Exam 1 study guide
Exam 1 study guide CHMY 321-001
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This 14 page Study Guide was uploaded by Rebeka Jones on Monday September 19, 2016. The Study Guide belongs to CHMY 321-001 at Montana State University taught by Holmgren, Steven in Fall 2016. Since its upload, it has received 207 views.
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Date Created: 09/19/16
Exam 1 Study Guide The Lewis dot structure of an individual atom indicates the number of valence electrons. Step 1: Determine the number of valance electrons Step 2: Place on valance electron by itself on each side of the atom. Step 3: If the atom has more than 4 valance electrons the remaining electrons must be paired with the already existing electrons. Lewis Dot Structures of atoms can be combined to create small molecules. When doing so follow the octet rule. Step 1: Draw all individual atoms. Step 2: Connect atoms that form more than one bond. Step 3: Connect single bonding elements. Step 4: pair any paired electrons to form an octet. A formal charge is associated with any atom that does not exhibit the appropriate number of valence electrons. When this exists in Lewis dot structures the charge must be drawn. Step 1: Determine the appropriate number of valance electrons for an atom. Step 2: Determine whether an atom exhibits the appropriate number of electrons. If there is an atom that has more electrons than it should assign a formal charge. Electronegativity generally increases as you go across the periodic table left and up. If the difference in the electronegativity of the two atoms is less than 0.5 the electrons are said to be equally shared resulting in a covalent bond. If the difference in electronegativity is between 0.5 and 1.7 the electrons are not shared equally resulting in a polar covalent bond. The withdrawal of electrons toward one atom is called induction and is indicated by an arrow showing where the electrons are coming from and where they are going. Induction causes areas of positive changes and negative charges. This is symbolized by delta. If the difference in electronegativity is greater than 1.7 the electrons are not shared at all resulting in an ionic bond. Ionic bonds are the result of the force of attraction between two oppositely charged ions. These values are simply guide lines many bond fall between groups. Step 1: Identify all polar covalent bond. Step 2: Determine the direction of each dipole. Step 3: Indicate the location of partial charges. H 2.1 Li Be B C N O F 1.0 1.5 2.0 2.5 3.0 3.5 4.0 Na Mg Al Si P S Cl 0.9 1.2 1.5 1.8 2.1 2.5 3.0 2 K Br 0.8 2.8 Orbitals with the same energy level are called degenerate orbitals. The order in which the orbitals are filled by electrons is determined by just three simple principals. 1. The Aufbau Principle – the lowest energy orbital is filled first. 2. The Pauli Principle – each orbital can accommodate a maximum of two electrons that have opposite spins. 3. Hund’s Rule – When dealing with degenerate orbitals such as p orbitals, one electron is placed in each degenerate orbital first, before electrons paired up According to the valence bond theory, a bond is simply the sharing of electron density between two atoms as a result of the constructive interface of their orbitals. Molecular orbital theory also describes a bond in terms of constructive interference between two overlapping atomic orbitals. But it also says that the orbitals are mathematically combined to produce new orbitals called molecular orbitals. Atomic orbitals are regions of space associated with an atom whereas molecular orbitals are associated with the whole molecules. When p orbitals overlap they create a pi bond. The bond forms above the plane of the molecule and below the plain. Sigma bonds experience free rotation at room temperature whereas pi bonds do not. A triple bind is formed by sp-hybridized carbon atoms one s orbital is mathematically averaged with only one p orbital. This lea ves 2 p orbitals 3 unaffected. Therefore, there is 2 sp orbitals and two p orbitals. The sp orbitals can form sigma bonds and the two p orbitals form pi bonds. A triple bond between two carbon atoms is therefore the result of three separate bonding interaction one sigma bond and two pi. A carbon atom with four single bond will be sp hybridized. A carbon atom 2 with three sigma bonds and one pi bond will be sp hybridized. A carbon atom with two sigma bonds and two pi bond will be sp hybridized. In order to predict the geometry of a bond we must count the number of sigma bonds and lone pairs. This is the steric number. It represents the number of electron pairs that are repulsing each other. This is VESPER theory. There are 3 different types of geometry arising from sp hybridization tetrahedral, trigonal pyramidal, bent. In all cases the electrons were arranged in a tetrahedron but the lone pair were ignored when describing geometry. 