Periodic Trends, Atomic Structure, and Nomenclature: General Chemistry Study Guide

Periodic Trends

Atomic Radius

The atomic radius is the distance from the nucleus to the outermost electron shell of an atom. It is a key periodic property that varies predictably across the periodic table.

• Trend Down a Group: Atomic radius increases as you move down a group due to the addition of electron shells.

• Trend Across a Period: Atomic radius decreases from left to right across a period because of increasing nuclear charge, which pulls electrons closer to the nucleus.

• Order of Atomic Radius (Example): For the elements Mg, Na, P, Si, Ar: $\text{Ar} < \text{P} < \text{Si} < \text{Mg} < \text{Na}$

Example: The atomic radius of potassium (K) is larger than that of calcium (Ca), which is larger than that of germanium (Ge), arsenic (As), and krypton (Kr).

Ionization Energy

Ionization energy is the energy required to remove an electron from a gaseous atom or ion. It reflects how strongly an atom holds onto its electrons.

• Trend Down a Group: Ionization energy decreases as you move down a group because electrons are farther from the nucleus and more shielded.

• Trend Across a Period: Ionization energy increases from left to right across a period due to increasing nuclear charge and decreasing atomic radius.

• Order of First Ionization Energies (Example): $\text{Cl} > \text{S} > \text{P} > \text{Si} > \text{Al}$

• Second Ionization Energy: The second ionization energy is always higher than the first, and is especially high for atoms that achieve a noble gas configuration after losing one electron.

Example: Neon (Ne) has a much higher ionization energy than sodium (Na) or magnesium (Mg).

Electron Affinity

Electron affinity is the energy change when an atom gains an electron to form a negative ion. It is most negative for elements that readily accept electrons.

• Trend Across a Period: Electron affinity becomes more negative from left to right across a period.

• Trend Down a Group: Electron affinity generally becomes less negative down a group.

• Most Negative Electron Affinity: Halogens (e.g., chlorine) have the most negative electron affinities.

Example: Chlorine has a more negative electron affinity than oxygen or sulfur.

Effective Nuclear Charge (Zeff)

The effective nuclear charge is the net positive charge experienced by valence electrons. It is affected by the number of protons and the shielding effect of inner electrons.

• Formula: $Z_{\text{eff}} = Z - S$ where $Z$ is the atomic number and $S$ is the shielding constant.

• Trend: Zeff increases across a period and slightly increases down a group.

Electron Configuration

Ground-State Electron Configuration

Electron configuration describes the arrangement of electrons in an atom's orbitals. The ground-state configuration follows the Aufbau principle, Pauli exclusion principle, and Hund's rule.

• Example: Argon (Ar, atomic number 18): $[\text{Ne}]3s^2 3p^6$

• Silver (Ag): $[\text{Kr}]4d^{10}5s^1$

• Oxide Ion (O2−): $1s^2 2s^2 2p^6$

Isoelectronic Series

An isoelectronic series consists of atoms and ions with the same number of electrons but different nuclear charges.

• Order of Increasing Radius: In an isoelectronic series, the species with the lowest nuclear charge has the largest radius.

• Example Table:

Species	Nuclear Charge (Z)	Relative Radius
O2−	8	Largest
F−	9	Smaller
Na+	11	Smallest

Periodic Table Organization

Periods and Groups

The periodic table is organized into horizontal rows called periods and vertical columns called groups.

• Periods: Indicate the principal energy level of valence electrons.

• Groups: Elements in the same group have similar chemical properties.

Classification of Elements

• Alkali Metals: Group 1A (e.g., Na, K), highly reactive, form 1+ ions.

• Alkaline Earth Metals: Group 2A (e.g., Mg, Ca), form 2+ ions.

• Halogens: Group 7A (e.g., Cl, Br), form 1− ions.

• Chalcogens: Group 6A (e.g., O, S), form 2− ions.

• Noble Gases: Group 8A (e.g., Ne, Ar), inert, full valence shell.

• Metalloids: Elements with properties intermediate between metals and nonmetals (e.g., Si, Ge, As).

Chemical Nomenclature

Binary Ionic Compounds

Binary ionic compounds consist of a metal cation and a nonmetal anion. The name is formed by stating the cation first, followed by the anion with an '-ide' ending.

• Example: Na2O is sodium oxide.

• Example: CaH2 is calcium hydride.

• Example: FeBr3 is iron(III) bromide.

Polyatomic Ions and Compounds

Polyatomic ions are charged species composed of two or more atoms covalently bonded.

• Common Polyatomic Ions:

Ion	Formula	Name
Sulfate	SO42−	Sulfate
Hydronium	H3O+	Hydronium
Ammonium	NH4+	Ammonium
Carbonate	CO32−	Carbonate

• Example: (NH4)2CO3 is ammonium carbonate.

• Example: Ca3(PO4)2 is calcium phosphate.

Acids and Molecular Compounds

• Acids: Named based on the anion present. For example, HNO3 is nitric acid, H2SO4 is sulfuric acid.

• Molecular Compounds: Use prefixes to indicate the number of atoms (e.g., N2O5 is dinitrogen pentoxide).

Atomic Structure

Subatomic Particles

Atoms are composed of protons, neutrons, and electrons.

• Protons: Positively charged, reside in the nucleus.

• Neutrons: Neutral, reside in the nucleus.

• Electrons: Negatively charged, occupy orbitals around the nucleus.

Rutherford Model

The Rutherford nuclear-atom model describes the atom as having a small, dense nucleus containing protons and neutrons, with electrons in the surrounding space.

• Key Point: Most of the atom's mass is concentrated in the nucleus.

Cathode Rays

Cathode rays are streams of electrons observed in vacuum tubes. They are deflected away from negatively charged plates because electrons are negatively charged.

Additional Properties and Examples

Reactivity and Storage of Elements

Some elements, such as sodium, are highly reactive and must be stored under oil to prevent reaction with air or water.

• Example: Sodium is solid at room temperature, forms oxides easily, reacts with water to release hydrogen gas, and must be stored submerged in oil.

Diatomic Molecules

Certain elements exist as diatomic molecules in their elemental form.

• Examples: H2, N2, O2, F2, Cl2, Br2, I2

Metalloids

Metalloids have properties intermediate between metals and nonmetals. Examples include silicon (Si), germanium (Ge), arsenic (As), and selenium (Se).

Summary Table: Key Periodic Trends

Property	Trend Down a Group	Trend Across a Period
Atomic Radius	Increases	Decreases
Ionization Energy	Decreases	Increases
Electron Affinity	Less negative	More negative

Additional info: Some context and examples have been inferred and expanded for clarity and completeness, including standard nomenclature rules, periodic trends, and atomic structure explanations.