Chemistry Semi Final Day 1

Basic structure of atoms. Atomic number, mass number. Elements, compounds. Isotopes. Radioactivity. A vogadro' s number. Definition of mole. Electronic structure of atoms. Building up/AUFBAU. Quantum numbers. Atomic orbitals. Pauli exclusion principle. Hund's rule.

-Atoms are divisible into subatomic particles protons, neutrons and electrons. -atoms of element are identical; the atoms of different elements are different in some fundamental way or ways. -Chemical compounds are formed when different atoms combine with each other. A given compound always has the same relative numbers and types of atoms -Atomic number (Z): number of protons in the nucleus. -Mass number (A): sum of the numbers of protons and neutrons. A=Z+N , N is # of neutrons -Isotope: Atoms w/ nuclei w same atomic number but diff mass numbers. In nature: most of the elements r mixture of isotopes the ratio of the isotopes is constant -Radioactivity: Spontaneous emission of radiation by unstable nuclei. alpha(Helium, over Z 83), beta(e-, positron, Zray photon), gamma decay(gamma photon) -ATOMIC MASS: unit amu (1 is 1/12th mass of carbon 12 isotope//12.0107 amu), avg atomic mass is the avg of the masses of all naturally occurring isotopes -1 MOLE: base unit, amount of substanceamount of substance that contains as many formula units (molecules, atoms, etc.) as the number of atoms in 12 g of carbon-12... which is 6.022x10^23 atoms MOLAR MASS: g/mol QUANTUM THEORY: Each electron is described by four quantum numbers. Three of these specify the wave function that gives the probability of finding the electron at various points in space. n& l(ang)= shape&sizeenergy of electron Ml magnestic quantum= orientation of orbital.. not E Ms magnetic spin orientations of e- Atomic orbital: the region of space where there is high probability of finding the electrons. The Aufbau principle used to determine the e conf of an atom. The principle postulates an atom is "built up" by progressively adding electrons. they assume their most stable conditions (electron orbitals) with respect to the nucleus and those e already there. The electrons fill up the orbitals in increasing order of energy. Building-up order of subshells!. # of electrons that can occupy each orbital is limited by the Pauli exclusion principle(No two electrons in an atom can have the same four quantum numbers). If multiple orbitals of the same energy are available, Hund's rule (Every orbital in a subshell is occupied by a single electron (having their spin in the same direction) before any orbital is doubly occupied.) states that unoccupied orbitals will be filled before occupied orbitals are reused

classification of dienes and polyenes based on the positions of the double bonds. Addition reactions of conjugated dienes. Structure of isoprene. Mevalonic acid. Isoprenes: terpenes, carotinoids. Vitamin A (retinol) and retinal. Molecular mechanism of vision.

...

Hydroxyl group-containing organic compounds: alcohols, enols and phenols. Classification, nomenclature and chemical properties, reactions of alcohols. Some important alcohols.

Alcohol: -OH on sp3 hybridized C. Denoted an kjsdbhfksf-OL. ie methanol from methane. Use prefixes to denote # of OH groups, ie Diol means two OH, trill etc. If its a single branch use hydroxy. Alcohols are typically thought of as (alkyl group+ alcohol) ie ethyl alcohol (aka ethanol) is C2H5OH. Benzyl alcohol is C6H5CH2- 1) Alcohols are classified by the number of alkyl groups bonded to the carbon bearing the -OH group. • In a primary (1°) alcohol, the carbon which carries the -OH group is only attached to one alkyl group (primary carbon) "propan-1-ol" • In a secondary (2°) alcohol, the carbon with the -OH group attached is joined directly to two alkyl groups, which may be the same or different (secondary carbon) "propan-2-ol, butane-2-ol" In a tertiary (3°) alcohol, the carbon atom holding the -OH group is attached directly to three alkyl groups, which may be any combination of same or different groups (tertiary carbon). 2) Alcohols are also classified by the number of -OH groups bonded in the molecule. a) Alcohols: one -OH group b) Diols: two -OH groups c) Triols: three -OH groups d) Polyhydroxy compounds: many -OH groups Properties: alcohols have higher BPs than alkanes.. also with more C's in the alcohol BP increases. BC of intermolecular attractions!!!!!!!!!!!!! Solubility of alcohols in water - Small alcohols are completely soluble in water. - Solubility decreases as the length of the hydrocarbon chain in the alcohol increases. Acidities of Alcohols • Simple alcohols are about as acidic as water • Alkyl groups make an alcohol a weaker acid • The more easily the alkoxide ion is solvated by water the more its formation is energetically favored • Steric effects are important • Electron-withdrawing groups make an alcohol a stronger acid by stabilizing the conjugate base (alkoxide RXNS 1) Acid-base properties Alcohols, like water, can act as either weak acids or weak bases (amphoteric) They are not acids (they do not change the pH of water, pKa = 16-19), but they react with Na: CH3CH2-OH + Na CH3CH2-O- Na+ + 1⁄2 H2 sodium ethoxide an „alkoxide" 2) Combustion reactions of alcohols Alcohols can burn with oxygen to produce water, carbon dioxide and heat (just like hydrocarbons) CH3CH2OH + 3O2 → 2CO2 + 3H2O + Heat 3) Dehydration • An alcohol can lose a water molecule to form an alkene using an acid catalyst such as H2SO4 and heat (an "elimination reaction"). This is the reverse of the addition of H2O to an alkene. • When more than one alkene can be formed, Zaitsev's rule states that the more substituted alkene will be the major product. Order of reactivity = 3° > 2° > (1° > methyl) (In fact, this reaction only works with 3° and 2° alcohols.) 4) Formation of alkyl halides (SN) (Replacing the -OH group by a halogen) 5) Oxidation of alcohols Oxidation is a loss of electrons Reduction is a gain of electrons. In organic terms: Oxidation: loss of H2; addition of O or O2; addition of X2 (halogens). Reduction: addition of H2 or H-; loss of O or O2; loss of X2. (Neither an oxidation nor reduction: Addition or loss of H+, H2O, HX) 6) Formation of esters a) Esters are produced when carboxylic acids are heated with alcohols in the presence of an acid catalyst. Reactivity order of alcohols: CH3OH > primary > secondary > tertiary H+ The most commonly discussed ester is ethyl ethanoate (or ethyl acetate). In this case, the hydrogen in the -COOH group has been replaced by an ethyl group Other Important Alcohols: 1) Ethanol: drinks, meths (industrial methylated spirits), fuel, solvent for organic compounds 2) Methanol As a fuel Methanol again burns to form carbon dioxide and water. Can be used as a petrol additive to improve combustion • Asanindustrialfeedstock 3) propan-2-ol Propan-2-ol is widely used in an amazing number of different situations as a solvent 4) Benzyl alcohol Benzyl alcohol (or phenylmethanol) is used in perfumes, flavors, soaps, cosmetics 5) Cinnamic alcohol Cinnamic alcohol (or cinnamyl alcohol) is found in cinnamon leaves, usually in the form of an ester. 6) Cyclohexanol Cyclohexanol is a cyclic, secondary alcohol.

Alkaline earth metals: electronic structure, physical and chemical properties. Oxides, hydroxides. and salts of alkaline earth metals. The electronic structure and chemical properties of aluminum. Acid-base properties of aluminum oxide. Aluminum chloride.

Alkaline Earth: 2ve, Ns2 These elements are found in the second group of the periodic table Stable noble-gas like cations: Mg2+, Ca2+, Ba2+ strong reducing agents. easily excitable ve (flame test!!!) They are metallic Having an oxidation number of +2 makes these elements very reactive They are not found free in nature They are present in the earth's crust but not in their basic form. They have high boiling and melting points. They have low density, electron affinity, and electronegativity. They react easily with halogens and water. They are softer and stronger than other metals (except the alkali metals). The oxides and hydroxides are white ionic solids. Magnesium oxide (MgO, magnesia) Sticky, water-insoluble powder. Usage: in tooth pastes, in some sports as a grip- improving agent. Calcium oxide (CaO, burnt lime) Reacts vigorously with water: CaO + H2O →Ca(OH)2 Calcium hydroxide (Ca(OH)2 Strong base. Moderately water-soluble. Calcium carbonate (CaCO3, limestone, chalk) In rocks, bones, eggshell. Decomposes at elevated temperature: CaCO3 →CaO + CO2 Water-soluble barium compounds are poisonous. Barium Barium sulfate (BaSO4, barite) Totally insoluble in water. Used in radiology as an X-ray contrast medium Aluminum: ns2np1 Stable noble-gas like cation: Al3+ Soft, silvery, durable metal of low density. Aluminum oxide layer protects its metal. Good conductor Amphoteric: reacts with acids and bases. -Aluminum oxide (Al2O3) hard, water-insoluble substance. Amphoteric -Aluminum hydroxide (Al(OH)3) Insoluble in water, Amphoteric -Aluminum chloride (AlCl3) Covalent, electron-deficient compound Cl Lewis acid

Arrhenius concept of acids and bases. Bronsted-Lowry concept of acids and bases. Conjugate acid-base pairs. Amphoteric compounds (ampholytes). Molecular structure and acid strength. Lewis concept of acids and bases.

Arrhenius:dissolved in water, acid produces protons (H+ or H3O+ ions). base gives OH- ions Bronsted Lowry: base accepts a proton, acid donates a proton (H+) Conjugate acid-Base pairs:for example HCl and Cl- or H2O and H3O+. Conjugate acid: in a conjugate acid-base pair, the species that DONATES PROTON Amphoteric compound (ampholyte): a species that can act either as an acid or a base. H2O and H3O+ Conjugate base: in a conjugate acid-base pair, the species that can accept a pro Molecular structure and acid strength: Two factors involved are POLARITY of bond, more polar means stronger acid bc proton more easily removed. Next factor is bond strength det by ATOMIC SIZE. BIGGER ATOM= weaker bond= stronger acid. Lewis Concept of acids and bases: Acid forms a covalent bond via accept electron pair and form a covalent bond. Base can form covalent bond by giving away e- pair... so formation of complex ions is a lewis acid//base reaction (lewis acid-base adduct) think of reactions of electrophile is always the base, acid attacks electrophilic base.

Molecular substances: chemical bonding. Describing ionic, covalent and metallic bonding. Formation of ions, ionization energy and electron affinity. Formation of an ionic solid. Metallic bonding.

CHEMICAL BONDING: goal octet. Atoms form bonds because the compound that results is more stable than the separate atoms IONIC: electrostatic attraction between + and - ions. term "ionic bonding" is given when the ionic character is greater than the covalent character - a bond in which a large electronegativity difference exists between the two atoms, causing the bonding to be more polar (ionic) than in covalent bonding where electrons are shared more equally. Bonds with partially ionic and partially covalent character are called polar covalent bonds. COVALENT: A chemical bond in which a pair of electrons is shared between two atoms. single sigma, double sigma pi, triple sigma 2 pi METALLIC: positively charged metal ions (cores of the atoms) and delocalized valence electrons (electron cloud or sea) Metallic bonding: the electrostatic attraction between them. IONIZATION E The first ionization energy is the energy needed to remove the highest- energy electron from the neutral atom in the gaseous state. highest top right. and ng Unit: J/mol E AFFINITY: The energy change for the process of adding an electron to a neutral atom in the gaseous state. Unit: J/mol. Electron affinity is either + or −, highest top right but low ng's

Describing covalent bonds. Molecular orbital theory. Single and multiple bonds. Bonding energies. Resonance description, delocalized bonding.

COVALENT: A chemical bond in which a pair of electrons is shared between two atoms. single sigma, double sigma pi, triple sigma 2 pi. small electronegativity difference. EACH COVALENT BOND HAS 2 e- Delocalized/Reasonance: Some molecules are have structures that cannot be shown with a single representation. Electrons of a bonding pair are spread over a number of atoms rather than localized between two.actual electron distribution is a composite of the resonance formulas.describing delocalized electrons within certain molecules or polyatomic ions where the bonding cannot be expressed by one single Lewis formula. A molecule or ion with such delocalized electrons is represented by several contributing structures Bonding energy: The energy needed to break a bond in the gaseous state. Reaction H2 → 2 H· requires 436 kJ/mol Reaction 2 H· → H2 releases 436 kJ/mol H2 has 436 kJ/mol less energy than two H atoms: H-H has bond strength of 436 kJ/mol. MO THEORY: The Molecular Orbital Theory, initially developed by Robert S. Mullikan, incorporates the wave like characteristics of electrons in describing bonding behavior. In Molecular Orbital Theory, the bonding between atoms is described as a combination of their atomic orbitals. While the Valence Bond Theory and Lewis Structures sufficiently explain simple models, the Molecular Orbital Theory provides answers to more complex questions. In the Molecular Orbital Theory, the electrons are delocalized. Electrons are considered delocalized when they are not assigned to a particular atom or bond (as in the case with Lewis Structures). Instead, the electrons are "smeared out" across the molecule. The Molecular Orbital Theory allows one to predict the distribution of electrons in a molecule which in turn can help predict molecular properties such as shape, magnetism, and Bond Order. ANTIBONDING BONDING ORBITALSEach pair of atomic orbitals interacts to form a set of bonding and antibonding molecular orbitals. orbital number conserved

Structure and nomenclature of cycloalkanes (cycloparaffins). The role of the number of carbon atoms of the ring in the stability. Conformations of cyclohexane: boat and chair. Axial and equatorial substituents. Isomerism of disubstituted cyclohexanes. Bi- and polycyclic alkanes: cis- and trans-decalin. Chemical properties of cycloalkanes.

CYCLOALKANES: cyclic single bonded Hydrocarbons Cycloalkanes have boiling points which are about 10 - 20 K higher than the corresponding straight chain alkane. Cycloalkanes are similar to alkanes in their general physical properties, but they have higher boiling points, melting points, and densities than alkanes. This is due to stronger London forces because the ring shape allows for a larger area of contact. Containing only C-C and C-H bonds, unreactivity of cycloalkanes with little or no ring strain (see below) are comparable to non-cyclic alkanes. Ring strain highest for cyclopropane, in which the carbon atoms form a triangle and therefore have 60 degree C-C-C bond angles. There are also three pairs of eclipsed hydrogens. Cyclobutane has the carbon atoms in a puckered square with approximately 90-degree bond angles; "puckering" reduces the eclipsing interactions between hydrogen atoms. Its ring strain is therefore slightly less, at around 110 kJ/mol. For a theoretical planar cyclopentane the C-C-C bond angles would be 108 degrees, very close to the measure of the tetrahedral angle. Actual cyclopentane molecules are puckered, but this changes only the bond angles slightly so that angle strain is relatively small. The eclipsing interactions are also reduced In cyclohexane the ring strain and eclipsing interactions are negligible because the puckering of the ring allows ideal tetrahedral bond angles to be achieved. As well, in the most stable chair form of cyclohexane, axial hydrogens on adjacent carbon atoms are pointed in opposite directions, virtually eliminating eclipsing strain. Explanations There isn't much EN difference between C and H, so little bond polarity. The molecules themselves also have very little polarity. A totally symmetrical molecule like methane is completely non-polar. This means that the only attractions between one molecule and its neighbours will be Van der Waals dispersion forces. These will be very small for a molecule like methane, but will increase as the molecules get bigger. That's why the boiling points of the alkanes increase with molecular size. THINK ...Where you have isomers, the more branched the chain, the lower the boiling point tends to be. Van der Waals dispersion forces are smaller for shorter molecules, and only operate over very short distances between one molecule and its neighbours. It is more difficult for short fat molecules (with lots of branching) to lie as close together as long thin ones. For example, the boiling points of the three isomers of C5H12 are: boiling point (K) pentane 309.2 2-methylbutane 301.0 2,2-dimethylpropane 282.6 The slightly higher boiling points for the cycloalkanes are presumably because the molecules can get closer together because the ring structure makes them tidier and less "wriggly"! BOAT & CHAIR(cyclohexane): six of the 12 carbon-hydrogen bonds end up almost perpendicular to the mean plane and almost parallel to the symmetry axis, with alternating directions, and are said to be axial. The other six C-H bonds lie almost parallel to the mean plane, and are said to be equatorial. At room temperature there is a rapid equilibrium between the two chair conformations of cyclohexane Two chairs are lowest E. The conformations involve following order of stability: chair form > twist boat form > boat form > half-chair form. The boat conformation (4, below) is also a transition state, allowing the interconversion between two different twist-boat conformations. Observations: -toxic but w/ added methyl (TOULENE) it becomes nontoxic (water soluble) - reacts slowly with Br2 to form bromobenzene: substitution product. - Addition products are not observed. C-C bond length: double shorter duh -Reasonance hybrid! How to tell if Aromatic:(HUCKNEL'S RULE): A planar cyclic molecule with alternating double and single bonds is aromatic if it has (4n+2) π electrons (n is a small integer) Benzene: n=1 ⇒ (4n+2)=6 electrons Benzene is stable and the electrons are delocalized.. ie benz 3 double bonds, six pi e-... if n=2, then 10 pi e- ---if there are 4n pi e-, compound is antiaromatic and behave likeordinary alkenes Nomenclature: alphabetically important, overall lowest numbers *if bonds are: 1,2: ortho (o). 1,3: meta(m). 1,4 para(p)... or you can just say like 1,4-dibromobenzene not p-dibromobenzene. All the 1,2-dichloro isomers are constitutional isomers of the 1,3-dichloro isomer Bicyclic Cyclohexanes: Fused ring systems that share more than two atoms are called bicyclic molecules. All the 1,2-dichloro isomers are constitutional isomers of the 1,3-dichloro isomers. In each category (1,2- & 1,3-), the (R,R)-trans isomer and the (S,S)-trans isomer are enantiomers. The cis isomer is a diastereomer of the trans isomers. Finally, all of these isomers may exist as a mixture of two (or more) conformational isomers, as shown in the table. The chair conformer of the cis 1,2-dichloro isomer is chiral. It exists as a 50:50 mixture of enantiomeric conformations, which interconvert so rapidly they cannot be resolved (ie. separated). Since the cis isomer has two centers of chirality (asymmetric carbons) and is optically inactive, it is a meso-compound. The corresponding trans isomers also exist as rapidly interconverting chiral conformations. The diequatorial conformer predominates in each case, the (R,R) conformations being mirror images of the (S,S) conformations. All these conformations are diastereomeric with the cis conformations. The diequatorial chair conformer of the cis 1,3-dichloro isomer is achiral. It is the major component of a fast equilibrium with the diaxial conformer, which is also achiral. This isomer is also a meso compound. The corresponding trans isomers also undergo a rapid conformational interconversion. For these isomers, however, this interconversion produces an identical conformer, so each enantiomer (R,R) and (S,S) has predominately a single chiral conformation. These enantiomeric conformations are diastereomeric with the cis conformations. The 1,4-dichlorocyclohexanes may exist as cis or trans stereoisomers. Both are achiral, since the disubstituted six-membered ring has a plane of symmetry. These isomers are diastereomers of each other, and are constitutional isomers of the 1,2- and 1,3- isomers. To name bicyclic alkanes, you follow these three steps: Count the total number of carbons in the entire molecule. This is the parent name (eg. ten carbons in the system would be decane) Count the number of carbons between the bridgeheads, then place in brackets in descending order. (eg. [2,2,1]) Place the word bicyclo at the beginning of the name.

