201 2l2g6u

51

CHEMISTRY NOTES

GENERAL CHEMISTRY

Dalton’s Atomic Theory Main postulates of Dalton atomic theory are as follows: Matter is composed of very tiny or microscopic particles called "atom". 1. Atom is an indivisible particle. 2. Atom can neither be created nor destroyed. 3. Atoms of an element are identical in size, shape, mass, and in other properties. 4. Atoms of different elements are different in their properties. 5. Atoms combine with each other in small whole numbers. 6. All chemical reactions are due to the combination or separation of atoms. Defects in Dalton’s Theory Postulates 2, 3, 4 and 6 are incorrect as described below: > Atom can be divided into a number of sub-atomic particles: electron, proton, and neutron. > Atoms of an element may be different in their masses. (1H1, 1H2, 1H3) > Not all compounds have small number of atoms. (Decane C10H22, Sugar C10H22O11) > Atoms can be destroyed by fission process (atom bomb, nuclear reactor) On the basis of above defects, Dalton's atomic theory has failed now.

Law of Conservation of Mass: The total mass of the products of a chemical reaction is the same as the total mass of the reactants (mass remains constant during chemical reaction) Law of Definite Composition (Law of Definite Proportions): The elemental composition of a pure compound is always the same, regardless of its source Law of Multiple Proportions: When two elements form more than one compound, the masses of one element in these compounds for a fixed mass of the other element are in ratios of small whole numbers Atomic Number: number of protons in the nucleus Mass Number: number of protons + number of neutrons Isotopes: atoms of the same element with different mass numbers (carbon-12 and carbon-13) Atomic Mass Unit (amu): exactly one-twelfth the mass of a carbon-12 atom Atomic Mass (weight): average mass of the atoms of an element in amu: mass in g of one mol of the element

Matter: anything that has mass and occupies space

Mole (mol): quantity of a given substance that contains as many molecules or formula units as the number of atoms in exactly 12 g of carbon-12; consists of 6.02 x 1023 particles

Mixtures: combinations of two or more substances in which each substance retains its own chemical identity and hence its own properties

Molar Mass: mass of one mol of substance; in g it is numerically equal to the formula mass (weight) in atomic mass units

Physical Change: change of state or phase

Mass % of A = (

𝑚𝑎𝑠𝑠 𝑜𝑓 𝐴 𝑖𝑛 𝑡ℎ𝑒 𝑤ℎ𝑜𝑙𝑒 ) × 100% 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑡ℎ𝑒 𝑤ℎ𝑜𝑙𝑒

Chemical Change: transformation in which substances are converted into other substances Compound: substance composed of two or more elements chemically combined in fixed proportions Element: substance that cannot be decomposed by any chemical reaction into simpler substances. Matter Heterogeneous Matter (variable composition)

by physical means into

Homogeneous Matter (variable composition)

by physical means into

Compounds

by chemical means into

Homogeneous Matter

Pure Substances (fixed composition)

Elements

Molecule: definite group of atoms that are chemically bonded together Ion: electrically-charged particle obtained from an atom or chemically bonded group of atoms by adding or removing electrons Chemical Formula: notation that uses atomic symbols with numerical subscripts to convey the relative proportions of atoms of the different elements in the substance Empirical Formula (Simplest Formula): formula with the smallest integer subscripts Molecular Formula: chemical formula that gives the exact number of different atoms of an element in a molecule Molecular Empirical Substance Formula Formula water H2O H2O hydrogen H2O2 HO peroxide ethane C2H6 CH3

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52 Molecular Mass / Weight (MW): the sum of the atomic masses/weights of all the atoms in a molecule Formula Mass / Weight (FW): the sum of the atomic masses/weights of all atoms in a formula unit of the compound, whether molecular or not Empirical Formula Mass or Empirical Formula Weight (EFW): the sum of the atomic masses divided by the weights of all the atoms in an empirical formula Molecular Formula = 𝑛 × 𝐸𝑚𝑝𝑖𝑟𝑖𝑐𝑎𝑙 𝐹𝑜𝑟𝑚𝑢𝑙𝑎 𝒎𝒐𝒍𝒆𝒄𝒖𝒍𝒂𝒓 𝑜𝑟 𝑓𝑜𝑟𝑚𝑢𝑙𝑎 𝑤𝑒𝑖𝑔ℎ𝑡 𝑒𝑚𝑝𝑖𝑟𝑖𝑐𝑎𝑙 𝑓𝑜𝑟𝑚𝑢𝑙𝑎 𝑤𝑒𝑖𝑔ℎ𝑡

where: 𝑛 =

Stoichiometry: Calculation of the quantities of reactants and products involved in a chemical reaction: N2(g) 1 molecule N2 1 mole N2 28.0g N2

+ + + +

3H2(g) 3 molecules H2 3 moles H2 3 X 2.02g H2

   

2NH3(g) 2 molecules NH3 2 moles NH3 2 X 17.0g NH3

chemical reaction

Atomic Size: within a period, decreases from left to right; within a group, increases from top to bottom Size of Cation: within a period, decreases from left to right; within a group, increases from top to bottom

molar interpretation mass interpretation

Ionization Energy (energy needed to remove an electron from a gaseous atom in the ground state): within a period, increases from left to right; within a group, decreases from top to bottom

Actual Yield: Amount of product actually obtained in a reaction 𝑥 100%

Atomic Structure and Periodic Table The Nuclear Atomic Model: Protons and neutrons in the nucleus; electrons outside the nucleus The Quantum Mechanical Model Quantum numbers: n, l, ml, ms > Principal Quantum Number: n = 1(K), 2(L), 3(M)…; identifies the shell or level to which the electron belongs > Azimuthal Quantum Number: l = 0(s), 1(p), 2(d), 3(f)…; n-1 identifies the subshell > Magnetic Quantum Number: ml = +1…0…-1 gives the shape of the subshell: orientation of atomic orbital spin quantum number 1 1 > Spin Quantum Number: ms =+ , − 2

Periodic Properties of the Elements

Size of Anion: within a period, decreases from left to right; within a group, increases from top to bottom

Theoretical Yield: Quantity of product that is calculated to form when all of the LR reacts

𝒂𝒄𝒕𝒖𝒂𝒍 𝒚𝒊𝒆𝒍𝒅 𝒕𝒉𝒆𝒐𝒓𝒆𝒕𝒊𝒄𝒂𝒍 𝒚𝒊𝒆𝒍𝒅

Aufbau Principle (Building-up Principle): A scheme used to reproduce the electronic configurations of atoms by successively filling subshells with electrons in a specific order (orbitals with lower energy first).

molecular interpretation

Limiting Reactant or Limiting Reagent (LR): reactant that is entirely consumed when a reaction completes

Percentage Yield =

Pauli’s Exclusion Principle: No two electrons in the same atom may have the same set of all four quantum numbers. Hund’s Rule of Multiplicity: In the ground state of an atom, electrons are distributed among the orbitals of a subshell in a way that gives the maximum number of unpaired electrons with parallel spins.

2

Heisenberg’s Uncertainty Principle: It is impossible to determine simultaneously the exact position and the exact momentum of a fast moving body like an electron.

Electron Affinity (energy change when an electron is added to a gaseous atom in the ground state): within a period, increases negative value from left to right; within a group, no clear trend Basicity of Metal Oxides: within a period, decreases from left to right; within a group, increases from top to bottom Acidity of Non-metal Oxides: within a period, increases from left to right; within a group, decreases from top to bottom Metallic Property (reducing property): within a period, decreases from left to right; within a group, increases from top to bottom Non-metallic Property (oxidizing property): within a period, increases from left to right; within a group, decreases from top to bottom Electronegativity (measure of the ability of an atom in a molecule to draw bonding electrons to itself): within a period, increases from left to right; within a group, decreases from top to bottom Chemical Bond Electrovalent Bond (ionic bond): formed by transfer of electrons to form ions Covalent Bond: formed by sharing of electrons between atoms > Lewis or electron-dot formulas

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53 > Normal and coordinate covalent bonds > Formal charge and oxidation number > Formal Charge (FC) = Group No. minus No. of unshared e’s minus No. of covalent bonds > The higher the bond order, the shorter the bond length and the greater the bond strength Increasing bond length: triple bond < double bond < single bond Decreasing bond length: triple bond > double bond > single bond Metallic Bond: a crystal of the metal consists of a regular arrangement of the positive ions, and the valence electrons move throughout the crystal attracted to the positive cores of all the positive ions. Polar and Non-Polar Bonds Non-Polar Covalent Bond: equal sharing of a pair of e’s in diatomic molecules consisting of identical atoms in identical atoms with identical neighbors Polar Covalent Bond: unequal sharing of a pair of e’s in identical atoms with different neighbors when two bonded atoms are dissimilar Polar and Non-Polar Molecules: consider polarity of bonds, then geometry of molecule > Non-polar bonds  non-polar molecules > Polar bonds but symmetrical  non-polar molecules > Polar bonds and not symmetrical  polar molecules Bonding e-pairs 6 5 4 3 2

Non-bonding e-pairs 0 1 2 3 4

Shape of molecule octahedral square pyramidal square planar T-shaped linear

Hybrid Orbitals: mixing of atomic orbitals to form a hybrid (e.g., sp, sp2, sp3). The number of hybrid orbitals is equal to the number of pure atomic orbitals that mixed. Phases of Matter: consider intermolecular forces of attraction 1. Ion-dipole 2. Dipole-dipole 3. London forces / Van der Waals forces 4. H-bonding

INORGANIC CHEMISTRY Gases Kinetic Molecular Theory 1. tiny molecules, far apart 2. constant, random, straight line motion: elastic collisions 3. average KE increases with increase in T 4. attractive forces are negligible Ideal Gas Law PV=nRT

“Carefree Boys & Girls Play Tennis & Volleyball” Law Charle’s Law Boyle’s Law Gay-Lussac’s Law

Constant Pressure Temperature Volume

Relationship ↑ V= ↑T ↑P=↓V ↑ P= ↑T

Liquids Vapor Pressure: pressure of vapor in equilibrium with a liquid at a given temperature Boiling Point: temperature at which the vapor pressure of a liquid is equal to the external pressure Viscosity: resistance of fluid to flow Surface Tension: measure of the inward force on the surface of the liquid caused by intermolecular forces of attraction Solids Molecular Solids: atoms or molecules held together by intermolecular forces Metallic Solids: positive cores of atoms held together by a “sea” of electrons (metallic bonding) Ionic Solids: cations and anions held together by the electrical attraction of opposite charges (ionic bonds) Covalent Network Solids: atoms held together in large networks or chains by covalent bonds Thermodynamics > Tells whether the reaction is spontaneous or not First Law of Thermodynamics: Energy can be converted from one form into another, but it can neither be created nor destroyed. Second Law of Thermodynamics: Every spontaneous change is accompanied by an increase in entropy. Gibb’s Free Energy: Reactions tend to seek a minimum in energy and a maximum in randomness. Third Law of Thermodynamics: At absolute zero the entropy of a perfect crystalline substance may be taken as zero.

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54 EXERGONIC ENDERGONIC Spontaneous Non-spontaneous ΔS=+ (positive) ΔS =- (negative) ↑ entropy (disorderliness) ↓ entropy (disorderliness) ΔG = - (negative) ΔG = + (positive) Gibb’s Free energy Gibb’s Free energy ΔH = - (negative) ΔH = + (positive) exothermic endothermic Ea = Activation Energy Ea exergonic < Ea endergonic Solutions Molarity, M: Number of moles of the solute per liter of the solution Molality, m: Number of moles of the solute per kilogram of the solvent Mole Fraction, X 𝑁𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑚𝑜𝑙𝑒𝑠 ∆ 𝑋∆ = 𝑇𝑜𝑡𝑎𝑙 𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 Equivalent Point: point in a titration when an equivalent amount of titrant has been added to the analyte solution Colligative Properties: depend on the number of particles and not on the kind of particle Raoult’s Law: Pertains to ideal solutions in which the intermolecular forces between molecules are all essentially the same. 1. Boiling Point Elevation, Tb = Kb m where: Kb for water is 0.520C per molal solution 2. Freezing Point Depression,Tf = Kf m where: Kf for water is 1.860C per molal solution Buffers: solutions capable of maintaining their pH at some fairly constant value even when small amounts of acid or base are added; could be a mixture of a weak acid and the salt of its conjugate base (e.g. HOAc-OAc-) or a mixture of a weak base and the salt of its conjugate acid (e.g. NH3-NH4-) Henderson-Hasselbalch Equation: [𝐴− ] 𝑝𝐻 = 𝑝𝐾𝑎 + log [𝐻𝐴] where pH = -log [H+] and pKa = - log Ka [A-] = molar concentration of a conjugate base [HA] = molar concentration of weak acid The equation can be rewritten to solve for pOH: [𝐻𝐵 + ] 𝑝𝑂𝐻 = 𝑝𝐾𝑏 + log [𝐵] [HB+] = molar concentration of the conjugate acid [B] = molar concentration of a weak base

pH = 7 pH > 7 pH < 7

SALT STRUCTURE +(BASE) (ACID ) Strong Base Strong Acid Strong Base Weak Acid Weak Base Strong Acid

HCl HBr HI

Strong Acid HClO4 HClO3 HBrO4 HBrO3 HIO4 HIO3

H2SO4 HNO3

Weak HClO2 HBrO2 HIO2 H2SO3 HNO2

Acid HClO HBrO HIO HNO

CH3COOH Acetic Acid (HAc) Strong Base

Weak Base

Amphoteric Base

IA

IIA

Heavy Metal

Li

Ca

NH4OH (ammonium hydroxide)

Pb

As

Na

Mg

R- NH2 (amines)

Al

Au

K

Sr

Hg

Ag

Ba Nuclear Chemistry Balancing Nuclear Reactions - Types of Nuclear Reactions: 1. Radioactive Decay: process in which an unstable nucleus is changed by the emission of radiation a. alpha decay

238 234 4 𝑈 → 𝑇ℎ + 𝐻𝑒 92 90 2

125 125 b. gamma radiation [ 𝑇𝑒] ∗ → 𝑇𝑒 + 𝛾 52 52 excited state  ground state c. -decay

131 131 0 𝐼 → 𝑋𝑒 + 𝑒 53 54 −1

d. positron decay

11 11 0 𝐶 → 𝐵+ 𝑒 6 5 1

e. electron capture 81 0 81 𝑅𝑏 + 𝑒(𝑜𝑟𝑏𝑖𝑡𝑎𝑙 𝑒𝑙𝑒𝑐𝑡𝑟𝑜𝑛) → 𝐾𝑟 37 −1 36 2. Transmutation: process in which one nucleus is transformed into another through bombardmen by various subatomic particles or ions 14 4 17 4 𝑁 + 𝐻𝑒 → 𝑂 + 𝐻𝑒 7 2 8 2 3. Nuclear Fission: process in which a heavy nucleus is split into lighter ones 1 235 142 91 1 𝑛+ 𝑈 → 𝐵𝑎 + 𝐾𝑟 + 3 𝑛 0 92 56 36 0 4. Nuclear Fusion: process in which light nuclei are fused into heavier one 1 1 2 0 𝐻+ 𝐻 → 𝐻+ 𝑒 1 1 1 1

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55 ORGANIC CHEMISTRY Organic chemistry is the study of the compounds of carbon, their structure, chemical and physical properties, as well as their preparation and source. Despite the limited number of atoms that consist organic compounds (C, H, O, N, S, P, and the halogens), their number is much greater than the inorganic ones. This is due to the unique versatility of carbon. Organic chemistry developed from one of the most fundamental theories, the structural theory which proposes that carbon has a valence of four and thus tetravalent, and that carbon can use one or more of its valences to form bonds with other atoms including carbon itself. The nature of chemical bonds in organic compounds is covalent (resulting from the sharing of electrons). Covalent bonds are formed by overlapping orbitals of two atoms in different ways: 1. Head-to-head overlap – leading to sigma (σ) bond formation. - by atomic orbitals (s, p, d) and/or hybrid orbitals (sp, sp2, sp3 and sp3d) 2. Side-to-side overlap – leading to pi (π) bond formation - by p and d orbitals only

Knowledge of the orbitals used in bond formation is needed to identify the types of bonds formed and the location of lone pairs of electrons (if any).

Characteristics of Covalent Bonds 1. Polarity – results when two atoms of different electronegatives (EN) form a covalent bond such that the electrons are not equally shared between them e.g. H - Cl or H – Cl δ+ δI - a dipole moment results from the partial positivity and partial negativity of a polar covalent bond - a molecule with no net dipole moment is a nonpolar molecule; a molecule with a net dipole moment is a polar molecule e.g. C - F > C - Cl > C - Br > C - I - there is also such thing as bond polarity (even if the two atoms held by the bond are the same) due to the difference in hybridization between the two atoms such as in the sp-sp3 and sp2-sp3 hybrid orbitals resulting to electron-withdrawing property due to the greater scharacter of one of the atoms. e.g. H–CC–CH3 > CH2  CH – CH3 >CH3 – CH2 – CH3 2. Bond Length – the average distance separating the nuclei of bonded atoms - bond length decreases with decreasing atomic radius/size: C–I > C–C > C–H - bond length decreases with increasing polarities: C–C > C–N > C–O > C–F - bond length decreases with increasing s character of the hybrid orbitals: H–CH2–CH2–CH3 > H–CHCH–CH3 > H–CC–CH3 - bond length decreases with increasing π bond electron density: H3C – CH3 > H2C = CH2 > HC  CH 3. Bond Energy/Strength – the energy needed to disrupt a covalent bond (in kJ/ mol)

- bond strength increases with increasing polarity: C–F > C – Cl > C – Br > C–I - bond strength increases with increasing s character of the hybrid orbitals: sp–sp > sp2 – sp > sp2 – sp2 > sp3 – sp2 > sp3 – sp3 - bond strength increases with increasing π bond electron density: C  C > C=C > C–C 198 kcal 163 kcal 88 kcal Structural Effects – show how electrons are distributed in organic molecules and how they affect the physical and chemical properties of the molecule 1. π -Electron Delocalization or Resonance - this concept generally applies to conjugated systems and is used to explain strengthening and shortening of bonds, nucleophilicity/electrophilicity of sites, and thermodynamic stability 2. Inductive Effects - results in the distortion of electron distribution due to the effect through sigma bond/s by strong electrondonating (EDG) or electron-withdrawing (EWG) atoms/ group of atoms - strong EDG which release electrons by resonance effect include: - NH2, -NHR and –NR2 - moderate EDG whose nonbonding electron pairs are less free include: -NH(C=O)R, -NR(C=O)R,-OH and -OR other EDG which release electrons by hyperconjugative effect include: -CH3, -CH2R, -CHR2, -CR3 and –C6H5 (the phenyl ring) - in organic compounds, EDG makes the electron pair more available, leading to ENHANCED BASICITY - strong EWG are due to electronegative atoms/ group of atoms: -NO2, -CX3 (where X = F, Cl, Br, I) - moderate EWG include: -CN, -(C=O)OH, -(C=O)OR, -(C=O)H, -(C=O)R, -SO3H - weak EWG include the halides: -F, -Cl, -Br, -I - in organic compounds, EWG enhance both the positivity of H and the dispersal of the negative charge (upon removal of the protonic H atom) stabilizing the conjugate base thus leading to ENHANCED ACIDITY 3. Steric effects - the presence of bulky substituents may result in the crowding of other substituents or atoms that may lead to weakening of bonds leading to acceleration of bond cleavage - steric effect near an acidic group favors the removal of protons regardless of whether the groups present are EDG or EWG (i.e. ENHANCED ACIDITY) - steric effect near the lone pair of electrons makes them less accessible or available for donation, thus DECREASING BASICITY 4. Intra- and Intermolecular Forces of Attraction (IMF) - boiling and melting substances involve disrupting IMFs existing between its particles - on the other hand, dissolution involves breaking of IMFs between similar molecules AND forming of IMFs between dissimilar molecules (solute and solvent) - the boiling point (Tb) of an organic compound is affected by several factors:

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56 a. Tb increases with increasing molecular size (due to larger surface area) b. Tb increases with decreased branching (branching decreases molecular surface area) c. Tb increases with increasing polarity of the molecule (strong dipole-dipole IMFs) d. Tb increases with decreasing molecular symmetry e. Tb increases with less intramolecular H-bonding (this minimizes IMF with the solvent) - the solubility of an organic compound in water is affected by several factors: a. it increases with decreasing carbon chain length b. it increases with branching c. it increases with increasing capacity to form intermolecular H-bonding Functional Groups Organic compounds are characterized by the presence of a certain arrangement of atoms called a functional group. This is the site in a molecule where a chemical reaction is most likely to occur and where physical properties are also determined. Organic compounds may be classified into aliphatic and aromatic compounds. I. ALIPHATIC HYDROCARBONS A. ALKANES and CYCLOALKANES- contains only C–C single bonds besides the C-H bonds IUPAC Nomenclature (substituents) + root + -ane Linear: also represented as CH4 methane CH3CH3 ethane CH3CH2CH3 propane CH3CH2CH2CH3 butane CH3CH2CH2CH2CH3 pentane CH3CH2CH2CH2CH2CH3 hexane Physical Properties Alkanes are nonpolar molecules. The C-H bond can be considered a nonpolar bond due to the almost equal EN of H and C (2.1 and 2.5 respectively). Alkanes are in the liquid state due to van der Waal’s attractive forces. Branching decreases the surface area of alkanes and results to lesser van der Waal’s interaction and thus lowers the leading to lower boiling point. Alkanes because of their weak intermolecular forces (IMF) can only be dissolved by similar nonpolar solvents with weak IMF of attraction. They are thus insoluble in water which possess strong IMF of attraction between them molecules. Chemical Properties Alkanes do not react with bases nor with acids because it has no H atom attached to an EN atom like oxygen and does not have an electron pair for donation. But it can easily catch fire, by the combustion reaction process: CH4 + 2O2  CO2 + 2H2O

This involves the rupture of C-H bonds by homolysis to form free radicals. Halogenation is another similar process to give monohalogenated products: CH4 + Cl2  CH3Cl + HCl The order of reactivities is: tertiary (3O) > secondary (2O) > primary (1O) > methane CH3 H H H l l l l H3CC–H H3CC–H H3CC–H H3CC–H l l l l CH3 CH3 H H B. ALKENES – contain at least one carbon-carbon double bond IUPAC Nomenclature (substituents) + root + -ene Physical Properties Alkenes are relatively nonpolar molecules. However, in the presence of alkyl substituents like the methyl group which is an electron-donating group, the π-bond electrons are slightly pushed creating a slight polarization. Electronegative atoms like the halogens also affect the dipole moment, this time by withdrawing electrons from the π-bond. Alkenes possessing larger dipole moments have relatively higher boiling point than the lower on zero dipole moment isomer. This is due to larger dipole-dipole IMF that holds the molecules together. On the contrary, alkenes with lower or zero dipole moments have higher melting point than isomers with higher dipole moment. A structure with zero dipole moment reflects its higher symmetry and therefore better and tighter packing in the solid crystal lattice together with more efficient IMF. Alkenes are relatively soluble in nonpolar solvents and insoluble in polar solvents. Chemical Properties The reactions of alkenes are mainly centered on the π bond which is a site of high electron density. They can therefore be considered as nucleophiles (seeking a positively-charged center, symbolized as Nu-) and the attacking reagents as electrophiles (seeking a negativelycharged center, symbolized as E+). One of the most common reactions of alkenes is catalytic hydrogenation Ni/Pd/Pt CH2 = CH2 + H2  CH3CH3 Electrophilic addition on alkenes is governed by Markovnikov’s rule, which says that the hydrogen of an attacking E+ will add to the alkene carbon containing the greater number of H substituents. C. ALKYNES – contain at least one carbon-carbon triple bond IUPAC Nomenclature (substituents) + root + -yne Linear: also represented as HCCH ethyne = HCC-CH3 propyne HCC-CH2CH3 1-butyne H3C-CC-CH3 2-butyne

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57 Physical Properties Alkynes have similar boiling and melting points to those of corresponding alkenes. They are soluble in nonpolar solvents but more soluble in water than alkenes or alkanes. Chemical Properties Alkynes are reduced (i.e. lower bond order) via the following reactions. Terminal alkynes are also characterized by acidity of the H atom. II. AROMATIC HYDROCARBONS Aromatic hydrocarbons are cyclic conjugated hydrocarbons which are specially stabilized and react by substitution instead of addition. According to Hückel’s Rule, a compound is aromatic if it satisfies all of the following: (1) it is planar, (2) fully conjugated, (3) monocyclic and (4) has (4n+2) electrons where n = 0, 1, 2, 3, etc. Application of this rule to reveals that the following are examples of aromatic compounds: H + l :+

benzene

cyclopropenyl cyclopentadienyl tropylium cation cation cation IUPAC Nomenclature Common names of frequently occurring parent compounds such as benzene are also considered IUPAC names. For monosubstituted benzene derivatives: substituent + benzene. For disubstituted benzene derivatives: positions of substituents relative to parent functional group are indicated by numbers. When the benzene ring is a substituent instead, the prefix phenyl- is used to indicate its presence. For common names, the positions of substituents relative to the parent functional group are indicated by the prefixes ortho- (o-) for groups on adjacent C atoms on the ring, meta- (m-) for groups separated by one C atom and para- (p-) for groups on C atoms directly opposite each other. Physical Properties Benzene is not as polar as the alkenes, and is soluble in relatively nonpolar to moderately polar solvents like hexane, ether, chloroform and ethyl acetate. Chemical Properties Benzene reacts by electrophilic aromatic substitution to maintain the highly stabilized phenyl ring in the following general manner H E H+

+ E+

+

The orientation of further substitution reactions of substituted benzene derivatives are also affected by the groups already attached to the phenyl ring. All EDG from weak to strong are ortho- and para- directors while all

EWG except for the weakly deactivating halides are metadirectors. The halides though they withdraw by inductive effect can also denote electrons by resonance effect and are thus ortho- and para- directors. III. ALKYL HALIDES – halogenated derivatives of alkanes with the general formula CnH2n+1X. Alkyl halides are usually prepared by halogenations of alkanes, alkenes or alkynes. Physical Properties Alkyl halides are slightly more polar than alkanes due to the C – X bond. Alkyl halides have higher boiling point than corresponding alkanes due to larger molecular weight and stronger dipole-dipole interactions. In of density, the order is R – I > R – Br > H2O > R – Cl > R – F Alkyl halides like alkanes are insoluble in water, although some alkyl fluorides exhibit slight solubility due to the occurrence of H-bonding: R – F...... H – OH δ+ δδ+ δChemical Properties Alkyl halides can either undergo nucleophilic substitution or elimination reactions, Nu:- +R – X  R – Nu + X – H X l l –C–C– + :B>C=C< + HB: + X– l l H H IV. ALCOHOLS and PHENOLS – hydrocarbon compounds with at least one of its hydrogens replaced by a hydroxyl group, and phenyl rings with a hydroxyl substituent IUPAC Nomenclature (substituents)+parent chain+ -ol The common way to name alcohols consist of identifying the parent chain: alkyl + alcohol Primary (1o): CH3CH2CH2CH2OH Secondary (2o):

OH l

1-butanol (n-butyl alcohol) 2-butanol (2-butyl alcohol)

OH l Tertiary (3o):

3-methyl-1,4-pentadiol OH

OH l 3,4-dimethyl-3-hexanol OH l

2-ethyl-3-methyl-1-penten-3-ol

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58 Physical Properties Alcohols have relatively higher boiling points than corresponding hydrocarbons because they are involved in intermolecular H-bonding. Since H-bonding is also possible with water, low-molecular weight alcohols (i.e. with C: OH ratio below 5) are soluble in water. Alcohols with more than one hydroxyl group may be involved in intra- and intermolecular H-bonding. Chemical Properties Generally, alcohols are more acidic than terminal alkynes but less acidic than water, with the decreasing acidity going from 1o to 3o alcohols. Phenols are much more acidic than alcohols, and even more acidic than water because of the large stabilization of the phenolate anion. EWG on the ortho- and para- positions especially enhance its acidity. Most of the reactions of alcohols consist of their conversions into aldehydes (with a special reagent called pyridinium chlorochromate, PCC) or carboxylic acids (from 1o alcohols), or ketones (from 2o alcohols) with oxidizing agents such as chromic acid (H2CrO4), KMnO4 and K2Cr2O7 in acid. O ll RCH2OH RCH aldehydes

RCH2OH

R2CHOH

O ll RC – OH O ll RCR

carboxylic acids

ketones

V. ETHERS – contain two hydrocarbon groups attached to a central oxygen atom, R – O – R’ IUPAC Nomenclature (substituents)+ -oxy- + parent chain Ethers may be named as alkyl group1+alkyl group2+ ether. Physical Properties Ethers contain polar C – O bonds are therefore more polar than alkanes. Ethers have higher boiling points than corresponding alkanes and alkenes, but much lower than alcohols due to absence of H-bonding. They are, however, as soluble in water as alcohols of the same molecular weight. As the alkyl groups of the ether increase in size, the ether becomes relatively more nonpolar. Chemical Properties The oxygen of ethers has free electron pairs available for proton donors or Lewis acids. Protonation of this oxygen will weaken the C – O bond and result in the partial positivity of the  - C which will be a good site for nucleophilic attack. The -H atoms are also readily abstracted by free radicals. H l H+ l –C– O – C– l l Nu:-

VI. AMINES – nitrogen-containing organic compounds that can be considered as alkyl or aryl (aliphatic or aromatic) derivatives of ammonia, NH3. Since ammonia contains three hydrogen atoms, there are three classes of amines. RNH2: 1o amines, characterized by an –NH2 group attached to one alkyl/ aryl group R2NH: 2o amines, characterized by an –NH group attached to two alkyl/ aryl groups R3N: 3o amines, characterized by full substitution of all H atoms by alkyl/ aryl groups IUPAC Nomenclature (substituents) + root + suffix –amine Amines can also be commonly named by placing the prefix amino- before the name of the parent chain. Physical Properties Amines are moderately polar compounds due to their C – N bond. In addition to this, the possibility of H-bonding (for 1o and 2o amines only) causes amines to have higher boiling points than alkanes of similar molecular weight. They have, however, lower boiling points than corresponding alcohols (stronger H-bonds due to more EN O atoms). All classes of amines can form H-bonds with water such that all low-molecular weight amines are very soluble in water. Chemical Properties Amines can act as Lewis bases due to the lone pair on N. However, the trend in basicity depends on the medium. In the aqueous phase, 3o < 2o < 1o < NH3 because the bulky and nonpolar alkyl/aryl groups on substituted amines hinder access to the lone pair, besides decreasing solvation of the amine. However in the gas phase, the trend is reversed because the alkyl groups can now enhance the amine’s basicity by inductive effect (EDG). The three classes of amines can form salts with mineral acids: CH3CH2NH2 + HCl CH3CH2NH3Cl VII. CARBONYL COMPOUNDS – a family of compounds whose chemical properties are predominated by the reactivity of the carbonyl moiety, a C atom doubly bonded to an O atom. Carbonyl compounds are also characterized by the acidity of the -H (even though bonded to C) A. ALDEHYDES and KETONES -aldehydes are organic compounds that are characterized by the formula O ll R–C–H - ketones have the remaining hydrogen substituted with the general formula O ll R–C–R Physical Properties Aldehydes and ketones have lower boiling points than corresponding alcohols, since they do not participate in Hbonding with similar molecules. However, they have higher boiling points that corresponding ethers and alkanes because of the more polar carbonyl group.

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59 Aldehydes and ketones are as soluble as their alcohol counterparts since they can form H-bonds with water. The same solubility restriction (i.e. C:O ratio should be less than 5) holds. Chemical Properties Most of these compounds’ reactions involve simple nucleophilic addition to the carbonyl bond wherein the C – O bond is reduced to a single bond: Aldehydes and ketones also undergo oxidation with KMnO4, H2CrO4 and K2Cr2O7 into carboxylic acids: O O ll ll RCH RC – OH O ll C – OH

=O

C – OH ll O B. CARBOXYLIC ACIDS and DERIVATIVES – organic compounds containing a hydroxyl group attached to a carbonyl O ll R – C – OH IUPAC Nomenclature (substituents) + root + suffix –oic + word “acid” The carbonyl C is assigned as C1 of the parent chain. They may also be identified with common names such as formic acid, butyric acid and caproic acid (found in ants, butter and goats, respectively). Carboxylic acids can also be named as derived from acetic acid. Examples are phenyl acetic acid and trimethyl acetic acid: O O ll ll CH2C – OH C – OH Illustrations: CH3CH2COOH O ll C – OH

propanoic acid (also methyl acetic acid)

O ll CH2C – OH

2,2-dimethylpropanoic acid phenylethanoic acid

Long-chain carboxylic acids are called fatty acids. They usually contain an even number of C atoms along the parent chain which may or may not contain a double bond. Saturated: CH3(CH2)4COOH caproic acid CH3(CH2)6COOH caprylic acid CH3(CH2)8COOH capric acid CH3(CH2)10COOH lauric acid CH3(CH2)12COOH myristic acid CH3(CH2)14COOH palmitic acid CH3(CH2)16COOH stearic acid

