Atoms, Elements & Compounds

Atomic Structure

• Atoms consist of a minute nucleus where all the positive charge and most of the mass are concentrated, surrounded by negatively charged electrons.
• The mass of the electron is so small it's often ignored.
• The nucleus contains:
  • Protons: positively charged particles.
  • Neutrons: neutral particles.
• Atoms are electrically neutral because the number of positive protons equals the number of negative electrons.

Sub-Atomic Particles

Atomic (Proton) Number (Z)

• The number of protons in the nucleus of an atom.

Mass (Nucleon) Number (A)

• Total number of protons and neutrons in the nucleus of an atom.

Electron Arrangement

• Electrons are arranged in a series of shells (energy levels) around the nucleus.
• Each shell can only contain a limited number of electrons.
• The shell nearest to the nucleus fills first, then the next, and so on.
• The first shell can hold only 2 electrons, the second 8, and the third 8 or 18.

Isotopes

• Different atoms of the same element with the same atomic number but different mass numbers due to a difference in the number of neutrons.
• Examples:
  • $^{235}_{92}U$ (uranium-235)
  • $^{238}_{92}U$ (uranium-238)
  • Carbon isotopes:
    • Carbon-12: 6 protons, 6 neutrons.
    • Carbon-13: 6 protons, 7 neutrons.
    • Carbon-14: 6 protons, 8 neutrons.

Calculating Relative Atomic Mass (Ar)

• $Ar = (Mass \space no \space for \space isotope \space 1 \space x \space its \space$

Elements, Compounds, and Mixtures

• Element: A substance made up of one type of atom (e.g., Aluminum, sulfur).
• Compound: A substance consisting of 2 or more elements chemically combined (e.g., water, carbon dioxide).
• Mixture: Consists of 2 or more elements or compounds which have not been chemically combined (e.g., water and ethanol).

Properties of Metals and Non-Metals

Chemical Bonding

• Noble gases (He, Ne, Ar) are unreactive (stable) because their atoms have full outer shells of electrons.
• Atoms of other elements combine/form chemical bonds to attain a stable electronic structure like that of the nearest noble gas.

Ionic Compounds

• An ionic bond is formed when one or more electrons is transferred from an atom of a metal to an atom of a non-metal.
• Ionic compounds are made up of ions.
• Formation of ions:
  • Sodium (Na) loses one electron to become a sodium ion (Na+), a positive ion (cation).
  • Chlorine (Cl) gains one electron to become a chloride ion (Cl-), a negative ion.
• Ions are electrically charged particles formed from an atom by the loss or gain of one or more electrons.

Dot and Cross Diagrams

• Magnesium oxide (MgO): Mg transfers 2 electrons to O.
• Lithium Fluoride (LiF): Li transfers 1 electron to F.
• Calcium Chloride (CaCl2): Ca transfers 2 electrons, one to each Cl.

Writing Formulas

• Calcium Chloride: $Ca^{2+}Cl^- \rightarrow CaCl_2$
• Calcium Oxide: CaO (not $Ca<em>2O</em>2$)

Giant Ionic Structure (Giant Ionic Lattice)

• A regular arrangement of alternating positive and negative ions in ionic crystals.
• Ions are held strongly together by electrostatic attraction between oppositely charged ions.
• Explains the high melting points of ionic compounds.

Properties of Ionic Compounds

1. Solid at room temperature with high melting and boiling points due to strong electrostatic attractions between ions; requires high energy to break (e.g., NaCl melting point is 800°C).
   • MgO has a higher melting/boiling point than NaCl because of the increased charge of ions, leading to greater forces holding the ions together.
2. Solid ionic compounds do not conduct electricity because ions cannot move.
3. Conduct electricity when melted or dissolved in water because they have free-moving ions.
4. Usually soluble in water but not in organic solvents like ethanol.

Molecules and Covalent Bonding

• In any bond, particles are held together by electrostatic attractions between oppositely charged particles.
• A covalent bond is formed when a pair of electrons is shared between two atoms (both non-metals).
• Covalent bonding is the electrostatic attraction between the bonding pair of electrons (negative) and the nuclei of the atoms involved (positive).

Dot and Cross Diagrams for Covalent Compounds

• A single covalent bond is formed when a pair of electrons is shared.

• Examples:

  • Hydrogen molecule (H2)
  • Hydrogen Chloride molecule (HCl)
  • Chlorine molecule (Cl2)
  • Methane molecule (CH4)
  • Water molecule (H2O)
  • Ethane molecule (C2H6)

• A double covalent bond is formed when two pairs of electrons are shared.

  • Oxygen molecule (O2) - O=O

• Carbon dioxide (CO2)

• Ethene (C2H6)

• A triple covalent bond is formed when three pairs of electrons are shared.

  • Nitrogen molecule (N2) - N≡N

Properties of Covalent Substances (Simple Molecular Structures)

1. Most covalent substances are gases and liquids; few are solids with lower melting and boiling points than ionic compounds.
   • Molecules are held together by weak intermolecular forces, requiring low energy to break.
2. Do not conduct electricity.
3. Less soluble in water but soluble in organic solvents such as ethanol.

Diatomic Molecules

• Molecules consisting of 2 atoms (e.g., H2, O2, N2, Cl2, Br2, I2, F2).

Comparison of Ionic and Covalent Bonding

Properties of Ionic and Covalent Compounds

Allotropes of Carbon

Diamond

• Tetrahedral arrangement of C atoms.
• Each carbon atom is bonded to 4 carbon atoms.
• High melting point: high energy needed to overcome strong covalent bonds.
• Hard: due to strong covalent bonds.
• Used in drillers.

Graphite

• Each carbon atom is bonded to 3 carbon atoms in repeated hexagons.
• Hexagons are arranged in layers joined by weak intermolecular forces.
• High melting point: high energy needed to overcome strong covalent bonds.
• Soft: due to weak intermolecular forces between molecules.
• Used in pencils.

C60 Fullerene

• Each C60 fullerene molecule consists of 60 carbon atoms.
• Molecules are joined together by weak intermolecular forces.
• Low melting point: low energy needed to overcome weak intermolecular forces.
• Soft: due to weak intermolecular forces.
• Poor electrical conductor.

Metallic Bonding

• Giant metallic lattice consisting of positive ions surrounded by a sea of free (mobile) electrons.
• Metallic bond: electrostatic attractions between positive ions and mobile (free) electrons.

Properties of Metals

1. High melting and boiling points: high energy needed to overcome electrostatic attractions between positive ions and negative electrons.
2. Good conductors of electricity: free-moving delocalized electrons.
3. Malleable and ductile: layers of positive ions can slide.