2 Sp -hybridization orbital achieves maximal separation in trigonal planer. Can also be bent. Sp hybridized orbitals have maximum separation when they are linear. If steric number is 4 If the steric If steric number is number is 3 2 3 2 Sp Sp Sp Tetrahedral arrangement of Trigonal Planar Linear electron pairs arrangement of arrangement of electron pairs electron pairs 4 No lone One Two One No lone Linear pairs lone pair lone lone pairs pairs pair Tetrahedral Trigonal Bent Trigonal planar planar Things are said to exhibit a dipole movement because the center of negative charge and the center of positive charge are separated from one another by a certain distance. The dipole movement is used as an indicat or of polarity, where it is defined as the amount of partial charge on either end of the dipole multiplied by the distance of separation. measuring the dipole moment allows us to calculate the % ionic character of that bond. When dealing with a compound that has more than one polar bond, it is necessary to take the vector sum of the individual dipole moments. This is called the molecular dipole moment, and it takes into account both magnitude and the direction of each individu al moment. The presence of lone pairs has a significant effect on the dipole moment. There is a dipole moment connected to each lone pair. Step 1: predict the molecular geometry – use steric number Step 2: Identify the direction of hall dipole moments. Step 3: Draw the net dipole moment. Intermolecular forces- attractive forces between individual molecules. All intermolecular forces are electrostatic meaning that they occur as a result of the attraction between opposite charges for neutral molecules there are 5 -dipole-dipole -hydrogen bonding -fleeting dipole dipole interactions In solid phase molecules align to attract each other whereas the liquid phase molecules are free and move more often to attract then to repel. More attraction = higher boiling point. Hydrogen bonding is a type of dipole-dipole interactions. Due to individual hydrogen has a partial positive charge. To understand fleeting dipole-dipole interactions consider the electrons to be in constant motion so the center of negative charge coincides with the venter of positive charge, resulting in a zero dipole moment. But the venters might not co-inside. This resulting transient dipole movement can create separate transient dipole moment in other molecules – this is called London dispersion forces. Larger molecules experience this more. Branching molecules have small surface areas. When comparing boiling points of compounds, we look for -are there any dipole-dipole interaction in either compound - will either compound form hydrogen interactions in either compound how much branching is in each compound. Polar compounds are solvable in polar solvents and vice versa. 6 Partially condensed structures – the C-H bonds are not all drawn explicitly – still only useful for small molecules Condensed structures – single bonds are not drawn. Instead groups of atoms are clustered together when possible. For example: if a molecule has two CH groups it would be written (CH ) . This still is impractical for large 3 3 2 molecules. Bond-line structures – simplify the drawing process and are easier to read. Most of the atoms are not drawn but with times these become user friendly -bond-line structures are drawn in a zig zag format and each corner or end point represents a carbon atom. -double bonds are shown with two lines and triple bonds are shown with three lines. *triple bonds are drawn in a linear fashion because triple bonds involve sp- hybridization which have linear geometry. -hydrogen atoms are not show because it is assumed each carbon has enough hydrogen atoms to achieve four bonds. How to draw bond line structures -carbon atoms in a straight line should be drawn in a zig zag format -when drawing double bonds, draw all bonds as far apart as possible - when drawing single bonds, the direction in which the bonds are drawn is irrelevant -all heteroatoms (atoms other than carbon and hydrogen) must be drawn and any hydrogen atoms attached to a heteroatom must be dr awn -never draw a carbon with more than 4 bonds 7 When given a Lewis structure -delete hydrogen atoms, except for those attached to heteroatoms -draw in a zig zag format -delete carbon atoms with bond line drawings it is easier to identify functional groups functional group – a characteristic group of atoms/bonds that possess a predictable chemical behavior The chemistry of every organic compound is determined by the functional groups present in the compound. Therefore, the classification of the compounds is based off of the functional groups. a carbon atom will generally have four bonds only when it doesn’t have a formal charge -when there is a positive formal charge carbon with form 3 bonds - when there is a negative formal charge carbon will form 3 bonds and hold a lone pair *formal charge must always be shown in bond line structure; lone pairs can be left out A bond-line structure is only clear if it contains either all of the lone pairs or all of the formal charges. Since there is normally more lone pairs than charges it is convention of always drawing formal charges and leaving off lone pairs 8 Patterns you will see for oxygen atoms -a negative charge corresponds with one bond and three lone pairs. -absence of charge corresponds with two bonds and two lone pairs -positive charge corresponds with three bonds and one lone pair Patterns you will see with nitrogen atoms -negative charge corresponds with two bonds and two lone pairs -absence of charge corresponds with three bonds and one lone pair -positive charge corresponds with four bonds and no lone pairs wedges – represent a group coming out of the page dash – represents a group going behind the page The way to deal with in adequacy of bond line structures is called resonance. According to this we draw one more bond-line and mentally meld them together. These are called resonance structures. They show how the positive charge is spread over two locations *we separate them with a straight two headed arrow and place brackets around them This metal meld is called a resonance hybrid Curved arrows are tools necessary to draw resonance structures properly. Every curve arrow has a tail and head. *They do not represent the movement of electrons The arrows treat electron AS IF they were moving even though they actually are not moving at all. 9 The tail shows where the electrons are coming from and the head shows where they are going. There are two rules to follow when drawing curved arrows for resonance structures. 