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CHAPTER 1: STRUCTURE BONDING
Every living organism is made of organic chemicals. The proteins that make up your hair, skin, and muscles; the DNA that controls your genetic heritage; the foods that nourish you; and the medicines that heal you are all organic chemicals. Little was known about chemistry at that time, and the behavior of the “organic” substances isolated from plants and animals seemed different from that of the “inorganic” substances found in minerals. Organic compounds were generally low-melting solids and were usually more difficult to isolate, purify, and work with than high-melting inorganic compounds. To many chemists (difference of organic and inorganic) ● was that organic compounds contained a peculiar “vital force” as a result of their origin in living sources. Because of this vital force, chemists believed, organic compounds could not be prepared and manipulated in the laboratory as could inorganic compounds. ● As early as 1816 , however, this vitalistic theory received a heavy blow when Michel Chevreul found that soap, prepared by the reaction of alkali with animal fat, could be separated into several pure organic compounds, which he termed fatty acids. For the first time, one organic substance (fat) was converted into others (fatty acids plus glycerin) without the intervention of an outside vital force. ● Little more than a decade later , the vitalistic theory suffered further when Friedrich Wöhler discovered in 1828 that it was possible to convert the “inorganic” salt ammonium cyanate into the “organic” substance urea, which had previously been found in human urine. ● By the mid-1800s , the weight of evidence was clearly against the vitalistic theory and it was clear that there was no fundamental difference between organic and inorganic compounds. ● The only distinguishing characteristic of organic compounds is that all contain the element carbon. ● Organic chemistry, then, is the study of carbon compounds. But why is carbon special? ● It is because of carbon’s electronic structure and its consequent position in the periodic table. ● As a group 4A element, carbon can share four valence electrons and form four strong covalent bonds. ● Carbon atoms can bond to one another, forming long chains and rings. ● Carbon, alone of all elements, is able to form an immense diversity of compounds, from the simple methane, with one carbon atom, to the staggeringly complex DNA, which can have more than 100 million carbons.
1.1 ATOMIC STRUCTURE: THE NUCLEUS
ATOM
● consists of a dense, positively charged nucleus surrounded at a relatively large distance by negatively charged electrons. ● The dense, positively charged nucleus contains most of the atom’s mass and is surrounded by negatively charged electrons. ● The three dimensional view on the right shows calculated electron-density surfaces. Electron density increases steadily toward the nucleus and is 40
times greater at the blue solid surface than at the gray mesh surface. ● The nucleus consists of subatomic particles called ○ protons , which are positively charged, and; ○ neutrons , which are electrically neutral. ● Because an atom is neutral overall: Number of protons (in the nucleus) = Number of electrons (surrounding the nucleus) ● Although extremely small—about 10 -14^ to 10 - meters (m) in diameter— the nucleus nevertheless contains essentially all the mass of the atom. ● Electrons have negligible mass and circulate around the nucleus at a distance of approximately 10 -10^ m. Thus, the diameter of a typical atom is about 2x10-10m, or 200 picometers (pm), where 1 pm = 10 -12^ m. ● A specific atom is described by its atomic number (Z) , which gives the number of protons (or electrons ) it contains, ● and its mass number (A) , which gives the total number of protons and neutrons in its nucleus.
EXAMPLE:
All the atoms of a given element have the same atomic number—1 for hydrogen, 6 for carbon, 15 for phosphorus, and so on— but they can have different mass numbers depending on how many neutrons they contain. ISOTOPES ● Atoms with the same atomic number but different mass numbers. ATOMIC MASS/ ATOMIC WEIGHT ● The weighted-average mass in atomic mass units (amu) of an element’s naturally occurring isotopes— EXAMPLES: ● 1.008 amu for hydrogen ● 12.011 amu for carbon ● 30.974 amu for phosphorus, and so on. 1.2 ATOMIC STRUCTURE: ORBITALS WAVE EQUATION ● according to the quantum mechanical model, the behavior of a specific electron in an atom can be described by a mathematical expression called a wave equation— ● the same type of expression used to describe the motion of waves in a fluid. ● The solution to a wave equation is called a wave function , or orbital , and is denoted by the Greek letter psi ( 𝜓 ).
By plotting the square of the wave function, 𝜓^2 , in
three-dimensional space, an orbital describes the volume of space around a nucleus that an electron is most likely to occupy. ● the orbital would appear as a blurry cloud, indicating the region of space where the electron has been. ● This electron cloud doesn’t have a sharp boundary , but for practical purposes we can set its limits by saying that an orbital represents the space where an electron spends 90% to 95% of its time.
