Chapter 5

 Scientific Models  Models are things used to represent real phenomena.  simplify and explain complex realities.  can take many forms  scale models, e.g. a globe  mathematical models, e.g. P/V = k  computer models, etc., e.g. weather predictions

 It explained much about the structure  Nucleus: positive, very dense, most of atom’s mass  Electrons: outside the nucleus  Empty space: most of the volume of the atom  It could not explain chemical behavior of elements, such as….  Why did elements give off light when heated?  Why did one element react with another to form a new compound?

Rutherford’s Model Rutherford’s model could not explain why matter gave off light when heated

The Bohr Model Neils Bohr Danish Physicist 1913: Proposed new model of the atom

 Bohr proposed that an electron is found only in specific circular paths, or orbits, around the nucleus.  Each possible electron orbit in Bohr’s model has a fixed energy.  The fixed energies an electron can have are called energy levels.  Higher energy levels are farther away from the nucleus  A quantum of energy is the amount of energy required to move an electron from one energy level to another energy level.

 Energy levels are like rungs on a ladder  Higher energy levels are closer together  Takes less energy to change between higher levels

The Bohr Model  When electrons absorb exactly the right quanta of energy….  They jump to higher energy level  When it jumps back down…  It gives off (emits) the same energy as light.

ERWIN SCHRÖDINGER  Bohr’s planetary model only worked for hydrogen  But it could not explain motion of electrons  Schrödinger and others developed a new mathematic model of the atom….  Called the quantum mechanical model

 Like Bohr’s model, electrons are restricted to certain energy levels  Unlike Bohr’s model, the exact pathway of the electron is uncertain  Locations of electrons are uncertain, and described terms of probability….  i.e. the likelihood of finding the electron at a given point in time

 Electrons are found within an “electron cloud” outside the nucleus

The electron cloud is more dense where the probability of finding the electron is high.

 A spinning fan blade  Forms a ‘fuzzy’ image  You know the fan blade is within the fuzzy region, but at any point in time you don’t know exactly where it is  Electrons are located in regions of probability called “ orbitals ”

Quantum Number DefinesDescribesValues 1 st PrincipalEnergy Level-----n = 1 to 7 2 nd Angular Momentum Energy Sub- level Shapes, p, d, f 3 rd MagneticOrbital3-D orientationx, y, z, etc. 4 th Spin-----Magnetic spin+1/2 or -1/2

AN “ s ” ORBITAL  Orbitals are represented by “electron density maps”  Probability is represented by the density of color  The more probable location of the electron is in the darker blue region

 Regions of space in which there is a high probability of finding an electron  Various types of orbitals exist, depending upon the sublevel  S sublevels have one orbital  P sublevels have 3 orbitals each

 d sublevels have 5 orbitals each  f sublevels have 7 orbitals each

Energy Level # Sublevels# Orbitals Electron capacity nnn2n2 2n22n

 Each orbital can contain up to 2 electrons! Sublevel# Orbitals per sublevel Electron capacity per sublevel s 12 p 36 d 510 f 714

 In most natural phenomena, change trends toward lower energy  Systems are more stable when they have less energy.  Electrons also tend to arrange themselves in their lowest energy states.  The arrangement of electrons within an atom is called an electron configuration.

 Three rules are used to determine electron configurations  Aufbau Principle  Pauli Exclusion Principle  Hund’s Rule

 Electrons occupy the lowest energy level first  This diagram is known as an electron orbital diagram

 4 th quantum number is the “spin” number  Electrons “spin”, either clockwise & counter- clockwise  Spin is symbolized ↑ or ↓  PEP says….  Two electrons in the same orbital must have opposite spins. Therefore ….  No two electrons in an atom can have the same identical set of 4 quantum numbers.

 Electrons fill orbitals within a sublevel such that they have maximum number of unpaired spins  This is because they have the lowest energy this way

 Determine the number of electrons in the diagram. How?  Begin filling orbitals at the lowest energy level (Aufbau principle)  Continue filling, applying Hund’s rule  All “up” spins  Follow by “down” spins  Stop when you have assigned all the electrons to orbitals

 A shorthand way for writing orbital diagrams  Write the energy level, sublevel, and number of electrons in the sublevel  Li1s 2 2s 1  C1s 2 2s 2 2p 2  N1s 2 2s 2 2p 3  O1s 2 2s 2 2p 4  F1s 2 2s 2 2p 5  Ne1s 2 2s 2 2p 6  Na 1s 2 2s 2 2p 6 3s 1

 Periods (rows) in the PT correspond to energy level  Certain groups (columns) correspond to the sublevels (s, p, d, f) (see page 166)

 Also called “noble gas notation”  An element’s electron configuration contains the e-config of a noble gas (group VIIIA, 18)  Begin with the preceding noble gas  Then complete the e-config

 Transition elements (groups 3-12) tend to prefer half-filled or completely filled d-orbitals at the expense of the s-orbital.  For chromium, you would expect  ….4 s 2 3 d 4,  but in fact one of the 4 s electrons is promoted to 3 d, resulting in  ….4 s 1 3 d 5  Try copper….

 Much of what is known about the atom is due to the study of light  Light has properties of waves  Waves have amplitude, wavelength, and frequency

 Inversely proportional  c = speed of light = 3.00 x 10 8 m/s  (constant)  lambda = wavelength (meters)  nu = frequency (Hertz, Hz, s -1 )

 Visible light is a small portion of the electro- magnetic spectrum

 All EM radiation travels at the same speed  c = 3.00 x 10 8 m/s  EM radiation varies in wavelength and frequency  Longer wavelength → Lower frequency  Shorter wavelength → Higher frequency

 Light separates into different colors (wavelengths) when it passes through a prism  It is a continuous spectrum

 Electrons of an element can absorb energy and emit the energy as EM radiation  These emission spectra are not continuous

 Each element has a unique emission spectra  Like a bar code for an element

 Electron at ground state absorbs a quantum of energy  Excited electron returns to ground state, emitting the quantum as light  Frequency of the light is directly proportional to the energy change of the electron

 Lyman Series is in the UV range  Balmer series is visible  Paschen series is in IR range

 Einstein determined that light behaved like a particle  “Particle” of light is the photon  Photon is a quantum of light  So light can behave as a wave and a particle, which is it?  Both

 If light (a wave) can behave as a particle, can a particle behave as a wave?  Yes  So electrons can be thought of as waves.

 Uncertainty principle  It is not possible to know the location and momentum (speed) of an electron at the same time

 Schrodinger equation  Mathematically described sublevels and orbitals