GACE Chemistry (6–12) (712)

Dmitri Mendeleev organized elements by increasing atomic mass in 1869. He noticed periodic repetition of properties and arranged elements into rows and columns. He left gaps for undiscovered elements and predicted their properties. His predictions proved remarkably accurate when gallium, scandium, and germanium were discovered.

Henry Moseley reorganized the periodic table by atomic number in 1913. X-ray experiments revealed that atomic number, not atomic mass, determines elemental properties. This resolved anomalies such as tellurium and iodine being out of order by mass.

The modern periodic table arranges elements by increasing atomic number. Horizontal rows are called periods and numbered 1 through 7. Vertical columns are called groups or families and numbered 1 through 18. Elements in the same group have similar chemical properties due to identical valence electron configurations.

Group 1 elements are alkali metals. They have one valence electron and are highly reactive. They form +1 cations and react vigorously with water. Examples include lithium, sodium, and potassium.

Group 2 elements are alkaline earth metals. They have two valence electrons and are reactive but less so than alkali metals. They form +2 cations. Examples include magnesium, calcium, and barium.

Groups 3 through 12 are transition metals. They have partially filled d orbitals and exhibit variable oxidation states. They often form colored compounds and can act as catalysts. Examples include iron, copper, and zinc.

Group 17 elements are halogens. They have seven valence electrons and are highly reactive nonmetals. They form -1 anions and exist as diatomic molecules. Examples include fluorine, chlorine, and bromine.

Group 18 elements are noble gases. They have full valence shells and are chemically inert under normal conditions. They exist as monatomic gases. Examples include helium, neon, and argon.

The lanthanides are the first row of the f-block, elements 57-71. The actinides are the second row of the f-block, elements 89-103. These inner transition metals are placed below the main table for compactness.

Atomic radius measures the size of an atom from the nucleus to the outermost electron orbital. Trends across a period show decreasing atomic radius from left to right. This occurs because protons increase while electrons add to the same shell, increasing effective nuclear charge and pulling electrons closer.

Trends down a group show increasing atomic radius. This occurs because each period adds a new principal energy level, placing electrons farther from the nucleus. The shielding effect of inner electrons also increases.

Cations are smaller than their parent atoms because removing electrons reduces electron-electron repulsion and the remaining electrons are pulled closer to the nucleus. Anions are larger than their parent atoms because adding electrons increases electron-electron repulsion and expands the electron cloud.

Ionization energy is the energy required to remove an electron from a gaseous atom or ion. First ionization energy removes the outermost electron. Successive ionization energies remove additional electrons.

Trends across a period show increasing ionization energy from left to right. Electrons are held more tightly due to higher effective nuclear charge. Exceptions occur at Groups 3 and 16 due to orbital stability considerations.

Trends down a group show decreasing ionization energy. Valence electrons are farther from the nucleus and more shielded. They require less energy to remove.

Electron affinity is the energy change when an electron is added to a gaseous atom. Negative values indicate energy is released and the process is favorable. Positive values indicate energy must be added.

Trends across a period generally show increasingly negative electron affinity from left to right. Atoms more readily accept electrons when approaching a full valence shell. Group 17 elements have the most negative electron affinities.

Trends down a group generally show less negative electron affinity. Larger atoms have valence orbitals farther from the nucleus. Added electrons experience less attraction.

Exceptions include Group 2 and Group 18 elements, which have near-zero or positive electron affinities. Their filled s subshells or full valence shells make electron addition unfavorable.

Electronegativity measures an atom's ability to attract electrons in a chemical bond. The Pauling scale ranges from 0.7 for cesium to 4.0 for fluorine. Higher values indicate stronger electron attraction.

Trends across a period show increasing electronegativity from left to right. Nonmetals attract bonding electrons more strongly than metals. Fluorine is the most electronegative element.

Trends down a group show decreasing electronegativity. Larger atoms have valence electrons farther from the nucleus. Bonding electrons experience less attraction.

Electronegativity differences between atoms determine bond character. Differences less than 0.5 indicate nonpolar covalent bonds. Differences between 0.5 and 1.7 indicate polar covalent bonds. Differences greater than 1.7 generally indicate ionic bonds.

Metallic character refers to how readily atoms lose electrons to form positive ions. Metals exhibit luster, conductivity, malleability, and ductility.

Trends across a period show decreasing metallic character from left to right. Elements transition from metals through metalloids to nonmetals.

Trends down a group show increasing metallic character. Valence electrons are more easily lost from larger atoms with lower ionization energies.

Metalloids border the staircase line on the periodic table. They exhibit properties intermediate between metals and nonmetals. Examples include boron, silicon, germanium, arsenic, antimony, and tellurium.