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Modern Atomic Theory and Chemical Bonding. Copyright Wayne Suits 2004. Rutherford's Atom. Rutherford showed that:An atom is composed of a positive nucleus surrounded by electrons (negative).The nucleus contains protons (positive) and neutrons (neutral).The nucleus is very small compared to the si
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1. For practice problems on the material covered in this power-point presentation visit the following address http://faculty.ccri.edu/wsuits/Sample%20Problems.htm
2. Modern Atomic Theory and Chemical Bonding Copyright Wayne Suits 2004
3. Rutherford’s Atom Rutherford showed that:
An atom is composed of a positive nucleus surrounded by electrons (negative).
The nucleus contains protons (positive) and neutrons (neutral).
The nucleus is very small compared to the size of the entire atom.
Questions left unanswered:
How are the electrons arranged around the nucleus and what are their travel paths like? Slide #8 is Problems with Rutherford’s Nuclear Model of the Atom.Slide #8 is Problems with Rutherford’s Nuclear Model of the Atom.
4. The Rutherford atom – a tiny, dense nucleus that contains all of the positive charge and most of the mass of the atom surrounded by electrons
5. Electromagnetic Radiation (Light) as a Key to Understanding Electron Paths Early scientists discovered that Electromagnetic radiation (light) is given off by atoms of an element when they have been excited by some form of energy
Furthermore, atoms of different elements give off different colors of light when they are excited.
6. When salts containing Li, Cu, and Na dissolved in methyl alcohol are set on fire, brilliant colors result.
7. Show Clip From Brown and LeMay
8. Spectral Analysis of Emitted Light from Excited Atoms When the emitted light from excited atoms was passed through a prism a curious spectrum of discrete lines of separate colors, separate energies was observed rather than a continuous spectrum of ROY G BIV.
Furthermore, different elements show totally different line spectra.
In fact, line spectra are used to identify the presence of different elements
10. Interpretation of Line Spectrum of Elements Atoms which have gained or absorbed extra energy from some excitation energy source (flame, electric discharge etc.) release that energy in the form of light
The light atoms give off contain very specific wavelengths called a line spectrum
light given off = emission spectrum
Each element has its own line spectrum which can be used to identify it
12. Figure: 06-12a,b
Figure: 06-12a,b
13. Interpretation of Atomic Spectra The line spectrum must be related to energy transitions in the atom.
Absorption = atom gaining energy
Emission = atom releasing energy
Since all samples of an element give the exact same pattern of lines, every atom of that element must have only certain, identical energy states
The energy of an atom is quantized – limited to discrete values
If the atom could have all possible energies, then the result would be a continuous spectrum instead of lines
14. Interpretation of Atomic Line Spectra in terms of Electron Paths electrons may be thought of as traveling in concentric shells or energy levels about the nucleus.
the energy of the shells increase as one proceeds away from the nucleus.
When an atom absorbs energy, electrons are promoted from an inner, low energy, shell to an outer, higher energy shell.
Conversely, when an excited atom emits energy, electrons drop down from an excited outer higher energy shells to an inner lower energy shells
16. An excited lithium atom emitting a photon of red light to drop to a lower energy state.
17. A sample of H atoms receives energy from an external source) (b) The excited atoms (H) can release the excess energy by emitting photons.
18. When an excited H atom returns to a lower energy level, it emits a photon that contains the energy released by the atom.
19. When excited hydrogen atoms return to their lowest energy state, the ground state, they emit photons of certain energies, and thus certain colors.
20. Hydrogen atoms have several excited-state energy levels.
21. Each photon emitted by an excited hydrogen atom corresponds to a particular energy change in the hydrogen atom.
22. (a) Continuous energy levels. Any energy value is allowed. (b) Discrete (quantized) energy levels. Only certain energy states are allowed.
23. The difference between continuous and quantized energy levels can be illustrated by comparing a flight of stairs with a ramp.
24. Show Clip from Brown and LeMay
25. Niels Hendrik David Bohr (1885-1962)
26. The Bohr model of the hydrogen atom represented the electron as restricted to certain circular orbits around the nucleus.
27. I. Atomic Structure (Review) Atoms are primarily composed of 3 sub atomic particles.
28. An atom is neutral if # e-’s = # p’s. If a neutral atom gains extra electron(s) then it becomes a negatively charged species called an anion.
If a neutral atom loses electron(s) then it becomes a positively charged species called a cation.
29. An atom is completely characterized by two numbers; the atomic #(Z) and the atomic mass # (A). Atomic # (Z) - the # of protons in the nucleus - responsible for identity of the element.
Mass # (A)- the total # of protons plus neutrons.
30. Representing Atoms of an Element An atom may be represented as its Symbol preceded by its subscripted atomic number, Z, and its superscripted atomic mass number, A.
