1. What is the atomic number of an element?
Answer: The atomic number is the number of protons in the nucleus of an
atom. It uniquely identifies an element. For example, carbon has atomic
number 6, meaning every carbon atom has exactly 6 protons.
2. How is mass number defined?
Answer: Mass number (A) = number of protons + number of neutrons. It
represents the total count of nucleons (protons and neutrons) in the nucleus.
For example, carbon-12 has 6 protons and 6 neutrons, so A = 12.
3. What are isotopes?
Answer: Isotopes are atoms of the same element (same number of protons)
that differ in the number of neutrons. For example, ¹²C, ¹³C, and ¹⁴C are all
carbon isotopes. They have the same chemical properties but different
masses.
4. How is average atomic mass calculated?
Answer: Average atomic mass = Σ (fractional abundance × isotopic mass).
For example, if chlorine has ³⁵Cl (75.77%, 34.97 amu) and ³⁷Cl (24.23%, 36.97
amu): avg mass = (0.7577)(34.97) + (0.2423)(36.97) = 35.45 amu.
5. What is the difference between an atom and an ion?
Answer: An atom is electrically neutral (protons = electrons). An ion has
gained or lost electrons: a cation has lost electrons (positive charge) and an
anion has gained electrons (negative charge). For example, Na⁺ is a cation
(lost 1 electron).
6. Describe the Bohr model of the hydrogen atom.
Answer: Bohr proposed that electrons orbit the nucleus in fixed circular
energy levels (shells). Electrons absorb energy to jump to higher levels and
emit photons when falling to lower levels. The energy of each level is En = -
13.6/n² eV, where n is the principal quantum number.
,7. What are the four quantum numbers and what do they describe?
Answer: n (principal): energy level/shell size (n = 1,2,3...). l (angular
momentum): sublevel shape (0 to n-1). ml (magnetic): orbital orientation (-l to
+l). ms (spin): electron spin (+½ or -½). Together they uniquely describe every
electron in an atom.
8. State the Pauli Exclusion Principle.
Answer: No two electrons in an atom can have the same set of all four
quantum numbers. This means each orbital can hold a maximum of 2
electrons, and those two electrons must have opposite spins (+½ and -½).
9. State Hund's Rule.
Answer: When filling degenerate orbitals (orbitals of equal energy), electrons
occupy separate orbitals with parallel spins before pairing up. This minimizes
electron-electron repulsion and results in maximum multiplicity (spin).
10. What is the Aufbau principle?
Answer: Electrons fill orbitals in order of increasing energy: 1s, 2s, 2p, 3s, 3p,
4s, 3d, 4p, 5s, 4d, 5p... The name comes from German for 'building up.' The
lowest energy orbitals are filled first.
11. Write the electron configuration for Fe (Z=26).
Answer: Fe: [Ar] 3d⁶ 4s². The noble gas core [Ar] = 1s²2s²2p⁶3s²3p⁶. Then 4s
fills before 3d (lower energy), followed by 3d with 6 electrons. Note: When
forming Fe²⁺, the 4s electrons are lost first: [Ar] 3d⁶.
12. What is the trend in atomic radius across a period?
Answer: Atomic radius decreases left to right across a period. As we add
protons without adding a new shell, the effective nuclear charge (Zeff)
increases, pulling electrons closer to the nucleus. Thus Li > Be > B > C > N >
O > F > Ne.
13. What is the trend in atomic radius down a group?
Answer: Atomic radius increases going down a group. Each successive
period adds a new electron shell, increasing the distance between the nucleus
and the outermost electrons. For example, Li < Na < K < Rb < Cs.
14. Define ionization energy (IE).
Answer: Ionization energy is the minimum energy required to remove the
outermost electron from a gaseous atom: X(g) → X⁺(g) + e⁻. First IE increases
, across a period (greater Zeff) and decreases down a group (electrons further
from nucleus).
15. Why is the second ionization energy always greater than the first?
Answer: After removing the first electron, the resulting cation has fewer
electrons but the same nuclear charge. The remaining electrons experience
greater effective nuclear charge and are held more tightly, requiring more
energy to remove.
16. Define electron affinity.
Answer: Electron affinity is the energy change when a neutral gaseous atom
gains an electron: X(g) + e⁻ → X⁻(g). Generally becomes more negative (more
energy released) across a period. High electron affinity means the atom
strongly attracts an extra electron.
17. Define electronegativity.
Answer: Electronegativity is the ability of an atom in a molecule to attract
electron density toward itself. Fluorine is the most electronegative element
(3.98 on Pauling scale). Electronegativity increases left to right across a
period and decreases down a group.
18. What is effective nuclear charge (Zeff)?
Answer: Zeff = Z – S, where Z is atomic number and S is the shielding
constant. Core electrons shield outer electrons from the full nuclear charge.
Zeff increases across a period (more protons, similar shielding), causing the
trends in atomic radius, IE, and EN.
