Modern Atomic Theory and the Periodic Table Chapter 10 the exam would have to be given earlier Hein and Arena Version 1.1 Eugene Passer Chemistry Department Bronx Community 1 College John Wiley and Sons, Inc.
A Brief Review 2
Composition of the Atom 1. The atom has a dense Nucleus containing Protons( p) and Neutrons( n). 2. n have no charge. p have a relative charge of +1. 3. The atomic number of an element is the same as the number of p in the Nucleus. 4. Electrons( e - ) are located outside of the Nucleus. e - have a relative charge of -1. 3
Electromagnetic Radiation 4
Energy can travel through space as Examples electromagnetic radiation. 5
light from the sun x-rays microwaves radio waves television waves radiant heat All show wavelike behavior. Each travels at the same speed, c, in a vacuum. c = 3.00 x10 8 m/s 6
Characteristics of a Wave 7
Wavelength (λ) 8
Light has the properties of a wave. Wavelength (nm) Wavelength (nm) (measured from (measured from peak to peak) trough to trough) 10.1 9
Frequency (ν) 10
Frequency is the number of wavelengths that pass a particular point per second (i.e. s -1 or Hz). 10.1 11
Speed (v) 12
Speed is how fast a wave moves through space (i.e. a non-vacuum environment). 10.1 13
Light also exhibits the properties of a particle. Light particles are called photons, which have discrete energies (i.e. E = hn, where h is Plank s Constant). The wave model and the particle model are both used to explain the properties of light. 14
The Electromagnetic Spectrum 15
visible light is part of the electromagnetic spectrum Each color in the visible region has a very specific wavelength (l); red light has a l =~650 nm while blue light has a l =~ 450 nm. 10.2 16
At high temperatures, or voltages, gaseous atoms are excited; excited atoms in the gaseous state emit light of different colors upon relaxation. When the light is passed through a prism or diffraction grating a line spectrum results. 17
Each element has its own unique set of spectral emission lines that distinguish it from other elements. These colored lines indicate that light is being emitted only at certain wavelengths. 10.3 Line spectrum of hydrogen. Each line (i.e. color) corresponds to the wavelength of the energy emitted when the electron of a hydrogen atom, which has absorbed energy falls back to a lower principal energy level. 18
The Bohr Atom 19
Electrons An electron revolve has a around discrete the energy nucleus when in it orbits occupies that an are orbit. located at fixed distances from the nucleus. 10.4 20
When The color an electron of the falls light from emitted a higher corresponds energy level to to one a lower of thenergy lineslevel of the a quantum hydrogenof spectrum. energy in the form of light is emitted by the atom. 10.4 21
Different lines of the hydrogen spectrum correspond to different electron energy level shifts. 10.4 22
Light is not emitted continuously. It is emitted in discrete packets called quanta. 10.4 23
E 1 E 2 E 3 Bohr s H atom with quantized energy levels; e - energies: E 3 > E 2 > E 1. 10.4 24
Bohr s calculations succeeded very well in correlating the experimentally observed spectral lines with electron energy levels for the hydrogen atom. Bohr s methods did not succeed for heavier atoms (i.e. atoms with several e - ). More theoretical work on atomic structure was needed. 25
Erwin Schröedinger created a mathematical model that showed electrons as waves; this lead to the duality of matter (i.e. e - :wave or particle). Schröedinger s work led to a new branch of physics called wave or quantum mechanics. Using Schröedinger s equation, the probability of finding an electron in a given region of space around the nucleus can be determined. The actual location of an electron within an atom cannot be determined. 26
Based on wave mechanics it is clear that electrons are not revolving around the nucleus in orbits. Instead of being located in orbits, the electrons are located in orbitals. An orbital is a region around the nucleus where there is a high probability (~90%) of finding an electron. 27
Energy Levels of Electrons 28
The wave-mechanical model of the atom also predicts discrete principal energy According to Bohr the energies of levels for e - of an atom. electrons in an atom are quantized. 29
As n increases, the energy of the electron increases. The first four principal energy levels of the hydrogen atom. Each level is assigned a principal quantum number n. 10.7 30
10.7, 10.8 Each principal energy level is subdivided into sublevels. 31
Within sublevels the electrons are found in orbitals. An s orbital is spherical in shape. The spherical surface encloses a space where there is a 90% probability that the electron may be found. 10.10 32
An atomic orbital can hold a maximum of two electrons. An electron can spin in one of two possible directions represented by (spin up) or (spin down). Two electrons that occupy the same atomic orbital must have opposite spins. 10.10 This is known as the Pauli 33 Exclusion Principal.
