Chapter 4

 Why aren’t e- drawn into nucleus?  Why do atoms of some elements behave the way they do?

Niels Bohr studied under Rutherford Worked w/ him on gold foil experiment Bohr refined Rutherford's idea by adding that the e- were in orbits like planets orbiting the sun each orbit only holding a set number of e-

Bohr’s Atom electrons in orbits nucleus

 to aid in finding out answers, scientists analyzed light samples from elements heated in flames  began to see pattern where chemical behavior of elements related to arrangement of e- in atoms  study of light necessary

 EMG- electromagnetic radiation  visible light, x rays, infrared, UV waves  described by 1. wavelength ( )- distance between 2 consecutive crests 2. frequency ( )- # of waves that pass a certain pt per second Hertz (Hz)- SI unit 1 wave/second 3. amplitude- wave’s height form origin to crest

 all EMG radiation travels at the speed of light (c) 3.00 x 10 8 m/s  speed = wavelength x frequency  higher the frequency, the higher the energy

 light can be explained as moving thru space in form of waves, but not in its interactions w/ matter  early 1900s, scientists conducted 2 exp that didn’t line up w/ light being a wave  photoelectric effect - emission of e- from metal when light of certain frequency shines on metal  prob: light as a wave should knock off e- regardless of light’s frequency, but not the case  Max Planck- studied hot obj & their emission of light  did not emit light continuously if light was only in form of waves

 Planck suggested obj emits energy in small packets called quanta  proposed quantum theory  quantum - minimum amount of energy that can be lost or gained by an atom  1905, Albert Einstein proposed EMG radiation has dual wave-particle nature  sometimes acts like wave, sometimes acts like particles  Einstein suggested light is made up of streams of particles carrying a quantum of energy  particles called photons

 photon- particle of EMG radiation w/ no mass carries quantum of energy  Einstein’s explanation of photoelectric effect  2 nd exp involved hydrogen  when gases have elec current pass thru them, some of their atoms will increase in energy  they will go from their ground state (lowest energy state) to the excited state ( state of higher PE)  when excited atom returns back to its ground state, it emits colored light energy (neon lights)

 worked w/ hydrogen- emitted pink light  passed light thru prism, it separated into 4 specific colors of visible spectrum  these 4 bands of light were hydrogen’s line- emission spectrum  spectrum - pattern of radiant energy (fingerprint)  scientists had expected to see a continuous spectrum if light was in form of waves, but didn’t  looking for explanation of the specific energy states of H e-

 enter Bohr’s theory of e- circling the nucleus in only allowed paths or orbits (planetary model)  e- absorb energy, they move into larger orbit (excited state)  when they emit energy, e- return to original orbit (ground state)  Bohr assigned value to each orbit & calculated radius  mathematically speaking, Bohr’s calculated values for the orbits matched & explained the observed spectral lines of H

 thought all atoms would follow same pattern, but not the case  remember that light has ability to act like waves & particles  using Einstein’s formula & Planck’s quantum theory, Louis de Broglie proposed that it was possible for e- to have the same properties  lead to Wave-particle model of e-  confirmed by exp

 in 1924, Edwin Schrodinger devised an equation that treated e- as moving about the nucleus as waves  equation laid foundation for quantum theory  QT- describe mathematically the wave properties of e- & other small particles  e-, like light waves, can be bent or diffracted & they can interfere w/ each other  So where are e- in atoms?

 theory was only accepted after Werner Heisenberg proposed his uncertainty principle - it is impossible to know both the exact position and the velocity of an object at the same time  e- detected by photons, any attempt to locate e- knocks it off its course  quantum numbers- used to describe e- behavior  e- move about the nucleus at extremely high speeds filling the entire area in e- cloud

 quantum numbers used to describe e configuration  atomic orbital- a region of space in which the probability of finding an e- is high  4 quantum numbers:

 describes the energy level an e- occupies  can only be whole numbers  as n increases, the distance of main energy levels from the nucleus increases & energy increases  known elements utilize main energy levels 1-7

 indicates the shape of the region in e- cloud the e- occupies  regions referred to as sublevels or a specific kind of atomic orbital  the # of possible orbital shapes is equal to the value of n  a letter has been designated to represent each different kind of sublevel

 s = sphere shaped  p = dumbell/ p-nut shaped  d = double p-nut  f = flower shaped  1 st energy level has 1 sublevel (s)  2 nd energy level has 2 sublevels (s, p)  3 rd e.l. has 3 sublevels (s,p,d)  4 th e. l. has 4 sublevels (s, p, d, f)

 indicates the orientation of an orbital about the nucleus  indicated by n 2  s sublevel has only 1 possible orientation, therefore only 1 s orbital in each sublevel  p sublevel has 3 different orientations, lobes extend along x, y, z axes  designated as p x p y p z  these have equal energy

 d sublevel has 5 different d orbitals  d 1, d 2, d 3, d 4, d 5  f sublevel has 7 different f orbitals  f 1 –f 7  each orbital can hold a max of 2 e-  with increasing main energy levels, there are larger # of orbitals

 spin of e- and the orientation of the magnetic field produced by the motion of the e-  indicates also the maximum number of e- the energy level can hold  represented by 2(n 2 )  only 2 possible values -1/2 or +1/2

 the arrangement of e- in atoms  e- tend to arrange themselves in ways that give them the lowest possible energy  3 rules that govern e- conf: 1. Aufbau principle- an e- occupies the lowest energy orbital that can receive it

 aids in applying Aufbau principle

2. Hund’s Rule - orbitals of equal energy are occupied by 1e- before any orbitals are occupied by 2e- sharing room w/ sibling (e- to e- repulsion is minimized therefore lowering energy they have)

3. Pauli exclusion principle- no 2 e- in the same atom can have the same 4 quantum #s the 2 values of the spin quantum # permits 2 e- of opposite spins to occupy the same orbital 3 notations used to indicate e- conf: 1 st 2 indicates ground state e- conf

 unoccupied orbital is represented by a line ____  arrows are used to denote the e- in the orbital  label lines w/principle quantum #s & sublevels  write name or symbol of element before lines

 eliminates arrows & lines  # of e- in sublevel is shown by adding superscripts to the sublevel designation

 can abbreviate by using the noble gas e- configuration that comes before it

 **most imp for representing how atoms form compounds***  shows the e- in the outermost main e- level  known as valence e-  use dots arranged around the element’s symbol to represent these valence e-

The end.