Chemistry
Chemistry December-05-11 2:56 PM
Quantum numbers: principle quantum # (n)- represents main energy level of an atom's electrons by a number. *increasing value of n corresponds to increasing values of energy for the electron {n n: I1,| 2, n>0} 3, 4 Letter code: K, L, M, N Angular momentum (l): tells about shape of electron cloud/orbital; l=0-sperical, l=1-dumbell shaped, l=2= 5 variations of shapes Letter code: S p d l can equal values from 0--> (n-1) ex. n=2 l=0,1 Magnetic quantum # ): determines the # of permitted orbitals in any group orientations in space of electron cloud, tells how many orbitals exist in sub shells Spin quantum # ( define orbital
values range from -l to +l; l=2
Defines orbital
-2, -1, 0,1, 2; specifies permitted
): not from Shrodinger' s equations, limits # of spin energies for an electron to 2 values 1/2; defines electrons while other Q#'s
Defines electron
Pauli exclusion principle- no 2 electrons can have all 4 Q#'s the same; 3 can be the same 1 must be different Hund's rule- electrons distributed in sub shell to yield maximum unpaired electrons Spin pair- an orbital is full- 2 electrons spinning opposite directions ( 1/2) 1s2 electrons AUFBAU principle- generalization, building up principle- electron structure Angular momentum # (l) determined by adding electrons 1 at a time until none left- MUST OBEY HUND'S Principle quantum # RULE, orbitals arranged by increasing energies; exception: 4s 3d14s2 NOT 4s23d1 Exception: Cr, Mo, W, and Cu, Ag, Au Cr--> [Ar] 3d44s2-->[Ar]3d54s1 Cu--> [Ar] 3d94s2--> [Ar] 3d104s1 For electron configuration of ions- start with neutral atom than adjust according to charge
Magnetic Properties: Diamagentism- not attracted or slightly repelled by magnetic fields Paramagnetism- drawn into by magnetic fields from unpaired electrons; paired electrons cancel each other out; electron behaves like magnets,larger # of unpaired electrons the larger magnetic moment
Periodic Trends:
Increasing
Atomic radius- L-R decrease in size (same principle shell and increase nuclear charge) T-B increase in size (increase principle shell) Decreasing Ionization energy- energy required to remove an electron; L-R increase ionization energy T-B decrease ionization energy Electron affinity- energy released when neutral atom gains an electron in gas phase; NOT COVALENT. Measures tendency to gain electrons large and -ve for non-metals follows trend of ionization energy Size and charge (ion vs. atom) removal of electron from neutral atoms results in remaining electrons more tightly held Na>Na+; addition of electrons overall attraction of nucleus for electrons decreases anion larger than atom ClF->Ne>Na+>Mg2->Al3+ Ionic radius = 1Ȧ= m=1 Cm Electronegativity-power atom has for attracting electrons COVALENT L-R: increases T-B: decreases
Decreasing Increasing
decreasing Increasing
