Atomic Structure

Question 1

Q: What is the atomic number (Z)?

Answer 1

A: The number of protons in an atom’s nucleus.

Question 2

Q: What is the mass number (A)?

Answer 2

A: The sum of protons and neutrons in an atom.

Question 3

Q: What is an ion?

Answer 3

A: An atom or molecule with a charge due to electron gain or loss.

Question 4

Q: What is an isotope?

Answer 4

A: Atoms of the same element with different numbers of neutrons.

Question 5

Q: What do electron configurations describe?

Answer 5

A: The arrangement of electrons in atomic orbitals.

Question 6

Q: What are the three rules for electron configuration?

Answer 6

A: Aufbau principle, Pauli exclusion principle, Hund’s rule.

Question 7

Q: How do electrons fill orbitals according to Hund’s rule?

Answer 7

A: Electrons occupy orbitals singly before pairing.

Question 8

Q: What is an exception to the expected electron configuration?

Answer 8

A: Chromium and copper have unusual configurations due to energy levels.

Question 9

Q: What are the four types of subshells in an atom?

Answer 9

A: s, p, d, and f.

Question 10

Q: How many electrons can each atomic orbital hold?

Answer 10

A: A maximum of two electrons.

Question 11

Q: What is ionisation energy?

Answer 11

A: The energy needed to remove an electron from an atom.

Question 12

Q: What are the three rules for filling electron orbitals?

Answer 12

A: Aufbau principle, Pauli exclusion principle, and Hund’s rule.

Question 13

Q: Why does successive ionisation energy increase?

Answer 13

A: Fewer electrons remain, increasing nuclear attraction.

Question 14

Q: What distinguishes isotopes of the same element?

Answer 14

A: They have the same number of protons but different numbers of neutrons.

Question 15

Q: What is the electron configuration of Na⁺ (sodium, 11)?

Answer 15

A: 1s² 2s² 2p⁶

Question 16

Q: According to the Pauli exclusion principle, how do electrons occupy an orbital?

Answer 16

A: No two electrons in the same orbital can have the same spin.

Question 17

Q: Why is the second ionisation energy higher than the first?

Answer 17

A: The remaining electrons experience stronger attraction to the nucleus.

Question 18

Q: What are isotopes?

Answer 18

A: Atoms of the same element with different numbers of neutrons.

Question 18

Q: How are isotopes represented?

Answer 18

A: Using nuclear notation (𝑍^𝐴X) or element-A (e.g., ²H or Hydrogen-2).

Question 19

Q: How do isotopes affect chemical and physical properties?

Answer 19

A: They have the same chemical properties (same electron configuration) but different physical properties due to mass differences.

Question 20

Q: What is the atomic mass unit (amu) based on?

Answer 20

A: 1 amu is defined as 1/12 the mass of a carbon-12 atom.

Question 21

Q: Why is relative atomic mass (Ar) not a whole number?

Answer 21

A: It is a weighted average of an element’s isotopes based on natural abundance.

Question 22

Q: What is the formula for calculating relative atomic mass (Ar)?

Answer 22

A:
r.a.m.=
(x/100 × mass isotope-a )+ (100 -x/100× mass isotope-b)

Question 23

Q: How are elements represented?

Answer 23

A: By chemical symbols (e.g., H for Hydrogen, O for Oxygen).

Question 24

Q: What is the periodic table based on?

Answer 24

A: Increasing atomic number (number of protons).

Question 25

Q: What are the four blocks of the periodic table?

Answer 25

A: s, p, d, and f blocks, corresponding to electron sub-levels.

Question 26

Q: What does the group number indicate?

Answer 26

A: The number of valence electrons.

Question 27

Q: What does the period number indicate?

Answer 27

A: The number of electron shells.

Question 28

Q: What happens to atomic radius across a period?

Answer 28

A: It decreases due to increased nuclear charge pulling electrons closer.

Question 29

Q: What happens to atomic radius down a group?

Answer 29

A: It increases due to additional electron shells.

Question 30

Q: How does ionisation energy change across a period?

Answer 30

A: It increases because electrons are held more tightly by the nucleus.

Question 31

Q: How does electronegativity change down a group?

Answer 31

A: It decreases as outer electrons are farther from the nucleus.

Question 32

Q: How do metals react in terms of electrons?

Answer 32

A: They lose electrons and form cations.

Question 33

Q: How do non-metals react in terms of electrons?

Answer 33

A: They gain electrons and form anions.

Question 34

Q: Why do alkali metals become more reactive down the group?

Answer 34

A: Outer electrons are further from the nucleus, making them easier to lose.

Question 35

Q: Why do halogens become more reactive up the group?

Answer 35

A: Smaller atoms attract electrons more strongly.

Question 36

Q: What are the four types of oxides?

Answer 36

A: Basic, acidic, amphoteric, and neutral oxides.

Question 37

Q: What is an amphoteric oxide?

Answer 37

A: An oxide that reacts with both acids and bases (e.g., Al₂O₃).

Question 38

Q: What does mass spectrometry measure?

Answer 38

A: The isotopic composition, relative atomic mass, and isotope abundances.

Question 39

Q: What is the difference between absorption and emission spectra?

Answer 39

A: Absorption spectra show energy absorbed; emission spectra show energy emitted.

Question 40

Q: What does atomic absorption spectroscopy (AAS) detect?

Answer 40

A: The concentration of metallic ions in a solution.

Question 41

Q: What is metallic bonding?

Answer 41

A: A lattice of cations surrounded by a sea of delocalized electrons.

Question 42

Q: What type of elements form ionic bonds?

Answer 42

A: Metals (which lose electrons) and non-metals (which gain electrons).

Question 43

Q: What type of elements form covalent bonds?

Answer 43

A: Non-metals, by sharing electrons.

Question 44

Q: Why do transition metals form multiple ions?

Answer 44

A: They have variable oxidation states due to unfilled d-orbitals.

Question 45

Q: What does the periodic law state?

Answer 45

A: Elements arranged by atomic number show recurring properties.

Question 46

Q: Why does first ionisation energy increase across a period?

Answer 46

A: Increased nuclear charge pulls electrons closer, requiring more energy to remove them.

Question 47

Q: Why are alkali metals more reactive as you go down the group?

Answer 47

A: Outer electrons are further from the nucleus and more easily lost.

Question 48

Q: What does a mass spectrometer determine about an element?

Answer 48

A: Its isotopic composition and relative atomic mass.

Question 49

Q: What type of oxide is Al₂O₃ and why?

Answer 49

A: Amphoteric, because it reacts with both acids and bases.

Question 50

Q: Which elements can have an expanded octet?

Answer 50

A: Sulfur (S) and Phosphorus (P) in Period 3 and beyond.

Question 51

Q: What does it mean when an atom has an expanded valence shell?

Answer 51

A: It has more than 8 electrons because extra orbitals are available.

Question 52

Q: Which elements can have fewer than 8 electrons in a stable structure?

Answer 52

A: Beryllium (Be) and Boron (B).

Question 53

Q: What are the general steps to draw a skeletal structure of a molecule?

Answer 53

1. Place the least electronegative atom in the center. 2. Connect atoms with bonding pairs of electrons. 3. Subtract used electrons from the total valence count. 4. Distribute remaining electrons to satisfy the octet rule. 5. Use multiple bonds if necessary.

Question 54

Q: When drawing a skeletal structure, where should hydrogen (H) be placed?

Answer 54

A: Always terminal (never the central atom).