31. Arrangement of the subatomic particles within the atom At the center is the nucleus which contains the protons and neutrons.
electrons may be thought of as traveling in concentric shells or energy levels about the nucleus.
the energy of the shells increase as one proceeds away from the nucleus.
32. There is a max. # of e-’s that can be accommodated in each shell.
33. Shell diagram for neutral atom of Phosphorus (P)
34. Further development of atomic model. Each shell is composed of 1 or more subshells.
Each shell has as many subshells as its own number.
1st shell has 1 subshell.
2nd shell has 2 subshells.
3rd shell has 3 subshells.
4th shell has 4 subshells.
35. There are only four different kinds of subshells. These subshells are labeled, in order of increasing energy, by the letters s, p, d & f. Each subshell can accomodate a different # of e-’s
36. Thus the total capacity of shell is distributed amongst its subshells.
38. Atomic subshells in order of increasing energy, filling order. NOTE:
Although the 4th shell is higher in energy than the 3rd shell, not all subshells of the 4th shell are higher in energy than all subshells of the 3rd shell. In fact, the highest subshell of the 3rd shell (3d) is higher in energy than the lowest subshell of the 4th shell (4s)
39. Further development of atomic model
40. Further development of atomic model
41. Now, to fully develop our theory of atomic structure we must understand that the subshells (s, p, d, f) of our earlier atomic model consist of orbitals that are not all concentric in shape. Furthermore, any one orbital can only accommodate 2 e-’s. Consequently, the number of orbitals that comprise a subshell can easily be calculated by simply dividing the subshell capacity by 2.
42. Number of Orbitals in each Subshell Any s subshell has a capacity of 2 e-’s
The number of orbitals that comprise any s subshell is 1.
Any p subshell has a capacity of 6 e-’s
The number of orbitals that comprise any p subshell is 3.
Any d subshell has a capacity of 10 e-’s
The number of orbitals that comprise any d subshell is 5.
Any f subshell has a capacity of 14 e-’s
The number of orbitals that comprise any f subshell is 7.
43. Orbitals (s p d + f) All orbitals of the same kind have the same 3 dimensional shape but different sizes. The size increases with the energy level. All s subshells consist of one s orbital that is spherically symmetrical about the nucleus. An s orbital can accommodate 2 e- This accounts for the 2e- capacity of the s subshell
44. s Orbitals
45. Figure: 06-19
Figure: 06-19
46. Figure: 06-16
Figure: 06-16
47. Each p subshell actually consists of a set of three p orbitals of equal energy; px py pz. Each of the three p orbitals is dumbbell shaped and all are oriented in space perpendicular to one another.
The max. capacity of each p orbital is 2e-.
This accounts for the total capacity of the p subshell as being 6 e-’s.
48. Figure: 06-20
Figure: 06-20
49. Shown together the three p orbitals look like this:
50. The d subshell actually consists of a set of five d orbitals of equal energy. Each d orbital can hold a maximum of 2e-. This accounts for the total capacity of the d subshell as being 10 e-’s. The d orbitals do not play as important a role in the chemistry that we will be discussing therefore their shapes and names need not be memorized.
51. Figure: 06-21
Figure: 06-21
52. Electron Spins Electrons spin on their axis
Physics tells us that any charged species that spins, generates a magnetic moment. That is to say, it acts like a tiny bar magnet with a North and a South Pole.
Furthermore, the “Right Hand Rule” tells us that if we wrap the fingers of our right hand around the spinning species, in the direction of the spin, then our thumb will be pointing to the magnetic north.
53. Figure: 06-23
Figure: 06-23
54. Represeanting Electrons Therefore, because of their magnetic moments, we generally represent electrons using a single barbed arrow. The tip of the arrow points to the magnetic north of the electron.
55. Atomic Orbitals in order of Increasing Energy
56. Figure: 06-22
Figure: 06-22
57. Ground - state electron configurations This refers to the lowest energy arrangement of e-’s in orbitals about the nucleus.
To obtain this ground - state electron configuration electrons are assigned to the orbitals of the previous slide according to the three rules.
58. Rules for Filling Orbitals Always fill the lowest energy orbitals first.
The two electrons that occupy any orbital must have opposite spins.
When filling orbitals of equal energy (those of the p,d,or f subshells) put one electron in each orbital with their spins parallel until all are half filled, then go back and pair them.
59. Orbital Electron Configurations Write the orbital electron configuration for P
Write the orbital electron configuration for O
60. When liquid oxygen is poured between the poles of a magnet, it “sticks” until it boils away. Oxygen is magnetic because of its unpaired electrons
61. Using the periodic table to write electron configurations The P.T. is arranged such that each horizontal row (period) represents the filling of orbitals in their proper order.
62. Figure: 06-27
Figure: 06-27
63. More information from the Periodic Table
The term valence electron refers to the # of e-’s in the outermost energy level or shell of an atom.