Lewis Structures: - Use valence electrons Electron Electron pair 1. Determine # electrons 2. Draw skeleton structure with single bonds 3. Subtract electrons used from total valence and distribute leftover electrons 4. Too few electrons use double and triple bonds 5. Too many add leftover to centre atom as lone pairs Formal charges- =ve on free atom--> 1/2(# of electrons in bond)-(#of electrons in lone pairs); when you can write several Lewis structures choose 1 with lowest magnitude of formal charge, if same choose 1 with -ve on more electronegative atom Resonance- can't write single Lewis structure- true state of bonding not correctly represented by ANY Lewis diagram; exists between extremes Octet failure- satisfy for more electronegative atom
Chemistry 1040 exam review Page 1
VESPR Theory: - Framework for molecular geometry predictions - Steric # count σ bonds and lone pairs
Polar and Non-polar molecules:
Steric # Framework
Shape
Angles
hypridization
2
Linear
Linear
180⁰
sp
3
Triangular planar
Triangular planar Bent (angular)
120⁰ change in# of moles of gases (products-reactants), when adding eqns x k values Le Chatlier's Principle-when chemical equilibrium is disturbed by increase/decrease yield of product by changing [], partial pressure, volume, or temp., the equilibrium shifts to counteract the change pH Scale: pH=-log[H3O+], pOH=-log[OH-], [H3O+] x [OH-]= Kw=1.0x10-14, pH + pOH= pKw=14.00 Weak acids and weak bases- react only partially with water; weak acid:
Mathematical treatment: HA(aq) +H2O(l) H3O+(aq) + A-(aq) I: CHA 0 0 C: -x x x E: CHA -x x x
*Ka corresponds to weak acid + water Weak acid types • organic weak acids- carboxylic acids in general R-COOH, some alcohols R-OH = x2/(CHA -x) • Inorganic weak acids (no carbon or hydrogen) HF, HCN, B(OH) 3 + • Conjugate acids of weak bases NH 3 (weak acid) NH4 (conjugate W.A.) • Small highly charge metal ions Al 3+, Cr3+, Fe3+,Fe2+ All follow equilibrium: HA(aq) +H2O(l) H3O+(aq) + A-(aq); Ka(HA) OR BH+(aq) +H2O(l) H3O+(aq)+ B(aq); Ka(BH+) Weak acids partially dissociate, strong acids completely dissociate, SA have more ions SA >WA for conductivity, free H3O+ [H3O+] from SA > [H3O+] from WA- for similar concentrations pH (SA) < pH (WA) 5% Rule: find the value of x/CHA- if the value is 5% or less, the assumption that (CHA -x) equation for accurate x Chemistry 1040 exam review Page 2
CHA is valid, if >5% assumption fails, must use quadratic
equation for accurate x *Note x/CHA x 100% is also known as %ionization, %dissociation, % deprotonation, or %protonation depending on situation Percent Dissociation (and % ionization)- how much an acid dissociated = x/CHA x 100% *% dissociation (WA) INCREASES with dilution, for any weak acid x= [H3O+] = ) Weak Bases types - Ammonia and its organic derivatives; in general B(aq) + H 2O(l) BH +(aq) + OH-(aq); Kb(B) - Conjugate bases of weak acids; in general A -(aq)+ H2O(l) HA(aq) + OH -(aq); Kb(A-) Mathematical treatment same as weak acid Ka and Kb: Ka x Kb = Kw therefore any acid-base conjugate pair: Ka x Kb = Kw = 1.0x1014 at 25⁰C and Ka = Kb /Kw and Kb = Ka /Kw and pKw = pKa + pKb= 14.00 at 25⁰C
Aqueous Solutions: Cations:
Anions:
Acidic
Examples
Acidic
Examples
Conjugate acids of weak bases
C5H5NH+, C6H5NH3+, NH4+, CH3NH3+, etc
Very few
HSO4-, H2PO4-
Small, highly charged metal ions Fe3+, Cr3+, Al3+, Fe2+, Cu2+, Ni2+
Neutral Anions of strong acids
Cl–, Br–, I–, NO3–, ClO4-
Neutral Li+, Na+,...; Mg2+, Ca2+, …; Ag+, Cu+, etc.
Group 1 & 2 cations +1 metals
Basic Conjugate bases of weak acids F–, O2–, OH–, S2–, HS–, CN–, CO3 2–, HCO3–, HPO4–, etc.
Common ion effect: the shift in an ionic equilibrium caused by the addition of a solute that contains an ion that is part of the equilibrium Buffers: a solution which resists change in pH when either acid or base is added; it must have both a weak acid and its conjugate base (or vise versa) present in substantial amounts (conjugate pair ratio must be in range of 1/10 to 10--> buffer) Mathematical treatment: HA(aq) +H2O(l) H3O+(aq) + A-(aq) I: CHA 0 CA C: -x x x E: CHA -x x CA+x = x(CA+x)/(CHA -x) Henderson-Hasselbalch equation: pH= pKa + log(CA/CHA) Buffer limits- pH= When SA is added to buffer, acid reacts with base; When SB is added to buffer, base reacts with acid; Making a buffer solution: Add conjugates to water (within ratio), generate conjugate base from weak acid by adding OH-, generate conjugate acid from weak base (by adding H3O+) Buffer of desired pH- select acid with pKa close to required pH, fine tune by adjusting conjugate acid or base
Titrations: Indicators- used to indicate when reaction is complete (reacted equivalent # of mols of acid and base; organic dyestuffs that are weak acids HIn and In- must be brightly and differently coloured Equivalence point- the point in titration when equivalent moles of acid and base have been added or mixed together (aka stoichiometric point) End point- pH at which the indicator changes colour (ranges from **equivalence point and end point have nothing to do with eachother unless we intentionally choose the proper indicator- choose indicator so that pKa indicator is close to pH of anticipated equivalence point--> end point and equivalence point pHs should be as close as possible Strong acid- Strong base titration- H3O+(aq) + OH-(aq) 2H2O(l); K=1.0x1014 pH at equivalence point will be 7 bc we have a neutral solution
WA + SB HA(aq) +OH-(aq)
H2O(l) + AC-(aq)
WB + SA:
Chemistry 1040 exam review Page 3
Organic Chemistry: - Chemistry of carbon compounds - Carbon has unique ability to form 4 strong covalent bonds with itself or 4 covalent bonds Constitutional or structural isomers: compounds have same molecular formula but different structures and therefore different properties Organization of organic chemistry- need to classify compounds, based on chemical properties by certain reactive parts of molecule. Classify carbons according to number of carbon atoms to which they are attached 1⁰- attached to 1 other carbon, 2⁰ attached to 2 other carbons, 3⁰ attached to 3 other carbons Hydrogens classified in same way- based on carbon they are attached to Classify compounds based on chemical properties ie functional groups Conformations- free rotation about C-C single bonds, gives rise to infinite # of arrangements of atoms Stereochemistry- structural or constitutional isomers, steroisomers- different isomers in space: Geometric (Cis-Trans) - due to lack of free rotation about double bond or within ring, substitutions have different spatial arrangements Cis- same side, trans- opposite side; Optical isomers- chiral compounds (4 different substituents, react to polarized light, enantiomer- non-superimposable mirror images,
Intermolecular forces: London forces (van der Waals)- moving electrons produce a fleeting instantaneous dipole moment, partial charges on different molecules attract one another and molecules stick together, strength increases with molar mass, strength depends on shape Dipole-Dipole- polar molecules have permanent partial charges on top of instantaneous partial charges, strength depends on magnitude of dipoles and shape Hydrogen bonding- O, N, F only, occurs when a hydrogen atom lies between 2 small, strongly electronegative atoms with lone pairs of electrons, lone pair and partial charge attract each other strongly and form a bond, strongest when the H is on a straight line between 2 oxygens Melting and boiling points- reflects strength of intermolecular forces H-bonding> dipole> london Water solubility- "LIKE DISOLVES LIKE" H-bonding> dipole> london Chemical properties: Chemically reactive sites- electrophilic agent-H+ likes to attack π bond, nucleophilic agent- a negative or neutral species wanting to donate a pair of electrons to a positive charged carbon near an electronegative species like a halogen Free radicals- a neutral species having an unpaired electron Reaction intermediates- species that are produced and consumed during a reaction but do not appear in the overall chemical equation
Organic Reactions: Alkanes- generally lacking in reactivity due to strength and non-polar nature of bonds- inert to: common strong acids, common strong bases, oxidising agents, reducing agents; reactions mostly involve free radicals Types of reactions: pyrolysis- decomposition of compound by heating usually in of catalyst; combustion- burn in the presence of O2; Halogenation- react with a halogen in the presence of UV light or heat (free radical mechanism) - substitution reactions, very difficult to control, product separated by fractional distillation F2>Cl2>Br2>I2 preferred order of reactivity with H's in molecule: 3⁰>2⁰>1⁰ therefore major and minor products Alkenes- electrophilic addition- π source of electrons, attracts electrophilic agent, 2 step process, preferred position of B- 3⁰>2⁰>1⁰ (Markovnikov's Rule) Catalytic hydrogenation- or reduction (not electrophilic addition), H2 adds to π bond in the presence of a catalyst addition to same side of C=C bond Polymerization- polymers- compounds made up of many repeating subunits either naturally or synthetically made, monomer- smaller molecule units making up the polymer, polymerization- process by which monomers are combined to form polymers, alkenes are starting materials for large class of polymers (addition polymers), free radical sometimes used as initiator, n determines length of chain Organic Halogen compounds- important starting materials, carbon-halogen bond has large dipole, carbon site for nucleophilic attack Nucleophilic substitution (Nu:- or Nu:) replace halogen with nucleophile where Nu:- = OH-, NH2-, CN-, SH-, OR-, used to produce other functional groups (alcohol, amides, ethers, etc., most general and useful types of rxns Aromatic hydrocarbons: reactivity- structure is planar with bond length b/w a single and double - resonance hybrid, benzene ring very stable, under forcing conditions can hydrogenate, need strong catalyst for halogenation with all 6 hydrogens acting equivalently Alcohols and phenols: acid-base properties give nucleophile formation; oxidation reactions Ethers- prepared by nucleophilic substitution, more hydrocarbon- like alcohols, relatively inert, useful solvents for organic reactions Amines- prepared by nucleophilic substituion, basic in nature (acid base chemistry), used to form amides Aldehydes and ketones- carbon-oxygen bond has a large dipole, carbon site for nucleophilic attack, π bond source for electrons- electrophilic attack, aldehydes readily oxidise to carboxylic acids, ketones, under same conditions- not oxidized, polarized bond that is a resonance hybrid, dipole-dipole interactions, no H-bonding, water solubility decreases with # of carbon atoms; Reduction (electrophilic attack) aldehydes reduce to 1⁰ alcohols, ketones reduce to 2⁰ alcohols, reducing agents NaBH4 +H+/H2O, LiAlH4 + H+/H2O, Nucleophilic additions- carbon end attacked by nucleophiles, oxygen end attacked by electrophiles; Addition of water- forms geminal diol- hydroxyl groups bonded to the same atom, addition of alcohols- forms hemiacetals, equilibrium usually favours the carbonyl compound and the alcohol; test for aldehydes vs. ketones- Tollens reagent (Ag+ in NH3) precipitate of Ag with aldehydes none in ketones, only aldehyde can further reduce to carboxylic acid Carboxylic acids and derivatives- carbon-oxygen bond has a large dipole, carbon site for nucleophilic attack, carboxylic acids are weak acids and when placed
Chemistry 1040 exam review Page 4
Carboxylic acids and derivatives- carbon-oxygen bond has a large dipole, carbon site for nucleophilic attack, carboxylic acids are weak acids and when placed in water will produce acidic solutions, reacts with strong bases; Preparation of esters- direct esterification (acid + alcohol) or acid anhydride + alcohol; Hydrolysis of esters- formation of salts (aka saponification) Preparation of amides ester+amine--> amide + alcohol, carboxylic acid + amine--> amide + water Condensation polymers (Remove water by heat), polyesters and polyamides, monomers have 2 functional groups (diols, dicarboxylic acids), polyesters- ester formation reaction: dicarboxylic acid + diol--> polyester + water; can also use molecules having bothe functional groups: diacid halide + halide--> polyester + HX Organic acids and bases- acids- H's are attached to electronegative atoms (oxygen mainly), electron withdrawing effects increase the acidity of a hydrogen Bases- atoms with lone pair electrons can act basic if willing to donate them to a proton (H+), organic bases tend to be nitrogen compounds Salt formation: organic acid/base + acid/base--> salt, salt + acid/base--> organic acid/base
Chemistry 1040 exam review Page 5
Quantum numbers: principle quantum # (n)- represents main energy level of an atom's electrons by a number. *increasing value of n corresponds to increasing values of energy for the electron {n n: I1,| 2, n>0} 3, 4 Letter code: K, L, M, N Angular momentum (l): tells about shape of electron cloud/orbital; l=0-sperical, l=1-dumbell shaped, l=2= 5 variations of shapes Letter code: S p d l can equal values from 0--> (n-1) ex. n=2 l=0,1 Magnetic quantum # ): determines the # of permitted orbitals in any group orientations in space of electron cloud, tells how many orbitals exist in sub shells Spin quantum # ( define orbital
values range from -l to +l; l=2
Defines orbital
-2, -1, 0,1, 2; specifies permitted
): not from Shrodinger' s equations, limits # of spin energies for an electron to 2 values 1/2; defines electrons while other Q#'s
Defines electron
Pauli exclusion principle- no 2 electrons can have all 4 Q#'s the same; 3 can be the same 1 must be different Hund's rule- electrons distributed in sub shell to yield maximum unpaired electrons Spin pair- an orbital is full- 2 electrons spinning opposite directions ( 1/2) 1s2 electrons AUFBAU principle- generalization, building up principle- electron structure Angular momentum # (l) determined by adding electrons 1 at a time until none left- MUST OBEY HUND'S Principle quantum # RULE, orbitals arranged by increasing energies; exception: 4s 3d14s2 NOT 4s23d1 Exception: Cr, Mo, W, and Cu, Ag, Au Cr--> [Ar] 3d44s2-->[Ar]3d54s1 Cu--> [Ar] 3d94s2--> [Ar] 3d104s1 For electron configuration of ions- start with neutral atom than adjust according to charge
Magnetic Properties: Diamagentism- not attracted or slightly repelled by magnetic fields Paramagnetism- drawn into by magnetic fields from unpaired electrons; paired electrons cancel each other out; electron behaves like magnets,larger # of unpaired electrons the larger magnetic moment
Periodic Trends:
Increasing
Atomic radius- L-R decrease in size (same principle shell and increase nuclear charge) T-B increase in size (increase principle shell) Decreasing Ionization energy- energy required to remove an electron; L-R increase ionization energy T-B decrease ionization energy Electron affinity- energy released when neutral atom gains an electron in gas phase; NOT COVALENT. Measures tendency to gain electrons large and -ve for non-metals follows trend of ionization energy Size and charge (ion vs. atom) removal of electron from neutral atoms results in remaining electrons more tightly held Na>Na+; addition of electrons overall attraction of nucleus for electrons decreases anion larger than atom ClF->Ne>Na+>Mg2->Al3+ Ionic radius = 1Ȧ= m=1 Cm Electronegativity-power atom has for attracting electrons COVALENT L-R: increases T-B: decreases
