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Chemistry Study Notes: Electrons, Bonding, and Intermolecular Forces, Lecture notes of Chemistry

Dakota State University (DSU)Chemistry

These comprehensive study notes cover the fundamental concepts of electronic structure, bonding, and intermolecular forces in chemistry. They delve into the wave nature of light, quantized energy, and the bohr model, providing a detailed explanation of atomic orbitals and electron configuration. The notes also explore different types of chemical bonds, including ionic, covalent, and metallic bonds, and discuss the factors influencing bond polarity and electronegativity. This resource is ideal for students seeking a thorough understanding of these essential chemical concepts.

Typology: Lecture notes

2023/2024

Uploaded on 11/08/2024

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Unit 9: Electrons, Bonding, and Intermolecular Forces

Electronic Structure of Atoms

- The Wave Nature of Light

- Electromagnetic spectrum/Electromagnetic radiation = radiant energy

- Scale ranges from gamma rays to radio waves

- Visible light = the only section humans can see (colors)

- The speed of all light is constant (in a vacuum)

- c (speed of light) = 3.00 x 108 meters/second

- Wavelength (λ) = the distance from any point on one wave to the same point on the next wave

- Short wavelength = high frequency + high energy

- Long wavelength = low frequency + low energy

- Frequency (v) = the number of wave fronts that pass a given point in an allotted amount of time

- Hertz (s-1) = the amount of wave fronts that pass a given point in 1 second

- c = λv

- Diagram of a Wave

- Crest = top of the wave

- Trough = bottom of the wave

- Amplitude = distance from the crest/trough to the x-axis of the wave

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Unit 9: Electrons, Bonding, and Intermolecular Forces Electronic Structure of Atoms

The Wave Nature of Light
- Electromagnetic spectrum/Electromagnetic radiation = radiant energy
  - Scale ranges from gamma rays to radio waves
  - Visible light = the only section humans can see (colors)
  - The speed of all light is constant (in a vacuum)
    - c (speed of light) = 3.00 x 10^8 meters/second
- Wavelength (λ) = the distance from any point on one wave to the same point on the next wave
  - Short wavelength = high frequency + high energy
  - Long wavelength = low frequency + low energy
- Frequency (v) = the number of wave fronts that pass a given point in an allotted amount of time
  - Hertz (s-1) = the amount of wave fronts that pass a given point in 1 second
- c = λv
- Diagram of a Wave
  - Crest = top of the wave
  - Trough = bottom of the wave
  - Amplitude = distance from the crest/trough to the x-axis of the wave

Quantized Energy and Photons
- Energy is quantized, meaning it can only be released in specific amounts (quantum or photon)
- E = hv
  - E = energy
  - h (Planck’s constant) = 6.62607004 × 10-34^ m^2 kg / s
    - Planck’s constant links the amount of energy a photon carries with the frequency of its electromagnetic wave
  - v = frequency
- Radiant energy is quantized
  - Light possesses both wave-like and particle-like behavior (photon vs. wave behavior)
Line Spectra and the Bohr Model
- Monochromatic light = light of a single wavelength
- Spectrum = when radiation from a source is separated into its different wavelengths
  - Continuous spectrum = rainbow of colors, containing light of all wavelengths
  - Line spectra = an emission spectrum consisting of separate isolated lines (can be used to identify an element, because elements have specific line spectra) - Emission spectra = a spectrum of the electromagnetic radiation emitted by a source - Absorption spectra = a spectrum of electromagnetic radiation transmitted through a substance, showing dark lines or bands due to absorption of specific wavelengths
- Bohr’s Model
  - Depicts the atom as a small, positively charged nucleus surrounded by electrons that travel in circular orbits around the nucleus
  - n = quantum energy level (7 on the periodic table)
Quantum Mechanics and Atomic Orbitals
- Quantum Mechanics
  - Quantum mechanics = the study of electrons and their locations within an atom (takes into account the particle and wave behaviors of electrons)
  - Heisenberg Uncertainty Principle = it is inherently impossible to simultaneously know both the momentum and location of an electron with any precision - Scientists can locate where the electrons have been (general area/map) and where the electrons are in their resting/ground state
  - Modern model = the electron is a particle, whose behavior is described in terms appropriate to waves, the energy of an electron can be described while discussing its location in terms of probabilities
- Electrons
  - Principle Energy Levels
    - Each electron can be found within a principal/primary/quantum energy level that is arranged from lowest energy(closer to the nucleus) to highest energy (farther from nucleus) - 1st level→quantum level 1 (n=1) - 2nd level→quantum level 2 (n=2) - 3rd level→quantum level 3 (n=3)