1. Avoid breaking single bonds 2. Never exceed an octet for second row elements (octet rule) *second curved arrows can remove violations Structures are not complete without drawing formal charges. There are five patterns to drawing resonance structures. 1) An allylic lone pair – lone pair by pi bond a. Lone pair to create pi bond – two curved arrows 2) An allylic positive charge – positive charge in allylic position a. One arrow – tail of curved arrow is placed on lone pair and the head is placed to form pi bond between lone pair and positive charge 3) Lone pair adjacent to a positive charge a. One arrow – tail of curved arrow is placed on a lone pair and the head is placed to for pi bond between lone pair and positive charge 4) Pi bonds between two atoms of difference electronegativity a. Move pi bond up onto the electronegative atom to become a lone pair 5) Conjugate pi bonds enclosed in a ring a. Conjugate pi bonds are enclosed in a ring of alternating double and single bonds…we push all of the pi bond over by one position 10 A compound could have three resonance structures; the three structures might not contribute equally to the overall resonance hybrid. One might be major significance while another is insignificant 1) Minimizes charges a. Structures with more than two charges should be avoided 2) Electronegative atoms such as N, O, and Cl, can bear a positive charge but only if they possess an octet *the most significant resonance structures are generally those in which all the atoms have an octet. 3) Avoid drawing a resonance structure in which two atoms bear opposite charges -they are usually insignificant a lone pair that is allylic to a pi bond is said to be delocalized. They occupy a p orbital rather than a hybridized orbital. A localized lone pair does not participate in resonance. Whenever an atoms possess both a pi bond and a lone pair, they will not both participate in resonance. Functional Groups 1) Alkyl Halide: has C to halogen bonded to C-H a. X = F, Cl, Br, or I 2) Alkene: C double bond C 3) Alkyne: C triple bond C 4) Alcohol: C bonded to OH 5) Ether: C bonded to O bonded to C 6) Thiol: C bonded to SH 7) Sulfide: C bonded to S bonded to C 8) Aromatic (Arene): cyclic array of single and double C bonds 11 The definition of a Bronsted-Lowery acids and bases is based on the transfer of a proton (H+) An acid is a proton donor A bas is a proton acceptor The products of a proton transfer reaction are all conjugate based and the conjugate acid -conjugate base: what remains of the acid *water can be an acid or a base All reactions are accomplished via a flow of electron density. The flow of electron density is illustrated with curved arrows. Even though they look just like the arrows used for resonance structures these ones actually represent the movement of electrons. The arrows show the reaction mechanism – they show how the reaction occurs in reference to the electrons *always involves at least two arrows In order to compare acids without the use of pka values we have to look at the conjugate base of each acid. - - If A is very stable (weak base) than HA must be a strong acid. If A is very unstable (strong base) the HA must be a weak acid. Electronegative atoms stabilize a negative charge, Because carbon is not very electronegative it doesn’t stabilize negative charges well. 12 *by determining the more stable conjugate base we can identify the stronger acid 1. Which atom bears the charge a. Comparing the atoms that bear a negative charge in the conjugate base b. To determine which of these is more stable we must consider where these elements are on the periodic table i. When two atoms are in the same row, electronegativity is dominant ii. When two atoms are in the same column size is dominant 2. Resonance a. If your negative charge is on the same atom check for resonance. When there is resonance on one of the atoms the charge is more stabilized therefore that atom is more acidic. i. These compounds are called carboxylic acids 3. Induction a. When factor 1 and 2 do not show a difference look for difference in the other atoms in the molecules. i. For example, if one molecule has 3 Cl atoms while the other has none. Cl has an induction effect and therefore stabilizes the molecule. 4. Orbitals a. The differences before will not explain the differences in acidity between identified protons on an atom i. The answer comes from looking at the hybridization state of the orbitals that hold the charge 1. The closer the electrons are to the nucleus the more stable it is. 2 a. Sp-hybridized is more stable than sp - hybridized 13 The equilibrium will always favor the more stable negative charge. A Lewis acid is defined as an electron acceptor A Lewis base is defined as an electron donor 14
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