Rule 1 ● The lowest-energy orbitals fill up first, according to the order 1s → 2s → 2p → 3s → 3p → 4s → 3d. Note that the 4s orbital lies between the 3p and 3d orbitals. Rule 2 ● Electrons act in some ways as if they were spinning around an axis, somewhat like how the earth spins. ● This spin can have two orientations, denoted as up (↑) and down (↓). Only two electrons can occupy an orbital, and they must be of opposite spin , a statement called the Pauli exclusion principle. Rule 3 ● If two or more empty orbitals of equal energy are available, one electron occupies each with spins parallel until all orbitals are half-full, a statement called Hund’s rule Some examples of how these rules apply are shown in Table 1-1. Hydrogen, for instance, has only one electron, which must occupy the lowest-energy orbital. Thus, hydrogen has a 1s ground-state configuration. Carbon has six electrons and the ground-state configuration
1s^2 2s^2 2px^1 2py^1 , and so forth. Note that a superscript is
used to represent the number of electrons in a particular orbital. Problem 1- Give the ground-state electron configuration for each of the following elements: (a) Oxygen Answer: 1s^2 2s^2 2p^4
(b) Nitrogen Answer: 1s^2 2s^2 2p^3 (c) Sulfur Answer: 1s^2 2s^2 2p^6 3s^2 3p^4 Problem 1- How many electrons does each of the following elements have in its outermost electron shell? (a) Magnesium Answer: 2 electrons (b) Cobalt Answer: 2 electrons (c) Selenium Answer: 6 electrons
1.4 DEVELOPMENT OF CHEMICAL BONDING
THEORY
By the mid-1800s, chemists had begun to probe the forces holding compounds together. ● August Kekulé and Archibald Couper independently proposed that, in all organic compounds, carbon is tetravalent —it always forms four bonds when it joins other elements to form stable compounds. August Kekulé ● Said that carbon atoms can bond to one another to form extended chains of linked atoms. ● Suggested that carbon chains can double back to form rings of atoms.
oxygen has six valence electrons (2s2 2p4), needs two more, and forms two bonds; and the halogens have seven valence electrons, need one more, and form one bond. Valence electrons that are not used for bonding are called lone-pair electrons , or nonbonding electrons. The nitrogen atom in ammonia, NH3, for instance, shares six valence electrons in three covalent bonds and has its remaining two valence electrons in a nonbonding lone pair. nonbonding electrons are often omitted when drawing line-bond structures Problem 1- Draw a molecule of chloroform, CHCl3, using solid, wedged, and dashed lines to show its tetrahedral geometry. Solution: (a) Step 1 (b) Step 2 Problem 1- Convert the following representation of ethane, C2H6, into a conventional drawing that uses solid, wedged, and dashed lines to indicate tetrahedral geometry around each carbon (gray 5 C, ivory 5 H). Solution: (a) Step 1
(b) Step 2 (c) Step 3 1.5 DESCRIBING CHEMICAL BONDS: VALENCE BOND THEORY VALENCE BOND THEORY ● a covalent bond forms when two atoms approach each other closely and a singly occupied orbital on one atom overlaps a singly occupied orbital on the other atom. The electrons are now paired in the overlapping orbitals and are attracted to the nuclei of both atoms, thus bonding the atoms together ● Example: In H 2 molecule the H-H bond is from the overlap of two singly occupied hydrogen 1s orbitals HYBRIDIZATION ● Combining atomic orbitals to make hybrid orbitals ● Example: sp^3 = s + p + p + p HYBRID ORBITALS ● Developed by Linus Pauling in 1930s ● Model of bonding used to explain the 3D placement of atoms in a molecule SIGMA BONDS ( 𝛔 ) ● formed by the head on overlap of two atomic orbitals along a line drawn between the nuclei ● Represents a single bond occupied by a single pair of electrons. ● Made