General Chemistry Unit 2 PDF

https://youtu.be/2ekjGl8yWZk?si=51hG0v-Zst4HBszu https://youtu.be/7YlJhPLocLA?si=9v-jCrUbcVBb1wxN GENERAL CHEMISTRY UNIT 2 → Rutherford’s Gold Experiment: → fired alpha particles at a thin sheet of gold foil to study atomic structure...

https://youtu.be/2ekjGl8yWZk?si=51hG0v-Zst4HBszu https://youtu.be/7YlJhPLocLA?si=9v-jCrUbcVBb1wxN GENERAL CHEMISTRY UNIT 2 → Rutherford’s Gold Experiment: → fired alpha particles at a thin sheet of gold foil to study atomic structure → most particles passed through the foil, but some were deflected at large angles → atoms are mostly empty space with a small, dense nucleus that has a large positive charged and where nearly all of the atom’s mass is concentrated → developed nuclear model of the atom → Isotopes: atoms of the same element that have the same number of protons but different number of neutrons → different atomic masses → ex) carbon-12 vs carbon-14 → Periodic Law: the properties of elements are periodic functions of their atomic numbers→ elements with similar properties occur at regular intervals → periodic table arranges elements in increasing order of their atomic number → Coulomb’s Law: interaction between two charged particles, potential energy (V) depends on the product of the charges (q1,q2) and the distance ® between them → opposite charges attract, and the closer they are, the stronger the attraction → crucial for understanding how protons and electrons interact within an atom → Moseley: established that the atomic number (number of protons) is fundamental property of elements, not atomic mass → experiment with X-ray Spectra: linked X-ray frequency to atomic number → elements should be arranged by atomic numbers → Noble Gases: He, Ne, Argon: very unreactive → Ionization energy: energy required to remove an electron from an atom, converting it into an ion → strength of the attraction between an electron and the nucleus is described by Columbs Law: electrons closer to nucleus → stronger attraction→ more energy required to remove → ionization energy increase across periods as atomic number increases → decrease in IE after noble gas bc electrons in alkali metals are farther from nucleus → Valence electrons: electrons in the outermost shell → determine reactivity and bonding behaviors → Successive ionization energy: (IE1, IE2, IE3) describe the energy required to remove more electrons from an atom → large jumps in IE show when an electron is being removed from a full inner shell, as opposed to the valence shell → ex) Na has an easy removable valence electron (shown by the large increase in energy required to remove the second electron from a more tightly bound inner shell → each successive electron removal requires more energy, since closer to nucleus, experiencing stronger reaction → Electron Shell Model: atomic properties based on valence electrons, shell size, nuclear charge → Hydrogen Spectrum: it emits light at specific frequencies corresponding to visible, ultraviolet and infrared regions → Rydberg equation: → pattern of emission highlights that only certain energy transitions are allowed, electron energy levels in hydrogen are quantized → Photoelectric Effect: demonstrates that light is quantized into energy packets called photons → when light of a certain frequency strikes a metal surface, it ejects electrons → kinetic energy of e- depends on lights frequency, not its intensity → energy of a photon is related to its frequency E=hv H: Planck’s constant → Waves: transfer energy without matter moving along with them → electromagnetic waves: oscillating electring and magnetic fields that can travel at the speed of light Wavelength(λ): distance between two consecutive peaks in waves Frequency (v): measured in hertz (Hz) → wavelength and frequency are inversely proportional → as wavelength increases, frequency decreases C = λv C: speed of light → Interference: when two waves meet → constructive: when wave peaks align → increasing wave’s amplitude → destructive: when peaks and troughs align → reducing wave’s amplitude → Quantum energy level model: describes specific energy levels in hydrogen atoms → Electron orbitals: → n: energy levels → quantum numbers: → n: size of orbital → l : shape of orbital ( s: spherical, p: lobes) → m: orientation of orbital → Photoelectron Spectroscopy: → measures ionization energies of electrons in various orbitals by analyzing the energy needed to eject them from atoms → electron configuration (1s2, 2s2, 2p6 etc) → Electron-electron repulsion: electrons repel each other, and shielding effects reduce the attraction between the nucleus and outer electrons, leading to trends in atomic size, ionization energy, and electron affinity across the periodic table. → Effective nuclear charge: net positive charge felt by electrons after accounting for repulsion from other electrons → number of electrons in each shell is limited by the Pauli Exclusion Principle and energy constraints → states that no two electrons can occupy the same quantum state within an atom, combined with energy constraints→ limits number of electrons in each orbital → elements in the same group have similar valence