Chemical equilibrium, equilibrium constant and the law of mass action. Le Chatelier's principle. Changing the reaction conditions (pressure, temperature and concentration). Application of Le Chatelier's principle.

Chemical reactions in general do not go to completion. Can be made to go in either one direction or to the other. When forward and reverse reactions occur at the same rate, so there is no net change, equilibrium is reached. Defined: rates of forward and reverse reactions equal, net change no longer occurs. Equillib constant: product cxns root molar coefficient over product cans root molar coefficients. at equillib cxns of abcd are constant. Kc is eqillib constant for a given temperature. Kp for gases Law of mass action: states that the values of Kc are constant for a particular reaction at a given temperature Le Chat& applying it:Henri le Chat, 1884If a reaction in equilibrium is disturbed, the composition of the system will change until that new equilibrium state is attained. The new equilibrium state is one that reduces the effect of the change. If a system in equilibrium is altered by the change of some conditions, chemical reaction occurs to shift the equilibrium composition in a way that attempts to reduce that change of conditions ie add reactant/ remove product, shift right Reaction Condition Change: Three ways to alter the equilibrium composition: 1. Changing the concentrations of reactants or products. 2. Changing the temperature:always an exo/endo forward reverse visav. If the forward reaction is exothermic, the reverse reaction is endothermic and vica versa. The effect of an increase in temperature can be reduced in the way that the endothermic reaction is taking place. The effect of a decrease in temperature can be reduced in the way that the exothermic reaction is taking place.3. Changing the pressure: A pressure change alters an equilibrium if the reaction involves a change of total moles of gases. The pressure can be changed by changing the volume of the vessel in which the equilibrium is taking place. equal moles on both sides= no change from altering pressure equlib moves to side with less moles in P increased!!! bc partial pressure proportional to # moles of gases as fewer moles of gases have smaller partial pressure. decrease in p shifts to less moles. A catalyst cannot alter the equilibrium.

chemical reactions of aromatic compounds. Electrophilic substitutions: halogenation, sulfonation, nitration, Friedel-Crafts alkylation and acylation. Mechanism of electrophilic aromatic substitution. Effect of substituents on aromatic substitution.

Chemical reactions of aromatic hydrocarbons: aromatic electrophilic substitutions: Halogenation, Nitration Sulfonation Friedel-Crafts alkylation, Friedel-Crafts acylation aromatic HC +E(+)-N(-)-> E shift out H and HN forms separately Electrophilic substitutions: halogenation, forms anti mark prod with FeBr3 catalyst sulfonation: Heat catalyst, SO3H replaces H, H2O byproduct nitration: NO2 replaces H, H2o also produced (SA catalyst) Friedel-Crafts alkylation: Alkyl Benzene formed (Benzene w/ R group... HCl byproduct.. R-Cl adds to benzene w/ AlCl3 catalyst) and acylation: R-C=O added to benzene with HCl byproduct, AlCl3 cata Mechanism of electrophilic aromatic substitution. Effect of substituents on aromatic substitution. Driving force is reconstruction of aromatic system. Activating groups promote substitution (faster than benzene) OH, NH2, alkyl, direct substitution to o- and p-positions.direct substitution to o- and p-...Donate electrons to the ring Deactivating groups promote substitution... slower than benzene..NO2, COOH, CN, SO3H positions. The others direct substitution to meta position.... withdraw e- from ring Halogens Deactivate ring Substitution is slower than for benzene Groups direct substitution to o/p (ortho/para) position

Colligative properties of solutions. Vapor pressure of a solution. Boiling-point elevation and freezing-point depression. Determination of molecular weights based on the colligative properties. Background of chromatography.

