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General Chemistry
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:
General Chemistry
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: