Unit 4 Chemistry Test

Periodic Relationships Among the Elements

Periodic Table

A tabular arrangement of the chemical elements, organized by increasing atomic number.

Dmitri Mendeleev

Russian chemist who created the first widely recognized periodic table in 1869.

Henry Moseley

Scientist who rearranged the periodic table by atomic number in 1913, leading to the modern periodic table.

Periodic Law

When elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties.

Atomic Radius

Half the distance between the nuclei of two identical atoms in a diatomic molecule.

Cations

Positively charged ions that result from an atom losing one or more electrons — smaller than neutral atom

Anions

Negatively charged ions that result from an atom gaining one or more electrons — larger than neutral atom

Isoelectronic

Describes atoms or ions having the same number of electrons as a noble gas.

Metallic Character

The level of reactivity of a metal (ex. sodium is a very reactive metal); increases down a group and decreases across a period.

Nuclear Charge

The total charge of the nucleus, increases with every proton it gains.

Shielding Effect

The phenomenon where inner electrons block the effective nuclear charge experienced by outer electrons.

Electron Configuration

The arrangement of electrons in an atom's orbitals, which influences the chemical behavior of the element.

Periodic Trends

Patterns in the properties of elements that occur across periods and down groups in the periodic table.

Electronegativity

The tendency for an element to attract electrons when chemically combined with another element — decreases down a group and increases across a period.

How does a cation’s size change when its number of electrons changes?

As more electrons are lost, the ion becomes smaller (e.g., Mg > Mg+1 > Mg+2).

How does an anion’s size change when its number of electrons changes?

As more electrons are gained, the ion becomes larger (e.g., O < O-1 < O-2).

Which element has the highest electronegativity

Fluorine (4.0)

Ionization Energy (I.E.)

The amount of energy required to remove an electron from a gaseous atom — decreases down a group due to valence electrons being further from the nucleus, increases as you move across a period due to increased effective nuclear charge

1st Ionization Energy

The energy required to remove the first electron from an atom.

2nd Ionization Energy

The energy required to remove the second electron from an atom.

3rd Ionization Energy

The energy required to remove the third electron from an atom.

Electron Affinity

The amount of energy released or spent when an electron is accepted by an atom in the gaseous state to form an anion — increase across a period and decrease down a group.

Exceptions in Electron Affinity

The electron affinities of elements B through F are less negative than those of the elements immediately below them due to electron-electron repulsions.