Colligative Properties depend on #particles in solution NOT nature of particles (physical properties of solutions) Vant Hoffs factor: NaCl, i = 2, CaCal2 is 3 1. VP Lowering Non-volatile solutes reduce the ability of the surface solvent molecules to escape the liquid. Therefore, vapor pressure is lowered. The amount of vapor pressure lowering depends on the amount of solute. 2. Freezing pt depression Addition of a solute lowers the freezing point of a liquid and shifts phase diagram as nonvolatile solute decreases FP of a liquid. Solute molecules make it more difficult for the solvent molecules to come together and freeze. Lower temperature is needed to freeze the solution. 3. BP elevation Addition of a non volatile solute increases the boiling point of a liquid. Solute molecules make it more difficult for the solvent molecules to vaporize. Higher temperature is needed to boil the solvent. 4 Osmotic pressure: Osmosis is the selective passage of solvent molecules through a semipermeable membrane from a dilute solution to a more concentrated one. Osmotic pressure (π) is the pressure required to stop osmosis. solvent moves both ways, - solute molecules are too large to pass through membrane, and - net movement is to try to equalize the solute concentrations on both sides of the membrane. Molecular weights: Molar mass of a nonelectrolyte can be determined from ∆Tf and ∆Tb. (Tb is the boiling point of the solution) m= moles solute/ mass of solvent where m and kg are known moles solute= mass solution/ mm of solute Chroma Bckgrnd: Partition: separation of a substance between two immiscible solvents (A and B) based on its different solubility in the solvents. partition coeff gives cxn is certiain solutions, comparing solubility. is [I2] chloroform/ [I2] h2o= 130 means iodine is 130* more soluble in chl. this means we can use EXTRACTION CHROMATOGRAPHY=way to separate mixtures. The separation is based on differential partitioning of a substance between the mobile and stationary phases. THREE types: 1.Based on the physical state of the mobile phase: liquid and gas chroma. 2. Based on the bed shape: Column chromatography, Thin-layer chromatography, and Planar chromatography 3. Based on the separation mechanism: Ion-exchange chromatography and Size-exclusion chromatography

Electrodes and voltaic cells. Types of electrodes. Electrode potential, standard potential. Electromotive force. Electrical work and free energy change. Nernst equation. Redox reactions in the living states.

Components of an electrochemical cell that dip into an electrolyte. They are generally solid metals connected with a wire. A voltaic cell is an electrochemical cell in which a spontaneous reaction generates an electric current. An electrolytic cell is an electrochemical cell in which an electric current drives an otherwise nonspontaneous reaction. A REDOX RXN OCCURS. Salt bridge connects two half runs. Types of electrodes includes a metal electrode, gas electrode, metal/ppt electrode, and redox electrode. A standard electrode is an electrode in which the molarities of ions and the pressure of gas (in atmospheres) are equal to 1 at 25 °C. The Standard H electrode is our reference. The potential of one electrode can not be measured directly. Only the potential difference between two electrodes can be measured. We need a reference value to which the potentials of other electrodes can be compared. Reference: standard hydrogen electrode (SHE) Standard Reduction Pot: Standard reduction potential is the voltage associated with a reduction reaction at an electrode when all solutes are 1 M and all gases are at 1 atm and the temperature is 298 K. Positive means its easily reduced. Negative means its easily oxidized. Electromotive force (emf) is the maximum potential difference between the electrodes in a voltaic cel. driving force that propels electrons from the anode to the cathode. It is positive in voltaic cells. Ecell=Ec-Ean Where the anode has a lower E(standard electrode pot) Nernst equation used to calculate the emf of a voltaic cell under nonstandard conditions. ie molarities not equal and also electrode potential of an electrode under nonstandard conditions. Some examples include conversion of NAD+ to NADH. NADP+ to NADPH. Acetate to acetaldehyde. Concentration Cell: The cell has two identical electrodes but different electrolyte concentrations. The electromotive force is determined by the concentration difference.

The acidity of phenols. Influencing effects of substituents on the acidity. Nomenclature of phenols. Chemical reactions. Derivatives of phenols. Oxidation of phenols. Some important phenols and phenol-derivatives. Esters formed from alcohols and inorganic acids: nitrites, nitrates, sulfates and phosphates.

