Electrons in Atoms Chapter 5 SAVE PAPER AND INK!!! When you print out the notes on PowerPoint, print "Handouts" instead of "Slides" in the print setup. Also, turn off the backgrounds (Tools>Options>Print>Uncheck "Background Printing")!

Section 5.1 Objectives Define a quantum of energy, and explain how it is related to an energy change of matter. Compare the wave and particle natures of light Contrast continuous electromagnetic spectra and atomic emission spectra. Review Radiation: the rays and particles —alpha particles, beta particles, and gamma rays—that are emitted by radioactive material The Rutherford’s model doesn’t explain why negatively charged electrons aren’t pulled into the positively charged nucleus

Electromagnetic Radiation Visible light is a type of electromagnetic radiation, a form of energy that exhibits wave-like behavior as it travels through space Radio waves, microwaves, visible light, and x rays are all examples of electromagnetic waves that differ from each other in wavelength and frequency Most subatomic particles behave as Particles and obey the physics of waves

Wave Characteristics Hz = 1/sec Frequency () – number of waves that pass a point per second Units are “cycles per sec” or Hertz Hz = 1/sec Wavelength () – distance from crest to crest Amplitude – is the wave’s height from the origin to a crest All radiation: c =  •  c = velocity of light = 3.00 x 108 m/sec Crest Amplitude Wave Cycle Trough

Electromagnetic Spectrum White light (sunlight) is made up of Continuous Spectrum (7 colors of rainbow) ROY G BIV = Red, Orange, Yellow, Green, Blue, indigo, violet Each color has a specific Frequency & wavelength Electromagnetic spectrum

Wavelength vs. Frequency As frequency increases the wavelength decreases (indirectly proportional) Long wavelength --> small frequency --> low energy Short wavelength --> high frequency --> High energy

Bright-line Spectrum Bright-line Spectrum – distinct separate lines of color for an element Atoms emit an atomic emission spectra, characteristic set of discrete wavelengths not a continuous spectrum Atomic Emission spectrum can be used as a “fingerprint” for an element

H, Hg, and Ne Spectrums Light can be grouped into packets called Photons Like matter is divided into atoms Energy of Photon uses the eqn: E = h  E = energy (Joules, J) H = Planck constant 6.6262 x 10-34 J·sec

Section 5.2 & 5.3 Objectives Identify the relationship among an atom’s energy levels, sublevels, and atomic orbital’s Apply the Pauli exclusion principle, the aufbau principle, and Hund’s rule to write electron configurations using orbital diagrams and electron configuration notation Define Valence electrons, and draw electron-dot structures representing an atom’s valence electrons

Terms The wave model of light cannot explain all of light’s characteristics A beam of light has wave like and particle like particles Matter can gain or lose energy only in small, specific amounts called quanta A quantum is the minimum amount of energy that can be gained or lost by an atom A photon is a particle of electromagnetic radiation with no mass that carries a quantum of energy

Electron Configuration Is the arrangement of electrons around the nuclei A listing showing how many electrons are in each orbital or subshell in an atom or ion Example: Bohr atomic model/Quantum Mechanics Bohr's model was rather simplistic and as scientists made more discoveries about more complex atoms, Bohr's model was modified and eventually was replaced by more sophisticated models. The Quantum Mechanical Model of the atom presents a more accurate model of the atom. It is a more sophisticated model based on complex mathematical calculations and interpretations 1s2 2s2 2p6 3s1

Electron Configuration Quantum mechanics is the study of the relationship between energy quanta (radiation) and matter, in particular that between valence shell electrons and photons Electrons in atoms are arranged (or located) as: LEVELS (n) Energy level of electron – region electron is likely to be found SUBLEVELS ORBITALS

Bohr’s Model of the Atom Bohr suggested that an electron moves around the nucleus only in certain allowed circular orbits Energy Added (A quantum) Hydrogen’s single electron is in the n=1 orbit in the ground state When energy is added, the electron moves to the n=2 orbit

How are photons emitted? Electron starts at 2nd energy level grounded state Excited electron steps from current to higher level Quantum energy is the amount of energy required to move an electron from its present energy level (grounded state) to the next higher level (excited state) Energy is absorbed and released by atoms in certain fixed amounts known as quanta quanta = bundles of energy Excited State n=3 Quantum Energy Grounded state n=2 n=1

How are photons emitted? Electron at excited state Electron will “fall” back to 2nd energy level Photon released At a specific color Due to frequency and wavelength The photon is equal in energy to the quantum energy added Excited State n=3 Grounded state n=2 n=1

How are photons emitted? Electron starts at 2nd energy level grounded state 3. Excited again This time adding more Quantum energy Result: higher energy level reached 4. When electron “falls” Photon different color Higher frequency, shorter wavelength More energy = higher energy color Excited State n=4 Quantum Energy n=3 Grounded state n=2 n=1

How are photons emitted? We can only see electrons that start at a grounded state at level 2. Why?? We can only see a certain wavelength and frequency Electrons returning to other levels do emit frequency’s but they are out of the visible light spectrum n=3 n=2 n=1

Electron Configurations Sublevel 2p4 Number of electrons in the sublevel Energy Level 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14… etc.

Principal Quantum Numbers 1s2 2s2 2p6 3s1 Energy Levels are called Principal quantum numbers (n) Assigned values in increasing energy: n=1,2,3,4,5,6,7 (think of n as an address for electrons) corresponds with rows Energy Level n Max # of electrons 1 2 8 3 18 4 32 5 n = 1 n = 2 n = 3 n = 4

Sublevels 1s2 2s2 2p6 3s1 Sublevels are in each Energy level # of sublevels equals the principal quantum number (n) Max # of sublevels is 4 The sublevels are: s, p, d, f Sublevel Max # of electrons s 2 p 6 d 10 f 14

Orbitals 1s2 2s2 2p6 3s1 In each sublevels there are orbitals: S subshell (1 orbital) P subshell (3 orbitals) d subshell (5 orbitals) F subshell (7 orbitals) Each orbital holds 2 electrons, that spin in opposite direction Total # of orbitals in a n level n2 Total # of electrons in a n level 2n2 An atomic orbital is a mathematical function that describes the wave-like behavior of an electron in an atom.

Orbitals and the Periodic Table Orbitals grouped in s, p, d, and f orbitals s orbitals d orbitals p orbitals f orbitals

Rules for electron configuration 1s2 2s2 2p6 3s1 Aufbau Principle Electrons enter orbitals of lowest energy first, starting with the 1s orbital Pauli Exclusion Principle An orbital is filled when it contains 2 electrons.  After that, you have to put the electrons in a different orbital Hund’s Rule Within a subshell, the electrons will occupy the orbitals singly first, and will only pair up when there are no longer any empty orbitals available in that subshell

Valence Electrons Valence Electrons are defined as electrons in the atom’s outermost orbitals – those associated with the atoms highest principal energy level Electron-dot structure consists of the elements symbol representing the nucleus, surrounded by dots representing the element’s valence electrons The octet rule states that atoms with 8 electrons in their outer shell are stable.

Extra Info

An excited lithium atom emitting a photon of red light to drop to a lower energy state.

Orbital occupancy for the first 10 elements, H through Ne.

Order to fill orbital’s

Aufbau principle The Aufbau principle states that each electron occupies the lowest energy orbital available

Dot Diagram order for Valence electrons Below you will see an example of the order of filling in the dots on a dot diagram for an element with eight valence electrons.   Please note that you can place the first two dots on any side, but the rest of the dots should be placed in either a clockwise or counter clockwise manner, with no side receiving two dots until each side gets one. Step 1 Step 2 Step 3 Step 4 Step 5 Step 6 Step 7 Step 8

Quantum Mechanical Model Quantum numbers are used to describe certain aspects of the locations of electrons. The shape, size, and energy of each orbital is a function of 3 quantum numbers n (principal quantum number) energy level, position of the electron with respect to the nucleus Sublevels shape of orbital s,p,d,f Orbitals describes the direction of the electron’s spin within a given orbital (clockwise or counterclockwise) 1s2 2s2 2p6 3s1