• EMR: energy that exhibits wavelike behavior and travels thru space at the speed of light in a vacuum

• Ex: light from the sun, X-rays
• All EMR travel at speed of light
• Each form of EMR is only different in wavelength
• Photon: tiny particle of light that acts as a carrier of energy

• Shorter wavelength = higher frequency = more energy
• So there is an inverse relationship between wavelength and frequency

Essential Formulas

• Formula:

3 Characteristics of Waves

 Wavelength Frequency Speed of light ● Distance between two consecutive peaks ● Unit = Meters       ○ Is often given in nm but must always beconverted           ■ 1m = 10⁹ nm ● The number of waves(cycles) per second ● Unit = s⁻¹ (Hertz) ● 2.9979 X 10⁸

Albert Einstein

• The intensity of light is a measure of the number of photons in a beam
• Greater intensity = more photons are available to release electrons
•
• Mass (kg) of a particle:
• v = velocity
• Can also use equation to calculate wavelength of a particle

The Dual Nature of Light: EMR can show both wave properties and particulate matter properties

• Electrons seem to move in an interference pattern (like a wave) and have the ability to carry energy and momentum when in motion (like a particle) → an electron is both a wave and particle

The Bohr Model

• Predicted that electrons orbit the nucleus at fixed radii
• Atoms absorb energy in the form of electromagnetic radiation → electrons move to higher energy lvl
• As electrons become more tightly bound, its energy becomes more negative
• As the electron is brought closer to the nucleus, energy is released from the system
• Is wrong bcuz does not take into account for sublevels (s, p, d, f), orbitals, or electron spin → electrons aren’t actually locked into orbits

The Quantum Mechanical Model of the Atom

• Quantum Mechanical Model: specifies the probability of finding an electron in the 3D space

• Wave function of an electron: the amplitude where an electron can stay in an atom

• Square of the wave function = the probability of the electron in the orbital

 Shell Subshell Orbital Pathway followed by an electron around the atom’s nucleus Given n Can hold max 32 e Pathway an electron moves inside the shell Given l Maximum e- depends on type of subshell Most probable location to find an electron and its spatial distribution Given m Can hold max 2 e- Energy of a orbital is determined by its value of n → all orbitals with the same value of n have the same energy (are said to be degenerate) Orbital configurations determine the shape of a molecule → determines their properties and how they behave

Heisenberg Uncertainty Principle

• Says that it’s impossible to know the exact motion/position of an electron as it moves around the nucleus
• Two types of electron motions in an atom
• Orbit motion around a nucleus
• Magnetic motion created by the spin on its axis
• The probability of finding the electron at a specific position is greatest the closer it is to the nucleus

Different Types of Wave Functions

 S Sublevel P Sublevel D Sublevel F Sublevel 1 orbital = max 2 electrons → the simplest 3 orbitals = max 6 e- There are 3 P orbitals because are talking about 3D space → can be one orbital on the x-axis, y-axis, and z-axis      ● Each axis canhold a total of 8electrons (s=2 &p=6)→ octet rule 5 orbitals → max 10electrons 7 orbitals → max 14electrons

Quantum Numbers

• Each of these orbitals is characterized by a series of numbers called quantum numbers
 Name Symbol Allowed Values Notes Principle quantum Number n 1, 2, 3, 4, … Describes the size and energy level of an orbital + relative distance from nucleus Is equal to the number of sublevels Ex: n = 2 (2nd electron shell) → 2s & 2p orbitals n increases → orbital becomes larger = higher energy electrons Angular momentum/azimuthal quantum number 1 0 ≤ L ≤ n-1 Describes the shape of an orbital L = 0 → s orbital L = 1 → p orbital L = 2 → d orbital L = 3 → f orbital L ≤ n-1 ● Ex: n = 2 L = 0 → s sublvl L = 1 → p sublvl Magnetic quantum number m L ≤ m ≤ L Describes the orientation of the orbitalTo determine value, draw orbitals → place electron → ans will be where the last electron is Ex: ● S sublvl has 1 orbital and l = 0→ m = 0 ● P sublvl has 3 orbitals & l = 1→ -1 ≤ m ≤1 Spin quantum number m +½. -½ Describes the spin of an e-→ an e- can only spin in a clockwise or counter-clockwise direction + ½ ( up arrow) or -½ (down arrow) Answer will be the spin of the last electron placed

Example: 2p⁵ → find n, l , m◻, m◻

• N =2
• L =1
• m◻→ 0
• Ms → -½

Important Principles

• Pauli Exclusion Principle: two electrons which share an orbital cannot have the same spin → have different values of m◻

• Aufbau principleelectrons are placed in orbitals, shells, and subshells of increasing energy

• Hund’s rule: Every orbital in a sublevel is singly occupied before any orbital is doubly occupied + All of the electrons in singly occupied orbitals have the same spin

• Because of the symmetrical distribution of electrons, orbitals in which the subshell is exactly half-filled or filled are more stable → so removing an electron from these atoms requires more energy

Radial Probability, Penetration, and Electron Repulsion

• Electrons are attracted to the nucleus at the same time as electrons repel each other.
• Penetration: The ability of an electron to get close to the nucleus

• Penetration depends on the shell (n) and subshell (ml)
• Within same shell value (n), penetrating power follows trend in subshells (ml)
• s>p>d>f
• For different shall values and subshell (l), penetrating power follows this trend
• 1s>2s>2p>3s>3p>4s>3d>4p>5s>4d>5p>6s>4f….
• The closer an electron comes to the nucleus/more it penetrates → the stronger its attraction to the nucleus
• Penetration and shielding result in an Effective force that holds the outer electrons to the atom

Electron Shielding

• Shielding: describes the blocking of the attraction between valence electrons and nucleus by the inner-shell electrons

• Results from balance between attractive and repulsive forces

Writing electron configuration

Electron configuration: tell us where electrons are approx

○ For D block elements, count starting from S block

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

○ Use as shortcut → on AP exam can always use

● Ex:

Ion Configurations

• Electrons are removed/added to the valence energy lvl first → only then can e- be removed/added from the d sublvl
• Ex:

Write out configuration for pure element first and then remove/add electrons

• Can’t just keep noble gas in brackets → have to back up to noble gas before