A p sublevel is made up of three p orbitals. Each p atomic orbital has two lobes. Each p atomic orbital can hold a maximum of two electrons. A p sublevel can hold a maximum of 6 electrons. 10.10 34
p z p y p x The three p orbitals share a common center, the nucleus. The three p orbitals point in different directions along x, y and z in 3-space. 10.10 35
A d sublevel is made up of five d orbitals. The five d atomic orbitals all point in different directions. Each d atomic orbital can hold a maximum of two electrons. A d sublevel can hold a maximum of 10 electrons. 36 10.11
n=1 1s 1 n=2 2s 1 n = 3 3s 1 n = 4 4s 1 Distribution of Sublevels and Orbitals by Principal Energy Level 2p 2p 2p 3 3p 3p 3p 3d 3d 3d 3d 3d 3 5 4p 4p 4p 4d 4d 4d 4d 4d 4f 4f 4f 4f 4f 4f 4f 3 5 7 Note: The number in RED indicates the number of orbitals in a given sublevel 37
The Hydrogen Atom In the ground state hydrogen s single electron lies in the 1s orbital. Hydrogen can absorb energy and the electron will move to an excited state. 38 10.12
Electron Configurations of the First 18 Elements 39
Electron Configuration Step by Step 40
1. Maximum of two electrons per orbital 10.10 41
1 s orbital 2 s orbital 2. Lowest energy orbitals occupied first; higher energy orbitals occupied only after the lower energy orbitals are filled. 3. Orbital energies: s < p < d < f for a given value of n. 10.10 42
10.10 4. Each orbital in a sublevel is occupied by a single electron before a second electron enters. For example, all three p orbitals must contain one electron before a second electron enters a p orbital. 43
Electron Configuration Arrangement of electrons within their respective sublevels. 2p Principal energy level Number of electrons in sublevel orbitals Type of orbital 6 44
Orbital Filling 45
The following diagrams boxes represent orbitals. Electrons are indicated by arrows: or. Each arrow direction represents one of the two possible electron spin states. 46
Filling the 1s Sublevel 47
H 1s 1 Hydrogen has 1 electron. It will occupy the orbital of lowest energy which is the 1s. He 1s 2 Helium has two electrons. Both helium electrons occupy the 1s orbital with opposite spins. 48
Filling the 2s Sublevel 49
Li 1s 2 2s 1 1s 2s The 1s orbital is filled. Lithium s third electron will enter the 2s orbital. Be 1s 2s 1s 2 2s 2 The 2s orbital fills upon the addition of beryllium s third and fourth electrons. 50
Filling the 2p Sublevel 51
B 1s 2 2s 2 2p 1 1s 2s 2p Boron has the first p electron. The three 2p orbitals have the same energy. It does not matter which orbital fills first. C 1s 2s 2p 1s 2 2s 2 2p 2 The second p electron of carbon enters a different p orbital than the first p electron so as to give carbon the lowest possible energy. N 1s 2 2s 2 2p 3 1s 2s 2p The third p electron of nitrogen enters a different p orbital than its 52 first two p electrons to give nitrogen the lowest possible energy.
O 1s 2s 2p 1s 2 2s 2 2p 4 There are four electrons in the 2p sublevel of oxygen. One of the 2p orbitals is now occupied by a second electron, which has a spin opposite to that of the first electron already in the orbital. F 1s 2s 2p 1s 2 2s 2 2p 5 There are five electrons in the 2p sublevel of fluorine. Two of the 2p orbitals are now occupied by a second electron, which has a spin opposite to that of the first electron already in the orbital. 53
Ne 1s 2 2s 2 2p 6 1s 2s 2p There are 6 electrons in the 2p sublevel of neon, which fills the sublevel. This electronic configuration is referred to as a full octet (ns 2 np 6 ). All noble gases, except He, have a full octet! It is this electron configuration that make noble gases un-reactive. The valence electrons occupy the valence shells (i.e. ns 2 np 6 )
Filling the 3s Sublevel 55
Na 1s 2 2s 2 2p 6 3s 1 1s 2s 2p 3s The 2s and 2p sublevels are filled. The next electron enters the 3s sublevel of sodium. Mg 1s 2s 2p 3s 1s 2 2s 2 2p 6 3s 2 The 3s orbital fills upon the addition of magnesium s twelfth electron. The valence electrons occupy the valence shells (i.e. ns 2 np 6 ) 56
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Which atom above has the largest number of unpaired electrons? P 58
Electron Configurations and the Periodic Table 59
Elements in the same group, or family, have similar chemical properties because the valence electron (i.e. ns 2 np 6 ) configuration is the same in each group, or family. 10.15 60
With the exception of helium which has a filled s orbital, the nobles gases have filled p orbitals. 10.15 61
The electron configuration of any of the noble gas elements can be represented by the symbol of the element enclosed in square brackets; the core. B 1s 2 2s 2 2p 1 [He]2s 2 2p 1 Na 1s 2 2s 2 2p 6 3s 1 [Ne]3s 1 Cl 1s 2 2s 2 2p 6 3s 2 3p 5 [Ne]3s 2 3p 5 62
The electron configuration of argon is Ar 1s 2 2s 2 2p 6 3s 2 3p 6 The elements after argon are potassium and calcium. Instead of entering a 3d orbital, the valence electrons of these elements enter the 4s orbital. K 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 [Ar]4s 1 Ca 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 [Ar]4s 2 63
d orbital numbers are 1 less than dthe orbital period filling number 10.16 Arrangement of electrons according to sublevel being filled. 64
f orbital numbers are 2 less than the f orbital period filling number 10.16 Arrangement of electrons according to sublevel being filled. 65
Period number corresponds with the highest principle energy level occupied by electrons in that period. 10.17 66
10.17 The The elements group numbers of a group for have the representative the same valence elements electron are equal configuration to the total except number that of the electrons valence electrons are in different in the principle atoms of energy the group. levels. 67
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