2.1: Types of Chemical Bonds🔗

electronegativity: measure of an atom(or a group of atoms) to attract shared electrons

• Fluorine is the most electronegative element

partial charges: created in a polar covalent bond when one atom is more electronegative than another

• region with high electron density will have a partial negative charge while a region with low density will have a partial positive charge

• higher electronegativity=partial negative

  lower electronegativity=partial positive

Ionic bonds:

• Ionic interactions occur between metal and nonmetal atoms when they lose or gain electrons to form ions

• coulombic or electrostatic attraction🧲

• Stronger when the charges are larger and the ions are smaller

• Properties:

  • forms crystals🔮

  • high melting/boiling points

  • hard & brittle

  • conducts electricity when dissolved

  • good insulators as a solid

    

Metallic bonds:

• Occurs between metal atoms🪙

• Due to multiple metallic cations being attracted to a delocalized sea of valence electrons

• IMF is stronger💪 when there are smaller metallic cations and more valence electrons

• Properties:

  • shiny✨

  • malleable & ductile

  • conducts heat🔥 & electricity

  • metallic oxides are basic and ionic

  • lose electrons to form cations

    

Covalent bonds:

• Electrons are shared between two or more atoms(typically nonmetals)

• Properties:

  • non-lustrous, various colors

  • brittle, hard or soft

  • poor conductors

  • nonmetallic oxides are acidic and covalent

  • form anions by gaining electrons

• Polar covalent bond🐻‍❄:

  • when atoms are shared unequally in a covalent bond

    

  Nonpolar covalent bond🚫🐻‍❄:

  • when atoms are shared equally in a covalent bond

    

2.2: Intramolecular Force & Potential Energy🔋

Covalent bonds:

• Can be single, double, triple(or an average if there are resonance structures)

• Occur at the lowest energy state

• Happens when the attraction between the nuclei is greatest for the shared electrons and repulsions between electrons and nuclei is the least

• If atoms are too close the nuclei will repel, if atoms are too far apart the attraction is not enough to hold them together

  

Bond enthalpy: the energy required to break a bond, or the energy released when a bond is formed

• larger radii increases the bond length

  longer bond length decreases the bond energy

• more electrons & shorter bond length = greater coulombic attraction

Lattice energy: the energy to separate ions in ionic compounds

• larger charges = higher lattice energy

  bigger radii = smaller lattice energy

• increasing the bond order = increasing the bond energy

2.3: Structure of Ionic Solids🪨

Ionic solids:

• Consist of cations and anions

• molecules are held together by coulombic forces

• higher ion charges = stronger bonds

  larger atoms = weaker bonds

• properties:

  • nonvolatile & high melting points

    • ionic bonds need to be broken to melt the solid, which separates oppositely charged particles

  • do not conduct electricity🚫⚡

    • charged ions are fixed in place

    • when melted or dissolved, ions are free to move, which enables electrical conduction⚡

  • many are soluble in polar solvents and insoluble in nonpolar

    

2.4: Structure of Metals & Alloys🪙

Metals:

• composed of cations that are embedded in delocalized sea of valence electrons

• electrons do not stay with one atom, move throughout the entire substance

• cations and electrons are attracted through coulombic attraction

• # of valence electrons determines amount of electrons in the delocalized sea of electrons

• increased charge & increased # of electrons = greater attractions

  decreased ionic radius = increased attractions

• Alloys:

  • Mixtures of metals🪙

  • interstitial alloys: small atoms added to the metal

    that fit in between the metal atoms(often H,B,C,N)

  • substitutional alloys: atoms added to the metal have smaller radii so they replace atoms in the lattice

    

2.5: Lewis Diagrams📊

lewis structures:

• Covalent bonds are formed between atoms sharing electrons

• lewis structures are a simple way of representing covalent bonds

2.6: Formal Charge & Resonance⚡︎

formal charge: method that helps determine what resonance structure is the most valid

• sum of lone electrons & bonds connected to the atom - valence electrons

  

2.7: VSEPR Theory & Bond Hybridization💡

VSEPR theory: predicts the geometrics of molecules and polyatomic ions

• the shape, or geometry, of a molecule is determined by lone pairs or bonds on the central atom of a molecule

• electron-electron repulsions are minimized by positioning themselves as far apart as possible

• lone pairs tend to repel more than bonds

  

Bond hybridization:

• to explain molecular geometries, we assume that the atomic orbitals on an atom mix to form hybrid orbitals

• the shape of a hybrid orbital is a mix of the shapes of the original atomic orbitals such as s(spherical) and p(dumbbell)

• the total # of atomic orbitals on an atom remains constant, so the # of hybrid orbitals on an atom equals the # of atomic orbitals that are mixed

• in CH4, the 2s and 3 2p orbitals of carbon mix to form 4 sp3 hybrid orbitals

Sigma and Pi bonds:

• Sigma bonds(σ) are always the first bond between two atoms - a single bond

• Pi bonds(π) are second and third bonds resulting from the overlap of p orbitals🥧