Decreasing Increasing
decreasing Increasing
Lewis Structures: - Use valence electrons Electron Electron pair 1. Determine # electrons 2. Draw skeleton structure with single bonds 3. Subtract electrons used from total valence and distribute leftover electrons 4. Too few electrons use double and triple bonds 5. Too many add leftover to centre atom as lone pairs Formal charges- =ve on free atom--> 1/2(# of electrons in bond)-(#of electrons in lone pairs); when you can write several Lewis structures choose 1 with lowest magnitude of formal charge, if same choose 1 with -ve on more electronegative atom Resonance- can't write single Lewis structure- true state of bonding not correctly represented by ANY Lewis diagram; exists between extremes Octet failure- satisfy for more electronegative atom
Chemistry 1040 exam review Page 1
VESPR Theory: - Framework for molecular geometry predictions - Steric # count σ bonds and lone pairs
Polar and Non-polar molecules:
Steric # Framework
Shape
Angles
hypridization
2
Linear
Linear
180⁰
sp
3
Triangular planar
Triangular planar Bent (angular)
120⁰ change in# of moles of gases (products-reactants), when adding eqns x k values Le Chatlier's Principle-when chemical equilibrium is disturbed by increase/decrease yield of product by changing [], partial pressure, volume, or temp., the equilibrium shifts to counteract the change pH Scale: pH=-log[H3O+], pOH=-log[OH-], [H3O+] x [OH-]= Kw=1.0x10-14, pH + pOH= pKw=14.00 Weak acids and weak bases- react only partially with water; weak acid:
Mathematical treatment: HA(aq) +H2O(l) H3O+(aq) + A-(aq) I: CHA 0 0 C: -x x x E: CHA -x x x
*Ka corresponds to weak acid + water Weak acid types • organic weak acids- carboxylic acids in general R-COOH, some alcohols R-OH = x2/(CHA -x) • Inorganic weak acids (no carbon or hydrogen) HF, HCN, B(OH) 3 + • Conjugate acids of weak bases NH 3 (weak acid) NH4 (conjugate W.A.) • Small highly charge metal ions Al 3+, Cr3+, Fe3+,Fe2+ All follow equilibrium: HA(aq) +H2O(l) H3O+(aq) + A-(aq); Ka(HA) OR BH+(aq) +H2O(l) H3O+(aq)+ B(aq); Ka(BH+) Weak acids partially dissociate, strong acids completely dissociate, SA have more ions SA >WA for conductivity, free H3O+ [H3O+] from SA > [H3O+] from WA- for similar concentrations pH (SA) < pH (WA) 5% Rule: find the value of x/CHA- if the value is 5% or less, the assumption that (CHA -x) equation for accurate x Chemistry 1040 exam review Page 2
CHA is valid, if >5% assumption fails, must use quadratic
equation for accurate x *Note x/CHA x 100% is also known as %ionization, %dissociation, % deprotonation, or %protonation depending on situation Percent Dissociation (and % ionization)- how much an acid dissociated = x/CHA x 100% *% dissociation (WA) INCREASES with dilution, for any weak acid x= [H3O+] = ) Weak Bases types - Ammonia and its organic derivatives; in general B(aq) + H 2O(l) BH +(aq) + OH-(aq); Kb(B) - Conjugate bases of weak acids; in general A -(aq)+ H2O(l) HA(aq) + OH -(aq); Kb(A-) Mathematical treatment same as weak acid Ka and Kb: Ka x Kb = Kw therefore any acid-base conjugate pair: Ka x Kb = Kw = 1.0x1014 at 25⁰C and Ka = Kb /Kw and Kb = Ka /Kw and pKw = pKa + pKb= 14.00 at 25⁰C
Aqueous Solutions: Cations:
Anions:
Acidic
Examples
Acidic
Examples
Conjugate acids of weak bases
C5H5NH+, C6H5NH3+, NH4+, CH3NH3+, etc
Very few
HSO4-, H2PO4-
Small, highly charged metal ions Fe3+, Cr3+, Al3+, Fe2+, Cu2+, Ni2+
Neutral Anions of strong acids
Cl–, Br–, I–, NO3–, ClO4-
Neutral Li+, Na+,...; Mg2+, Ca2+, …; Ag+, Cu+, etc.
Group 1 & 2 cations +1 metals
Basic Conjugate bases of weak acids F–, O2–, OH–, S2–, HS–, CN–, CO3 2–, HCO3–, HPO4–, etc.