up of hybridized orbitals ● Example: H-H bond from previous example exhibits a sigma bond PI BONDS (π) ● covalent bond formed between two neighboring atom's unbonded p-orbitals. ● involve a parallel overlapping of the atomic orbitals ● Unhybridized orbitals makes pi bonds NOTE: ● Single bond = 1 sigma bond and 0 pi bond ● Double bond = 1 sigma bond and 1 pi bond ● Triple bond = 1 sigma bond and 2 pi bonds BOND STRENGTH ● strength with which a chemical bond holds two atoms together. ● conventionally measured in terms of the amount of energy, in kilocalories per mole, required to break the bond. BOND LENGTH ● Optimum distance between nuclei that leads to maximum stability
● has a length of 154 pm ● has a strength of 377 kJ/mol Bond Angles of Ethane ● near but not exactly equal to the tetrahedral value of 109.5o 1.8 sp^2 HYBRID ORBITALS AND THE STRUCTURE OF ETHYLENE Single Bonds ● a result from sharing of one electron pair between bonded atoms ● examples: methane and ethane Double Bonds ● sharing two electron pairs between atoms ● example: ethylene with a structure H 2 C = CH 2 , contains a carbon-carbon double bond Triple Bonds ● sharing three electron pairs ● example: acetylene with a structure HC ≡ CH, contains carbon-carbon triple bond sp^2 hybrid orbitals ● 2s orbital combines with only two of three available 2p-orbitals will result to 3 sp^2 hybrid orbitals and one 2p orbital remains unchanged ● are unsymmetrical about the nucleus ● are strongly oriented in a specific direction to form strong bonds ● the three sp^2 orbitals lie in a plane that angles of 120o to one another ● the remaining p orbital is perpendicular to the sp^2 plane Carbon-carbon Double Bond ● two carbons with sp^2 hybridization approach each other, forming a strong ø bond by sp^2 -sp^2 head-on overlap ● the unhybridized p orbitals interact by overlapping sideways to form a pi ( 𝛑 ) bond ● the combination of an sp^2 -sp^2 ø bond and a 2p-2p 𝛑 bond results in sharing four electrons and the formation of the carbon-carbon double bond ● the electrons in the ø bond occupy the region centered between the nuclei ● the electrons in the 𝛑 bond occupy the regions above and below a line drawn between the nuclei The Structure of Ethylene ● when four hydrogen atoms form ø bonds with the remaining four sp2 orbitals will complete the structure of ethylene ● Ethylene has a planar structure with H–C–H and H–C–C bond angles of approximately 120o Bond Angle of H – C – H ● 117.4o Bond Angle of H – C – C ● 121.3o C – H Bond ● has a length of 108.7 pm ● has a strength of of 464 kJ/mol (111 kcal/mol) Ethylene vs Ethane ● the carbon-carbond double bond of ethylene is shorter and stronger than the single bond in ethane because it has four electrons bonding in the nuclei together, rather than just two ● Ethylene - 134 pm (C=C bond length)
- 728 kJ/mol or 174 kcal/mol (strength) ● Ethane - 154 pm (C-C bond length)
- 377 kJ/mol (strength) ● the carbon-carbon double bond is less than twice as strong as the singlebond because the sideways overlap in the 𝛑 part of the double bond is not as great as the head-on overlap in the ø part
1.9 sp HYBRID ORBITALS AND THE STRUCTURE OF ACETYLENE CARBON can also form triple bonds by sharing six e-^ in addition to forming single (2e-) and double (4e-) bonds sp hybrid ● third kind of orbital accounting for the triple bond in a molecule such as acetylene, H–C≡C–H ● two carbon s orbital hybridizes with only a single p orbital resulting in two sp orbitals (oriented 180° apart on the x -axis), and two p orbitals (perpendicular on the y -axis and the z -axis) remain unchanged In the figure below, the two sp hybrid orbitals are oriented 180° away from each other, perpendicular to the two remaining p orbitals (red/blue).