electron configurations → leading to similar chemical properties → Line spectra: atoms or molecules emit light at specific wavelengths when heated or exited, producing discrete line spectra → each energy line corresponds to a quantized energy level in atom → Orbital energy: As the principal quantum number increases, the energy of orbitals generally increases → in larger atoms, orbital overlaps occur (4s orbital fills before 3d) caused by the effects of shielding and penetration → shielding reduces attraction between outer electrons and the nucleus → penetration allows s orbitals to be closer to nucleus → making them more stabilized than p,d,f orbitals → Electron affinity: energy change when an electron is gained, becomes more negative across a period, indicating atoms are more readily gain electrons → Hund’s Rule: electrons fill degenerate orbitals (orbitals of same energy) before pairing up → minimizes electron-electron repulsion → leading to more stable electron configurations by maximizing number of unpaired electrons → Emission: where an electron in an atom releases energy as a photon wen it transitions from a higher to lower energy state → Absorption: when an electron in an atom absorbs energy (typically from a photon) causing it to transition from lower to higher energy level → measuring energy Nf, ni : initial and final quantum numbers → Bohr Model of Hydrogen: → describes the hydrogen atom as a single electron orbiting a positively charged nucleus in specific, fixed circular orbits, with each orbit corresponding to a distinct energy level → electrons have quantized energy levels → electrons orbit the nucleus in fixed orbits → energy is emitted or absorbed between levels → Limitations: fails for multi electron atoms because it doesnt account for electron-electron interactions, like not accounting for subshells. It simplifises electron orbits as fixed circular path with no sub shell consideration. → Heisenberg’s Uncertainty Principle: it is impossible to know the exact position and momentum of an electron simultaneously → wave like nature of electrons → the more precisely the position is known, the less precisely momentum is known → Electron Orbitals: describes region around nucleus where an electron is likely to be found → Quantum numbers: describe properties of these orbitals and the distribution of electrons → Principal quantum numbers (n) : defines the size and energy of the orbital, higher n values correspond to orbitals that are larger and have higher energy → ex) n=1 (first shell) , n=2 (second shell) → Angular Momentum quantum number (l): defines the shape of orbital, ranges from 0 to n-1 → 0: s-orbital (spherical), equal probability of finding electron in all directions from nucleus → 1: p-orbital (dumbell shape), exists in 3 orientations (x,y,z axes) → 2: d-orbital (complex shape), five orientations → 3: f-orbital (even more complex shape), seven orientations → Magnetic Quantum number (ml): describes the orientation of orbital in space → ex) p orbitals can be -1, 0, +1, resulting in 3 orbitals → Spin Quantum number (ms): describes electrons spin → +½ spin up → -½ spin down → Electron Configuration Notation: Describes the distribution of electrons among the orbitals. For example, the configuration of oxygen is 1s² 2s² 2p⁴, meaning it has two electrons in the 1s orbital, two in the 2s orbital, and four in the 2p orbitals → Degeneracy: Orbitals that have the same energy level are degenerate. For example, the three 2p orbitals (px, py, pz) are degenerate because they have the same energy. → Multi-Electron Atoms: In atoms with more than one electron, electron-electron repulsions break degeneracy, making different subshells (s, p, d, f) within the same energy level have different energies. → Valence: The number of bonds an atom can form, based on the number of electrons needed to complete its valence shell. ○ Noble Gases: Valence = 0 (they do not form bonds). ○ Oxygen: Valence = 2 (forms two bonds, e.g., H₂O). ○ Hydrogen: Valence = 1 (forms one bond, e.g., H₂O). ○ Carbon: Valence = 4 (forms four bonds, e.g., CH₄). ○ Nitrogen: Valence = 3 (forms three bonds, e.g., NH₃). ○ Fluorine: Valence = 1 (forms one bond, e.g., HF). → Octet Rule: Atoms in groups 4–8 aim to have eight electrons in their valence shell by forming bonds. For example, carbon forms 4 bonds to complete its octet. → Ionic Bonds: Involves the transfer of electrons between atoms, leading to the formation of cations (positive ions) and anions (negative ions). → Covalent Bonds: Atoms form covalent bonds by sharing pairs of electrons. Each bond consists of one shared pair. → Double Bonds: Two pairs of electrons are shared (e.g., O₂, C₂H₄). → Stronger and Shorter than single bonds. → Triple Bonds: Three pairs of electrons are shared (e.g., N₂, C₂H₂). → Stronger and Shorter than double bonds. → Lone Pairs: Electrons not involved in bonding, crucial for determining molecular geometry and reactivity. Exceptions to the Octet Rule: → Electron-Deficient Molecules: Some atoms, like beryllium (Be) and boron (B), can form stable compounds even though their central atoms have fewer than 8 electrons. → Example: ○ BeH₂: Beryllium forms two bonds with hydrogen, giving it only 4 electrons around the central Be atom (instead of 8).  ○ BF₃: Boron forms three bonds with fluorine, giving it only 6 electrons around the central B atom. → These compounds are stable despite not following the octet rule due to the nature of the elements involved. → Hypervalent Molecules: Atoms in the 3rd period or below (like phosphorus, sulfur) can hold more than 8 electrons because they have access to empty d-orbitals. → Example: ○ PCl₅: Phosphorus forms five bonds with chlorine, resulting in 10 electrons around the central phosphorus atom. ○ SF₆: Sulfur forms six bonds with fluorine, resulting in 12 electrons around the central sulfur atom. → Hypervalent molecules expand their octet due to the availability of extra orbitals, allowing more bonding. → Isomers: Compounds with the same molecular formula but different atomic arrangements (e.g., ethanol and dimethyl ether with C₂H₆O). → Isomers have distinct chemical and physical properties (e.g., ethanol is a liquid at room temperature, while dimethyl ether is a gas). → Structural Isomers: Differ in how atoms are bonded (e.g., C-O-C vs. C-C-O). → Spatial Isomers: Differ in the spatial arrangement of atoms (e.g., carvone isomers have different smells). → Resonance: Occurs when a molecule can be represented by multiple valid Lewis structures. → The actual structure is a resonance hybrid, averaging the characteristics of all the structures. → Example: Benzene (C₆H₆) has alternating double and single bonds, but the actual structure has equal bond lengths due to delocalized electrons. → Ozone (O₃) is another resonance example, where both O-O bonds are equivalent due to electron delocalization. → Formal Charge is a way to predict the most stable structure when multiple Lewis structures are possible. ○ Formal charge = # of valence electrons (free atom) – # lone pair electrons – ½ bonding electrons. → The best structure minimizes formal charges and places negative charges on more electronegative atoms (e.g., oxygen in nitrite NO₂⁻). → Example: For carbon dioxide (CO₂), the structure with no formal charges is preferred over those with charge imbalances. - Rydberg equation: 2 versions, 1 over wavelength or 1 over frequency - Use either - 2s and 2p graphs: 2p is closer to nucleus, but since its closer to 1s it repels with the electrons in the 1s - 2s is lower in PE (fill it up first) - Potential energy and kinetic energy graphs - Total energy= E+KE - Use wave functions to expect how far from nucleus we expect electrons to be - Quantum numbers, for ms u can choose between -½ and +½ - Highest energy electron is always in the furthest shell - Xray photoelectron spectroscopy: put enough energy to reject an electron out - Photo emission: enough energy to excite electrons to emit light/photons - Photoelectron (ionization energy), prove subshells: saw more ionization energies than the number of shells - XPS taking any electron at any point - Number of ionization energies: number of subshells and then the energy jumps represent the shells - Wave particle duality: defraction patterns, shows that it cant just be a particle, know this experiment - Successive IE vs XPS: taking electrons from the same atom over(in order) and over vs one electron from any subshell (out of order) - In XPS probably still electron electron repulsion since not in order vs Successive IE: no repulsion - Copper and chromium - Electron configuration: AR= 4s1 3d10 - Energies are close, so they are able to jump, since they are close, 4s doesnt need to be filled first - Learn how to draw energy diagram - Why nitrogen has a positive electron affinity: - Periodic table, right side, core charge of +5 and small radius - Determining electron affinity: trends across periodic table - Core charge increases, affinity increase - Further from nucleus, less attraction - Core Charge # protons (atomic number) - # inner shell electrons - Noble gases have 0 electron affinity (power set) - Photoelectric effect: https://byjus.com/jee/photoelectric-effect/ - What happens when ur shooting photons to try and eject an electron - - All the experi ments and what they told us - Draw orbital lobes (s,p) (s just circles, and get bigger and bigger, draw nodes) - Lower formal charge is better  - Preferred to have -ve charge on atom w higher electron affinity CDS Video 15: Electron wave motion; uncertainty; orbitals (16:53) Rutherford’s Experiment: 1.) Particles went straight through (no deflection) (most alpha particles) → most of the atom is empty space 2.) Particles “bounced back” (large deflection) (small amount of alpha particles) → a small part of the atom is more massive than the alpha particle, thus very dense 3.) Shallow deflections (even small amount of alpha particles) → the small part of the atom is +ve charged - Atom is mainly empty space with a dense, +ve charged nucleus. Moseley’s X-Ray Experiment: Bombarded atoms with energy and then measured X-ray frequencies emitted → each element has a unique x-ray frequency (unique, physical property) → assigned an atomic number to each element-integer ranking of ordered mass → when graphing x-ray