Chemistry Unit 2: the 2 nd half! Electrons and their Properties

I. Why isn’t the electron pulled into the nucleus? A. Remember the nucleus is positive and the electron is negative. B. However, there is a special relationship between light (electromagnetic radiation) and the electrons of an atom

II. Properties of Light A. made of many particles B. described as a wave

C. Electromagnetic Radiation  1. Includes: x rays, ultraviolet, infrared, microwaves, radio waves, visible light  2. Constant speed: 3.0 x 10 8 m/s

D. Wavelength (λ)  1. distance between 2 waves  2. unit varies depending on type of EM ray  3. List the types of EM radiation in order from longest wavelength to shortest wavelength.

E. Frequency (v)  1. number of waves passing a point in a given amount of time  2. Unit: waves per second 1 wave/second = 1 Hz (Hertz)  3. Arrange EM radiation types form highest frequency to lowest frequency

F. Mathematical Relationship c = λv c : speed of light λ : wavelength v : frequency

III. Photoelectric Effect A. Emission of electrons from a metal when light shines on the metal B. Wave model could NOT explain this  1. light below a minimum frequency would not knock off any electrons

C. Max Planck  1. Objects emit energy in small specific amounts know as quanta  2. Quantum – minimum amount of energy that can be lost or gained by an atom

3. Mathematical Relationship  E quantum = hv E: energy (Unit is Joules) v: frequency (Hz) h: Planck’s Constant h = x J s

D. Einstein  1. Light has dual nature  2. wave and particle-like properties  3. made of photons 1. photons have zero mass and a quantum of energy  4. Mathematical relationship E photon = hv

IV. Atomic Emission Spectrum A. When energy passes through a gas at low pressure  the potential energy of some of the atoms increases

B. Ground State  Excited State  1. Ground state: lowest energy state of an atom  2. Excited state: atom has higher potential energy than ground state

C. Excited State  Ground State  1. gives off energy gained  2. produces colors  3. Example: Neon signs

D. Atomic Emission Spectra  1. Set of frequencies emitted by the element  2. Unique to each element Ex. Strontium – Red Hydrogen – Pink  3. Known as a line-emission spectrum

Hydrogen line emission spectrum

V. Bohr’s Model & Atomic Emission Spectrum A. Bohr proposed that electrons can circle the nucleus in fixed paths known as orbits  1. Electron in orbit = fixed energy  2. Therefore, the lowest energy state is the orbit closest to the nucleus a. energy increases with each successive orbit away from the nucle

B. Compare Bohr’s model to the rungs of a ladder. (Think about potential energy)

C. Bohr’s model can be used to explain atomic emission lines  1. in orbit, energy is neither gained nor lost  2. Gain energy  move to higher orbit (Excited state)  3. Electron drops down  photon emitted E photon = higher orbit energy – lower orbit energy

D. Bohr’s Model only worked for the line- emission spectrum of hydrogen  1. Did not fully account chemical behavior of atoms  2. Electrons DO NOT move in fixed circular orbits

VI. Quantum Model of the Atom A. Louis de Broglie  1. Electrons could behave as waves at specific frequencies  2. Electrons could be bent and diffracted like light waves  3. Electron waves can also interfere with one another a. when waves overlap b. reduces energy in some areas, increases it in others

Electron Diffraction Light Diffraction

4. De Broglie’s equation for electron behavior: λ = h__ mv

B. Heisenberg Uncertainty Principle  1. Dealt with detection of electrons  2. Photons have about the same energy as electrons a. When photons are used to locate electrons, they knock them off. b. There is always uncertainty when trying to locate electrons

3. Heisenberg Uncertainty Principle states: it is impossible to determine simultaneously both the position and velocity of an electron or any other particle  a. fundamental to the foundation of scientists understanding of light and matter

C. Schrodinger Wave Equation  1. developed equation that treated electrons as waves  2. Along with Heisenberg Principle, laid the foundation of Quantum Theory  3. Quantum Theory: describes mathematically the wave properties of electrons and other very small particles a. Proposed probable locations of electrons b. Exist in regions known as orbitals (3-dimensional region around the nucleus that indicates the probable location of an electron)

VII. Atomic Orbitals & Quantum Numbers A. Quantum Numbers  1. Specify the properties of atomic orbitals and the properties of electrons in the orbitals  2. Indicate the following for an orbital: Main energy level The shape The orientation

3. Types of Quantum Numbers  Principal Quantum Number (n)  Angular Momentum Quantum Number (l)  Magnetic Quantum Number (m)  Spin Quantum Number (+1/2 or -1/2)

a.Principal Quantum Number (n)  1. main energy level  2. n = a positive integer (1, 2, 3, 4, etc)  3. Often referred to as shells

b. Angular Quantum Numbers (l)  1. Exist at each Main energy level AFTER the first one  2. Known as sublevels – orbitals of different shapes  3. Indicates the shape of the orbital s, p, d, f  4. Each orbital is represented by a principal quantum number and the letter of the sublevel

c. Magnetic Quantum Numbers (m)  1. Indicates orientation around the nucleus  2. s sublevel has only one orientation, but p, d, and f have multiple possibilities.

3 orientations for the p sublevel

5 orientations for the d sublevel

d. Spin Quantum Number (+1/2 or -1/2)  1. electrons spin on an internal axis  2. single orbital holds 2 electrons and each has an opposite spin Pauli Exclusion Principle

VIII. Electron Configuration A. method of indicating the arrangement of electrons around a nucleus B. Consists of the following:  1. A number indicates energy level (n)  2. A letter indicates type of orbital: s, p, d, f (l)  3. A superscript indicating the number of electrons in the orbital  Example: 1s 2 (This is Helium)

C. Rules for Electron Configuration 1. Aufbau Principle  Shows the order in which electrons occupy orbitals  States: an electron occupies the lowest-energy orbital that can receive it

2. Pauli Exclusion Principle  Importance of the spin quantum number is reflected here  States: no two electrons in the same atom can have the same 4 quantum numbers  How can an orbital hold two electrons then? OPPOSITE SPINS!!!!

3. Hund’s Rule  Placing unpaired electrons in separate orbitals in the same sublevel  States: orbitals of equal energy are each occupied by a one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin.

D. Methods 1. There are three methods for writing electron notation  A. Orbital Notation – Shows spin direction of each electron  B. Electron-Configuration Notation – Uses the superscript with the orbital letter to represent number of electrons in each  C. Noble Gas Notation – shorthand method used to represent larger elements Noble gases are in Group 18!

E. How do we do this? (STEPS) 1. Determine the total number of electrons to be represented 2. Use the Aufbau process to fill the orbitals with electrons 3. The sum of the superscripts should equal the total number of electrons

F. Practice (Assume atoms are neutral) Hydrogen  Orbital Notation:  EC Notation: Helium  Orbital Notation:  EC Notation Boron  Orbital Notation:  EC Notation:

G. Noble Gas Notation 1. Noble Gases – Group 18 Elements  Helium, neon, argon, krypton, xenon, and radon 2. Method of simplifying the electron configuration of larger elements

H. Practice Problems Write the complete electron configuration for Iron (Fe) and the noble-gas notation.

Write both the electron configuration and noble gas notation for an atom of Barium.