Periodicity

Main Group Elements

Elements that belong to the s and p blocks of the periodic table

Second Ionization Energy, I2

Energy required to remove the second electron from a gaseous ion: Na+(g) → Na2+(g) + e-

Qualities of Nonmetals

- More diverse in their behavior than metals - Nonlustrous, are poor conductors of heat and electricity, and exhibit lower melting points than metals - When nonmetals react with metals, nonmetals tend to gain electrons - Compounds composed entirely of nonmetals are molecular substances. - Most nonmetal oxides are acidic: - Nonmetal oxides react with bases to form salts and water:

Metalloids

- Properties intermediate of metals and nonmetals - Found fame in semiconductor industry

Qualities of Metals

- Shiny, Lustrous, Malleable, and Ductile - Solids at room temperature with very high melting temperatures - Metals tend to be oxidized when they react and compounds of metals with nonmetals tend to be ionic substances. - Most metal oxides are basic and are able to react with acids to form salts and water.

Exceptions to ionization energy

A IIIA element (ns2 np1) has smaller ionization energy than the preceding IIA element (ns2). Apparently, the np electron of the IIIA element is more easily removed than one of the ns electrons of the preceding IIA element. Also note that a VIA element (ns2 np4) has smaller ionization energy than the preceding VA element. As a result of electron repulsion, it is easier to remove an electron from the doubly occupied np orbital of the VIA element than from a singly occupied orbital of the preceding VA element.

Isoelectronic Series

A group of ions that all have the same number of electrons

Core Electrons

All Electrons in Inner Energy Levels including completely filled d and f subshells.

First Ionization Energy, I1

Amount of energy required to remove an electron from a gaseous atom: Na(g) → Na+(g) + e-

Isoelectronic Series Trend

As nuclear charge increases in an isoelectronic series the ions become smaller: O2- > F- > Na+ > Mg2+ > Al3+

Trends for Atomic Radius

As the principal quantum number increases (i.e., we move down a group), the distance of the outermost electron from the nucleus becomes larger. Hence the atomic radius increases. As we move across the periodic table, the number of core electrons remains constant, however, the nuclear charge increases. Therefore, there is an increased attraction between the nucleus and the outermost electrons. This attraction causes the atomic radius to decrease.

Atomic Number

Characteristic of elements based on table location

Apparent Radius

Determined by the closest distances separating the nuclei during such collisions.

In which direction on the periodic table does metallic character increase?

Down and to the left

What is true about the increasing ionization energy?

Each successive ionization requires increasing amounts of energy.

Which of the following was an element unknown to, but correctly predicted by, Mendeleev?

Eka-Silicon (Germanium)

Trends for electron affinity

Electron affinity increases upward for the groups and from left to right across periods of a periodic table

Why are anions larger than their parent atoms?

Electrons have been added to the most spatially extended orbital which means total electron-electron repulsion has increased.

Why are Cations smaller than their parent atoms?

Electrons have been removed from the most spatially extended orbital and the effective nuclear charge has increased.

Trend of Ionic Size

For ions with the same charge, ionic size increases down a group.

Why do Halogens tend to form anions?

Gaining electrons will fill their octet faster than losing them.

Under which process does an atom demonstrate the largest increase in size in one step?

Gaining its first electron to become a -1 anion.

Ionization Energy Trends

Generally increases across a period since as we move across a period, Zeff increases, making it more difficult to remove an electron. It also decreases down a group since the outermost electron is more readily removed as we go down a group and the atom gets bigger, it becomes easier to remove an electron from the most spatially extended orbital.

Exceptions for Electron Affinity

Group 2A has less electron affinity than Group 1A as there is now a new p orbital created which causes the outermost electrons to be farther from the nucleus, causing even more repulsion. Group 5A has less electron affinity than Group 4A as the electron gets added to an already occupied orbital.

Which group are halogens found in?

Group 7A

Covalent Radius

Half of the bond distance if the two atoms that make up the molecule are the same

trends in electronegativity

Increases from left to right across a period on the periodic table and increase from bottom to top of a group on the periodic table.

Trends for metallic character

It increases down a group and decreases from left to right across a period.

