Study Guide: Atoms, Elements, and Ionic/Molecular Compounds (Chapters 4 & 6, Sections 1-3)

Atoms, Elements, and the Periodic Table

Elements and Symbols

An element is a pure substance that cannot be broken down into simpler substances by chemical means. Each element is represented by a unique one- or two-letter chemical symbol (e.g., H for hydrogen, O for oxygen). The periodic table organizes all known elements by increasing atomic number.

• Element: Fundamental substance; cannot be decomposed further.

• Chemical Symbol: One- or two-letter abbreviation (e.g., Na for sodium).

• Compound: Combination of two or more elements (e.g., H2O).

• Atomic Number: Number of protons in the nucleus; uniquely identifies the element.

• Example: Element #16 is sulfur (S).

The Periodic Table

The periodic table is a chart of elements arranged by increasing atomic number. It consists of 7 periods (horizontal rows) and 18 groups (vertical columns). Elements are classified into metals, nonmetals, and metalloids, each with distinct properties.

• Periods: Horizontal rows (1-7).

• Groups/Families: Vertical columns (1-18).

• Common Group Names:

  • Group 1: Alkali metals

  • Group 2: Alkaline earth metals

  • Group 17: Halogens

  • Group 18: Noble gases

  • Transition metals: Groups 3-12

  • Representative elements: Groups 1, 2, 13-18

  • Lanthanides and actinides: Separate rows below main table

• Element Categories:

  • Metals: Left side; shiny, malleable, good conductors.

  • Nonmetals: Right side; dull, brittle, poor conductors.

  • Metalloids: Border staircase; intermediate properties.

• Example: Aluminum (Al) is a metal; chlorine (Cl) is a nonmetal; silicon (Si) is a metalloid.

The Atom

An atom is the smallest unit of an element that retains its chemical properties. Atoms consist of a nucleus (protons and neutrons) and electrons in surrounding orbitals. The number of atoms is conserved in chemical reactions (conservation of mass).

• Proton: 1+ charge, 1 amu, in nucleus.

• Neutron: 0 charge, 1 amu, in nucleus.

• Electron: 1- charge, ~1/1850 amu, outside nucleus.

• Like charges repel; opposite charges attract.

• Neutral atom: Number of protons = number of electrons.

• Atomic mass unit (amu): Used for microscopic masses; grams for macroscopic amounts.

• Example: Carbon atom: 6 protons, 6 neutrons, 6 electrons.

• Additional info: The atom is mostly empty space; most mass is in the nucleus; electrons occupy most of the volume.

Atomic Number and Mass Number

The atomic number is the number of protons in an atom and determines the element's identity. The mass number is the sum of protons and neutrons. The average atomic mass (with decimals) is the weighted average of all isotopes.

• Atomic Number (Z): Number of protons.

• Mass Number (A): Number of protons + number of neutrons.

• Number of neutrons: $\text{Neutrons} = \text{Mass Number} - \text{Atomic Number}$

• Average atomic mass: Weighted average of isotopes, in amu.

• Example: Oxygen-18 ($^{18}\text{O}$): 8 protons, 10 neutrons, mass number 18.

Isotopes and Atomic Mass

Isotopes are atoms of the same element with different numbers of neutrons. Most elements have multiple isotopes, which contribute to the average atomic mass.

• Isotope: Same number of protons, different number of neutrons.

• Example: $^{18}\text{O}$ has 8 protons, 10 neutrons.

Electron Energy Levels

Electrons occupy quantized energy levels in atoms. Electromagnetic radiation (light) is both a wave and a particle (photon). The energy of a photon is proportional to its frequency. Atomic spectra show discrete lines, corresponding to electron transitions between energy levels.

• Photon: Particle of light; energy is quantized.

• Energy of photon: $E = h\nu$ (where $h$ is Planck's constant, $\nu$ is frequency)

• Atomic spectrum: Discrete lines; each line corresponds to a specific electron transition.

