Chemistry 1050- Chapter 8: Periodic Properties of the Elements

Periodic Property

A property of an element that is predictable based on an element’s position in the periodic table. Aka. Elements in the same column have similar properties

Periodic Law

When elements are arranged by increasing atomic mass, their chemical properties repeat in a regular pattern.

Henry Moseley

Discovered that arranging elements according to their whole numbers (atomic number) better correlates elements according to elemental properties

Core Electrons

Electrons that are in complete principal energy levels (inner e-) and those in complete d and f sublevels. Held most strongly by the nucleus, are the most stable, and never react.

Valence Electrons

Outermost electrons found in the outermost principal energy level. Held loosely by the nucleus and are the most important bonding electrons

Row Number

Equal to number of the highest principal level

Inner Electron Configuration

The elections of the noble gas that preceded the element of interest (core electrons)

Outer electron configuration

The electrons beyond the previous noble gas of the element of interest (valance electrons)

Noble Gases

Eight valance electrons (two for He)- full quantum level and stable (lowest energy) and electrons cannot lower their energy by reaction with other substances so are particularly unreactive (inert).

Non-bonding Atomic Radius (Van der Waals Radius)

The size of an atom when it is touching but not bonded to another atom; it is half the distance between two nonbonding atoms that are in direct contact.

Bonding Atomic Radius (Covalent Radius)

The size of an atom when it is bonded to another atom; it is half the distance between the nuclei of two bonded atoms.

Atomic Radius

An average value based on many measured bonding radii across different elements and compounds.

Effective Nuclear Charge (Zeff)

The net positive charge an electron actually feels in an atom. In one-electron atoms, the electron only feels the nucleus pulling on it. In multi-electron atoms, inner electrons repel outer electrons, reducing how much nuclear pull they feel.

Slater’s Rules (John C. Slater)

A method for calculating the shielding constant (S) to find how much of the nuclear charge an electron actually feels.

Cations

Smaller than their corresponding atoms as they lose electrons. The nucleus still has the same number of protons so the current electrons feel the nucleus more strongly.

Anions

Larger than their corresponding atoms as they gain electrons. More electrons repelling each other, spreading out and making the atom blow up (like a puffer fish).

Isoelectronic

Iona with the same number of electrons though they are different elements.

Ionization Energy

The energy required to remove an electron from the atom or ion in the gaseous state and it is always positive.

Electron Affinity

The energy released (exothermic) when an atom gains an e- in the gas phase.

Alkali Metals (Group 1)

Highly reactive metals (Li- Cs) with one loosely held valence electron, known for low ionization energies, vigorous reactions (especially with water), and increasing reactivity down the group.

Alkaline Earth Metals (Group 2)

Metals with two valence electrons that become larger and easier to ionize down the group; they are denser with higher melting points than alkali metals and show moderate reactivity—reacting with nonmetals and slowly with water (with Be not reacting, Mg reacting only when heated, and Ca reacting readily but still far less violently than Group 1 metals). As well as reacting with oxygen forms quicklime (CaO) or lime.

Halogens (Group 17)

Nonmetals with seven valence electrons that are powerful oxidizing agents; their reactivity decreases from fluorine to iodine, and they commonly react with metals to form metal halides, with hydrogen to form hydrogen halides (which become acidic in water), and with each other to form interhalogen compounds. They increase in size and density down the group and exist in different physical states at room temperature (F₂ and Cl₂ as gases, Br₂ as a liquid, I₂ as a solid).

Noble Gases (Group 18)

Elements with completely full valence shells, giving them extremely high ionization energies and making them almost entirely unreactive; they exist as gases at room temperature, increase in size and density down the group, and only form compounds under very extreme conditions (mainly xenon and krypton reacting with fluorine). Their chemical inertness is useful for preventing unwanted reactions.