Unsaturated: palmitoleic acid CH3(CH2)5CH = CH(CH2)7COOH oleic acid CH3(CH2)7CH = CH(CH2)7COOH linoleic acid CH3(CH2)4CH=CHCH2CH=CH(CH2)7COOH linolenicacid CH3(CH2)4CH=CHCH2CH=CHCH2CH=CH(CH2)7COOH Physical Properties Only carboxylic acids with five or fewer carbons are soluble in water. They are soluble because they can form H-bonds with water. As the alkyl portion of the acid increases, the nonpolar property and hydrophobicity of the compound predominate. Due to intermolecular H-bonding, carboxylic acids such as acetic acid are found as dimmers in the gas state and in nonpolar solvents. Chemical Properties Carboxylic acids react with bases to form the carboxylate salts, RCOOH+ NaOH RCOO-Na+ + H2O The presence of EWG on the R substituent enhances the acidity of carboxylic acids because of dispersion of electron density and consequent stabilization of the –COOgroup. On the other hand, substituents on the -C (branched acids) are usually less acidic than the unbranched acids since their – COO- groups are shielded from solvent molecules and cannot be stabilized by solvation as effectively. Carboxylic acids can be readily derivatized to give other functional groups: 1. Formation of ESTERS (esterification) - reaction with alcohols in the presence of an acid catalyst lower rates are O ll RC – OH +

R’ – OH

O ll RC – OR’

+

OH2

IUPAC Nomenclature (substituents) + parent chain + suffix –oate The chain to which the carbonyl group belongs is considered as the parent chain, with the carbonyl C as C1. The other group R’ attached to O is considered as the substituent, with the O- linked C as C1. O ethyl ethanoate (also called ethyl acetate) OCH2CH3 O methyl benzoate OMe Physical Properties of Carboxylic Acid Derivatives Esters, anhydrides and acyl chlorides do not form Hbonds with themselves unlike carboxylic acids and alcohols. They, therefore, have lower boiling points than carboxylic acids and alcohols for comparable molecular weight. Their boiling points are similar to those of aldehydes and ketones of corresponding molecular weight.

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60 They are soluble in water as long as they have 5 C or less because of H-bonding for the esters; and because of reaction with water (hydrolysis) leading to the formation of carboxylic acids for the acyl chlorides and anhydrides. Amides can form H-bonds with themselves and thus have high boiling points as well as melting points. The N,N-disubstituted amides, however, do not form H-bonds with themselves and therefore have lower boiling and melting points. Amides with 5 C or less are soluble in water due to formation of H-bonds between solute and solvent.

c. cellulose: insoluble CHO abundant in the tough outer wall of plant cells. In humans, cellulose does not serve as f 6. Carbohydrates are formed from CO2, H2O and chlorophyll, in the presence of sunlight, by the process of photosynthesis. chlorophyll : sunlight 6 CO2 + 6 H2O C6H12O6 + 6 O2 7. Hydrolysis of a polysaccharide monosaccharides that is made up of.

yields

the

Chemical Properties of Carboxylic Acid Derivatives The reactions undergone by these families of compounds are predominantly centered on the partially positive carbonyl carbon being attacked by nucleophiles. The relative reactivities of the carbonyl compounds is in the order O O O O O ll ll ll ll ll RC – Cl > RC – O – CR’ > RC – OR’ > RC – NH2

8. Carbohydrates undergo fermentation to produce alcohols. Sucrose and maltose will ferment with yeast but not lactose.

most reactive

10. Proteins are the main constituents of living cells. They are made up of amino acids (acids that contain the amine group as well as the carboxyl group) linked through the nitrogen atom to form very large molecules with molecular weights as high as 50 million.

least reactive

The relative reactivity depends on (1) relative electron density or withdrawing effect of a group –X towards the carbonyl C, (2) the relative ability of –X as a leaving group, and (3) both the electronic and steric effect of the R-group. Nucleophilic substitution takes place readily if the incoming group (Nu:-) is a stronger base than the leaving group (-X) or if the final product is resonance-stabilized BIOCHEMISTRY 1. Carbohydrates are polydroxy aldehydes and ketones (called aldose and ketose) or, compounds yielding these substances upon hydrolysis; they include sugars, starches and cellulose. 2. Carbohydrates can be monosaccharides, (containing one sugar unit) or polysaccharides (containing two or more units); those containing two units are often called disaccharides. M.S. + M.S. + energy D.S. (dehydration synthesis involving removal of water) hydrolysis D.S. M.S. + M.S. + energy Lactose glucose + galactose Maltose 2 glucose sucrose glucose (C6H12O6) + fructose (C6H12O6) 3. Glucose is one of the most common monosaccharides; it is the sugar found in the blood and hence is often referred to as blood sugar. (also known as dextrose or grape sugar) 4. Sucrose: (table sugar) most common disaccharides; it consists of a glucose unit and a fructose unit. 5. Starch (an important nutrient for man) is a polysaccharide made up of many glucose units. a. starch: storage carbohydrate found in plants b. glycogen: storage carbohydrate of animals

9. Reducing sugars are identified by becoming oxidized by Cu2+ (which is reduced and hence blue color disappears) in the FEHLING’S and Benedict’s tests. Maltose and lactose show reducing properties except sucrose. Polysaccharides: (-) tests for reducing sugars.

11. All the amino acids that have been isolated from proteins are alpha amino acids and hence have the general formula however, they can also have another NH2 O R – CH – C – OH; -NH2 or –COOH as part of the R group. Also, the R group can contain an –OH group S atoms, and it can be aliphatic or aromatic. 12. The essential amino acids are those required by man but cannot be synthesized by the body. 13. The amino acids in a protein are ed together through the nitrogen atom on one amino acid and the carboxyl group on another, to form a peptide linkage. Hence proteins are sometimes called polypeptides. 14. The kind and order of amino acids in a protein determine the primary structure, and the type of chain and hydrogen bonding within the molecule determines the secondary structure. The tertiary structure has to do with the overall shape of the molecule, and the quaternary structure refers to interaction between protein molecules. 15. Since proteins contain both an amine group (-NH2, which is basic) and the acid or carboxyl (-COOH) also form a type of double ion, called a Zwitterion. The pH value where it is neutral is called the isoelectric point. 16. Proteins undergo hydrolysis to yield the amino acids that the protein is composed of; they can be denatured by heat, acid, base, or alcohol. This breaks the hydrogen bonds and thus changes the structure and reactivity of the protein.

