Organic Chemistry 1 Chapter 1-3 – Flashcards
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Methane
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1 C; CH4; CH4
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Ethane
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2 C; C2H6; CH3CH3
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Propane
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3 C; C3H8; CH3CH2CH3
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Butane
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4 C; C4H10; CH3CH2CH2CH3
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Pentane
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5 C; C5H12; CH3(CH2)3CH3
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Hexane
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6 C; C6H14; CH3(CH2)4CH3
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Heptane
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7 C; C7H16; CH3(CH2)5CH3
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Octane
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8 C; C8H18; CH3(CH2)7CH3
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Nonane
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9 C; C9H20; CH3(CH2)8CH3
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Decane
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10 C; C10H22; CH3(CH2)8CH3
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Ionic bond
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To obtain a noble gas configuration, atoms may transfer electrons from one atom to another. Electron density completely transferred from one atom to another.
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Covalent bond
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Electron density is shared between two atoms, where each atom has a full valence shell.
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Nonpolar Covalent Bonds
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Electrons are shared equally by two bonded atoms.
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Polar Covalent Bonds
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Electrons are shared unequally by two bonded atoms.
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Electronegativity
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Is the relative ability of an atom in a particular molecule to attract shared electrons to itself. It is used to determine whether a given bond with be non polar covalent, polar covalent, or an ionic bond.
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Electronegativity range
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values range from 0.7-4.0; larger values means the element attracts electrons more strongly (negative end), while the smaller values attract electrons weakly (positive end).
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Electronegativity scale on the Periodic Table
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It increases as you go from left to right and as you go from bottom to top.
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How are bonds classified?
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Classified based upon the difference in electronegativity between the bonding atoms.
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Pure Covalent Bonds (non polar) Classification
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Difference is 0-0.4; Shared equally. C-H 2.5-2.1=0.4
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Polar Covalent Bonds Classification
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Difference is >0.4-2.0; H-F 2.1-4.0=1.9
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Ionic Bonds Classification
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Difference is >2.0; electron has been completely transferred at this point. NaCl 0.9-3.0=2.1
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G.N Lewis
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American chemist who built the UC Berkeley Chem. Dept. Lewis symbols and structures and Lewis acid-base theory.
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Octet Rule
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When atoms bond together to form molecules, the atoms transfer or share electrons in such a way as to attain a filled shell of electrons.
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Drawing Lewis Structures 1
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Determine the number of electrons in the structure by: summing the valence electrons for each atom in the formula, then add 1 for each negative charge or subtract 1 for each positive charge.
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Drawing Lewis Structures 2
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Draw skeletal structure- central atom generally least electronegative element. Hydrogens are always outer atoms. Connect the atoms with lines (bonds).
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Drawing Lewis Structures 3
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Count the number of electrons used and subtract from the total available.
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Drawing Lewis Structures 4
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With the remaining electrons start with the outer atoms and add pairs of electrons to complete the octets for atoms bonded to the central atom.
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Drawing Lewis Structures 5
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If any electrons remain after step 4, place them on the central atom.
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Drawing Lewis Structures 6
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If any atoms lack an octet at this point, move non-bonding pairs of electrons to make double or triple bonds until atoms have an octet.
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Lone Pairs
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AKA Nonbonding electrons, are valence shell electrons that are not shared between atoms.
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Condensed Structures
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All atoms are drawn in, but the bond lines are omitted (double and triple bonds are usually drawn). Atoms drawn next to the atoms to which they are bonded. Parentheses are used around similar groups bonded to the same atom. Lone pairs are omitted. CH3(CH2)2CH3
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Heteroatom
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Any element other than carbon or hydrogen.
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Skeletal Structures
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AKA Line structures: don't label carbon atoms, they are at the intersection of 2 or more lines and at the end of lines. Don't draw hydrogens attached to carbon (it is a given that they are there). Draw in all heteroatoms and hydrogens directly bound to them.
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Formal Charge
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A hypothetical charge on an atom. They are not real or measured.
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How to calculate formal charge
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+( # of valence electrons of atom) - ( # of unshared electrons of atom) - ( 1/2 # of shared electrons.
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Expanded Octets
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Some elements form more than 4 bonds. Such as Sulfur and Phosphorous form up to 6 bonds, meaning they share 12 electrons.
