Chapter 9: Periodic Properties of Elements

periodic property

A property of an element that is predictable based on an element’s position in the periodic table.

electron configuration

A notation that shows the particular orbitals that are occupied by electrons in an atom.

ground state

The lowest energy state in an atom, ion, or molecule.

orbital diagram

A diagram that gives information similar to an electron configuration but symbolizes an electron as an arrow in a box representing an orbital, with the arrow’s direction denoting the electron’s spin.

Pauli exclusion principle

The principle that no two electrons in an atom can have the same four numbers.

degenerate

A term describing two or more electron orbitals with the same value of n that have the same energy.

Coulomb's law

A scientific law stating that the potential energy between two charged particles is proportional to the product of the charges divided by the distance that separates the charges.

shielding

The effect on an electron of repulsion by electrons in lower-energy orbitals that screen it from the full effects of nuclear charge.

effective nuclear charge (Z_eff)

The actual nuclear charge experienced by an electron, defined as the charge of the nucleus plus the charge of the shielding electrons.

penetration

The phenomenon of some higher-level atomic orbitals having significant amounts of probability within the space occupied by orbitals of lower energy level. For example, the 2s orbital penetrates into the 1s orbital.

Hund's rule

The principle stating that when electrons fill degenerate orbitals, they first fill them singly with parallel spins.

aufbau principle

The principle that indicates the pattern of orbital filling in an atom.

valence electrons

Those electrons that are important in chemical bonding. For main-group elements, the valence electrons are those in the outermost principal energy level.

core electrons

Those electrons in a complete principal energy level and those in complete d and f sublevels.

van der Waals radius (nonbonding atomic radius)

(nonbonding atomic radius) One-half the distance between the centers of adjacent, nonbonding atoms in a crystal.

covalent radius (bonding atomic radius)

In nonmetals, one-half the distance between two atoms bonded together, and in metals one-half the distance between two adjacent atoms in a crystal of the metal.

atomic radius

The average bonding radius of an atom determined from measurements on a large number of elements and compounds.

paramagnetic

The state of an atom or ion that contains unpaired electrons and is, therefore, attracted by an external magnetic field.

diamagnetic

The state of an atom or ion that contains only paired electrons and is, therefore, slightly repelled by an external magnetic field.

ionization energy (IE)

The energy required to remove an electron from an atom or ion in its gaseous state.

electron affinity (EA).

The energy change associated with the gaining of an electron by an atom in its gaseous state.

Which ion is pumped into the cell during nerve signal transmission?

K⁺ – Potassium ions

Mendeleev arranged elements primarily by:

Atomic mass

Which orbital penetrates most into the nucleus?

2s

The spin quantum number (ms) can have which of the following values?

½ or –½

What causes sublevels in the same energy level to have different energies in multielectron atoms?

Penetration and shielding

True or False: In hydrogen, all sublevels within a principal energy level are degenerate.

True

True or False: Paramagnetic species have unpaired electrons and are attracted to magnetic fields.

True

True or False: Noble gases are highly reactive due to their stable electron configurations.

False

True or False: Electron affinity is always an exothermic process.

False

True or False: The first ionization energy generally increases across a period due to increasing effective nuclear charge.

True

According to Coulomb’s Law, the potential energy is ___________ when opposite charges are brought closer together.

Lower (or decreases)

The electrons in the outermost principal energy level of an atom are called ___________ electrons.

Valence

___________ states that no two electrons in the same atom can have the same set of four quantum numbers.

Pauli Exclusion Principle

___________ metals are very reactive, have low ionization energies, and are found only in compounds in nature.

Alkali

The shorthand electron configuration uses the symbol of the previous ___________ gas in brackets.

Noble

Explain why 2s electrons experience a greater attraction to the nucleus than 2p electrons.

2s electrons penetrate closer to the nucleus, experiencing less shielding and a stronger attraction to the nucleus than 2p electrons, which stay further out and are more shielded by inner electrons.

Why does the atomic radius increase down a group but decrease across a period?

Down a group, atoms get larger because they have more energy levels (shells). Across a period, atoms get smaller because increasing nuclear charge pulls electrons closer without adding more shielding.

Describe the difference in electron configurations between a transition metal cation and its neutral atom.

A transition metal cation loses electrons first from the outermost s orbital, not the d orbital. For example, Fe: [Ar] 4s² 3d⁶ → Fe²⁺: [Ar] 3d⁶ (loses the 4s electrons first).

Why do elements in the same group generally have similar chemical properties?

Elements in the same group have the same number of valence electrons, so they undergo similar types of chemical reactions and form similar compounds.

What happens to ionization energy as you remove successive electrons from an atom? Why?

Ionization energy increases with each successive electron removed because you’re removing electrons from a more positively charged ion, and eventually from a core shell, which requires much more energy.