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61 17. Proteins can be simple (only contain amino acids) or conjugated protein (contain phosphoric acid, carbohydrates, or nucleic acids). They can be classified according to their function as enzymes, hormones, antibodies, structural (in skin, hair and so on), contractile (in skeletal muscles), or blood proteins. 18. Lipids are fatlike substances that are esters, acids or alcohols. Fats, oils and waxes are examples of lipids. 19. Fatty acids are long chain monocarboxylic acids that can be saturated or unsaturated. 20. Fats and oils are esters of glycerol and fatty acids; all R groups (from the acids) are the same in a simple glyceride, and they are different in a mixed glyceride. Fats have saturated R group and oils the R groups contain multiple bonds; an oil can be converted to a fat by hydrogenating it to form saturated R groups. 21. The number of double bonds in a fat or oil can be determined by the iodine number, which tells how many grams of iodine will react with 100 grams of fat or oil. The saponification number is used to determine the molecular weight of a fat; by measuring the number of milligrams of KOH that will react with one gram of fat, to form a soap (an alkali metal salt of a fatty acid). 22. Waxes are esters of high molecular weight monohydroxy alcohols. 23. Simple lipids are esters of fatty acids; compound lipids contain other groups (such as phosphoric acid, carbohydrates, and nitrogen compounds) along with the acid and alcohol; derived lipids are obtained by the hydrolysis of simple or compound lipids – for example, fatty acids and sterols (high molecular weight cyclic alcohols – such as cholesterol) are derived lipids. 24. Most enzymes are proteins and will undergo all the reactions of proteins. The enzymes in the body function best about 400C. Temperatures above or below body temperature will decrease the activity of enzymes. Each enzymes has a certain pH at which it can function best. An increase in the amount of enzyme will

increase the rate of reaction. An increase in the amount of substrate will increase the rate of reactions.

Compounds that increase the activity of an enzyme are called activators. Compounds that interfere with the activity of an enzyme are called inhibitors. Enzymes contain an “active site” that binds to the substrate to form an enzyme substrate complex. This complex yields the products and regenerates the enzyme. Many enzymes contain two parts – a protein and a nonprotein part. The protein part of an enzyme is called the

apoenzyme.

Some enzymes require the presence of a substance called a coenzyme before they can act effectively. Coenzymes frequently contain the B vitamins or compounds derived from the B vitamins. Under the older system of naming enzymes the substrate was not mentioned, the newer system indicates the substrate being acted on. The names of enzymes

under this system end in-ase. Enzymes can be classified as oxidoreductases (enzymes that catalyze oxidation – reduction reactions between two substrates), transferases (which catalyze the transfer of a functional group between substrates), hydrolases (which catalyze hydrolysis reactions), lyases (which catalyze the removal of groups from substrates by means other than hydrolysis), isomerases (which catalyze the interconversion of cis-trans isomers), and ligases (which catalyze the coupling of two compounds with the braking of pyrophosphate bonds). Some hydrolytic enzymes are found in the lysosomes of the cytoplasm. The cytoplasm also contains mitochondria. These structural and functional units contain most of the oxidative enzymes and are deeply involved in the electron transport system of oxidation reaction. The mitochondria also produce ATP, the cells’ chief source of energy. Abnormal plasma enzyme concentrations are of clinical significance in the diagnosis of certain diseases. Isozymes are enzymes with the same function but slightly different structural features. Allosteric enzymes are key metabolic enzymes whose activity can be changed by molecules other than the substrate. Zymogens are the precursors of enzymes. 25. Enzymes are proteins that act as catalysts for the biochemical reactions take place in the body. There is a specific one for each reaction. Without these, the biochemical reaction will not take place fast enough to sustain life. 26. The main constituent of genes and viruses in the body are nucleo-proteins, which consist of proteins and nucleic acids. These are two classes of nucleic acids – ribonucleic acids, RNA and deoxyribonucleic acids, DNA. These are responsible for the genetic code; transfer RNA’s (tRNA) select specific amino acids from a cell and transfer them to messenger RNA’s (mRNA), which transmit the genetic code from the DNA in protein synthesis. 27. Food taken into the body first undergoes digestion. which breaks up the complex molecules into simple ones, and then metabolism, in which the digested food undergoes chemical changes in the tissues. There are two types of metabolism – catabolism produces energy from the nutrients and anabolism either stores the nutrients or produces proteins and other substances from them. 28. Carbohydrates, which usually consumed as starches, are one of the main sources of energy in the body. By a series of processes they are converted to glucose and either used to produce energy (by forming carbon dioxide and water, plus energy, in the Krebs or citric acid cycle) or stored for later use. 29. In the metabolism of proteins, they are hydrolyzed to amino acids, which are then either used for tissue building or broken down to ammonia, carbon dioxide and water with the production of energy. The ammonia, which is toxic, is removed by converting it to urea in the liver. In adults there is a nitrogen balance between the amount

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62 taken in and amount excreted; however in children this is not the case since some is used in tissue building. 30. Fats and oils are oxidized to produce energy in the body, being an even richer source than carbohydrates. Fats can also be stored in the body; these not only serve as a source of reserve food but also cushion the body organs. 31. The food we eat consists of carbohydrates, proteins, fats, minerals, water, and vitamins (compounds that are essential to an organism but cannot be synthesized by the organism) 32. Nutrients can be classified as according to the following functions;  provision from energy  body-building materials  regulation of the body processes. The three main food types  carbohydrates, proteins and fats  proteins, calcium, and phosphorus are used  elements such as sodium, potassium, chlorine, magnesium, sulfur and iron are necessary, along with traces of copper, cobalt, iodine, zinc, molybdenum, and fluorine. 33. Drugs fall into two categories: pharmacodynamic agents, which alter body functions by acting on the nervous system and chemotherapeutic agents, which destroy organisms that are harmful to the body. Additional Notes in Proteins Colloidal Nature: Proteins form colloidal dispersions with water. Being colloidal, protein will thru filter paper but not membrane. Proteins present in the bloodstream cannot through the capillaries and should remain in the blood. Since proteins cannot through membranes, there should be no protein material present in the urine. The presence of protein in the urine indicates damage to the membranes in the kidneys-possibly nephritis. Denaturation of a protein refers to the unfolding and rearrangement of the secondary and tertiary structures of a protein without breaking the peptide bonds. A protein that is denatured loses it biological activity. Reversible denaturation occurs when the conditions for denaturation are mild, the protein can be restored to the original conformation by carefully reversing the conditions that caused their denaturation. Reagents that cause denaturation 1. Alcohol coagulates (precipitates) all types of proteins except prolamines. 70% alcohol is used as disinfectant because its ability to coagulate the protein present in bacteria. Alcohol denatures proteins by forming H-bonds that complete with the naturally occurring H-bonds in the protein. 2. Salts of heavy metals such as mercuric chloride or silver nitrate, precipitate proteins. The antidote for mercuric