• Energy levels: Numbered $n=1,2,3,...$; higher $n$ means higher energy and larger orbital.

• Sublevels: Each energy level has $n$ sublevels: s, p, d, f.

• Orbital shapes: s (sphere), p (dumbbell), d (cloverleaf), f (complex).

• Maximum electrons per orbital: 2 (with opposite spins).

• Electron transitions: Absorbing a photon moves electron up; emitting a photon moves electron down.

• Example: n=2 has one s and three p orbitals.

Electron Configurations

Electron configuration describes the arrangement of electrons in orbitals. Orbital diagrams use boxes and arrows; condensed configurations use noble gas notation.

• Orbital diagram: Boxes for orbitals, arrows for electrons (spin up/down).

• Electron configuration: e.g., $1s^2 2s^2 2p^6$

• Condensed configuration: e.g., $[\text{He}] 2s^2 2p^6$

• Example: Sodium (Na): $1s^2 2s^2 2p^6 3s^1$ or $[\text{Ne}] 3s^1$

Trends in Periodic Properties

Periodic trends describe how properties change across periods and groups. Valence electrons are the outermost electrons, important for chemical bonding. Atomic radius, ionization energy, and metallic character vary predictably across the table.

• Valence electrons: Outer-shell electrons; Group 1 has 1, Group 2 has 2, etc.

• Lewis structure: Shows valence electrons as dots around element symbol.

• Atomic radius: Increases left and down.

• Ionization energy: Increases right and up.

• Metallic character: Increases left and down.

• Example: Potassium (K) has 1 valence electron; Lewis structure: K•

Ionic and Molecular Compounds

Ionic vs. Molecular Compounds

Ionic compounds consist of metal cations and nonmetal anions, held together by ionic bonds (electrostatic attraction). Molecular compounds consist of nonmetals sharing electrons via covalent bonds.

• Ionic compound: Metal + nonmetal; electrons transferred.

• Molecular compound: Nonmetals only; electrons shared.

• Cation: Positive ion; formed by losing electrons (metals).

• Anion: Negative ion; formed by gaining electrons (nonmetals).

• Octet rule: Atoms gain/lose electrons to achieve noble gas configuration.

• Common ion charges:

  • Group 1: 1+

  • Group 2: 2+

  • Group 3: 3+

  • Group 15: 3-

  • Group 16: 2-

  • Group 17: 1-

  • Group 18: No ions (noble gases)

• Example: Sodium (Na) forms Na+; chlorine (Cl) forms Cl-.

Forming Ionic Compounds

Ionic compounds are formed by combining cations and anions in ratios that result in a neutral compound. The charges must balance.

• Neutral compound: Total positive charge = total negative charge.

• Crisscross method: Use ion charges to determine subscripts.

• Examples:

  • NaCl: Na+ + Cl- → NaCl

  • MgCl2: Mg2+ + 2Cl- → MgCl2

  • Na2O: 2Na+ + O2- → Na2O

Naming and Writing Ionic Formulas

To name ionic compounds, the cation retains its element name, and the anion is named with the suffix "-ide." Transition metals may have multiple charges, indicated by Roman numerals in the name.

• Cation: Element name (e.g., sodium).

• Anion: Element root + "-ide" (e.g., chloride).

• Transition metals: Use Roman numerals to indicate charge (e.g., Iron(II) chloride = FeCl2, Iron(III) chloride = FeCl3).

• Example: CuBr = copper(I) bromide; CuBr2 = copper(II) bromide.

Table: Common Ion Charges by Group

Group	Common Ion Charge	Example Element
1	1+	Na+
2	2+	Mg2+
3	3+	Al3+
15	3-	N3-
16	2-	O2-
17	1-	Cl-
18	0	Ne (no ion)

Summary

• Understand the structure and properties of atoms and elements.

• Use the periodic table to classify elements and predict properties.

• Describe electron arrangements and periodic trends.

• Distinguish between ionic and molecular compounds.

• Name and write formulas for ionic compounds, including transition metals.