Atomic Structure and the Periodic Table

Topic 1 - Atomic Structure and the Periodic Table

Atoms

  • Definition: All substances are composed of atoms, which are extremely small, undetectable even with a typical microscope.

  • Size Example: A 50 pence coin comprises approximately 77.4\cdot10^{18} atoms, indicating the massive quantity even in a small object.

Composition of Atoms

General Structure

  • Atoms consist of three principal subatomic particles: proton-s, neutrons, and electrons.

  • Atom Size:

    • Radius of an atom: approximately 1\cdot10^{-10} m (0.1 nanometres).

    • Radius of the nucleus: around 1\cdot10^{-14} m, about 1/10,000th of the atom's radius.

Components

  1. The Nucleus

    • Located at the center of the atom.

    • Contains protons (positively charged) and neutrons (neutral charge).

    • Carries nearly the entire mass of the atom.

    • Nucleus charge: positive due to protons.

  2. Electrons

    • Revolve around the nucleus in shells.

    • Charge: negatively charged and extremely lightweight; often considered with negligible mass.

    • The volume of the electron orbits influences the atom's overall size.

    • Movement creates a shell structure like:

      • Electron - Relative mass: very small (often approximated as zero) - Charge: -1

      • Proton - Relative mass: 1 - Charge: +1

      • Neutron - Relative mass: 1 - Charge: 0

Neutral Atoms and Ions

  • Neutral Atoms: The atom has no overall charge as the number of electrons equals protons.

  • Ions: Formed when atoms lose/gain electrons, leading to an imbalance between protons and electrons. E.g., an ion with a 2- charge has two more electrons than protons.

Atomic and Mass Numbers

  • Definitions:

  1. Atomic Number (Z): Number of protons in an atom.

  2. Mass Number (A): Total number of protons and neutrons in an atom.

  3. Finding Neutrons: N = A - Z (where N = number of neutrons, A = mass number, Z = atomic number).

Example of Nuclear Symbol: Sodium

  • Notation: ^{23}_{11}Na

  • Mass Number = 23, Atomic Number = 11.

Elements

  • Definition: Substances containing atoms with the same number of protons.

  • Element Characteristics: Different number of neutrons/electrons do not change the atomic identity but affect physical properties.

  • There are roughly 100 known elements.

  • Each element symbol represents atoms of that element succinctly using a shorthand format (e.g., C for Carbon, Na for Sodium).

Isotopes

  • Definition: Variants of the same element; same number of protons but different neutrons.

  • Example:

    • Carbon-12: 6 protons, 6 neutrons, mass number 12.

    • Carbon-13: 6 protons, 7 neutrons, mass number 13.

  • Relative Atomic Mass: Average mass considering the abundance of all isotopes of an element. Computed using:

    • \text{Relative Atomic Mass} = \frac{\text{sum of (isotope abundance} \times \text{isotope mass number)}}{\text{sum of abundances of all isotopes}}

Example Calculation

  • For Copper (Cu) with isotopes Cu-63 (69.2% abundance) and Cu-65 (30.8% abundance):

    • = \frac{(69.2 \times 63) + (30.8 \times 65)}{100} = 63.616 \text{ (round to 63.6)}

Compounds

  • Definition: Substances formed from two or more elements, where atoms bond chemically.

  • Atoms combine in fixed proportions.

  • Bonds involve electron sharing (covalent) or electron transfer (ionic).

Types of Bonding

  1. Ionic Bonding: Generally between metals and non-metals, involves ion formation; e.g., NaCl.

  2. Covalent Bonding: Involves sharing electrons, typical for nonmetals; e.g., HCl.

Properties of Compounds

  • Compounds have vastly different properties from their constituent elements (e.g., iron (shiny metal) when combined with sulfur (yellow powder) forms dull iron(II) sulfide).

Representation of Compounds

  • Chemical Formula: Indicates the types and numbers of atoms in compounds (e.g., CO₂ for carbon dioxide, H₂SO₄ for sulfuric acid).

  • Brackets indicate multiplicative relationships (e.g., Ca(OH)₂ implies 1 calcium, 2 oxygen, and 2 hydrogen).

Examples of Common Compounds

  1. Carbon Dioxide (CO₂): 1 Carbon, 2 Oxygens.

  2. Water (H₂O): 2 Hydrogens, 1 Oxygen.

  3. Sodium Chloride (NaCl): 1 Sodium, 1 Chlorine.

Chemical Changes and Equations

  • Word Equations: Describe chemical reactions using names (e.g., methane + oxygen → carbon dioxide + water).

  • Symbol Equations: Provide shorthand for reactions, balancing is crucial to show conservation of mass.

  • Example: 2Mg + O_2 \rightarrow 2MgO

Balancing Equations

  • Must have equal numbers of each atom on both sides. Strategies involve adjusting coefficients in front of the formulas.

  • Practice aids proficiency in balancing equations efficiently.

The History of Atomic Theory

  • Early Interpretations: John Dalton (1810s): Atoms as solid spheres; J.J. Thomson (1897): Discovered electrons, forming the Plum Pudding Model.

  • Rutherford's Model: Alpha-particle scattering demonstrated a nucleus, leading to the modern nuclear model of the atom.(1909)

  • Bohr's Model: Electrons in defined shells; stability led to advancements in understanding atomic structure.

Electronic Structure

  • Electron Shell Rules:

  1. Electrons occupy shells (energy levels).

  2. Fill lowest energy shells first.

  3. Max electrons: First shell - 2, Second shell - 8, Third shell - 8.

  • Electron Configurations: Illustrated as shell numbers (2,8,1 for Sodium).

The Periodic Table

  • Structure: Elements organized by increasing atomic number, demonstrating periodic trends in properties (e.g., metallic vs non-metallic).

  • Groups: Vertical columns determined by the number of electrons in the outer shell influence reactivity.

  • Periods: Horizontal rows denote new electron shells.

Metals and Non-Metals

  • Metals: Typically located on the left, form positive ions, exhibit metallic properties.

  • Non-Metals: Located to the right, generally do not form positive ions or exhibit metallic properties.

  • Transition Metals: Located in the center, exhibit typical metallic traits, can have multiple oxidation states.

Group 1: Alkali Metals

  • Characteristics: Highly reactive, low density, softer than most metals, forms ionic compounds.

  • Reactivity Trend: Increases down the group; easier to lose the outer electron.

CReactions of Alkali Metals

  1. With Water (e.g., 2Na + 2H2O \rightarrow 2NaOH + H2)

  2. With Chlorine (e.g., 2Na + Cl_2 \rightarrow 2NaCl)

  3. With Oxygen: Varying metal oxide forms.

Group 7: Halogens

  • Properties: Reactive, non-metals, form colored vapors and diatomic molecules.

  • Reactivity Trend: Decreases down the group. Harder to gain electrons

Group 0: Noble Gases

  • Characteristics: Inert, non-reactive, colorless gases, full outer electron shells.

  • Boiling Points: Increase down the group due to greater electron count leading to stronger intermolecular