Bonding Notes FCP
Electron Shells and Energy Levels
Bohr Energy Levels
Drawings of energy levels for hydrogen (H), lithium (Li), and sodium (Na).
Atomic Orbitals: Mathematical descriptions of likely electron locations in atoms/molecules.
1s Orbital:
First energy level.
Spherically symmetric around the nucleus.
Shape & Structure:
Each orbital can hold 2 electrons.
p Orbitals:
Begin at the second energy level (2s, 2p).
Three equivalent orbitals: px, py, pz at each energy level.
Atomic Orbitals and Electron Configuration
Higher Energy Levels:
3rd level: 3s, 3p (px, py, pz), and 5 d orbitals (total of 9 orbitals).
4th level: 4s, 4p, 4d, and 7 f orbitals (total of 16 orbitals).
Shorthand Electron Configuration:
Use noble gas symbols to simplify lengthy configurations.
Energy Level Transitions
Movement of Electrons:
Electrons can transition between energy levels through energy gain/loss.
Atomic Line Spectrum:
Unique color signatures for different atoms (e.g., H2, O2, He).
Lewis Dot Symbols and Valence Electrons
Lewis Dot Symbols:
Symbol of an element with dots representing valence electrons.
Valence Electrons:
Electrons in the last energy level, represented on four sides of the atomic symbol.
Maximum of 2 electrons per side.
Octet Rule:
Atoms gain/lose electrons to achieve a full outer shell of 8 electrons (similar to noble gases).
Ionic Compounds and Bonds
Ionic Compounds:
Composed of cations and anions with an overall net charge of zero.
Ionic Bond:
Complete transfer of electrons between atoms (examples: NaCl, MgCl2).
ΔEN > 1.7.
Properties of Salts:
Formed from ionic bonds between metals and nonmetals.
Characteristics: soluble in water, conduct electricity, high boiling, and low freezing points.
Naming Ionic Compounds
Binary Compounds:
Compounds containing only two elements.
Named by stating the metal first, followed by the non-metal with the suffix -ide (e.g., NaCl as sodium chloride).
Practice Naming:
Exercise for writing names from formulas (Li2O, KI, Na2S).
Oxidation Numbers and Cross-Over Rule
Oxidation Number:
Charge of an atom when involved in a bond.
Cross-Over Rule for Ionic Bonds:
Method to determine subscripts in formulas:
Write symbols and charges.
Cross-over charges without signs for subscripts.
Find the highest common factor.
Divide crossed charges by this factor.
Example: MgO, Li2S, K3P.
Transition Metals and Naming
Transition Metals:
Can have multiple charges, necessitating Roman numerals in names (e.g., iron (II) chloride for FeCl2).
Naming Rules:
Always place the metal first in the name.
Determine charges to assign Roman numerals.
Examples:
Osmium (III) sulfide from Os2O3 due to charge balancing.
Covalent Bonds
Covalent Bonds:
Formed by sharing pairs of electrons between atoms.
Types of Covalent Bonds:
Pure Covalent Bonds: Equal sharing of electrons (ΔEN = 0).
Polar Covalent Bonds: Unequal sharing (ΔEN between 0.1 - 1.6).
Lewis Dot Structures and Molecular Polarity
Drawing Lewis Structures:
Sum valence electrons and find electron pairs.
Organize atoms around a central atom.
Distribute electron pairs through bonds.
Molecular Polarity:
Molecules can be nonpolar even if they contain polar bonds if the bond dipoles cancel out.
Example: Water (polar due to bent geometry and O-H polar bonds).
Naming Covalent Compounds
Naming Binary Covalent Compounds:
Use prefixes to indicate quantity (e.g., di- for two).
Examples:
Name compounds like PH3, H2O, SO.
Write formulas for names like sulfur dioxide or dihydrogen oxide.
Metallic Bonds and Alloys
Metallic Bonds:
Formed by a 'sea of electrons' surrounding metal nuclei.
Electrons are mobile, allowing layers of atoms to slide past one another without breaking bonds.
Alloys:
Mixtures of two or more metals or metals and non-metals.
Exhibits different properties than individual components (e.g., steel from iron and carbon).
Common Alloys: steel, brass, sterling silver.