Atomic structure and properties flashcards

Atomic structure and properties covers the fundamental makeup of atoms, including protons, neutrons, and electrons, and how these determine an element's identity and behavior. Students must understand isotopes, mass number, and how to calculate average atomic mass from isotopic abundances. Electron configuration using the aufbau principle, Hund's rule, and the Pauli exclusion principle is essential, as is understanding periodic trends such as atomic radius, ionization energy, and electronegativity. The quantum mechanical model of the atom, including the four quantum numbers and the shapes of orbitals, forms the theoretical backbone. Students should also connect electron configuration to an element's position on the periodic table and predict properties based on that position.

1. Card 1

   What defines the identity of an element?

   The number of protons in the nucleus, called the atomic number, uniquely identifies each element.

2. Card 2

   What are isotopes?

   Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons, giving them different mass numbers.

3. Card 3

   How is average atomic mass calculated?

   It is the weighted average of all naturally occurring isotopic masses, each multiplied by its percent abundance expressed as a decimal.

4. Card 4

   State the Pauli exclusion principle.

   No two electrons in an atom can have the same set of four quantum numbers; each orbital holds at most two electrons with opposite spins.

5. Card 5

   State Hund's rule.

   Electrons fill orbitals of equal energy one at a time before pairing up, maximizing the number of unpaired electrons in a subshell.

6. Card 6

   What are the four quantum numbers?

   Principal (n), angular momentum (l), magnetic (ml), and spin (ms) quantum numbers together describe the state of any electron in an atom.

7. Card 7

   What is the trend in atomic radius across a period?

   Atomic radius decreases across a period because increasing nuclear charge pulls electrons closer without adding a new shell.

8. Card 8

   What is the trend in first ionization energy across a period?

   First ionization energy generally increases across a period as greater nuclear charge holds electrons more tightly, making removal harder.

9. Card 9

   Why does ionization energy decrease down a group?

   Electrons are farther from the nucleus and experience more shielding from inner electrons, reducing the effective nuclear charge felt by valence electrons.

10. Card 10

    What is electronegativity and how does it vary on the periodic table?

    Electronegativity measures an atom's pull on shared electrons; it increases across a period and decreases down a group, peaking at fluorine.

11. Card 11

    What is the aufbau principle?

    The aufbau principle states electrons occupy the lowest available energy orbitals first, filling 1s before 2s, 2s before 2p, and so on.

12. Card 12

    What is effective nuclear charge (Zeff)?

    Effective nuclear charge is the net positive charge experienced by a valence electron after accounting for shielding by inner-shell electrons.

• Students often confuse mass number with average atomic mass, forgetting that mass number is a whole-number count of protons plus neutrons for one specific isotope.
• Many students forget that ionization energy dips at Group 13 and Group 16, missing that half-filled and fully filled subshells create slight stability exceptions.
• Students frequently misapply the aufbau order by forgetting that 4s fills before 3d but empties before 3d when forming transition metal cations.

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