Bohr Model of Atom - Bohr 2nd Postulate

  • Bohr’s postulates explained the behavior of electrons in an atom
  • Second postulate states that electrons move in fixed, circular orbits around the nucleus
  • These orbits are called energy levels or shells
  • Each energy level has a specific energy associated with it
  • Electrons can only exist in these energy levels, not in between
  • Electrons can transition between energy levels by absorbing or emitting energy
  • The energy absorbed or emitted is quantized, meaning it is discrete and specific
  • The energy difference between energy levels determines the frequency and wavelength of light emitted or absorbed
  • The smaller the energy difference, the longer the wavelength of light emitted

Energy Levels and Quantum Numbers

  • Energy levels are denoted by the principal quantum number, n
  • Each energy level can have a maximum number of electrons, given by 2n^2
  • The first energy level (n=1) can hold a maximum of 2 electrons
  • The second energy level (n=2) can hold a maximum of 8 electrons
  • The third energy level (n=3) can hold a maximum of 18 electrons
  • Each energy level is further divided into subshells or orbitals
  • Subshells are denoted by the angular momentum quantum number, l
  • The number of subshells in an energy level is given by n
  • Each subshell can hold a maximum number of electrons, given by 4l + 2

Electron Configurations

  • Electron configurations describe the arrangement of electrons in an atom
  • Aufbau principle: Electrons fill the lowest energy levels first
  • Hund’s rule: Electrons fill degenerate orbitals (same energy) one by one before pairing up
  • Pauli exclusion principle: No two electrons can have the same set of quantum numbers
  • Electron configurations are written using the noble gas notation
  • The noble gas notation represents the previous noble gas before the element, followed by the electron configuration of the remaining electrons

Electron Affinity

  • Electron affinity is the amount of energy released when an electron is added to a neutral atom
  • It is a measure of the atom’s tendency to gain electrons
  • Electron affinity values can be positive or negative
  • Positive values indicate that energy is required to add an electron, making the process unfavorable
  • Negative values indicate that energy is released when an electron is added, making the process favorable
  • Group trend: Electron affinity generally decreases down a group
  • Periodic trend: Electron affinity generally increases across a period, with some exceptions

Ionization Energy

  • Ionization energy is the amount of energy required to remove an electron from a neutral atom
  • It is a measure of the atom’s tendency to lose electrons
  • Ionization energy values can vary depending on the energy level from which the electron is removed
  • First ionization energy refers to the energy required to remove the first electron
  • Second ionization energy refers to the energy required to remove the second electron, and so on
  • Group trend: Ionization energy generally decreases down a group
  • Periodic trend: Ionization energy generally increases across a period, with some exceptions

Electron Configuration and Periodic Table

  • Electron configurations can help determine the position of an element in the periodic table
  • Elements in the same group have similar electron configurations in their outermost energy level
  • Elements in the same period have sequential addition of electrons to the energy levels
  • Group 1 elements (alkali metals) have a configuration of ns1
  • Group 2 elements (alkaline earth metals) have a configuration of ns2
  • Transition metals have a configuration of (n-1)d^1-10 ns^1-2
  • Group 17 elements (halogens) have a configuration of ns2 np5
  • Group 18 elements (noble gases) have a full outermost energy level (ns2 np6) and are stable

Orbital Diagrams

  • Orbital diagrams show the arrangement of electrons in an atom using boxes to represent orbitals and arrows to represent electrons
  • Each box represents an orbital, and each arrow represents an electron
  • Orbitals within a subshell are sometimes shown as lines (s) or as boxes (p, d, f)
  • Hund’s rule is followed when filling degenerate orbitals: electrons fill orbitals singly before pairing up
  • The spin of electrons is represented by arrows pointing up or down
  • Orbital diagrams can be used to determine electron configurations and to predict the reactivity of elements