64. For all main group elements the # of the column (family) of the Periodic Table in which the symbol for the element occurs = the # of valence electrons.
65. Lewis Structures These are shorthand techniques for emphasizing the outer shell or valence e-’s of an atom by representing an atom as its symbol surrounded by its valence e-’s, the e-’s in the atoms outermost shell. Note that the symbol of the element represents the nucleus plus all inner shell e-’s.
66. Write Lewis dot structures for carbon, hydrogen, oxygen, nitrogen and chlorine.
67. Why do atoms react together to form compounds? Atoms react with one another to form compounds in an attempt to achieve the e- configuration of their nearest noble gas neighbor (family 8). The reason for this is that the e- configuration of the noble gases represents an extremely stable situation.
68. There are two ways in which atoms can bond together so as to achieve the e- configuration of their nearest noble gas neighbor. They can loose or gain the necessary e-’s and thereby become ions and ultimately form ionic bonds.
Two or more atoms can share e-’s and form covalent bonds.
69. Ionic Bonds These are formed when ions anions/cations of opposite charge come together. Generally ionic compounds are formed between metals (left of step) and nonmetals (right of step).
70. Consider the formation of the ionic compound magnesium bromide. Magnesium (Mg ) could achieve the e- config. of Neon by loosing 2e- .
71. Bromine could achieve the e- config. of krypton by gaining one e-. Consequently one magnesium combines with two bromine atoms to form MgBr2.
72. Note all atoms in MgBr2 are isoelectronic with their nearest noble gas neighbor. Mg+2 + Br-1 = MgBr2
73. Covalent Bond A covalent bond results from the sharing of an electron pair between two atoms.
Whenever two atoms share a pair of e-’s, it is as if each member of the bonded pair of atoms has gained an extra electron.
As atoms bond together to become isoelectronic with their nearest noble gas neighbors, covalent bonds generally occur when two or more nonmetallic elements (right of step) bond together because the nearest noble gas neighbors for these elements lies ahead of them. Consequently, they all need to gain electrons to become isoelectronic with their nearest noble gas neighbors.
75. Let’s Look at the Water Molecule
76. Kekulé structure for water molecule.
77. Now let’s build the Ammonia Molecule
78. The Covalent Bond and Electronegativity The sharing of an e- pair between two atoms may be equal .
If this is the case then the resulting covalent bond is a nonpolar covalent bond.
If, on the other hand the sharing is unequal then a polar covalent bond results.
79. The reason for this variance in bond polarity is due to the fact that different elements have different tendencies to attract to themselves extra electrons. In other words, each element has a different electronegativity
80. Electronegativity The tendency of an atom, when in combination with other atoms, to attract to itself the bonded (extra) e-’s.
81. Electronegativity values increase from left to right across any horizontal row (period) of the P.T. and they decrease going down any vertical column (family) of the P.T.
Consequently the most electronegative elements are
N, O, F, Cl, Br
82. Figure: 08-06
Figure: 08-06
83. Electronegativity values for selected elements.
84. If two atoms are covalently bonded and one has a high electronegativity and the other has a low electronegative then the electron pair comprising that bond is not shared equally but spends more of its time closer to the more electronegative atom. The immediate result of this unequal sharing is that the more electronegative atom gains a partial negative charge (?-) while the less electronegative element gains a partial positive charge (? +). This type of bond is called a polar covalent bond.
85. The degree to which a covalent bond is polarized is indicated by the electronegativity difference between the two bonded atoms. Refer to next slide for electronegativity values of elements. If the electronegativity difference is greater than .5 but less than 2.0 then the covalent bond is polar.
If the electronegativity difference is less than .5 then the covalent bond is nonpolar.
86. Polar Covalent Bonds in H2O
87. A molecule typical of those found in petroleum. The bonds are not polar.
88. Electronegativity values for selected elements.
89. Ionic Bond and Electronegativity Consideration of electronegativity can demonstrate that ionic bonds are nothing more than an extreme case of a polar covalent bond. In fact…
if the electronegativity difference between two atoms is greater than 2.0, then any bond between these two atoms would be ionic.
90. Molecular Polarity If a molecule contains polar bonds, and if those polar bonds are located such that the ? + charges are at one end of the molecule and the ? - charges are at the other end, then the molecule is a polar molecule. The measure of molecular polarity is a quantity called the dipole moment (D).
91. Like Dissolves Like Polar molecules dissolve in Polar Solvents
Nonpolar molecules dissolve in nonpolar solvents
Polar molecules do not dissolve in nonpolar solvent
Nonpolar molecules do not dissolve in polar solvents
92. : An oil layer floating on water. The oil is nonpolar and the water is polar
93. Polar water molecules interact with the positive and negative ions of a salt. Ionicly bonded materials are the extreme case of polar substances
