Periodic Table and Periodic Properties – Lecture Review

A comprehensive set of Q&A flashcards covering electronic configurations, periodic table history and structure, periodic trends (atomic/ionic radii, ionization energy, electron gain enthalpy), shielding, penetration, and oxide nature to aid exam preparation.

How many horizontal rows are in the modern periodic table and what are they called?

Seven horizontal rows called periods.

How many vertical columns are in the modern periodic table and what are they called?

Eighteen vertical columns called groups.

What determines the block (s, p, d, f) to which an element belongs?

The subshell in which its valence (outer-most) electron is being added.

What is the maximum number of electrons that can be accommodated in the nth shell?

2n² electrons.

How many electrons can an individual orbital hold and what is special about their spins?

Two electrons with opposite spins (Pauli Exclusion Principle).

How many electrons can the s subshell hold in total?

Two electrons (one orbital).

How many electrons can the p subshell hold in total?

Six electrons (three orbitals).

How many electrons can the d subshell hold in total?

Ten electrons (five orbitals).

How many electrons can the f subshell hold in total?

Fourteen electrons (seven orbitals).

State Aufbau’s Principle in one sentence.

Electrons occupy orbitals of lowest available energy first before filling higher-energy orbitals.

What does Hund’s Rule of Maximum Multiplicity say about filling degenerate orbitals?

Every degenerate orbital is singly occupied before any pairing occurs.

According to Pauli’s Exclusion Principle, what must differ for two electrons in the same orbital?

Their spins must be opposite.

Give the expected and actual ground-state electronic configuration of chromium (Z = 24).

Expected: [Ar] 4s² 3d⁴; Actual: [Ar] 4s¹ 3d⁵ (half-filled d subshell stability).

Give the expected and actual ground-state electronic configuration of copper (Z = 29).

Expected: [Ar] 4s² 3d⁹; Actual: [Ar] 4s¹ 3d¹⁰ (completely filled d subshell stability).

Why are half-filled and fully-filled subshells especially stable?

They have extra stability from symmetry and exchange energy.

What does the term ‘exchange energy’ refer to in atomic structure?

Energy lowering due to possible exchanges of electrons with parallel spins in degenerate orbitals, favouring half-filled or filled subshells.

Which suffix is used in IUPAC temporary systematic element names?

The name ends with “-ium.”

What IUPAC root corresponds to the digit 0?

Nil.

What IUPAC root corresponds to the digit 5?

Penta.

State Döbereiner’s Law of Triads in one line.

In a triad, the atomic mass of the middle element is approximately the average of the other two and its properties are intermediate.

What was the main limitation of Döbereiner’s triads?

Only a few elements fit into triads because few elements were known and atomic masses were insufficiently precise.

What pattern did Newlands observe in his ‘Law of Octaves’?

Every eighth element (arranged by increasing atomic mass) showed similar properties, analogous to musical octaves.

Who produced a graph relating atomic volume to atomic mass that showed periodic repetition but published after Mendeleev?

Lothar Meyer.

List two key features of Mendeleev’s periodic table.

(1) Nine groups divided into A and B sub-groups, (2) Left intentional gaps predicting undiscovered elements (eka-elements).

Give an example of an element predicted by Mendeleev and its modern name.

Eka-aluminium predicted by Mendeleev is gallium.

What are ‘inverted pairs’ in Mendeleev’s table, and give one example.

Pairs where an element with higher atomic mass precedes a lower one to preserve chemical similarity, e.g., Ar (39.9) before K (39.1).

Who introduced the modern periodic law and what does it state?

Henry Moseley; properties of elements are periodic functions of their atomic numbers (and electronic configurations).

Define covalent radius.

Half the internuclear distance between two identical atoms joined by a single covalent bond.

Define metallic radius.

Half the distance between nuclei of two adjacent atoms in a metallic lattice.

Define van der Waals radius.

Half the distance between nuclei of two identical non-bonded atoms in neighbouring molecules in the solid state.

Explain the general trend of atomic radius down a group.

Atomic radius increases because additional electron shells are added.

Explain the general trend of atomic radius across a period (left to right).

Atomic radius decreases as effective nuclear charge increases while the principal shell number remains the same.

Why do noble gases deviate from the usual atomic-radius trend across a period?

Their radii are measured as van der Waals radii (no bonding) which are larger than covalent radii of neighbouring atoms.

Arrange the isoelectronic species N³⁻, O²⁻, F⁻, Na⁺, Mg²⁺, Al³⁺ in order of increasing radius.

Al³⁺ < Mg²⁺ < Na⁺ < F⁻ < O²⁻ < N³⁻ (radius decreases with increasing nuclear charge for same electron count).

What is meant by ‘shielding (screening) effect’?

Inner electrons partially block the attraction between the nucleus and outer-shell electrons, reducing effective nuclear charge on the outer electrons.

Which subshell provides the best shielding and which the worst?

s subshell shields best; f subshell shields worst.

Give the order of orbital penetration power toward the nucleus.

s > p > d > f (s penetrates most, f least).

Define first ionization enthalpy.

Minimum energy required to remove the most loosely bound electron from a neutral isolated gaseous atom.

State two factors that increase ionization enthalpy.

Higher effective nuclear charge and smaller atomic radius.

Why is the first ionization energy of oxygen slightly lower than that of nitrogen?

Oxygen’s 2p⁴ configuration has one paired electron causing repulsion; removing it gives a more stable half-filled 2p³ state.

How does successive ionization enthalpy change for a given element?

Each successive IE increases, with a large jump when an electron is removed from a new (inner) shell or a stable configuration.

Describe the trend of ionization enthalpy across a period.

Generally increases left to right due to rising nuclear charge and decreasing atomic size.

Describe the trend of ionization enthalpy down a group.

Generally decreases because atomic size increases and effective nuclear charge felt by valence electrons decreases (greater shielding).

Define electron gain enthalpy (electron affinity).

The energy released when an electron is added to a neutral isolated gaseous atom to form an anion.

Why is electron gain enthalpy of chlorine more negative than that of fluorine?

In F the very small size leads to greater electron-electron repulsion in the compact 2p shell, reducing energy released; Cl’s larger size accommodates the extra electron more easily.

Why do noble gases have positive (endothermic) electron gain enthalpies?

They possess completely filled shells, so adding an electron requires input of energy to enter a higher energy level.

State the general trend of electron gain enthalpy across a period up to the halogens.

Becomes increasingly negative from left to right, reaching maximum negativity at the halogen group.

Classify the oxides of metals and non-metals in terms of acid-base character.

Metal oxides tend to be basic; non-metal oxides tend to be acidic; some oxides are amphoteric (both acidic and basic); a few are neutral.

In size comparison, place the radii of a cation, its neutral atom, and its anion in order.

R(cation) < R(atom) < R(anion).