frequency against atomic number→ perfect curve → comparing the square root of x-ray to atomic number = linear relationship → x-ray frequency is unique for each atom which corresponds to a unique physical property of each element → there is a perfect relationship between x-ray frequency and atomic number, so the atomic number must also correspond to a unique physical property of the atom → atomic numbers are integers → From Rutherford’s experiment, we know that an atom contains a dense, positive nucleus → atomic number counts the number of protons (units of positive charge) in the nucleus of an atom Coulomb’s Law: → quantifies the force between two objects Potential energy: → PE is proportional to Coulombs law → PE will always be negative → Large magnitudes of PE are lower numbers than high magnitudes of PE → Lower PE = more stable state (electron experiencing greater Coulombic interactions) Ionization Energy: → amount of energy required to remove an electron from the atom → for an electron to be removed, then the Coulombic attraction between the electron and the atom must be zero → Low PE= large IE → High PE= small IE → IE= -PE → To determine PE: → Look at Coulombs Law 1.) Core Charge: greater core charge = lower PE= higher IE 2.) Radius: smaller radius= lower PE= higher IE 3.) Electron-electron repulsion: like negative charges= repulsion=raise PE= lower IE Core Charge: → effective nuclear charge felt by an electron → number of protons - number of shielding electrons First IE vs. Atomic Number: 1.) Small gradual increase → within elements in the same period (across a row) 2.) Huge decrease → between two different periods (going to a new row) 3.) Overall decrease → within elements in the same group (down a row) Conclusion: → electron shell model → we already know that the protons in the center of an atom (nucleus) → group of electrons at a similar radius from are in a “shell” → we can figure out how many electrons are in a shell by looking at the big drops Shell Model to Explain Data: 1.) Small gradual increases → more protons = core charge increase → shielding electrons remain the same as electrons are in the same shell → radius also decreases because core charge increases, but not the main reason  2.) Huge decreases → happens from one period to the next. Electrons go into a new shell, greatly increasing the radius → core charge decreases because there are more shielding electrons Mention BOTH 3.) Overall decrease → when you go down a group, while the core charge remains the same, there are more electron shells, so there is a greater radius Successive Ionization Energy: → energy needed to remove each electron from the same time one after the other in ORDER Number of Valence Electrons: → small increases most of the time → electron-electron repulsion is reduced as an electron is removed → lowers the PE of all other electrons → increasing IE of the next electron → large increases → A larger increase indicates we have moved to a closer electron shell → OVERALL: counting the number of electrons before a large increase can tell us the number of valence electrons Light: → we can relate the speed of light to its frequency and wavelength using the equation: C = λv C: speed of light (3.00 x 10^8 m/s) Λ: wavelength of light (m) 1m = 10^9 nm V: frequency of light (1/s or Hz) Photoelectric Effect: → Shine light on the surface of the element (typically metal), which ejects electrons → light frequency can impact kinetic energy (increases) but not electron current → light intensity can impact electron current (increases), but not kinetic energy CHECK GRAPHS KE of electrons Current: number of electrons ejected → because a threshold frequency (V0) exists to eject electrons, it shows that the energy of the light is not additive  → which would be expected from a wave → In the KE vs Intensity graph, where increasing the amount of light doesn't increase how much energy each electron has → light is quantized and comes in discrete packets called photons Particle-wave duality: → ability of light to act as both a particle and a wave Photon= particle of light energy E=hv E: energy (J) H: Planck’s constant (6.6262 x 10^-34) Js V: frequency (Hz) X-ray Photoelectron Spectroscopy (XPS): → input a large amount of energy into the same and measure the KE of the ejected electrons → relate the KE of electrons ejected to the IE required to remove them from the atom Ephoton = hv = IE + KE Conclusion: → from XPS data: some elements have multiple groups of electrons with different KEs → different KEs can relate to different thresholds of IE → New IE threshold = new shell → New IE threshold of SIMILAR magnitude= new subshell → we can count the number of electrons in a subshell by counting the number of elements before moving to a new IE threshold → the lowest IE threshold corresponds to the first IE in the successive IE because the right-most/ lost IE is the IE of the valence electron Successive IE vs. XPS: → SIE: IE of an electron removed consecutively from the same atom (valence → inner shells) → XPS: IE of any electron, even inner shell, without needing to remove the valence electrons → thus measured IEs are NOT the same