Unit 3 Periodic Table and Atomic Structure Review

Comprehensive vocabulary flashcards covering periodic trends, orbital shapes, electron configurations, family groups, and historical developments of the periodic table.

Dmitri Mendeleev

The scientist who first arranged the periodic table, organizing the elements in order of increasing atomic mass.

Henry Moseley

The scientist who discovered atomic number and organized the elements of the periodic table in order of increasing atomic number.

Periodic Law

The law stating that when elements of the periodic table are arranged in order of increasing atomic number, there are periodic patterns in their physical and chemical properties.

Groups

The vertical columns down the periodic table; elements in the same group have similar chemical properties because they have the same number of outer (valence) electrons.

Periods

The horizontal rows going across the periodic table; elements in the same period have the same number of occupied electron shells or energy levels.

Valence Electrons

The electrons located in the outermost shell or energy level of an atom.

Energy Level Capacities

The maximum electron capacities of the first three energy levels: the first energy level can hold up to $2$ electrons, the second up to $8$ electrons, and the third up to $18$ electrons.

Alkali Metals

A group of extremely reactive metals having $1$ valence electron that are never found in nature in their pure elemental form.

Alkaline Earth Metals

The group of metallic elements located in Group $2$ of the periodic table, which includes elements such as Barium ($Ba$).

Transition Metals

The large central block of metallic elements spanning Groups $3$ through $12$ on the periodic table.

Halogens

A group of highly reactive nonmetals in Group $17$ with $7$ valence electrons that tend to form salts.

Noble Gases

Non-reactive nonmetals located in Group $18$ that have their outermost shell completely filled with $8$ valence electrons.

Metalloids

The elements that touch the 'staircase' on the periodic table, possessing intermediate properties between metals and nonmetals.

Hydrogen

The only element on the periodic table that does not match the properties of any other group.

Ion

An atom with a positive or negative charge, which is formed when an atom loses or gains electrons.

Atomic Radius (Period Trend)

The size of an atom, which decreases across a period from left to right due to an increased nuclear charge; for example, Lithium ($Li$) has a larger radius than Fluorine ($F$) because Lithium has a weaker nuclear charge.

Positive Ion Size

An ion with a positive charge that is always smaller than the size of the neutral atom from which it was formed.

Negative Ion Size

An ion with a negative charge that is always larger than the size of the neutral atom from which it was formed.

Ionization Energy

The amount of energy required to remove the most loosely held electron from an atom; it increases across a period from left to right and decreases down a group.

Second Ionization Energy

The energy required to remove a second or third electron from an atom, which is always greater than the energy required to remove the first electron.

Potassium Ionization Energy Jump

A large increase in ionization energy observed when removing the $2\text{nd}$ electron from Potassium ($K$), because it already achieved the stable electron configuration of Argon ($Ar$) after losing the first electron.

Electronegativity

The ability of an atom to attract electrons to itself when the atom is in a chemical bond; it increases across a period and decreases down a group.

Fluorine Electronegativity

The element with the highest electronegativity value on the periodic table, possessing an electronegativity of $3.98$.

Group 18 Electronegativity

The elements belonging to Group $18$ (noble gases) that do not have defined values for electronegativity.

Relationship Between Ionization Energy and Electronegativity

Elements with high ionization energies typically have high electronegativity values because their nuclei strongly attract electrons, making them more likely to attract shared electrons in chemical bonds.

s Orbital

An atomic orbital characterized by a spherical ('sphere') shape.

p Orbital

An atomic orbital characterized by a 'peanut' shape.

d Orbital

An atomic orbital characterized by a 'daisy' shape.

f Orbital

An atomic orbital characterized by a 'flower' shape.

Electron Configuration of Selenium

The complete arrangement of electrons for Selenium ($Se$): $1s^2\,2s^2\,2p^6\,3s^2\,3p^6\,4s^2\,3d^{10}\,4p^4$.