Compounds like alcohols and phenol which contain an -OH group attached to a hydrocarbon are very weak acids. Alcohols are so weakly acidic that, for normal lab purposes, their acidity can be virtually ignored. Physical Properties: The polar nature of the O-H bond (due to the electronegativity difference of the atoms ) results in the formation of hydrogen bonds with other phenol molecules or other H-bonding systems (e.g. water). The implications of this are: high melting and boiling points compared to analogous arenes high solubility in aqueous media The presence of intramolecular hydrogen bonding is believed responsible for the significantly lower boiling points of certain ortho-substituted phenols vs the meta- and para- analogs. However, phenol is sufficiently acidic for it to have recognisably acidic properties - even if it is still a very weak acid. A hydrogen ion can break away from the -OH group and transfer to a base With sodium hydroxide solution Phenol reacts with sodium hydroxide solution to give a colourless solution containing sodium phenoxide. With metallic sodium Acids react with the more reactive metals to give hydrogen gas. Phenol is no exception - the only difference is the slow reaction because phenol is such a weak acid. You can recognise phenol because: It is fairly insoluble in water. It reacts with sodium hydroxide solution to give a colourless solution (and therefore must be acidic). It doesn't react with sodium carbonate or hydrogencarbonate solutions (and so must be only very weakly acidic). Substituents, particularly those located ortho or para to the -OH group, can dramatically influence the acidity of the phenol due to resonance and / or inductive effects. Electron withdrawing groups enhance the acidity, electron donating substituents decrease the acidity. The resonance stabilization of o-nitrophenol is shown below:

Nomenclature and preparation of ethers, cyclic ethers and phenolethers. Chemical properties of ethers. Preparation, physical and chemical properties of thiols (mercaptans), oxidation reactions. Thioethers. Sulfoxides, sulfones, sulfonic acids.

