Organic compounds are compounds that contain carbon.
The atomic number of an atom equals the number of protons in its nucleus. The mass number of an atom is the sum of its protons and neutrons. Isotopes have the same atomic number, but different mass numbers.
An atomic orbital indicates where there is a high probability of finding an electron. The closer the atomic orbital is to the nucleus, the lower is its energy. Degenerate orbitals have the same energy.
Electrons are assigned to orbitals following the aufbau principle, the Pauli exclusion principle, and Hund’s rule.
The octet rule states that an atom will give up,accept,or share electrons in order to fill its outer shell or attain an outer shell with eight electrons. Electropositive elements readily lose electrons; electronegative elements readily acquire electrons.
The electronic configuration of an atom describes the orbitals occupied by the atom’s electrons. Electrons in inner shells are called core electrons; electrons in the outermost shell are called valence electrons. Lonepair electrons are valence electrons that are not used in bonding.
Attractive forces between opposite charges are called electrostatic attractions. An ionic bond is formed by a transfer of electrons; a covalent bond is formed by sharing electrons. A polar covalent bond has a dipole, measured by a dipole moment. The dipole moment of a molecule depends on the magnitudes and directions of the bond dipole moments.
Lewis structures indicate which atoms are bonded together and show lone pairs and formal charges. A carbocation has a positively charged carbon, a carbanion has a negatively charged carbon, and a radical has an unpaired electron.
According to molecular orbital (MO) theory, covalent bonds result when atomic orbitals combine to form molecular orbitals. Atomic orbitals combine to give a bonding MO and a higher energy antibonding MO.
Cylindrically symmetrical bonds are called sigma bonds; pi bonds form when p orbitals overlap side-toside. Bond strength is measured by the bond dissociation energy. A bond is stronger than a bond. All single bonds in organic compounds are bonds, a double bond consists of one bond and one bond, and a triple bond consists of one bond and two bonds. Triple bonds are shorter and stronger than double bonds, which are shorter and stronger than single bonds. To form four bonds, carbon promotes an electron from a 2s to a 2p orbital. C, N, and O form bonds using hybrid orbitals.
The hybridization of C, N, or O depends on the number of bonds the atom forms: No bonds means that the atom is hybridized, one bond indicates that it is hybridized, and two bonds signifies that it is sp hybridized. Exceptions are carbocations and carbon radicals, which are hybridized.
The more s character in the orbital used to form a bond, the shorter and stronger the bond is and the larger the bond angle is. Bonding and lone-pair electrons around an atom are positioned as far apart as possible.
An acid is a species that donates a proton, and a base is a species that accepts a proton. A Lewis acid is a species that accepts a share in an electron pair; a Lewis base is a species that donates a share in an electron pair.
Acidity is a measure of the tendency of a compound to give up a proton. Basicity is a measure of a compound’s affinity for a proton. The stronger the acid, the weaker is its conjugate base. The strength of an acid is given by the acid dissociation constant
Approximate values are as follows: protonated alcohols, protonated carboxylic acids, protonated water carboxylic acids protonated amines alcohols and water
The pH of a solution indicates the concentration of positively charged hydrogen ions in the solution.
In acid–base reactions, the equilibrium favors reaction of the strong and formation of the weak. The strength of an acid is determined by the stability of its conjugate base:
The more stable the base, the stronger is its conjugate acid.
When atoms are similar in size, the more acidic compound has its hydrogen attached to the more electronegative atom. When atoms are very different in size, the more acidic compound has its hydrogen attached to the larger atom. Inductive electron withdrawal increases acidity; acidity decreases with increasing distance between the electron-withdrawing substituent and the ionizing group.
Delocalized electrons are electrons shared by more than two atoms. A compound with delocalized electrons has resonance. The resonance hybrid is a composite of the resonance contributors,which differ only in the location of their lone-pair and electrons.
The Henderson–Hasselbalch equation gives the relationship between and pH:A compound exists primarily in its acidic form in solutions more acidic than its value and primarily in its basic form in solutions more basic than its value.
According to molecular orbital (MO) theory, covalent bonds result when atomic orbitals combine to form molecular orbitals. Atomic orbitals combine to give a bonding MO and a higher energy antibonding MO.
Cylindrically symmetrical bonds are called sigma bonds; pi bonds form when p orbitals overlap side-toside. Bond strength is measured by the bond dissociation energy. A bond is stronger than a bond. All single bonds in organic compounds are bonds, a double bond consists of one bond and one bond, and a triple bond consists of one bond and two bonds. Triple bonds are shorter and stronger than double bonds, which are shorter and stronger than single bonds. To form four bonds, carbon promotes an electron from a 2s to a 2p orbital. C, N, and O form bonds using hybrid orbitals.
The hybridization of C, N, or O depends on the number of bonds the atom forms: No bonds means that the atom is hybridized, one bond indicates that it is hybridized, and two bonds signifies that it is sp hybridized. Exceptions are carbocations and carbon radicals, which are hybridized.
The more s character in the orbital used to form a bond, the shorter and stronger the bond is and the larger the bond angle is. Bonding and lone-pair electrons around an atom are positioned as far apart as possible.
An acid is a species that donates a proton, and a base is a species that accepts a proton. A Lewis acid is a species that accepts a share in an electron pair; a Lewis base is a species that donates a share in an electron pair.
Acidity is a measure of the tendency of a compound to give up a proton. Basicity is a measure of a compound’s affinity for a proton. The stronger the acid, the weaker is its conjugate base. The strength of an acid is given by the acid dissociation constant
Approximate values are as follows: protonated alcohols, protonated carboxylic acids, protonated water carboxylic acids protonated amines alcohols and water
The pH of a solution indicates the concentration of positively charged hydrogen ions in the solution.
In acid–base reactions, the equilibrium favors reaction of the strong and formation of the weak. The strength of an acid is determined by the stability of its conjugate base:
The more stable the base, the stronger is its conjugate acid.
When atoms are similar in size, the more acidic compound has its hydrogen attached to the more electronegative atom. When atoms are very different in size, the more acidic compound has its hydrogen attached to the larger atom. Inductive electron withdrawal increases acidity; acidity decreases with increasing distance between the electron-withdrawing substituent and the ionizing group.
Delocalized electrons are electrons shared by more than two atoms. A compound with delocalized electrons has resonance. The resonance hybrid is a composite of the resonance contributors,which differ only in the location of their lone-pair and electrons.
The Henderson–Hasselbalch equation gives the relationship between and pH:A compound exists primarily in its acidic form in solutions more acidic than its value and primarily in its basic form in solutions more basic than its value.
from - Organic chemistry - Bruice - Chapter 1 summary
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