chloride when these poisons are taken internally is egg white. The heavy metal will react with the egg white and precipitate out. (The egg white colloid has a charge opposite to that of the heavy metal ion and so attracts it). The precipitate thus formed must be removed from the stomach by an emetic or the stomach will digest the egg white and return the poisonous material to the system. 3. Heat. Gentle heating causes reversible denaturation of protein, whereas vigorous heating causes irreversible denaturation by disrupting several types of bonds. 4. Alkaloidal Reagents such as picric acid or tannic acid denature protein by disrupting the salt bridges and hydrogen bonds. Tannic acid has been used extensively in the treatment of burns. When the substance is applied to the burned area, it causes the protein to precipitate as a tough covering, thus reducing the amount of water loss from the area. 5. Radiation. Proteins in cancer cells (rapidly dividing cells) are more susceptible to radiation than those present in normal cells so x irradiation is used to destroy cancerous tissues. 6. pH. Changes in pH can disrupt H-bonds and salt bridges, causing irreversible denaturation. Proteins are coagulated by such strong acids as concentrated HCl, sulfuric acids and nitric acids. Casein is precipitated from milk as a curd when it comes in with the HCl of the stomach. Heller’s ring test is used to detect the presence of albumin in the urine. A layer of conc. nitric acid is carefully placed in a sample of urine in a test tube. If albumin is present, it will precipitate out as a white ring at the interface of the two liquids. If acid or base remains in with protein for long period of time, the peptide bonds will break. 7. Oxidizing and reducing agents. Oxidizing agents such as bleach and nitric acid and reducing agents such as sulfites and oxalates denature protein irreversibly to disrupting disulfide bonds. 8. Salting out. Most proteins are insoluble in saturated salt solutions and precipitate out unchanged. To separate a protein from a mixture of other substances, the mixture is placed in a saturated salt solution (such as NaCl, or Na2SO4. The protein precipitates out and is removed by filtration. The protein can then be purified from the remaining salt by the process of dialysis. Biomolecules A. Carbohydrates (L. “carbos” charcoal, “hydros” water)  Most abundant organic molecule in nature  Monomer is monosaccharide, polymer is polysaccharide  Majority of these compounds contain carbon, hydrogen and oxygen with a common ration of Cn(H2O)n “hydrate of carbon”  Functions include energy sources and structural elements  They are classified as monosaccharides,

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63

     

disaccharides, oligosaccharides, and polysaccharides according to the number of simple sugars they contain. monosaccharides (glucose, fructose, galactose, mannose) disaccharides (lactose, maltose, sucrose) oligosaccharides (3-10 monosaccharide units, found in ER, GB) polysaccharides (starch, glycogen, cellulose, chitin, pectin, gum, peptidoglycan) Monosaccharides are linked together by a process called dehydration synthesis/ condensation (removal of water). Polysaccharides are broken down to their component monosaccharides through hydrolysis.

B. Proteins  Monomer is amino acid, polymer is polypeptide  There are currently 20 naturally-occurring amino acids  Amino acids are linked together by peptide bonds  Protein Denaturation=disruption of 3° structure  Of all the molecules encountered in living systems, proteins have the most diverse functions. Function 1. catalysis

Examples Enzymes (amylase,

cellulose, papase/ papain) Collagen (connective tissue matrix)

2. structure

Fibroin (silk), spider web Keratin (hair, horns, feathers, hooves)

Actin and tubulin (involved 3. movement

4. defense

5. regulation

6. transport

7. storage

in cell movements in

Amoeba and sperm cells; found in flagella and cilia) Skin keratin (first line of protection) Fibrinogen and thrombin (blood clot) Antibodies (secreted by lymphocytes in response to infection) Spines of sea urchin Hormones (insulin &

glucagon, growth hormone) Hemoglobin (oxygen and waste material transport) Myoglobin (transport of oxygen to muscles) Glucose transporters Zein (corn), gliadin (wheat), plastids, ovalbumin (egg white), casein (milk)

There are four levels of protein structure:  Primary (1°)  Secondary (2°)  Tertiary (3°)  Quaternary (4°) C. Lipids (L. “lipos” fat)  Unlike other biological polymers, lipids are not defined by specific, repeating monomeric units. Rather, they are defined by their water-repelling property.  They contain 3 fatty acids and 1 glycerol.  They also contain non-polar hydrocarbon groups (-CH3CH2)  Functions include transport, storage, and insulation.  3 types: oils (saturated and unsaturated), waxes, phospholipids  Oil is liquid at room temperature, while wax is solid.  Phospholipids usually assume two forms in an aqueous environment: micelle and phospholipids bilayer.  Oil is usually plant-derived, while fat (wax) is animalderived.  Lipids have the most amount of energy content: Lipids>Carbohydrates> Proteins D. Nucleic Acids  There are two types: DNA and RNA.  Monomer is nucleotide  The sugar component is a pentose (Nucleotide is made up of three components: nitrogenous base, sugar (ribose, deoxyribose) and phosphate group.  The nitrogenous bases are adenine, guanine, cytosine, thymine, uracil.  Types of bond found in Nucleic Acids: H-bonds (between complementary nitrogenous bases); Phosphodiester bond (connects nucleotides); and N-glycosidic bond (connects bases and sugars).  Main function is involved in heredity, protein synthesis, and as energy carriers.

Sugar Bases Strand Helix

DNA Deoxyribose A, G, C, T Double-stranded Helical

RNA Ribose A, G, C, U Single-stranded Non-helical

Examples

B-DNA, A-DNA, Z-DNA

mRNA, tRNA, rRNA, hnRNA, snRNA

* words in italics are specific names of proteins

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64 PRIORITY TABLE OF FUNCTIONAL GROUPS OF ORGANIC CHEMISTRY

Function

Suffix (main function)

Prefix (secondary function)

Example

Acid

-oic acid

carboxyl-

CH3-COOH Ethanoic acid

Ester

-yl, -oate

Amide

-amide

carboxamido-

CH3-CH2-CONH2 Propanamide

Nitrile

-nitrile (cyanide)

cyano-

CH3-CN Ethanenitrile Methyl cyanide

Aldehyde

-al

oxo-

CH3-CH2-CHO Propanal

Ketone

-one

oxo-

CH3-CO-CH3 Propanone

R-OH

Alcohol

-ol

hydroxi-

R-NH2

Amine

-amine

amino-

R-O-R’

Ether

-y … yl ether

oxa-

C=C

Alkene

-ene

-CC-

Alkyne

-yne

R-NO2

Nitro

nitro-

R-X

Halide

fluoro-, chloro-, bromo-, iodo-

-R

Radical

yl-

Formula O II R-C-OH O II R-C-O-R’ O II R-C-NH2 R-CN O II R-C-H O II R-C-R’

CH3-COO-CH3 Methyl ethanoate

CH3-CH2OH Ethanol CH3-CH2-NH2 Ethylamine CH3-O-CH2-CH3 Ethylmethyl ether CH3-CH=CH2 Propene CH3-CHCH Propyne CH3-CH2-NO2 Nitroethane CH3-CH2Br Bromoethane CH3 I CH3-CH-CH3 Methylpropane

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