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Paramagnetic
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Substances that are attracted to magnetic fields. They have one or more unpaired electrons. The more unpaired electrons, the stronger the attraction. Elemental Oxygen (O2) is paramagnetic, but the Lewis structure doesn't show unpaired electrons.
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Diamagnetic
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Substance that are weakly repelled by magnetic fields. They have no unpaired electrons.
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Molecular Orbital Theory
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The combination of atomic orbitals on different atoms forms molecular orbitals. Electrons in these orbitals belong to the molecule. Provides us with electron distributions, bond energies, and magnetic properties.
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Valence Bond Theory
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Bonds result from the sharing of electrons in overlapping orbitals of different atoms. Orbitals may be atomic or hybridized atomic. Electrons are localized in the bonds between the two atoms.
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Antibonding
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higher energy
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Bonding
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lower energy
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Pi Bonding
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The sideways overlap of two parallel p orbitals leads to pi bonding MO and an pi* anti bonding MO. It is not as strong as most sigma bonds.
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Boundary Surface Representation for s Orbitals
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circular (ball-like)
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Boundary Surface Representations for p Orbitals
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1 infinity sign. 2px, 2py, and 2pz
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Boundary Surface Representation for d Orbitals
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4 infinity signs together. dyz, dxz, dxy, dx2-y2, and dz2
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The overlap of s, p, d and f atomic orbitals often result in
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unpredictable predictions.
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Hybrid Orbitals
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s, px, py, and pz atomic orbitals can be combined to give four spy hybrid orbitals.
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VSEPR
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Each group of valence electrons around a central atom is located as far away from others as possible to minimize repulsions. The repulsions maximize the space that each object attached to the central atom occupies, result is 5 arrangements of electron-groups.
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The electron groups of VSEPR define the__ but the molecular shape is defined by the ___.
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object arrangement; relative portions of the atoms bonded to the central atom.
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Drawing in 3D
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Simple lines represent bonds that are in the plane of the paper. Wedge shaped lines represent bonds that project forward to atoms. Dashed lines indicate bonds that go back to atoms.
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Rotation of Single Bonds
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Ethane is composed of two methyl groups bonded by the overlap of the spy hybrid orbitals. There is free rotation along single bonds. The hydrogens arrange themselves around the central atoms when rotation occurs.
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Bonding in Ethylene
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Ethylene has three sigma bonds formed by its sp2 hybrid orbitals in a trigonal planar. The unhybridized p orbital of one carbon is perpendicular to its sp2 hybrid orbitals, and it is parallel to the unhybridized p orbital of the second carbon. Overlap of these 2 p orbitals will produce a pi bond (double bond) that is located above and below the sigma bond.
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Can rotation occur in double bonds?
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No they are permanent.
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Bond Dipole Moments
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Due to difference in electronegativity. They depend on the amount of charge and distance separation. Measured in debyes (D).
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Molecular Dipole Moments
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The vector sum of the bond dipole moments. Depends on bond polarity and bond angles. Lone pairs of electrons contribute.
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Intramolecular Forces
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bonding forces; exist within each molecule. They influence the chemical properties of the substance.
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Intermolecular Forces
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non-bonding forces; exist between molecules. They influence the physical properties of the substance. Intermolecular are much weaker than intramolecular. When a substance melts or boils the intermolecular forces are broken.
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Dipole-dipole interactions
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The attractive forces between the permanent dipoles of two polar molecules.
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Hydrogen Bonds
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They are very strong dipole-dipole forces. Interaction between a H covalently bonded to strongly electronegative atoms (Z=O, N or F) and nonbonding electron pairs on other strongly electronegative atoms (O, N or F). As electrons are pulled away from H by an electronegative atom, what is left is an unshielded proton that will strongly attract neighboring electrons. Represented by a dotted or dashed line.
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H-bond Acceptor between two water molecules
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An H is forming a hydrogen bond to the lone pair on the oxygen of the other molecule. This means that the molecule with the oxygen attached is the H-bond acceptor.
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H-bond Donor between two water molecules
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An H is forming a hydrogen bond to the lone pair on the oxygen of the other molecule. This means that it is the molecule with the hydrogen attached is the H-bond donor.
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Water is a molecule that is capable of being either the ___ or the ____ of a hydrogen bond.
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an acceptor or a donor.
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Some molecules have the ability to hydrogen bond, but they are only capable of playing the role of an acceptor.
