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Chapter 1- vitalism: the belief that natural products needed a vital force to create them- isotopes: atoms with the same number of protons but different numbers of neutrons- orbitals are where the electrons that are bound to nuclei are found- electron density: the relative probability of finding an electron in acertain region of space- is a function of the distance from the nucleus (less likely to be far away fromthe nucleus- the first electron shell can hold 2 electrons, the second can hold 8 electrons, and the third can hold 18 electrons- node: a region of zero electron density- 2p orbitals have higher density than 2s orbitals because the average location of theelectron in a 2p orbital is farther away from the nucleus- nodal plane: a flat (planar) region of space, including the nucleus, with 0 electron density- degenerate orbitals: orbitals with identical energies- Pauli exclusion principle tells us that each orbital can hold a max of 2 electrons, provided that their spins are paired- Aufbau principle says that we must fill up the lower energy orbitals before filling the higher energy orbitals- valence electrons: the electrons in the outermost shell- Hund’s rule states that when there are two or more orbitals of the same energy, electrons will go into different orbitals rather than pair up in the same orbital- G.N. Lewis proposed early ideas that electrons with a filled octet are especially stable and atoms transfer or share electrons in such a way to fill their octet, which we now call the octet rule- ionic bonding: ions with opposite charges that attract each other – common in inorganic compounds and pretty uncommon in organic compounds- covalent bonding: electrons are shared rather than transferred (most common way to bond in organic compounds)- nonpolar covalent bond: when electrons are shared equally between thetwo atoms- polar covalent bond: electrons unequally shared between two atoms- dipole moment: the amount of charge separation multiplied by the bond length- - means a small amount of negative charge and + means a smallδ δamount of positive charge- bond polarity is symbolized by an arrow with its head at the negative end ofthe polar bond and a plus sign at the positive end- Electronegativity… elements with higher electronegativities generally have more attraction for the bonding electrons – this means that in a bond between two different atoms, the atoms with the higher EN is the negative end of the dipole- Formal charge = [group number] – [nonbonding electrons] – ½ [shared electrons]- resonance structures/forms: different ways of drawing the same compound- the actual molecule is known as the resonance hybrid- delocalized charge: a charge which is spread out over two or more atoms; we usually draw resonance forms to show how the charge can appear on each of the atoms sharing the charge- resonance-stabilized cation – spreading the positive charge over two atoms makes the ion more stable than if the entire charge were localized only on the carbon or only on the nitrogen, so this term just puts that into words- remember – individual resonance forms do not exist; the molecule does not “resonate” between these structures. It is a hybrid with some characteristics of both- the more stable resonance form is the major contributor and the less stable form is the minor contributor – the structure of the actual compound resembles the major contributor more than it does the minor contributor- when defining the major contributor, look for full octets and no chargeseparation – it’s also the one with the lowest energy; also, negative chargesare more stable on more electronegative atoms, such as O, N, and S- when drawing resonance structures, the only thing that can move iselectrons… do not move nuclei, and bond angles must stay the same; thenumber of unpaired electrons must remain the same- resonance stabilization is most important when it serves to delocalize a charge over two or more atoms- structural formulas show which atoms are bonded to which- condensed structural formulas are written without necessary putting the lines there (so like CH3CH3) – nonbonding electrons rarely shown- line-angle formula aka skeletal structure is written as lines and where two lines meet and there is no other indication it implies that there is a carbon bond there- molecular formula is when it’s written like C4H10O- empirical formula = divide the # of grams by each element’s atomic weight, then divide each of these numbers by the smallest number gotten in the previous step- to find molecular formula, get the molecular weight from the empiricalformula, divide it by the molecular weight given in the problem, thenmultiple each little number by that number- Arrhenius theory defined acids as substances that dissociate in water to give hydronium ions… it said that stronger acids were assumed to dissociate to a greater degree than weaker acids…. Therefore it also said that bases are substances that dissociate in water to give hydroxide ions, and strong bases are said to dissociate more completely than weaker acids- Kw = [H3O+] [-OH] = 1.00 x 10-14 M2 (at 25 deg. Celcius)- [H3O+] = [-OH] = 1.00 x 10-7 M in a neutral solutiono acidic and basic solutions are defined by an excess of hydronium ion or hydroxide ion- pH = -log10 [H3O+]- Bronsted-Lowry acid is any species that can donate a proton and Bronsted-Lowry base is any species that can accept a proton- this includes the Arrhenius acids and bases, because compounds thatdissociate to give hydronium ion(s) are proton donors, and compounds thatdissociate to give hydroxide ions are proton acceptors (hydroxide ion acceptsa proton to form water)- conjugate acids and bases – example: when HCl donates a proton (H), Cl is called the conjugate base to HCl, which is an acid- Ka is the acid dissociation constant and it indicates the relative strength of the acid- Ka = ( [H3O+]] [A-] ) / [HA]- pKa = -log10Ka- strong acids have pKa values around 0 (or negative sometimes) and weak acids have pKa values greater than 4- in acid/base reactions, equilibrium generally favors the weaker acid and base- (Ka) (Ka) = 10-14- more electronegative elements bear a negative charge more easily, giving a more stable conjugate base and a stronger acid- as EN increases, so do stability and acidity- inductive effect: electron donation or withdrawal through the sigma bonds of a molecule- note – multiple electron drawing groups (so multiple high


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DREXEL CHEM 241 - Chapter 1

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