19. What is the photoelectric effect and its significance?
Answer: When light of sufficient frequency hits a metal surface, electrons are
ejected. Below the threshold frequency, no electrons are ejected regardless of
intensity. Einstein explained this by proposing light travels in photons with
energy E = hf. This proved the quantized nature of light.
20. What is the de Broglie wavelength?
Answer: De Broglie proposed that matter has wave-like properties. The
wavelength λ = h/mv, where h is Planck's constant, m is mass, and v is
velocity. This is significant for small particles like electrons but negligible for
macroscopic objects.
UNIT 2: Chemical Bonding
, 21. What is an ionic bond?
Answer: An ionic bond forms through the electrostatic attraction between
oppositely charged ions. It occurs when one atom transfers electron(s) to
another, typically between a metal and nonmetal. Example: Na⁺Cl⁻. The lattice
energy holds the crystal together.
22. What is a covalent bond?
Answer: A covalent bond forms when two atoms share one or more pairs of
electrons. It typically occurs between nonmetals. A single bond shares 2
electrons, double bond shares 4, triple bond shares 6. Example: H₂, O₂, N₂.
23. What is a polar covalent bond?
Answer: A polar covalent bond forms between atoms with different
electronegativities (ΔEN = 0.4 to 1.7). Electrons are shared unequally,
creating partial charges (δ+ and δ-). Example: H–Cl, where Cl is more
electronegative and has δ-.
24. How does electronegativity difference determine bond type?
Answer: ΔEN < 0.4: nonpolar covalent (e.g., H₂, Cl₂). ΔEN 0.4–1.7: polar
covalent (e.g., HCl, H₂O). ΔEN > 1.7: ionic (e.g., NaCl, MgO). These are
guidelines, not strict rules.
25. Explain the octet rule.
Answer: Main group atoms tend to form bonds to achieve 8 electrons in their
valence shell (like noble gases). H follows a duet rule (2 electrons).
Exceptions include incomplete octets (BF₃), expanded octets (PCl₅), and odd-
electron species (NO).
26. How do you draw a Lewis structure?
Answer: Steps: (1) Count total valence electrons. (2) Connect atoms with
single bonds. (3) Complete octets of outer atoms. (4) Place remaining
electrons on central atom. (5) Form multiple bonds if central atom lacks octet.
Check: total electrons used = total valence electrons.
27. What is formal charge and how is it calculated?
Answer: Formal charge = (valence electrons) – (lone pair electrons) –
½(bonding electrons). The best Lewis structure has formal charges closest to
zero. If nonzero, negative formal charge should be on the more
electronegative atom.
28. What is resonance?
Answer: The atomic number is the number of protons in the nucleus of an
atom. It uniquely identifies an element. For example, carbon has atomic
number 6, meaning every carbon atom has exactly 6 protons.
2. How is mass number defined?
Answer: Mass number (A) = number of protons + number of neutrons. It
represents the total count of nucleons (protons and neutrons) in the nucleus.
For example, carbon-12 has 6 protons and 6 neutrons, so A = 12.
3. What are isotopes?
Answer: Isotopes are atoms of the same element (same number of protons)
that differ in the number of neutrons. For example, ¹²C, ¹³C, and ¹⁴C are all
carbon isotopes. They have the same chemical properties but different
masses.
4. How is average atomic mass calculated?
Answer: Average atomic mass = Σ (fractional abundance × isotopic mass).
For example, if chlorine has ³⁵Cl (75.77%, 34.97 amu) and ³⁷Cl (24.23%, 36.97
amu): avg mass = (0.7577)(34.97) + (0.2423)(36.97) = 35.45 amu.
5. What is the difference between an atom and an ion?
Answer: An atom is electrically neutral (protons = electrons). An ion has
gained or lost electrons: a cation has lost electrons (positive charge) and an
anion has gained electrons (negative charge). For example, Na⁺ is a cation
(lost 1 electron).
6. Describe the Bohr model of the hydrogen atom.
Answer: Bohr proposed that electrons orbit the nucleus in fixed circular
energy levels (shells). Electrons absorb energy to jump to higher levels and
emit photons when falling to lower levels. The energy of each level is En = -
13.6/n² eV, where n is the principal quantum number.
,7. What are the four quantum numbers and what do they describe?
Answer: n (principal): energy level/shell size (n = 1,2,3...). l (angular
momentum): sublevel shape (0 to n-1). ml (magnetic): orbital orientation (-l to
+l). ms (spin): electron spin (+½ or -½). Together they uniquely describe every
electron in an atom.
8. State the Pauli Exclusion Principle.
Answer: No two electrons in an atom can have the same set of all four
quantum numbers. This means each orbital can hold a maximum of 2
electrons, and those two electrons must have opposite spins (+½ and -½).
9. State Hund's Rule.
Answer: When filling degenerate orbitals (orbitals of equal energy), electrons
occupy separate orbitals with parallel spins before pairing up. This minimizes
electron-electron repulsion and results in maximum multiplicity (spin).