Electron Pairing and Valence Electrons

  • Valence electrons are the electrons in the outermost energy level of an atom
  • They are involved in chemical bonding and determining the reactivity of elements
  • Electrons in the outermost energy level are added in pairs, resulting in paired (spin up and down) or unpaired electrons
  • Unpaired electrons make an atom more reactive because they can easily form bonds with other atoms
  • The number of valence electrons can be determined from the electron configuration
  • For main group elements (s and p block), the valence electrons are those in the outermost energy level
  • For transition metals (d block), the valence electrons are those in the outermost d orbitals
  • Atomic size refers to the size of the atom, typically measured as the atomic radius
  • Atomic radius is defined as half the distance between the nuclei of two adjacent atoms
  • Atomic size generally increases down a group
  • This is because each successive energy level further from the nucleus adds more electron-electron repulsion, causing the size to increase
  • Atomic size generally decreases across a period
  • This is because the effective nuclear charge (positive charge felt by valence electrons) increases, pulling the electrons closer to the nucleus, and thus reducing the size
  1. Bohr Model of Atom - Bohr 2nd Postulate
  • Electrons move in fixed, circular orbits around the nucleus
  • Orbits are called energy levels or shells
  • Each energy level has a specific energy associated with it
  • Electrons can only exist in these energy levels
  • Electrons can transition between energy levels by absorbing or emitting energy
  1. Bohr Model of Atom - Energy Levels and Quantum Numbers
  • Energy levels are denoted by the principal quantum number, n
  • Each energy level can have a maximum number of electrons, given by 2n^2
  • The first energy level (n=1) can hold a maximum of 2 electrons
  • Subshells are denoted by the angular momentum quantum number, l
  • Each subshell can hold a maximum number of electrons, given by 4l + 2
  1. Bohr Model of Atom - Electron Configurations
  • Electron configurations describe the arrangement of electrons in an atom
  • Aufbau principle: Electrons fill the lowest energy levels first
  • Hund’s rule: Electrons fill degenerate orbitals one by one before pairing up
  • Pauli exclusion principle: No two electrons can have the same set of quantum numbers
  • Electron configurations can be written using the noble gas notation
  1. Bohr Model of Atom - Electron Affinity
  • Electron affinity is the amount of energy released when an electron is added to a neutral atom
  • It is a measure of the atom’s tendency to gain electrons
  • Electron affinity values can be positive or negative
  • Positive values indicate that energy is required to add an electron
  • Negative values indicate that energy is released when an electron is added
  1. Bohr Model of Atom - Ionization Energy
  • Ionization energy is the amount of energy required to remove an electron from a neutral atom
  • It is a measure of the atom’s tendency to lose electrons
  • Ionization energy values can vary depending on the energy level
  • First ionization energy refers to the energy required to remove the first electron
  • Ionization energy generally increases across a period
  1. Bohr Model of Atom - Electron Configuration and Periodic Table
  • Electron configurations can determine the position of an element in the periodic table
  • Elements in the same group have similar electron configurations
  • Elements in the same period have sequential addition of electrons
  • Group 1 elements have a configuration of ns1
  • Group 17 elements have a configuration of ns2 np5
  1. Bohr Model of Atom - Orbital Diagrams
  • Orbital diagrams show the arrangement of electrons in an atom using boxes and arrows
  • Boxes represent orbitals and arrows represent electrons
  • Hund’s rule is followed when filling degenerate orbitals
  • Orbital diagrams can determine electron configurations and reactivity of elements
  • Example: Oxygen (O) configuration: 1s2 2s2 2p4
  1. Bohr Model of Atom - Electron Pairing and Valence Electrons
  • Valence electrons are in the outermost energy level of an atom
  • They are involved in chemical bonding and reactivity of elements
  • Unpaired electrons make an atom more reactive
  • Valence electrons can be determined from the electron configuration
  • Example: Carbon (C) configuration: 1s2 2s2 2p2, Valence electrons: 4
  1. Bohr Model of Atom - Periodic Trends in Atomic Size
  • Atomic size refers to the size of the atom, measured as the atomic radius
  • Atomic size generally increases down a group