The 3 orbitals in each energy level are set at 90° angles to each other
D orbitals = 4 orbitals have a “cloverleaf” shape, while the 5th orbital takes on a donut-like shape
F orbitals = too complicated
Electron Configuration
Electron configuration = the way in which electrons are distributed among the various orbitals of an atom
Aufbau Principle = electrons fill the energy levels one at a time beginning at the lowest level and moving to the higher levels (starts at n=1)
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d
Exceptions
Chromium→1s^2 , 2s^2 , 2p^6 , 3s^2 , 3p^6 , 4s^1 , 3d^5
4s gives one of its electrons to 4d, so that all of its orbitals have at least 1 electron
Sublevel 3d doesn’t want an empty orbital (unstable)
Copper→1s^2 , 2s^2 , 2p^6 , 3s^2 , 3p^6 , 4s^1 , 3d^10
4s gives one of its electrons to 4d, so that 4d has 2 electrons in each orbital (more stable)
Abbreviated Configurations
To shorten the configuration start with the noble gas that precedes it, then write the configuration
Iodine→[Kr] 5s^2 , 4d^10 , 5p^5
Valence Electrons
Valence electrons = the electrons found in the outermost p and s orbitals
Electrons can join, leave or be shared when in these orbitals
Electrons can Absorb Energy
Electrons absorb energy in quanta (small fixed amounts)
Excited state = when electrons absorb energy from their environment
When electrons absorb a quanta of energy, they become energized and jump to a higher energy level (discovered by Planck and Bohr)
The electron can’t hold on to the energy for long, so eventually they fall back to their original/ground state (normal energy levels), emitting the extra energy as light/photons (discovered by Einstein)
Often the wavelength of this emitted light is in the visible part of the electromagnetic spectrum (line spectra)
Photoelectron Spectroscopy
Photoelectron spectroscopy = the energy measurement of electrons emitted from solids, gases or liquids by the photoelectric effect, in order to determine the binding energies of electrons in a substance
Graph for Nitrogen (1s^2 2s^2 2p^3 ) Chemical Bonds

Types of Bonds
- Ionic bond = the electrostatic forces between oppositely charged ions
  - Typically when metals interact with nonmetals
- Covalent bond = chemical bond formed by sharing 1 or more pairs of electrons
  - Typically when non metals interact with other nonmetals
  - Lewis electron dot diagrams
- Metallic bond = when many metal nuclei share delocalized electrons
  - High melting point
  - Conducts electricity
- Network covalent solid = SiO 2 (quartz), Si, germanium, carbon (diamond and graphite)
  - High melting points
- Strongest to weakest bond: metallic, ionic, network covalent, covalent
Ionic Bonding
- Electron transfer
  - Ionic bonds typically form between elements with low ionization energy (readily lose an electron) and an element with high electron affinity (readily gain an electron)
- Lattice Energy
  - Lattice energy = the energy required to completely separate a mole of a solid ionic compound into its gaseous ions
  - LE = k(q + q)/d
    - k = constant
    - q = absolute value of the charge of the ion
    - d = distance between atomic nuclei (ionic radius)
  - The smallest molecule made with the largest charge = more stable = more lattice energy
  - The larger molecule made with the smallest charge = more unstable = less lattice energy
    - Charge is more important in determining lattice energy
- Ions of the Representative Elements
  - Octet Rule = atoms tend to gain, lose, or share electrons until they are surrounded by 8 valence electrons (many exception)
  - Group 1 elements form +1 ions
  - Group 2 elements form +2 ions
  - Group 15 elements form -3 ions
  - Group 16 elements form -2 ions
  - Group 17 elements form -1 ions
- Properties of Ionic Compounds
  - Brittle, with high melting points
  - Crystalline structure (rigid, well defined 3D arrangement)
Covalent Bonding
- Properties of Covalent Molecules
  - Low melting points
  - May vaporize readily
  - May be pliable in their solid forms
Bond Polarity and Electronegativity
- Bond polarity describes the sharing of electrons between atoms