Common ion effect: the shift in an ionic equilibrium caused by the addition of a solute that contains an ion that is part of the equilibrium Buffers: a solution which resists change in pH when either acid or base is added; it must have both a weak acid and its conjugate base (or vise versa) present in substantial amounts (conjugate pair ratio must be in range of 1/10 to 10--> buffer) Mathematical treatment: HA(aq) +H2O(l) H3O+(aq) + A-(aq) I: CHA 0 CA C: -x x x E: CHA -x x CA+x = x(CA+x)/(CHA -x) Henderson-Hasselbalch equation: pH= pKa + log(CA/CHA) Buffer limits- pH= When SA is added to buffer, acid reacts with base; When SB is added to buffer, base reacts with acid; Making a buffer solution: Add conjugates to water (within ratio), generate conjugate base from weak acid by adding OH-, generate conjugate acid from weak base (by adding H3O+) Buffer of desired pH- select acid with pKa close to required pH, fine tune by adjusting conjugate acid or base
Titrations: Indicators- used to indicate when reaction is complete (reacted equivalent # of mols of acid and base; organic dyestuffs that are weak acids HIn and In- must be brightly and differently coloured Equivalence point- the point in titration when equivalent moles of acid and base have been added or mixed together (aka stoichiometric point) End point- pH at which the indicator changes colour (ranges from **equivalence point and end point have nothing to do with eachother unless we intentionally choose the proper indicator- choose indicator so that pKa indicator is close to pH of anticipated equivalence point--> end point and equivalence point pHs should be as close as possible Strong acid- Strong base titration- H3O+(aq) + OH-(aq) 2H2O(l); K=1.0x1014 pH at equivalence point will be 7 bc we have a neutral solution
WA + SB HA(aq) +OH-(aq)
H2O(l) + AC-(aq)
WB + SA:
Chemistry 1040 exam review Page 3
Organic Chemistry: - Chemistry of carbon compounds - Carbon has unique ability to form 4 strong covalent bonds with itself or 4 covalent bonds Constitutional or structural isomers: compounds have same molecular formula but different structures and therefore different properties Organization of organic chemistry- need to classify compounds, based on chemical properties by certain reactive parts of molecule. Classify carbons according to number of carbon atoms to which they are attached 1⁰- attached to 1 other carbon, 2⁰ attached to 2 other carbons, 3⁰ attached to 3 other carbons Hydrogens classified in same way- based on carbon they are attached to Classify compounds based on chemical properties ie functional groups Conformations- free rotation about C-C single bonds, gives rise to infinite # of arrangements of atoms Stereochemistry- structural or constitutional isomers, steroisomers- different isomers in space: Geometric (Cis-Trans) - due to lack of free rotation about double bond or within ring, substitutions have different spatial arrangements Cis- same side, trans- opposite side; Optical isomers- chiral compounds (4 different substituents, react to polarized light, enantiomer- non-superimposable mirror images,
Intermolecular forces: London forces (van der Waals)- moving electrons produce a fleeting instantaneous dipole moment, partial charges on different molecules attract one another and molecules stick together, strength increases with molar mass, strength depends on shape Dipole-Dipole- polar molecules have permanent partial charges on top of instantaneous partial charges, strength depends on magnitude of dipoles and shape Hydrogen bonding- O, N, F only, occurs when a hydrogen atom lies between 2 small, strongly electronegative atoms with lone pairs of electrons, lone pair and partial charge attract each other strongly and form a bond, strongest when the H is on a straight line between 2 oxygens Melting and boiling points- reflects strength of intermolecular forces H-bonding> dipole> london Water solubility- "LIKE DISOLVES LIKE" H-bonding> dipole> london Chemical properties: Chemically reactive sites- electrophilic agent-H+ likes to attack π bond, nucleophilic agent- a negative or neutral species wanting to donate a pair of electrons to a positive charged carbon near an electronegative species like a halogen Free radicals- a neutral species having an unpaired electron Reaction intermediates- species that are produced and consumed during a reaction but do not appear in the overall chemical equation
