atoms
Atoms and Atomic Bonding
Atom Structure
Atoms consist of:
Nucleus: Central part containing:
Protons: Carry a positive charge (+).
Neutrons: Neutral particles that add mass but have no charge.
Electrons: Negatively charged particles (–) that orbit the nucleus in energy shells.
Protons define the atomic number, which is unique to each element.
Neutrons contribute to the atomic mass.
Electrons play a crucial role in chemical bonding.
Electron Shells
Electron shells are layers of electrons surrounding the nucleus.
Each shell can hold a maximum number of electrons:
The first shell holds 2 electrons.
The second shell can hold up to 8 electrons.
The outermost shell, known as the valence shell, is key to determining the atom's chemical behavior.
Valence Electrons
Valence electrons are the electrons in the outermost shell that participate in chemical bonding.
Atoms with incomplete valence shells are chemically reactive and tend to form bonds with other atoms to achieve stability.
Ion Formation
Atoms can become ions by gaining or losing electrons:
Positive Ions (Cations): Atoms with few valence electrons (e.g., sodium) donate electrons.
Negative Ions (Anions): Atoms with nearly full valence shells (e.g., chlorine) accept electrons.
When the number of protons equals the number of electrons, the atom is neutral.
Bond Types
Ionic Bonds:
Formed by the transfer of electrons between atoms.
Example: Sodium chloride (NaCl).
Weaker bonds that can be easily broken in water.
Covalent Bonds:
Formed by the sharing of electrons between atoms.
Example: Quartz (SiO₂).
Stronger and more stable compared to ionic bonds.
Periodic Table Use
The atomic number corresponds to the number of protons in the atom.
The electron configuration of an atom indicates the number of valence electrons, allowing predictions on whether the atom will donate, accept, or share electrons.
Example:
Oxygen has an atomic number of 8, which means it has 6 valence electrons.
Consequently, it tends to form covalent bonds, requiring 2 additional electrons to fill its valence shell.
Mineral Identification
Definition
A mineral is defined as a naturally occurring, inorganic solid with:
A definite chemical composition.
A crystalline structure.
Rocks are aggregates of multiple minerals.
Properties Used in Identification
Color:
The visible hue of a mineral.
It can be unreliable due to the presence of impurities.
Luster:
Refers to how light reflects off the mineral's surface.
Categories include:
Metallic: Shiny, resembling metal.
Non-metallic: Appears glassy, pearly, or earthy.
Metallic luster arises from free electrons reflecting light.
Streak:
The color of a powdered mineral when tested on a porcelain plate.
More reliable than the surface color.
Hardness:
The resistance of a mineral to scratching, measured using the Mohs scale (1–10).
Cleavage:
The tendency of a mineral to break along flat planes of weakness.
Fracture:
An irregular break occurring when no cleavage planes are present.
Specific Gravity:
The density of a mineral in comparison to water.
Reaction to Acid:
Carbonate minerals will fizz when exposed to dilute hydrochloric acid (HCl).
Mohs Hardness Scale Examples:
Talc = 1 (softest)
Fingernail ~2.5
Penny ~3
Steel Nail ~4.5
Glass ~5.5
Porcelain Plate ~6.5
Diamond = 10 (hardest)
Color Variation
Caused by impurities known as chromophores.
Example:
Copper can produce green/blue colors (e.g., malachite, azurite).
Iron/Magnesium leads to dark colored minerals (e.g., olivine, pyroxene).
Iron may yield red/yellow hues (e.g., hematite, limonite).
Cleavage vs. Fracture
Cleavage results in flat, smooth surfaces.
Fracture results in irregular or rough surfaces.
Specific Gravity
Defined as the ratio of the mineral's density to the density of water.
Heavy minerals like galena and gold exhibit high specific gravity.
Mineral Groups & Silicate Structures
Phosphorus Source
Apatite, a phosphate mineral, releases phosphorus upon weathering, which is essential for life.
Silicates
Characterized by SiO₄ tetrahedra, where a silicon atom is bonded to four oxygen atoms.
Carbonates
Composed of CO₃ groups (e.g., calcite, dolomite).
Other Mineral Groups:
Oxides: Contain oxide ions (O²⁻) bonded to metals (e.g., hematite).
Sulfates: Comprise sulfate ions (SO₄²⁻), such as gypsum.
Phosphates: Contain phosphate ions (PO₄³⁻), exemplified by apatite.
Halides: Composed of halide ions (Cl⁻, F⁻), e.g., halite and fluorite.
8 Most Common Crustal Elements
Oxygen
Silicon
Aluminum
Iron
Calcium
Sodium
Potassium
Magnesium
Silicate Tetrahedron
The silicate tetrahedron is the fundamental building block of silicate minerals with an overall charge of –4.
This charge is neutralized by bonding with cations such as Fe²⁺, Mg²⁺, Ca²⁺, Na⁺, and K⁺.
Progression of Silicate Structures:
Isolated Tetrahedra: Example – Olivine.
Single Chains: Example – Pyroxene.
Double Chains: Example – Amphibole.
Sheets: Examples – Micas, clay minerals.
Frameworks: Examples – Quartz, feldspar.
Trends with Si/O Ratio:
A higher Si/O ratio results in more covalent bonds, lower density, and greater stability.
Isolated tetrahedra possess a low Si/O ratio, yielding high density and less stability.
Framework silicates feature a high Si/O ratio, resulting in low density and very high stability.
Bonding & Stability
The presence of more covalent bonds enhances the strength and resistance of minerals to weathering.
Mafic vs. Felsic Silicates:
Mafic Silicates:
Characterized by dark, dense minerals (e.g., olivine, pyroxene, amphibole).
Felsic Silicates:
Comprised of lighter, less dense minerals (e.g., quartz, feldspar).
Non-Silicates Examples
Calcite
Halite
Gypsum
Hematite
Most Common Minerals on Earth’s Surface
1st: Feldspar
2nd: Quartz