- When two sp-hybridized carbon atoms approach each other, sp hybrid orbitals on each carbon overlap head-on to form a strong sp–sp σ bond. 2. At the same time, the pz orbitals from each carbon form a pz–pz π bond by sideways overlap, and the py orbitals overlap similarly to form a py–py π bond. 3. The net effect is the sharing of six electrons and formation of a carbon–carbon triple bond. 4. The two remaining sp hybrid orbitals each form a σ bond with hydrogen to complete the acetylene molecule. The Structure of Acetylene The two carbon atoms are joined by one sp–sp σ bond and two p–p π bonds. A comparison of sp , sp^2 , and sp^3 hybridization is given in the table. ● As suggested by sp hybridization, acetylene is a linear molecule with H–C–C bond angles of 180°. ● The C–H bonds have a length of 106 pm and a strength of 558 kJ/mol (133 kcal/mol). ● The C≡C bond length in acetylene is 120 pm, and its strength is about 965 kJ/mol (231 kcal/mol), making it the shortest and strongest of any carbon–carbon bond. 1.10 (^) HYBRIDIZATION OF NITROGEN, OXYGEN, PHOSPHORUS, AND SULFUR ● The valence-bond concept of orbital hybridization is NOT LIMITED to carbon. ● Covalent bonds formed by other elements can also be described using hybrid orbitals. METHYLAMINE (CH 3 NH 2 ) ● an organic derivative of ammonia (NH 3 ) ● substance responsible for the odor of rotting fish
Produced by some kind of bacteria, both can be described by approximate sp^3 hybridization around sulfur, although both have significant deviation from the 109.5° tetrahedral angle. 1.11 DESCRIBING CHEMICAL BONDS: MOLECULAR ORBITAL THEORY Molecular orbital (MO) theory describes covalent bond formation as arising from a mathematical combination of atomic orbitals (wave functions) on different atoms to form molecular orbitals , so called because they belong to the entire molecule rather than to an individual atom. atomic orbital ● describes a region of space around an atom where an electron is likely to be found, molecular orbital ● describes a region of space in a molecule where electrons are most likely to be found. (Both atomic orbital and molecular orbital have specific size, shape, and energy.) Two ways for the orbital combination to occur: Additive way
- additive combination leads to the formation of a molecular orbital that is lower in energy and roughly egg-shaped
- additive combination is a single, egg-shaped, molecular orbital; it is not the same as the two overlapping 1s atomic orbitals of the valence bond description.
- lower in energy than the two hydrogen 1s atomic orbitals and is called a bonding MO because electrons in this MO spend most of their time in the region between the two nuclei, thereby bonding the atoms together. Subtractive way
- subtractive combination leads to a molecular orbital that is higher in energy and has a node between nuclei
- subtractive combination is a single molecular orbital with the shape of an elongated dumbbell.
- higher in energy than the two hydrogen 1s orbitals and is called an antibonding MO because any electrons it contains can’t occupy the central region between the nuclei, where there is a node, and can’t contribute to bonding.
Just as bonding and antibonding σ molecular
orbitals result from the head-on combination of two s atomic
orbitals in H 2 , so bonding and antibonding π molecular
orbitals result from the sideways combination of two p atomic orbitals in ethylene.
● the lower-energy, π bonding MO has no node
between nuclei and results from the combination of p orbital lobes with the same algebraic sign.
● the higher-energy, π antibonding MO has a node
between nuclei and results from the combination of lobes with opposite algebraic signs. ● only the bonding MO is occupied; the higher-energy, antibonding MO is vacant.
1.12 DRAWING CHEMICAL STRUCTURES
Condensed Structures
- carbon–hydrogen and carbon–carbon single bonds aren’t shown; instead, they’re understood. Ex.: If a carbon has three hydrogens bonded to it, we write CH3; if a carbon has two hydrogens bonded to it, we write CH2; and so on.
- the horizontal bonds between carbons aren’t shown in condensed structures—the CH 3 , CH 2 , and CH units are simply placed next to each other—but the vertical carbon–carbon bond in the first of the condensed structures drawn above is shown for clarity.
- the two CH 3 units attached to the CH carbon are grouped together as (CH 3 ) 2 in the second of the condensed structures Skeletal Structures Rule 1 ● Carbon atoms aren’t usually shown. Instead, a carbon atom is assumed to be at each intersection of two lines (bonds) and at the end of each line. Occasionally, a carbon atom might be indicated for emphasis or clarity. Rule 2 ● Hydrogen atoms bonded to carbon aren’t shown. Because carbon always has a valence of 4, we mentally supply the correct number of hydrogen atoms for each carbon. Rule 3 ● Atoms other than carbon and hydrogen are shown. Note: Although such groupings as -CH3, -0OH, and -NH2 are usually written with the C, O, or N atom first and the H atom second, the order of writing is sometimes inverted to H 3 C- , HO- , and H 2 N- if needed to make the bonding connections in a molecule clearer. Larger units such as -CH 2 CH3 are not inverted, though; we don’t write H3CH 2 C- because it would be confusing. There are, however, no well-defined rules that cover all cases; it’s largely a matter of preference. Interpreting a Line-Bond Structure Worked Example ● Carvone, a substance responsible for the odor of spearmint, has the following structure. Tell how many hydrogens are bonded to each carbon, and give the molecular formula of carvone. Strategy ● The end of a line represents a carbon atom with 3 hydrogens, CH3; a two-way intersection is a carbon atom with 2 hydrogens, CH2; a three-way intersection is a carbon atom with 1 hydrogen, CH;