  • Atomic size generally decreases across a period
  • Effective nuclear charge increases across a period
  • Example: Atomic sizes of Li (largest) to F (smallest) increase across period 2
  1. Bohr Model of Atom - Summary
  • The Bohr model of the atom explains electron behavior using energy levels and quantum numbers
  • Electron configurations determine the arrangement of electrons in an atom
  • Electron affinity measures an atom’s tendency to gain electrons
  • Ionization energy measures an atom’s tendency to lose electrons
  • Valence electrons and orbital diagrams play a key role in chemical bonding and reactivity
  1. Quantum Mechanical Model of the Atom
  • Bohr’s model was unable to explain the behavior of atoms beyond hydrogen
  • Quantum mechanical model is a more accurate model
  • It considers electrons as both particles and waves
  • Electrons are described by wavefunctions and probability distributions
  1. Wave-particle Duality
  • Electrons exhibit both particle-like and wave-like properties
  • Electrons can be described by a wavefunction, Ψ
  • The probability of finding an electron at a certain position is given by |Ψ|^2
  • Electrons can exist in superposition states, where they are in multiple states simultaneously
  1. Heisenberg’s Uncertainty Principle
  • Heisenberg’s Uncertainty Principle states that it is impossible to simultaneously determine the exact position and momentum of an electron
  • ΔxΔp ≥ h/4π, where Δx is the uncertainty in position, Δp is the uncertainty in momentum, and h is Planck’s constant
  1. Quantum Numbers
  • Quantum numbers describe the energy, shape, orientation, and spin of an electron
  • Principal quantum number (n) determines the energy level and size of the orbital
  • Angular momentum quantum number (l) determines the shape of the orbital
  • Magnetic quantum number (ml) determines the orientation of the orbital in space
  • Spin quantum number (ms) determines the spin of the electron (+1/2 or -1/2)
  1. Electron Orbitals
  • Orbitals are regions in space where there is a high probability of finding an electron
  • s orbitals are spherical and have 1 orientation (l=0)
  • p orbitals are dumbbell-shaped and have 3 orientations (l=1)
  • d orbitals are complex and have 5 orientations (l=2)
  • f orbitals are even more complex and have 7 orientations (l=3)
  1. Electron-Filling Order and Stability
  • Electrons fill orbitals in order of increasing energy
  • Pauli Exclusion Principle: No two electrons in an atom can have the same set of quantum numbers
  • Hund’s Rule: Electrons fill orbitals singly with parallel spins before pairing up
  • Filled and half-filled orbitals are more stable than partially-filled orbitals
  1. Electron Configuration Notation
  • Electron configuration notation represents the distribution of electrons in an atom
  • The orbital notation uses arrows to represent electrons in each orbital
  • The electron configuration notation uses the principle, angular momentum, and magnetic quantum numbers to represent the electrons in each energy level and subshell (e.g., 1s^2 2s^2 2p^6)
  1. Valence Electrons and Chemical Properties
  • Valence electrons are the outermost electrons of an atom
  • They are responsible for the chemical properties and reactivity of elements
  • The number of valence electrons can determine the group number of an element in the periodic table
  • For example, group 1 elements have 1 valence electron, group 14 elements have 4 valence electrons
  1. Periodic Trends in Atomic Radius
  • Atomic radius is the size of an atom
  • Atomic radius generally increases down a group due to the addition of energy levels
  • Atomic radius generally decreases across a period due to the increasing effective nuclear charge (positive charge experienced by valence electrons)
  • Exceptions to the trend occur when electron-electron repulsion in larger atoms leads to an increase in atomic radius
  1. Summary
  • The quantum mechanical model provides a more accurate description of the behavior of electrons in atoms
  • Wave-particle duality describes the dual nature of electrons
  • Heisenberg’s Uncertainty Principle sets a limit on the precision of simultaneous position and momentum measurements
  • Quantum numbers describe the energy, shape, orientation, and spin of electrons
  • Orbitals represent regions of high probability of finding electrons
  • Electron configuration notation represents the distribution of electrons in an atom
  • Valence electrons determine the chemical properties and reactivity of elements
  • Atomic radius generally increases down a group and decreases across a period with some exceptions