Class 1 - 06/06/24:

• Atoms - smallest unit of any element - has protons, neutrons, and electrons

  • p = +1, mass = 1amu

  • e = -1, mass = 0amu

  • n = 0, mass = 1amu

• Atomic number = Z, the number of p*

• Mass number = p*+n

• The charge = p-e

  • Cations and anions are ions

    • C>0 = cation → +

    • C<0 = anion → -

    • C=0 = atom level

• Isotopes: two atoms of the same element that differ in their number of neutrons - determined by mass number

• Bohr model of the atom:

  • Electrons orbit at a fixed distance from the nucleus - the orbit decreases with distance from the nucleus - as we move away, the ends come closer and energy increases with distance from the nucleus

• Electrons absorb only specific allowed E(due to fixed quantities of E)

  • Current orbit = ground state

  • Higher E orbit = excited state

  • Ephoton = Ef-Ei

  • e- in an excited state can come to a lower level to emit a photon - when dropped, the e- becomes relaxed

• Hydrogen Absorption spectrum = dark bands on a bright background(absorb, so black line)

• Hydrogen Emission spectrum = bright bands on a dark background(emit, so bright lines)

• The Energy of a photon is related to its wavelength(λ lambda) and frequency(f)

  • E = hf = hc/λ → Wavelength and frequency are inversely related - when f is high, λ is low, E is high; when f is low, λ is high, E is low

  • 

• E- exists in 3D orbitals and 4 quantum number describe their structures(orbitals are the areas around the nucleus where an e- is most likely to be found)

  • s, p, d, f → s being lowest in E and F in highest E

  • ex. boron= 1s2 2s2 2p1 → goes by block and row of the periodic table

• 3 basic rules for Electron filling:

  • Pauli principle: there can be no more than 2 e- in any given orbital(spin up and down)

    • e- cannot be the same, hence, different spins

  • Aufbau principle: E- occupies the lowest orbitals first and is filled in increasing E

    • exception is 3d and 4s → 4s are removed before 3d

  • Hund’s Rule: E- first occupy an orbital singly then pair up(no orbital left empty)

• Anomalous electron configurations: when some elements prefer to be half-filled or filled by taking an e from 4s and putting it into 3d ex. Cr and Cu

• Paramagnetic = at least one unpaired e-

• Diamagnetic = all e- are paired ex. noble gases

• Ground state e-configuration: the ground state is the lowest e configuration → correct amount of e as the element has

• Excited state e-configuration: the element has jumped an orbital but it has not changed the amount of electrons as it originally had

• A half-filled shell is more stable than one that isn’t; filled one is the most stable ex. noble gases

• Valence shell configurations of elements determine the chemical reactivity of the elements → Elements in the same group show similar characteristics ex. noble gases as calm due to their octet shells, while halogens are reactive gases with 1 e-missing

• Valence shell electrons experience electrostatic attraction due to the nucleus → shielding effect

  • force of electrostatic attraction is proportional to Zeff + C/r²

• Atomic radius increases going down right to left

• Ionization E increase going up from left to right

• Electron affinity(negativity) increases going up from left to right

• Electronegativity increases going up from left to right

  • FON=ClBrISC=H

• Acidity increases going down left to right down

Class 2: 11/06/24:

• Molecular Structure

  • Drawing lewis structures:

    • count the number of valence e-

    • Arrange atoms with least e-neg in the center(C is always in middle and H is never in the middle) - Use FON=ClBrISC=H

    • Connect each outer atom to center(each line is two electrons)

    • add dots of pairs to each electron till none are remaining

    • Complete missing octets

    • Assign formal charges

      • FC = valence e -(1/2 bonding e - sticks)-(lone pair e - dots)

    • Always check the number of e, octet rule obeyed, formal charges add up to total charge, the smallest set of FC, + charges on fewer e-neg atoms and - charges on more e-neg atoms

  • Sets of electrons also determine hybridization → count each bond and lone pair as one group

    • 2 groups = sp linear ex. CO2

    • 3 groups = sp² trigonal planar

    • 4 groups = sp³ tetrahedral

  • The lone pairs on the molecule determine the molecular shape

    • AX4 = tetrahedral ex. CH4

    • AX3E = trigonal pyramidal ex. NH3

    • AX2E2 = bent ex. H2O

      • All have the same bond angles but different shapes

• Chemical bonding: these bonds form when electrons are shared between two atoms as their orbitals overlap

  • More electrons shares make a stronger bond

  • A shorter distance between atoms makes a stronger bond

  • Stronger bonds have higher dissociation energies

  • Breaking bonds is always endothermic

  • Ionic or covalent bond types

  • Electronegativity determines the bond → if higher EN, then the bond is polar and e are not shared equally; if small EN, then the bond is shared equally and non-polar

    • Covalent bonds: formed between non-metal-non-metal → high EN - these compounds are insulators

    • Metallic bonds are formed between atoms with low EN metal-metal - these compounds are conductors and malleable

    • Coordinate covalent bonds are formed between atoms with lone pairs and e-deficient - ligands ex. hemoglobin

    • Ionic bonds are formed with particles of opposite charges anions-cations - these are insulators and brittle

• Intermolecular Forces: when opposite charges attract each other - larger charges, stronger IMF and more tightly held together

  • Ion-dipole forces - between ions and polar molecules

  • Dipole-Dipole forces - between polar molecules and easier to cleave

  • Dipole-induced-dipole - between polar and non-polar - bigger e cloud and polarity from polar molecule but very easily cleaved

  • London Dispersion - temporary small dipoles formed by collisions - very weak

  • Hydrogen bonding - between very polar molecules with donors being (N-H, O-H, F-H) and acceptors being (O:, N:, F:)

    • Strongest bonds highest to lowest: Ionic>Covalent>Coordinate Covalent>Metallic

    • Strongest IMFs: Ion-dipoles>H-bond>Di-di>di-induced-di>LDF

• Chemical Thermodynamics

  • Enthalpy - The energy stored within chemical bonds or any attractive forces

  • the change in enthalpy of a reaction is teh difference between energy stored in reactant vs products - delta h reaction can be high or low

    • Low H is exothermic(-) Forming bonds

    • High H is endothermic(+)Breaking bonds

  • Enthalpies of formation is the amount of energy associated with forming One mole of compound = sum of products-sum of reactants

    • Enthalpy is independent of its pathways - Hess’s law requires a combination of two or more reactions to find the enthalpy for the overall reaction

  • Entropy = potential randomness - more particles, changing phases, increase temperature → all increases entropy

  • Gibbs free energy: energy available to do work

    • G = H - TS

      • Spontaneous process is exergonic → G is negative

      • non-spontaneous is endergonic → G is positive

Class 3 - 27/06/24: