A hydrogen atom is made up of a single positive-charged nucleus surrounded by a single electron. Hydrogen may be the most significant element of all because of its simple structure. In the Sun, hydrogen (H) nuclei combine to produce helium (He) nuclei, which accounts for roughly all of Earth's energy.

H atoms make up around 90% of all atoms in the universe, making it by far the most plentiful element. Only trace amounts of the free, diatomic element occur naturally on Earth, but hydrogen is plentiful in water when combined with oxygen.

• The term Nonpolar H2 refers to a colorless, odorless gas with extremely weak dispersion forces and a low molar mass due to its basic structure.

In the periodic table, hydrogen does not have a completely acceptable place (Figure 14.1). Hydrogen may fit better in Group 1A(1), 4A(14), or 7A(17) depending on the property: Hydrogen, like the Group 1A(1) elements, has an ns1 outer electron configuration and, most often, a +1 oxidation state.

Unlike alkali metals, however, hydrogen shares its single valence electron with nonmetals rather than transferring it to them.

Furthermore, hydrogen has the highest ionization energy (IE = 1311 kJ/mol) and electronegativity (EN = 2.1) of any alkali metal.

Lithium, on the other hand, has an IE of just 520 kJ/mol and an EN of 1.0, the highest among the alkali metals.

Hydrogen's valence level is half-filled, but with only one electron, and it possesses ionization energy, electron affinity, electronegativity, and bond energy values comparable to those of Group 4A(14) elements (14).

Hydrogen, like the elements of Group 7A(17), exists as diatomic molecules and fills its outer level either via electron sharing or by acquiring one electron from metal to create a 1 ion (hydride, H).

While monatomic halide ions (X) are abundant and stable, H is uncommon and reactive. Furthermore, hydrogen has a lower electronegativity (EN = 2.1) than halogens (whose ENs vary from 4.0 to 2.2) and lacks their three valence electron pairs.

Many covalent hydrides are formed when hydrogen interacts with nonmetals, including CH4, PH3, H2S, and HCl.

The majority are gases, however many boron and carbon hydrides are liquids or solids made up of considerably bigger molecules. Because the other nonmetal has a higher electronegativity, hydrogen has an oxidation number of +1 in most covalent hydrides.

• N2(g) + 3H2(g) ⟶ 2NH3(g) ΔH°rxn = −91.8 kJ

The conditions for producing covalent hydrides are determined by the reactivity of the other nonmetal. For example, hydrogen interacts at high temperatures (400°C) and pressures (250 atm) with stable, triple-bonded N2, and the reaction requires a catalyst to proceed at any realistic speed:

The peculiar behavior of hydrogen is due to its small size. It has a high IE because its electron is very close to the nucleus and there are no inner electrons to protect it from the positive charge.

Because it only has one proton to attract bonding electrons, it has a low EN (for a nonmetal). H will be treated in this chapter as part of either Group 1A(1) or 7A(17), depending on the property under consideration.

When hydrogen reacts with highly reactive metals, such as those in Group 1A(1) and the bigger members of Group 2A(2) (Ca, Sr, and Ba), it creates saltlike hydrides, which are white, crystalline solids consisting of the metal cation and the hydride ion:

• 2Li(s) + H2(g) ⟶ 2LiH(s)

• Ca(s) + H2(g) ⟶ CaH2(s)

In water, H− is a strong base that pulls H+ from surrounding H2O molecules to form H2 and OH−:

• NaH(s) + H2O(l) ⟶ Na+(aq) + OH– (aq) + H2(g)

The hydride ion is also a powerful reducing agent; for example, it reduces Ti(IV) to the free metal:

• TiCl4(l) + 4LiH(s) ⟶ Ti(s) + 4LiCl(s) + 2H2(g)

According to Pauli's exclusion principle and Hund's rule, electrons fill the one ns and three np orbitals.

Atomic size typically reduces when nuclear charge increases and electrons are added to orbitals of the same energy level (same n value), but initial ionization energy and electronegativity often rise (see bar graphs).

As elements go from metals to metalloids to nonmetals, their metallic character declines with increasing nuclear charge.

Except for the inert noble gas, reactivity is strongest at the left and right ends of the period because members of Groups 1A(1) and 7A(17) are just one electron away from achieving a filled outer level.

The bonding between an element's atoms evolve from metallic to covalent in networks, covalent in individual molecules, and none (noble gases exist as separate atoms).

Physical characteristics such as melting temperature, as predicted, change dramatically at the network/molecule border, which occurs between carbon (solid) and nitrogen (gas).

Each element's bond with an active nonmetal evolves from ionic to polar covalent to covalent. Each element's bond with an active metal evolves from metallic to polar covalent to ionic

As the connection between the element and oxygen evolves from ionic to covalent, the acid-base behavior of the common element oxide in water varies from basic to amphoteric to acidic (see the similar trend for the Period 3 elements in Figure 8.25).

Metals reduce strength, whereas nonmetals improve strength. Period 2 oxidation numbers (O.Ns) are equal to the A-group number for Li and Be and the A-group number minus 8 for O and F. Boron has numerous O.N.s, Ne has none, and C and N exhibit all conceivable O.N.s for respective groups.