10. What is the Aufbau principle?
Answer: Electrons fill orbitals in order of increasing energy: 1s, 2s, 2p, 3s, 3p,
4s, 3d, 4p, 5s, 4d, 5p... The name comes from German for 'building up.' The
lowest energy orbitals are filled first.
11. Write the electron configuration for Fe (Z=26).
Answer: Fe: [Ar] 3d⁶ 4s². The noble gas core [Ar] = 1s²2s²2p⁶3s²3p⁶. Then 4s
fills before 3d (lower energy), followed by 3d with 6 electrons. Note: When
forming Fe²⁺, the 4s electrons are lost first: [Ar] 3d⁶.
12. What is the trend in atomic radius across a period?
Answer: Atomic radius decreases left to right across a period. As we add
protons without adding a new shell, the effective nuclear charge (Zeff)
increases, pulling electrons closer to the nucleus. Thus Li > Be > B > C > N >
O > F > Ne.
13. What is the trend in atomic radius down a group?
Answer: Atomic radius increases going down a group. Each successive
period adds a new electron shell, increasing the distance between the nucleus
and the outermost electrons. For example, Li < Na < K < Rb < Cs.
14. Define ionization energy (IE).
Answer: Ionization energy is the minimum energy required to remove the
outermost electron from a gaseous atom: X(g) → X⁺(g) + e⁻. First IE increases
, across a period (greater Zeff) and decreases down a group (electrons further
from nucleus).
15. Why is the second ionization energy always greater than the first?
Answer: After removing the first electron, the resulting cation has fewer
electrons but the same nuclear charge. The remaining electrons experience
greater effective nuclear charge and are held more tightly, requiring more
energy to remove.
16. Define electron affinity.
Answer: Electron affinity is the energy change when a neutral gaseous atom
gains an electron: X(g) + e⁻ → X⁻(g). Generally becomes more negative (more
energy released) across a period. High electron affinity means the atom
strongly attracts an extra electron.
17. Define electronegativity.
Answer: Electronegativity is the ability of an atom in a molecule to attract
electron density toward itself. Fluorine is the most electronegative element
(3.98 on Pauling scale). Electronegativity increases left to right across a
period and decreases down a group.
18. What is effective nuclear charge (Zeff)?
Answer: Zeff = Z – S, where Z is atomic number and S is the shielding
constant. Core electrons shield outer electrons from the full nuclear charge.
Zeff increases across a period (more protons, similar shielding), causing the
trends in atomic radius, IE, and EN.
19. What is the photoelectric effect and its significance?
Answer: When light of sufficient frequency hits a metal surface, electrons are
ejected. Below the threshold frequency, no electrons are ejected regardless of
intensity. Einstein explained this by proposing light travels in photons with
energy E = hf. This proved the quantized nature of light.
20. What is the de Broglie wavelength?
Answer: De Broglie proposed that matter has wave-like properties. The
wavelength λ = h/mv, where h is Planck's constant, m is mass, and v is
velocity. This is significant for small particles like electrons but negligible for
macroscopic objects.
UNIT 2: Chemical Bonding
, 21. What is an ionic bond?
Answer: An ionic bond forms through the electrostatic attraction between
oppositely charged ions. It occurs when one atom transfers electron(s) to
another, typically between a metal and nonmetal. Example: Na⁺Cl⁻. The lattice
energy holds the crystal together.
22. What is a covalent bond?
Answer: A covalent bond forms when two atoms share one or more pairs of
electrons. It typically occurs between nonmetals. A single bond shares 2
electrons, double bond shares 4, triple bond shares 6. Example: H₂, O₂, N₂.
23. What is a polar covalent bond?
Answer: A polar covalent bond forms between atoms with different
electronegativities (ΔEN = 0.4 to 1.7). Electrons are shared unequally,
creating partial charges (δ+ and δ-). Example: H–Cl, where Cl is more
electronegative and has δ-.
24. How does electronegativity difference determine bond type?
Answer: ΔEN < 0.4: nonpolar covalent (e.g., H₂, Cl₂). ΔEN 0.4–1.7: polar
covalent (e.g., HCl, H₂O). ΔEN > 1.7: ionic (e.g., NaCl, MgO). These are
guidelines, not strict rules.
25. Explain the octet rule.
Answer: Main group atoms tend to form bonds to achieve 8 electrons in their
valence shell (like noble gases). H follows a duet rule (2 electrons).
Exceptions include incomplete octets (BF₃), expanded octets (PCl₅), and odd-
electron species (NO).
26. How do you draw a Lewis structure?
Answer: Steps: (1) Count total valence electrons. (2) Connect atoms with
single bonds. (3) Complete octets of outer atoms. (4) Place remaining
electrons on central atom. (5) Form multiple bonds if central atom lacks octet.
Check: total electrons used = total valence electrons.
27. What is formal charge and how is it calculated?
Answer: Formal charge = (valence electrons) – (lone pair electrons) –
½(bonding electrons). The best Lewis structure has formal charges closest to
zero. If nonzero, negative formal charge should be on the more
electronegative atom.
28. What is resonance?
