基础化学（全英语）

应叶

目录

  • 1 Nomenclature
    • 1.1 Nomenclature
    • 1.2 Inorganic compounds
    • 1.3 Organic compounds
  • 2 Atom
    • 2.1 Basic Atomic Theory
    • 2.2 Evolution of Atomic Theory
    • 2.3 Atomic Structure and Symbolism
    • 2.4 Isotopes
    • 2.5 Early development of the periodic table of elements
    • 2.6 Organization of the elements
  • 3 Atoms: the quantum world
    • 3.1 Wave Nature of Light
    • 3.2 Quantized Energy and Photons
    • 3.3 the Bohr Model
    • 3.4 Wave Character of Matter
    • 3.5 Atomic Orbitals
    • 3.6 3D Representation of Orbitals
    • 3.7 Electron Spin
    • 3.8 Electron Configurations
  • 4 Molecular Shape and Structure
    • 4.1 VSEPR theory
    • 4.2 Hybridization
    • 4.3 sp3 hybridization
    • 4.4 sp2 hybridization
    • 4.5 sp hybridization
    • 4.6 Other hybridization
    • 4.7 Multiple Bonds
    • 4.8 Molecular Orbitals
    • 4.9 Second-Row Diatomic Molecules
  • 5 Fundamentals of Thermochemistry
    • 5.1 Systems, States and Processes
    • 5.2 Heat as a Mechanism to Transfer Energy
    • 5.3 Work as a Mechanism to Transfer Energy
    • 5.4 Heat Capacity and Calorimetry
    • 5.5 The First Law of Thermodynamics
    • 5.6 Heats of Reactions - ΔU and ΔH
    • 5.7 Indirect Determination of ΔH - Hess's Law
    • 5.8 Standard Enthalpies of Formation
  • 6 Principles of Thermodynamics
    • 6.1 The Nature of Spontaneous Processes
    • 6.2 Entropy and Spontaneity - A Molecular Statistical Interpretation
    • 6.3 Entropy Changes and Spontaneity
    • 6.4 Entropy Changes in Reversible Processes
    • 6.5 Quantum States, Microstates, and Energy Spreading
    • 6.6 The Third Law of Thermodynamics
    • 6.7 Gibbs Energy
  • 7 Chemical equilibrium
    • 7.1 Equilibrium
    • 7.2 Reversible and irreversible reaction
    • 7.3 Chemical equilirbium
    • 7.4 Chemical equilibrium constant, Kc
    • 7.5 Le Chatelier's principle
    • 7.6 RICE table
    • 7.7 Haber process
  • 8 Acid–Base Equilibria
    • 8.1 Classifications of Acids and Bases
    • 8.2 The Brønsted-Lowry Scheme
    • 8.3 Acid and Base Strength
    • 8.4 Buffer Solutions
    • 8.5 Acid-Base Titration Curves
    • 8.6 Polyprotic Acids
    • 8.7 Exact Treatment of Acid-Base Equilibria
    • 8.8 Organic Acids and Bases
  • 9 Kinetics
    • 9.1 Prelude to Kinetics
    • 9.2 Chemical Reaction Rates
    • 9.3 Factors Affecting Reaction Rates
    • 9.4 Rate Laws
    • 9.5 Integrated Rate Laws
    • 9.6 Collision Theory
    • 9.7 Reaction Mechanisms
    • 9.8 Catalysis
Other hybridization



sp3d and sp3d2 Hybridization

To describe the five bonding orbitals in a trigonal bipyramidal arrangement, we must use five of the valence shell atomic orbitals (the s orbital, the three porbitals, and one of the d orbitals), which gives five sp3d hybrid orbitals. With an octahedral arrangement of six hybrid orbitals, we must use six valence shell atomic orbitals (the s orbital, the three p orbitals, and two of the d orbitals in its valence shell), which gives six sp3d2 hybrid orbitals. These hybridizations are only possible for atoms that have d orbitals in their valence subshells (that is, not those in the first or second period).


Three Lewis structures are shown along with designations of molecular shape. The left image shows a sulfur atom singly bonded to four fluorine atoms. The sulfur atom has one lone pair of electrons while each fluorine has three. Two fluorine atoms are drawn vertically up and down from the sulfur while the other two are shown going into and out of the page. The second structure shows one chlorine atom singly bonded to three fluorine atoms. The chlorine has two lone pairs of electrons while each fluorine has three. Two fluorine atoms are drawn vertically up and down from the sulfur while the other is shown horizontally. The right structure shows a chlorine atom singly bonded to four fluorine atoms. The chlorine atom has one lone pair of electrons and a superscript plus sign, while each fluorine has three lone pairs of electrons. Two fluorine atoms are drawn vertically up and down from the sulfur while the other two are shown going into and out of the page.

Figure 1 The three compounds pictured exhibit sp3d hybridization in the central atom and a trigonal bipyramid form. SF4 and 

ClF4+ have one lone pair of electrons on the central atom, and ClF3 has two lone pairs giving it the T-shape shown.


In a molecule of phosphorus pentachloride, PCl5, there are five P–Cl bonds (thus five pairs of valence electrons around the phosphorus atom) directed toward the corners of a trigonal bipyramid. We use the 3s orbital, the three 3p orbitals, and one of the 3d orbitals to form the set of five sp3d hybrid orbitals (Figure 1) that are involved in the P–Cl bonds. Other atoms that exhibit sp3d hybridization include the sulfur atom in SF4 and the chlorine atoms in ClF3 and in . (The electrons on fluorine atoms are omitted for clarity.)

Two images are shown and labeled “a” and “b.” Image a depicts a ball-and-stick model in a trigonal bipyramidal arrangement. Image b depicts the hybrid orbitals in the same arrangement and each is labeled, “s p superscript three d.”

Figure 2(a) The five regions of electron density around phosphorus in PCl5 require five hybrid sp3d orbitals. (b) These orbitals combine to form a trigonal bipyramidal structure with each large lobe of the hybrid orbital pointing at a vertex. As before, there are also small lobes pointing in the opposite direction for each orbital (not shown for clarity).

The sulfur atom in sulfur hexafluoride, SF6, exhibits sp3d2 hybridization. A molecule of sulfur hexafluoride has six bonding pairs of electrons connecting six fluorine atoms to a single sulfur atom. There are no lone pairs of electrons on the central atom. To bond six fluorine atoms, the 3s orbital, the three 3p orbitals, and two of the 3d orbitals form six equivalent sp3d2 hybrid orbitals, each directed toward a different corner of an octahedron. Other atoms that exhibit sp3d2hybridization include the phosphorus atom in , the iodine atom in the interhalogens , IF5, ,  and the xenon atom in XeF4.

Two images are shown and labeled “a” and “b.” Image a depicts a ball-and-stick model in an octahedral arrangement. Image b depicts the hybrid orbitals in the same arrangement and each is labeled, “s p superscript three d superscript two.”

Figure 3(a) Sulfur hexafluoride, SF6, has an octahedral structure that requires sp3d2 hybridization. (b) The six sp3d2 orbitals form an octahedral structure around sulfur. Again, the minor lobe of each orbital is not shown for clarity.


Summary of Hybrid Orbitals to Central Atoms

The hybridization of an atom is determined based on the number of regions of electron density that surround it. The geometrical arrangements characteristic of the various sets of hybrid orbitals are shown in the belowing Figure. These arrangements are identical to those of the electron-pair geometries predicted by VSEPR theory. VSEPR theory predicts the shapes of molecules, and hybrid orbital theory provides an explanation for how those shapes are formed. To find the hybridization of a central atom, we can use the following guidelines:

  1. Determine the Lewis structure of the molecule.

  2. Determine the number of regions of electron density around an atom using VSEPR theory, in which single bonds, multiple bonds, radicals, and lone pairs each count as one region.

  3. Assign the set of hybridized orbitals from the belowing Figure.  


A table is shown that is composed of five columns and six rows. The header row contains the phrases, “Regions of electron density,” “Arrangement,” (which has two columns below it), and “Hybridization,” (which has two columns below it). The first column contains the numbers “2,” “3,” “4,” “5,” and “6.” The second column contains images of a line, a triangle, a three sided pyramid, a trigonal bipyramid, and an eight-faced ocatahedron. The third column contains the terms, “Linear,” “Trigonal planar,” “Tetrahedral,” “Trigonal bipyramidal,” and “Octahedral.” The fourth column contains the terms “s p,” “s p superscript 2,” “s p superscript 3,” “s p superscript 3 d,” and “s p superscript 3 d superscript 2.” The last column contains drawings of the molecules beginning with a peanut-shaped structure marked with an angle of “180 degrees.” The second structure is made up of three equal-sized, rounded structures connected at one point with an angle of “120 degrees,” while the third structure is a three-dimensional arrangement of four equal-sized, rounded structures labeled as “109.5 degrees.” The fourth structure is made up of five equal-sized, rounded structures connected at “120 and 90 degrees,” while the fifth structure has six equal-sized, rounded structures connected at “90 degrees.”
Figure The shapes of hybridized orbital sets are consistent with the electron-pair geometries. For example, an atom surrounded by three regions of electron density is sp2 hybridized, and the three sp2 orbitals are arranged in a trigonal planar fashion.


It is important to remember that hybridization was devised to rationalize experimentally observed molecular geometries, not the other way around.

The model works well for molecules containing small central atoms, in which the valence electron pairs are close together in space. However, for larger central atoms, the valence-shell electron pairs are farther from the nucleus, and there are fewer repulsions. Their compounds exhibit structures that are often not consistent with VSEPR theory, and hybridized orbitals are not necessary to explain the observed data.

Three Lewis structures are shown. The left structure shows an oxygen atom with two lone pairs of electrons single bonded to two hydrogen atoms. The middle structure is made up of a sulfur atom with two lone pairs of electrons single bonded to two hydrogen atoms. The right structure is made up of a tellurium atom with two lone pairs of electrons single bonded to two hydrogen atoms. From left to right, the bond angles of each molecule decrease.

For example, we have discussed the H–O–H bond angle in H2O, 104.5°, which is more consistent with sp3 hybrid orbitals (109.5°) on the central atom than with 2p orbitals (90°). Sulfur is in the same group as oxygen, and H2S has a similar Lewis structure. However, it has a much smaller bond angle (92.1°), which indicates much less hybridization on sulfur than oxygen. Continuing down the group, tellurium is even larger than sulfur, and for H2Te, the observed bond angle (90°) is consistent with overlap of the 5p orbitals, without invoking hybridization. We invoke hybridization where it is necessary to explain the observed structures.