energy?

Answer 33

* Answer: Effective nuclear charge (Zeff) increases across a period because the number of protons in the nucleus increases while the shielding effect remains relatively constant. A higher Zeff results in a stronger attraction between the nucleus and the electrons, which increases the ionization energy since more energy is required to remove an electron.

Question 34

6. Why does ionization energy generally increase across a period?

Answer 34

* Answer: Ionization energy increases across a period because the effective nuclear charge increases while the atomic radius decreases. The electrons are held more tightly by the nucleus, so it requires more energy to remove an electron as you move from left to right across a period.

Question 35

7. Explain why there are irregularities in the trend of ionization energies.

Answer 35

* Answer: Irregularities in the ionization energy trend occur when electrons are added to a new sublevel with higher energy or when they are the first to pair in an orbital. In both cases, the added electron experiences more electron-electron repulsions, which lowers the ionization energy slightly compared to the general trend.

Question 36

8. Why does the second ionization energy of sodium (Na) exceed that of aluminum (Al)?

Answer 36

* Answer: The second ionization energy of Na is higher than that of Al because, after Na loses one electron, the next electron must be removed from a lower energy shell (closer to the nucleus), which is much harder to ionize. In contrast, Al’s second ionization energy involves removing an electron from the same shell, making it easier.

Question 37

9. What is electron affinity (EA), and how does it compare to ionization energy?

Answer 37

* Answer: Electron affinity is the energy released when an atom gains an electron to form an anion in the gas phase. Unlike ionization energy, which is always positive, electron affinity can be either positive or negative, depending on whether energy is released (exothermic) or required (endothermic) to add an electron. EA generally follows a similar trend to ionization energy across periods and groups.

Question 38

10. Why is the second electron affinity (EA2) always positive?

Answer 38

* Answer: The second electron affinity is always positive because adding a second electron to a negatively charged anion results in significant electron-electron repulsion. The attractive forces of the nucleus are not strong enough to completely overcome this repulsion, so energy must be added to force the second electron onto the anion.

Question 39

11. What is the main driving force behind chemical reactivity?

Answer 39

* Answer: The main driving force behind chemical reactivity is the tendency of atoms to achieve a complete valence shell, often known as the octet rule. Atoms react by either losing, gaining, or sharing electrons to achieve this stable configuration, which reduces their overall energy and makes them more stable.

Question 40

12. Differentiate between ionic and covalent bonds based on electronegativity differences.

Answer 40

* Answer: Ionic bonds form when there is a large difference (2.0 or more) in electronegativity between two atoms, typically between a metal and a non-metal. The more electronegative atom gains electrons to complete its valence shell, while the less electronegative atom loses electrons. Covalent bonds form when the electronegativity difference is smaller (less than 2.0), and atoms share electrons to fill their valence shells. Pure covalent bonds (ΔEN < 0.4) involve equal sharing, while polar covalent bonds (0.4 < ΔEN < 2.0) involve unequal sharing.

Question 41

13. Why do metals and non-metals typically form ionic bonds?

Answer 41

* Answer: Metals and non-metals typically form ionic bonds because metals have low electronegativity and tend to lose electrons, while non-metals have high electronegativity and tend to gain electrons. This large difference in electronegativity leads to the complete transfer of electrons from the metal to the non-metal, resulting in the formation of ions that are held together by electrostatic forces.

Question 42

14. Explain the concept of electronegativity and its trend across the periodic table.

Answer 42

* Answer: Electronegativity is the ability of an atom to attract electrons in a chemical bond. As you move from left to right across a period, electronegativity increases because the effective nuclear charge increases, pulling electrons more strongly. As you move down a group, electronegativity decreases because the added electron shells increase the distance between the nucleus and the bonding electrons, reducing the attraction.

Question 43

15. What is the difference between pure covalent and polar covalent bonds?

Answer 43

* Answer: In pure covalent bonds, the difference in electronegativity between the atoms is less than 0.4, so the electrons are shared equally. In polar covalent bonds, the electronegativity difference is between 0.4 and 2.0, meaning that one atom attracts the bonding electrons more strongly, leading to an unequal sharing of electrons.