Atoms, Electronic Structure & Bonding

Structure of an atom (3 parts)

• Electron clouds

• Nucleus

• Neutron

Principal Quantum Numbers (n)

• Distance of electron from nucleus

  • The main coefficient in electron configuration

  • 1s² 2s² 2p⁶ 3s² 3p²

• High n value means higher energy

• Valence steel - outermost (highest) n

Azimuthal Quantum Numbers (l)

• Shape of the orbital

• Range: 0 - (n-1)

  • 0 = s

  • 1 = p

  • 2 = d

  • 3 = f

• 1s² 2s² 2p⁶ 3s² 3p²

Magnetic Quantum Numbers (mᵢ)

• orientation of the orbital

• Range: (-1, +1) or (2l +1)

• For each sub-shell, how many electrons at most it can take

Spin Quantum Numbers (mₛ)

• Spin up or spin down

Octet Rule

• Atoms prefer to be in a state with access to 8 valence electrons

• Can drive periodic table trends

Orbital Filling vs Orbital Bonding

Even though we fill the 3d orbital later, the 4s orbital will bond first and dominate bonding behavior

• Radius of 4s > radius of 3d

D-shells and F-shells (2 properties)

1. Magnetism

   1. Elements with several unpaired electrons have higher magnetic spin

      1. Results in higher saturation magnetization and higher remanence

2. Fluorescence and Phosphorescence

   1. d and f orbitals can have split energy levels due to atom-atom interactions

   2. if light elevates an electron to a higher energy level (absorption), it will eventually relax back down releasing light (emission)

Primary Bonds (3)

1. Ionic

2. Covalent

3. Metallic

Secondary Bonds (2)

1. Van der Waals

2. Hydrogen Bonding

Electronegativity Trend

• Increases moving from bottom to top, and from left to right

• Atoms that want to receive an electron to achieve “Noble Gas” electron configuration will be more electronegative

• Atoms want to lose an electron to achieve “Noble Gas” configuration will be less electronegative

Linus Pauling Equation

χ₁ - χ₂ = ΔEN

• Electronegativity difference

• It is impossible to be 100% ionic!

ΔEN for Pure covalent

0

ΔEN for Non-polar covalent

0 - 0.4

ΔEN for Polar covalent

0.4 - 2.0

ΔEN for Ionic

> 2.0

Ionic Bonds

• Electrons are transferred between atoms

• Bonding is a result of a cation (+) and attracting anion (-)

• Oxide ceramics are typically considered ionic solids

Coulombic Attraction

Fa = ( |Z₁Z₂|q² ) / ( 4𝝅εₒx² )

Repulsive Force

Fᵣ = -B / x¹²

In ionic bonds, is there any difference in bond energy between the same atoms with different orientations?

No, ionic bonds are non-directional

Fracture in Ionic Bonds

• Basically no plastic deformation

• Very brittle

• Think ceramics

• Repulsive forces cause lattice to break

Covalent Bonds

• Electrons in the valence shell are shared between neighbor atoms

• Bonding is a result of orbital overlap between unfilled, valence orbitals

• Polymers and semiconductors usually have covalent bonds

• Generally strong bond strength

Hybridization

• Covalent bonds are STRONGLY directional

• Many common covalent solids rely on different types of orbitals overlapping to create “hybrid” orbitals

• These mixed orbitals have non-spherical electron density, which results in strongly directional bonds

Metallic Bonds

• Valence shell electron(s) are delocalized and shared across the entire solid

• Bonding is a result of the atomic cores “floating” in a “jellied (sea of electrons)”

• Metals have metallic bonding

Why are metals so ductile?

All atoms/bond/electrons “look the same” and are already shared, so atoms can slide past each other and form bonds in the new area of jellium

Density comparison between metallic solids, ionic solids, and covalent solids

• Metallic Solid: Packing spheres to maximize space, filling, packing efficiency

• Ionic Solid: Packing of spheres also maximize space, filling

  • Limited packing efficiency due to size mismatch of cations and anions

• Covalent Solids: More complex, “open” structures determined by fixed bonding angle

  • Directional Bonds

Van der Waals Bonding (3 categories)

• Much weaker than primary bonds

• Form due to electrostatic attraction

1. Induced dipole - induced dipole

   1. London Dispersion forces

   2. Weakest VdW

2. Dipole - induced dipole

3. Dipole (permanent) - dipole (permanent)

   1. Strongest VdW

London Dispersion Forces

Weak bond that forms between two instantaneous (induced) dipoles

1. Random fluctuation in electron cloud create an instantaneous dipole

2. Induces dipole in neighboring atom

3. A weak electrostatic attraction is formed

Dipole-Dipole Bonding

Strongest VdW

• Chlorine is more electronegative than hydrogen, therefor it “hogs” electrons (even though they are sharing) and creates slightly negative permanent dipole

• Hydrogen is less electronegative than chlorine, therefor it “gives” electrons (even though they are sharing) and creates slightly positive permanent dipole

Dipole - Induced Dipole

Permanent dipole induced dipole in neighboring atom

Why do polymers have such low elastic moduli?

Between the chains, there are Van der Waals bonds, which are very weak, letting the chains easily slide past each other

• Regular London Dispersion forces break apart very easily

How can we make polymers stronger?

By stretching C-C bonds, instead of L.D forces