Trigonal planar
- Central atom has 3 unpaired electron (Boron)
- Flat molecule
- Bond angle = 120°
Tetrahedral
- Central angle has 4 bonding spots (Carbon)
- Bond angle = 109.5°
Trigonal pyramidal
- Same as tetrahedral except the central angle has one unshared pair of electrons
- Bond = 107°
Trigonal bipyramidal
- Central atom has 5 bonding locations (has ability to hold more than an octet)
- Bond angles = 90° and 120°
Seesaw = central atom has 3 single bonds, 1 double bond, and a pair of unshared electrons
Octahedral
- Central atom has 6 bonding locations (has ability to hold more than an octet)
- Bonding angle = 90°
Electronic Geometry
Electronic geometry = molecular geometry *if the central has no unshared electrons Electron Domains Electron Geometry 2 Linear 3 Trigonal planar 4 Tetrahedral 5 Trigonal planar 6 Octahedral
VSERP (Valence Shell Electron Pair Repulsion)
Domains = how many things (paired electrons or atoms) are attached to the central atom
Electronic geometry (deals with domains)
Molecular geometry (deals with atoms attached to the central atom)
Bond polarity (if electronegativity difference between 2 elements is over .4 the bond is polar)
Molecular polarity
If bonds are nonpolar then molecule = nonpolar

If any of the bonds are polar then the shape of the molecule needs to be taken into account - Symmetrical = nonpolar - Asymmetrical = polar
Hybridization = the idea that atomic orbitals fuse to form newly hybridized orbitals, which in turn, influences molecular geometry and bonding properties
s = 1 domain
sp = 2 domains
sp^2 = 3 domains
sp^3 = 4 domains
sp^3 d = 4 domains
sp^3 d^2 = 6 domains
Resonance Structure
Sometimes a single Lewis structure is inadequate for describing a molecule or polyatomic ion
Sometimes it is possible to draw several correct, equivalent structures for the same molecule, in which only the placement of the electrons differs
These structures are called resonance structures
The “true” structure is somewhere in the middle (a blend of the equivalent structures)
Double and triple bonds can alter their location within the molecule
Strengths of Covalent Bonds
Bond enthalpy = the enthalpy change for the breaking of a particular bond in one mole of a gaseous substance
ΔHrxn = (enthalpies of the bonds broken) - (enthalpies of the bonds created)
The greater the bond enthalpy, the stronger the bond
A molecule with strong chemical bonds is less reactive than one with weak bonds
As the number of bonds between 2 atoms increases, the bond becomes shorter
Single bond is shorter than a double, which is shorter than a triple Intermolecular Forces
Intramolecular forces = covalent bond and ionic attractive force
Covalent = sharing of electrons
Ionic = transfer of electrons leading to an electrostatic attraction (positive-negative) between atoms
Intermolecular forces = the attractive or repulsive forces that exist between 2 or more molecules
Dispersion Forces/The London Dispersion Force
Deals with nonpolar molecules
There is no positive or negative end, therefore there is no attraction
Dispersion force = a temporary attractive force that results when the electrons in two adjacent atoms occupy positions that make the atoms form temporary dipoles (sometimes called an induced dipole-induced dipole attraction)
The larger the molecule, the larger the dispersion force
Hydrocarbons exhibit dispersion forces
Hexane is more attracted to hexane, then methane is to methane
In general gases exhibit dispersion forces
Weakest
Dipole-Dipole Force

Chemistry Study Notes: Electrons, Bonding, and Intermolecular Forces, Lecture notes of Chemistry

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