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For example, 2 molecules of Acetone cannot hydrogen bond because the oxygen does not have a hydrogen covalently bonded to it. However, Acetone could hydrogen bond with water where water is the H-donor and Acetone is the H-acceptor.
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London dispersion forces
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One instantaneous dipole can induce another instantaneous dipole in an adjacent molecule (or atom). London dispersion forces are the forces between instantaneous dipoles. AKA induced dipoles or Van der Waals forces. All molecules have them and th
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London dispersion forces characteristics
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All molecules have them and they are temporary. The size of the force depends on: The number of electrons and the shapes of the molecules.
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Surface area determines strength of Intermolecular Forces.
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The larger the surface area, the larger the attractive force between two molecules, and the stronger the intermolecular forces.
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Strength of Ion-Dipole
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Basis of Attraction: Ion-charge-dipole charge Energy: 40-600 kJ/mol
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Strength of H bond
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Basis of Attraction: Polar bond to H-dipole charge (high EN of N,F, O) Energy: 10-40 kJ/mol
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Strength of Dipole-Dipole
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Basis of Attraction: Dipole Charges Energy: 5-25 kJ/mol
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Strength of Ion-induced dipole
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Basis of Attraction: Ion-charge-polarizable electron cloud. Energy: 3-15 kJ/mol
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Strength of Dipole-induced dipole
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Basis of Attraction: Dipole-charge-polarizable electron cloud. Energy: 2-10 kJ/mol
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Strength of Dispersion (London)
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Basis of Attraction: Polarizable electron cloud. Energy: 0.05-40 kJ/mol
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3 States of Matter-> How many types of Solutions?
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7
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Gas as Solute:
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Solvents: Gas: (O2 in N2) Air Liquid: (CO2 in H2O) Carbonated Beverages Solid: (CH4 in Ice) Methane Hydrate
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Liquid as Solute
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Solvents: Gas: Not possible Liquid: (HC2H3O2 in H2O) Vinegar Solid: (Hg in Ag) Dental Fillings
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Solid as Solute
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Solvents: Gas: Not possible Liquid: (NaCl in H2O) Ocean Water Solid: (Zn in Cu) Brass
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The process of dissolution represents a competition:
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-Energy of the solvent-solute interactions. -Energy of the intermolecular forces of the solute itself. -Energy of the intermolecular forces of the solvent itself.
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If the solvent-solute interactions win:
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The material may dissolve (said to be soluble or miscible.)
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If the solute-solute interactions win:
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The material may not dissolve (insoluble or immiscible.)
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The solute-solvent interactions are greater than:
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the sum of the solute-solute and solvent-solvent interactions. Favorable for solution formation.
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The solute-solvent interactions are less than:
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the sum of the solute-solute and solvent-solvent interactions. Not favorable for solution formation.
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Resonance Structures
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When there is more than one Lewis structure for a molecule that differ only in the position of the electrons. Lone pairs and multiple bonds in different positions. One resonance contributor is converted by another by the use of curved arrows which show the movement of electrons.
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Resonance Hybrid
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The actual molecule is an average of all the resonance forms. It doesn't resonate between forms, although we draw it that way.
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Rules for Resonance:
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-Individual structures exist only on paper. -Individual structures must be valid Lewis Structures. -Only electrons are allowed to move between resonance structures. -position of nuclei must remain the same. -only electrons in pi bonds and lone pairs can be moved.
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Resonance forms can be compared using the following criteria, beginning with the most important:
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1. Has as many octets as possible. 2. Has as many bonds as possible. 3. Has the negative formal charge on the most electronegative atom. 4. Has as little formal charge separation as possible.
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Assessing unequal resonance forms with the use of formal charges:
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1. The fewer and smaller the formal charges of a Lewis structure, the better the structure. 2. Negative formal charges should reside on more electronegative atoms.
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The better resonance form for Sulfuric Acid
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One resonance structure for Sulfuric Acid contains more bonds (extending the octet rule on Sulfur with 12 electrons attached) but has less formal charges. We would think that this would be the better structure due to the formal charges and that Sulfur is capable of having an extended octet. This is not the case, the better resonance form would be to have less bonds (octet rule is being abided by sulfur with only 8 electrons attached) and having a larger formal charge. Octet rule is more important than a smaller formal charge.
