1: Review
- Page ID
- 451101
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\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)- 1.1: Atomic Structure - The Nucleus
- Atoms are comprised of protons, neutrons and electrons. Protons and neutrons are found in the nucleus of the atom, while electrons are found in the electron cloud around the nucleus. The relative electrical charge of a proton is +1, a neutron has no charge, and an electron’s relative charge is -1. The number of protons in an atom’s nucleus is called the atomic number, Z. The mass number, A, is the sum of the number of protons and the number of neutrons in a nucleus.
- 1.2: Atomic Structure - Orbitals
- An atomic orbital is the probability description of where an electron can be found. The four basic types of orbitals are designated as s, p, d, and f.
- 1.3: Atomic Structure - Electron Configurations
- The order in which electrons are placed in atomic orbitals is called the electron configuration and is governed by the aufbau principle. Electrons in the outermost shell of an atom are called valence electrons. The number of valence electrons in any atom is related to its position in the periodic table. Elements in the same periodic group have the same number of valence electrons.
- 1.4: Development of Chemical Bonding Theory
- Lewis Dot Symbols are a way of indicating the number of valence electrons in an atom. They are useful for predicting the number and types of covalent bonds within organic molecules. The molecular shape of molecules is predicted by Valence Shell Electron Pair Repulsion (VSEPR) theory. The shapes of common organic molecules are based on tetrahedral, trigonal planar or linear arrangements of electron groups.
- 1.5: Describing Chemical Bonds - Valence Bond Theory
- Covalent bonds form as valence electrons are shared between two atoms. Lewis Structures and structural formulas are common ways of showing the covalent bonding in organic molecules. Formal charge describes the changes in the number of valence electrons as an atom becomes bonded into a molecule. If the atom has a net loss of valence electrons it will have a positive formal charge. If the atom has a net gain of valence electrons it will have a negative formal charge.
- 1.6: sp³ Hybrid Orbitals and the Structure of Methane
- The four identical C-H single bonds in methane form as the result of sigma bond overlap between the sp3 hybrid orbitals of carbon and the s orbital of each hydrogen.
- 1.7: sp² Hybrid Orbitals and the Structure of Ethylene
- The C=C bond in ethylene forms as the result of both a sigma bond overlap between a sp2 hybrid orbital on each carbon and a pi bond overlap of a p orbital on each carbon
- 1.8: sp Hybrid Orbitals and the Structure of Acetylene
- The carbon-carbon triple bond in acetylene forms as the result of one sigma bond overlap between a sp hybrid orbital on each carbon and two pi bond overlaps of p orbitals on each carbon.
- 1.9: Noncovalent Interactions Between Molecules
- In contrast to intramolecular forces, such as the covalent bonds that hold atoms together in molecules and polyatomic ions, intermolecular forces hold molecules together in a liquid or solid. Intermolecular forces are generally much weaker than covalent bonds. The most common intermolecular forces in organic chemistry are from strongest to weakest are hydrogen bonds, dipole-dipole interactions, and London Dispersion (van der Waals) forces.
- 1.10: Drawing Chemical Structures
- Kekulé Formulas or structural formulas display the atoms of the molecule in the order they are bonded. Condensed structural formulas show the order of atoms like a structural formula but are written in a single line to save space. Skeleton formulas or Shorthand formulas or line-angle formulas are used to write carbon and hydrogen atoms more efficiently by replacing the letters with lines. Isomers have the same molecular formula, but different structural formulas
- 1.11: Functional Groups
- Functional groups are atoms or small groups of atoms (two to four) that exhibit a characteristic reactivity. A particular functional group will almost always display its characteristic chemical behavior when it is present in a compound. Because of their importance in understanding organic chemistry, functional groups have characteristic names that often carry over in the naming of individual compounds incorporating specific groups
- 1.12: Formal Charges
- A formal charge is the charge assigned to an atom in a molecule, assuming that electrons in all chemical bonds are shared equally between atoms, regardless of relative electronegativity.
- 1.13: Infrared (IR) Spectroscopy
- The infrared region of the electromagnetic spectrum causes asymmetric bonds to stretch, bend, and/or vibrate. This interaction can be measured to help elucidate chemical structures.
- 1.14: Interpreting IR Spectra
- The analysis and interpretation of the IR spectra for several compounds are explained.