Organic Reactions: Alkanes- generally lacking in reactivity due to strength and non-polar nature of bonds- inert to: common strong acids, common strong bases, oxidising agents, reducing agents; reactions mostly involve free radicals Types of reactions: pyrolysis- decomposition of compound by heating usually in of catalyst; combustion- burn in the presence of O2; Halogenation- react with a halogen in the presence of UV light or heat (free radical mechanism) - substitution reactions, very difficult to control, product separated by fractional distillation F2>Cl2>Br2>I2 preferred order of reactivity with H's in molecule: 3⁰>2⁰>1⁰ therefore major and minor products Alkenes- electrophilic addition- π source of electrons, attracts electrophilic agent, 2 step process, preferred position of B- 3⁰>2⁰>1⁰ (Markovnikov's Rule) Catalytic hydrogenation- or reduction (not electrophilic addition), H2 adds to π bond in the presence of a catalyst addition to same side of C=C bond Polymerization- polymers- compounds made up of many repeating subunits either naturally or synthetically made, monomer- smaller molecule units making up the polymer, polymerization- process by which monomers are combined to form polymers, alkenes are starting materials for large class of polymers (addition polymers), free radical sometimes used as initiator, n determines length of chain Organic Halogen compounds- important starting materials, carbon-halogen bond has large dipole, carbon site for nucleophilic attack Nucleophilic substitution (Nu:- or Nu:) replace halogen with nucleophile where Nu:- = OH-, NH2-, CN-, SH-, OR-, used to produce other functional groups (alcohol, amides, ethers, etc., most general and useful types of rxns Aromatic hydrocarbons: reactivity- structure is planar with bond length b/w a single and double - resonance hybrid, benzene ring very stable, under forcing conditions can hydrogenate, need strong catalyst for halogenation with all 6 hydrogens acting equivalently Alcohols and phenols: acid-base properties give nucleophile formation; oxidation reactions Ethers- prepared by nucleophilic substitution, more hydrocarbon- like alcohols, relatively inert, useful solvents for organic reactions Amines- prepared by nucleophilic substituion, basic in nature (acid base chemistry), used to form amides Aldehydes and ketones- carbon-oxygen bond has a large dipole, carbon site for nucleophilic attack, π bond source for electrons- electrophilic attack, aldehydes readily oxidise to carboxylic acids, ketones, under same conditions- not oxidized, polarized bond that is a resonance hybrid, dipole-dipole interactions, no H-bonding, water solubility decreases with # of carbon atoms; Reduction (electrophilic attack) aldehydes reduce to 1⁰ alcohols, ketones reduce to 2⁰ alcohols, reducing agents NaBH4 +H+/H2O, LiAlH4 + H+/H2O, Nucleophilic additions- carbon end attacked by nucleophiles, oxygen end attacked by electrophiles; Addition of water- forms geminal diol- hydroxyl groups bonded to the same atom, addition of alcohols- forms hemiacetals, equilibrium usually favours the carbonyl compound and the alcohol; test for aldehydes vs. ketones- Tollens reagent (Ag+ in NH3) precipitate of Ag with aldehydes none in ketones, only aldehyde can further reduce to carboxylic acid Carboxylic acids and derivatives- carbon-oxygen bond has a large dipole, carbon site for nucleophilic attack, carboxylic acids are weak acids and when placed
Chemistry 1040 exam review Page 4
Carboxylic acids and derivatives- carbon-oxygen bond has a large dipole, carbon site for nucleophilic attack, carboxylic acids are weak acids and when placed in water will produce acidic solutions, reacts with strong bases; Preparation of esters- direct esterification (acid + alcohol) or acid anhydride + alcohol; Hydrolysis of esters- formation of salts (aka saponification) Preparation of amides ester+amine--> amide + alcohol, carboxylic acid + amine--> amide + water Condensation polymers (Remove water by heat), polyesters and polyamides, monomers have 2 functional groups (diols, dicarboxylic acids), polyesters- ester formation reaction: dicarboxylic acid + diol--> polyester + water; can also use molecules having bothe functional groups: diacid halide + halide--> polyester + HX Organic acids and bases- acids- H's are attached to electronegative atoms (oxygen mainly), electron withdrawing effects increase the acidity of a hydrogen Bases- atoms with lone pair electrons can act basic if willing to donate them to a proton (H+), organic bases tend to be nitrogen compounds Salt formation: organic acid/base + acid/base--> salt, salt + acid/base--> organic acid/base
Chemistry 1040 exam review Page 5