Electronegativity

Measure of an atom's ability to attract electrons (Noble gases have no electronegativity).

Who developed the earliest form of our modern periodic table?

Mendeleev

Which of the groups of elements tend to form cations?

Metals

Ionization Energy

Minimum energy required to remove an electron from the ground state of the isolated gaseous atom or ion. Cl(g) --> Cl+(g ) + e-

Which of the following series of isoelectronic ions (Mg²⁺, N³⁻, F⁻, Si⁴⁺) has the largest ionic radii?

N³⁻

why is ionic size important?

Predicts lattice energy and determines the way in which ions pack in a solid. Ionic size is periodic.

What does the principal quantum number refer to?

Refers to the size and energy level of the outer orbital

Elements that are most similar in their properties are found in the same period or same group?

Same group

Periodic Law

Similar properties recur periodically when elements are arranged according to increasing atomic number

Why is sodium, NA, which has a nuclear charge of 11 have a nuclear effective charger near 2.667 rather than 1?

The assumptions that core electrons shield protons 1 for 1 and valence electrons do not shield are not completely valid.

In which direction in a given period of the periodic table does melting point tend to increase?

The closer an atom is to having a half-full valence shell, the higher the melting point. This yields a high melting point in the middle of the period and low melting points at the ends.

Why do noble gases have a low electron affinity?

The electron affinity is the change in energy when an electron is added to a neutral gaseous atom to form an ion. It would require a lot of energy to do this for neon because it is a noble gas with a full octet.

Electron Affinity

The energy change when a gaseous atom gains an electron to form a gaseous ion. It can be exothermic (release heat) or endothermic (absorb heat). (The value that is more negative is the more electron affinity.) Cl(g) + e- --> Cl-(g)

Metallic Character

The extent to which the element exhibits the physical and chemical properties of metals

Ionization Energy Changes

The larger the ionization energy, the more difficult it is to remove the electron and there is a sharp increase in ionization energy when a core electron is removed.

Which of the following types of electrons penetrates most toward the nucleus and therefore has the greatest shielding effect?

The s electrons penetrate most toward the nucleus and therefore have the greatest shielding effect.

S

The screening constant which represents the portion of the nuclear charge that is screened from the valence electron by other electrons in the atom; usually close to the number of core electrons in an atom.

Which group are metalloids found in?

The staircase of elements including Boron from Groups 3A to 6A

Typically, the ionization energy increases from left to right on the periodic table. However, there are minor irregularities. For example, the ionization of boron (Group 3A) is lower than beryllium (Group 2). Considering the electron configuration of these two atoms, choose the best explanation for this observation.

The valence electron removed from boron is in a 2p orbital rather than a 2s orbital. This is further from the nucleus and so it is not held as tightly.

Typically, the ionization energy increases from left to right on the periodic table. However, there are minor irregularities. For example, the ionization of oxygen (Group 6A) is lower than nitrogen (Group 5A). Considering the electron configuration of these two atoms, choose the best explanation for this observation.

There is more pairing energy in oxygen's electron configuration than in nitrogen's electron configuration.

Effective nuclear charge (Zeff) equation:

Z - S = Zeff

Aktive Metals

alkali metals (group 1A) and the alkaline earth metals (group 2A)

Most early attempts to organize the elements were based on atomic mass. Mendeleev occasionally switched elements so that apparently lighter elements occurred after a heavier element. He did this so that elemental properties properly recurred in the same group. This choice anticipated the discovery of which two features of elements?

atomic number and isotopes

Bonding atomic radius

distance between the two nuclei

Valence Electrons

outer shell electrons including partially filled d and f subshells; Each element in the same group contains the same amount of valence electrons

Effective nuclear charge (Zeff)

the charge experienced by an electron on a many-electron atom; depends on its distance from the nucleus and the number of electrons in the spherical volume out to the electron in question.

Periodic Table

the most significant tool that chemists use for organizing and recalling chemical facts

Factors of Atomic Radius

the principal quantum number, n, and the effective nuclear charge, Zeff.

Factors of Nuclear Effective Energy

the type of orbital an electron occupies, the number of core electrons, and the number of valence electrons