General Chemistry: Quantum Numbers & Electron Configuration PDF

**[Lesson 1: Quantum Numbers]** **Models of Electron Structure of Atoms (SiPoN PaQ -- JJ ENE)** 1. **Solid Sphere Model (by John Dalton 1803)** - All mater is made up of tiny, indivisible particles called atoms from Greek word 'atomos' - Atoms are solid, hard spheres that cannot be crea...

**[Lesson 1: Quantum Numbers]** **Models of Electron Structure of Atoms (SiPoN PaQ -- JJ ENE)** 1. **Solid Sphere Model (by John Dalton 1803)** - All mater is made up of tiny, indivisible particles called atoms from Greek word 'atomos' - Atoms are solid, hard spheres that cannot be created nor destroyed. - Atoms of the same element are identical, while atoms of different elements are different - Atoms combine to form compounds in fixed ratios 2. **Plum Pudding Model (J.J. Thomson 1904)** - Atoms are made up of a uniform sphere of positive charge - Negatively charged electrons are embedded in the sphere of positive charge, like plums in a plum pudding - The total negative charge of the electrons is equal to the total positive charge of the atom, so the atoms has no overall charge (- = +) 3. **Nuclear Model (Ernest Rutherford 1911)** - The atom must have a tiny, dense nucleus at the center, which contains all of the positive charge and most of the mass of the atom - The electrons must orbit the nucleus at a distance, in a similar way to how planets orbit the Sun - The positive charge and mass of the atom were evenly distributed throughout the atom 4. **Planetary Model (Niels Bohr 1913)** - The atom has a central nucleus containing protons and neutrons with the electrons in circular orbitals at specific distances from the nucleus - The electrons encircle the nucleus of the atom in specific allowable paths called orbits and can be moved into different orbits with the addition of energy 5. **Quantum Mechanical Model (Erwin Schrödinger 1926)** - The electron is a wave portrayed by a nucleus surrounded by an electron cloud - It predicts that electrons can only exist in certain fixed energy levels. Quantum numbers represent these energy levels - The electrons can only move between energy levels by absorbing or emitting energy in the form of photons of light **Quantum Numbers** -- describe electrons' properties in atoms, including energy levels, orbital shapes, orientations, and spin 1. **Principal (n)** -- energy level and orbital size 2. **Azimuthal/angular momentum (l)** -- type of orbital/sublevel and orbital shape 3. **Magnetic (m~l~)** -- specific orbital location and orbital orientation 4. **Spin (m~s~)** -- direction of electron and electron spin direction Formula for certain number of electrons that an energy level can hold: n^2^\*2 **Valence Electrons** -- electrons in the outermost energy level of an atom, which are responsible for the chemical properties of an atom **Sublevel** -- an energy level defined by quantum theory Types of sublevel: s, p, d ,f **Orbitals** -- regions of space around the nucleus where electrons are most likely to be found, are defined by four quantum numbers (principal, angular momentum, magnetic, spin) **Orbital Shapes** - **S orbital** -- spherical shape, can hold 2 electrons - **P-orbital** -- peanut or dumbbell shape, can hold 6 electrons - **D-orbital** -- cross peanut or cross dumbbell shape, can hold 10 electrons - **F orbital** -- tetrahedral shape, can hold 14 electrons **[Lesson 2: Electron Configuration]** **Electron Configuration** -- representation of the arrangement of electrons distributed among the orbital shells and subshell **Orbital Diagram** -- type of notation which illustrates an atom's electron distribution and electron spin within orbitals **3 rules in assigning electrons to orbitals:** 1. **Aufbau Principle** -- states that electrons fill lower-energy atomic orbitals before filling higher-energy ones, predicting the electron configurations for atoms or ions - **'Aufbau'** is German for "building-up" 2. **Pauli Exclusion Principle** -- states that no two electrons in an atom can have the same set of four quantum numbers. Since there are only two possible spin states, each orbital can only hold two electrons with opposite spins 3. **Hund's Rule** -- aka the **rule of maximum multiplicity**, states that electrons fill orbitals in a subshell with the same energy level in a way that maximizes the total spin. This means that electrons will first occupy all the available orbitals in a subshell with parallel spins before pairing up with other electrons **Noble Gas Configuration** -- consists of the elemental symbol of the last noble gas prior to that atom, followed by the configuration of the remaining electrons **Magnetic Property** -- unpaired electrons show that substance is paramagnetic; paired electrons show that substance is diamagnetic **[Lesson 3: Ionic and Covalent Bonding]** **Valence Electrons** -- determine how an atom interacts with others and its chemical properties **Noble Gases** -- called "noble" because they rarely bond, their complete valence electron shells make noble gases extraordinarily stable and unlikely to form chemical bonds **Octet Rule** - Refers to the tendency of atoms to prefer to have eight electrons in the valence shell - Allows us to determine the atomic structure of most chemicals **Lewis Structure** by Gilbert Lewis - Lewis electron dot structure, diagrams that show the bonding between a molecule's atoms and the lone pairs of electrons that may exist in the molecule **Ions** -- an atom or group of atoms with an electric charge - Cations -- ions with positive charge - Anions -- ions with negative charge **Ionic Compounds** -- compounds made of ions that form charged particles when an atom (or group of atoms) gains or loses electrons, referred to as electrovalent or saltlike compounds **Covalent Compounds** -- formed when atoms share electrons, creating molecules - Single bond -- two atoms share one pair of electrons - Double bond -- two atoms share two pairs of electrons - Triple bond -- two atoms share three pairs of electrons **[Lesson 4: Polarity of Compound]** **Polarity** -- distribution of electrical charge over the atoms joined by the bond - Electrons in a polar covalent bond are unequally shared between the two bonded atoms, which results in partial positive and negative charges - **Principle "like dissolves like"** -- polar dissolves polar, nonpolar dissolves nonpolar **[Lesson 5: Molecular Geometry of Compound]** **Molecular Geometry** -- 3D arrangement of atoms around a central atom, determined by VSEPR (Valence Shell Electron Pair Repulsion) theory, which minimizes repulsion between bonds and lone pairs **VSEPR Theory** -- Valence Shell Electron Pair Repulsion, states that electron pairs around a central atom repel each other, arranging themselves to be as far apart as possible, determining molecular shape **Types of Molecular Geometry (LBT TO)** 1. **Linear** - Atoms in a straight line around carbon - Features a central with two bonded atoms and no lone pairs, resulting in a straight-line shape with a 180 degree bond angle (CO2) 2. **Bent or Angular** - Lone pairs of oxygen create a bent shape - Has a central atom bonded to two atoms with one or more lone pairs, creating a v-shape with bond angles less than 120 or 109.5 degree (H2O) 3. **Trigonal Planar** - Atoms form a flat triangle around boron - Has a central atom bonded to three atoms with no lone pairs, forming a flat, triangular shape with 120 degree bond angles (BF3) 4. **Tetrahedral** - Hydrogen atoms form a 3D tetrahedron around carbon - Has a central atom bonded to four atoms with no lone pairs, creating a three-dimensional shape with bond angles of about 109.5 degree (CH4) 5. **Octahedral** - Has a central atom bonded to six atoms with no lone pairs, forming a three-dimensional shape with 90 degree bond angles (SF6)