ETHERS can be thought of as analogs to water -An ether has two organic groups (alkyl, aryl, or vinyl) bonded to the same oxygen atom: R-O-R′ -Diethyl ether is used industrially as a solvent -Tetrahydrofuran (THF) is a solvent that is a cyclic ether -Any ether is a constitutional isomer of an alcohol containing the same number of carbon atoms. Naming Ethers • Simple ethers are named by identifying the two organic substituents and adding the word ether • If other functional groups are present, the ether part is considered an alkoxy substituent Naming Cyclic Ethers Cyclic ethers are generally named by their common names (we will not study the IUPAC names) • A cyclic ether containing two carbons is called ethylene oxide (generally known as epoxides) • A cyclic ether containing 4 carbons (with 2 double bonds) is called a furan • A cyclic ether containing 5 carbons (with 2 double bonds) is called a pyran • A cyclic ether containing 4 carbons and 2 oxygens is called a dioxane Physical Properties of Ethers • They have a bent geometry around the O and are polar compounds • Ethers are only slightly soluble in water and are highly flammable • Ethers have considerably lower boiling points than those of alcohols Hydrogen bonding • No hydrogen bonds between the molecules in the ether: the boiling point is much lower than the corresponding alcohol - the ether is significantly less dense than the corresponding alcohol. • Ethers can act as a hydrogen-bond acceptor, but, they can't act as hydrogen-bond donors: ethers are less likely to be soluble in water than alcohols with the same molecular weight (eg. constitutional isomers) Ethers resemble alkanes more than alcohols (with respect to boiling point) • Intermolecular hydrogen bonding possible in alcohols; not possible in alkanes or ethers. Ethers resemble alcohols more than alkanes (with respect to solubility in water) - Hydrogen bonding to water possible for ethers and alcohols; not possible for alkanes Preparation of ethers 1) From two alcohols in the presence of heat and concentrated sulfuric acid: 2CH3CH2OH --H+--> CH3CH2OCH2CH3 Diethyl ether is prepared industrially by sulfuric acid - catalyzed dehydration of ethanol - also with other primary alcohols Secondary and tertiary alcohols do not form ethers when heated with an acid catalyst. The attached substituents prevent two molecules from approaching each other closely enough from intermolecular dehydration to occur. This effect is called STERIC HINDRANCE 2. Williamson Synthesis (SN2) is an important laboratory method for the preparation of symmetrical and unsymmetrical ethers Ethers containing substituted alkyl groups (secondary or tertiary) may also be prepared by this method. The reaction involves SN2 attack 3) By reacting a primary alkyl halide with an alkoxide ion CH3CH2CH2Br + CH3O- → CH3CH2CH2OCH3+ Br- 4) Silver oxide-catalyzed ether formation: Reaction of alcohols with Ag2O directly with alkyl halide forms ether in one step. Glucose reacts with excess iodomethane in the presence of Ag2O to generate a pentaether in 85% yield Reactions of Ethers • Ethers are unreactive towards acids, bases, and oxidizing agents. • Ethers, like alkanes, participate in combustion and halogenation reactions ONLY. 1) Acidic Cleavage Strong acids cleave an ether at elevated temperature. HI, HBr produce an alkyl halide from less hindered component by SN2 (tertiary ethers undergo SN1) 2) Combustion of ethers • Ethers can burn with oxygen to produce water, carbon dioxide and heat (just like hydrocarbons) CH3-O-CH3+ 3O2→2CO2+ 3H2O + Heat • Ethers are much more flammable than alcohols. Care should be taken when working with ethers in the laboratory (just a spark from static electricity can set off ether fumes) Use of ethers • Ethers are often used as solvents to carry out chemical reactions. • Diethyl ether, often known by the generic name "ether," was once used extensively as an anesthetic. (Because mixtures of diethyl ether and air explode in the presence of a spark, ether has been replaced by safer anesthetics.) Autooxidation - You should know about this if you ever work in the lab with ethers: O2 causes the formation of peroxides Cyclic Ethers • Cyclic ethers behave like acyclic ethers, except if ring is 3-membered (epoxides) • Dioxane and tetrahydrofuran are used as solvents Epoxides (Oxiranes) • Three membered ring ether is called an oxirane (root "ir" from "tri" for 3-membered; prefix "ox" for oxygen; "ane" for saturated) • Also called epoxides • Ethylene oxide (oxirane; 1,2-epoxyethane) is industrially important as an intermediate • Prepared by reaction of ethylene with oxygen at 300 °C and silver oxide catalyst Crown Ethers • Large rings consisting repeating -OCH2CH2- or similar units • Named as x-crown y-ether -x is the total number of atoms in the ring -y is the number of oxygen atoms - 18-crown-6 ether: 18-membered ring containing 6 oxygens atoms • Central cavity is electronegative and attracts cations • Complexes between crown ethers and ionic salts are soluble in nonpolar organic solvents • Creates reagents that are free of water that have useful properties • Inorganic salts dissolve in organic solvents leaving the anion unassociated, enhancing reactivity THIOLS (RSH) Nomenclature There are several ways to name the alkylthiols: • The preferred method (used by the IUPAC) is to add the suffix-thiol to the name of the alkane. (The method is nearly identical to naming an alcohol.) Example: CH3SH methanethiol • An older method: the word mercaptan replaces alcohol in the name of the equivalent alcohol compound. Example: CH3SH methyl mercaptan (just as CH3OH is called methyl alcohol). • As a prefix, the terms sulfanyl or mercapto are used. Example: CH3CH2SH mercaptoethane • The hydrogen bonding between thiol groups is weak, the main cohesive force is van der Waals interaction between the highly polarizable divalent sulfur centers. • Due to the lesser electronegativity difference between sulfur and hydrogen compared to oxygen and hydrogen, an S-H bond is less polar than the hydroxyl group. • Thiols have a lower dipole moment relative to the corresponding alcohol. Boiling points and solubility • Thiols show little association by hydrogen bonding with both water molecules and among themselves. They have lower boiling points and are less soluble in water and other polar solvents than alcohols of similar molecular weight. • Thiols and thioethers have similar solubility characteristics and boiling points. Odor • The odors of thiols are often strong and repulsive, particularly for those of low molecular weight. • Many thiols have strong odors resembling that of garlic. • The spray of skunks consists mainly of low-molecular-weight thiols and derivatives. • Human sweat also contains thiols. **Natural gas distributors add thiols (originally ethanethiol) to natural gas, which is naturally odorless, when processed. Most gas odorants utilized currently contain mixtures of mercaptans and sulfides with t-butyl mercaptan as the main odor constituent.** Formation of Thiols 1) From alkyl halides by displacement with a sulfur nucleophile such as −SH (The alkylthiol product can undergo further reaction with the alkyl halide to give a symmetrical sulfide) 2) In industry, methanethiol is prepared by the reaction of hydrogen sulfide with methanol in the presence of acidic catalysts. CH3OH + H2S → CH3SH + H2O 3) The other principal route to thiols involves the addition of hydrogen sulfide to alkenes. Such reactions are usually conducted in the presence of an acid catalyst or UV light. 4) Halide displacement, using the suitable organic halide and sodium hydrogen sulfide has also been utilized. 5) Laboratory methods - direct reaction of alkyl halides with sodium hydrosulfide is generally inefficient owing to the competing formation of thioethers: CH3CH2Br + NaSH → CH3CH2SH + NaBr CH3CH2Br + CH3CH2SH → (CH3CH2)2S + HBr Chemical properties • Thiols are weak acids but much stronger than alcohols, because the S‐H bond is weaker than the O‐H bond, (Sulfur is much larger than oxygen) However, thiols are weaker acids than phenols. • Thiols are easily oxidized to disulfides by many oxidizing agents: 1) Thiolates (RS−) are formed by the reaction of a thiol with a base • Thiolates react with primary or secondary alkyl halide to give sulfides (R-S-R') • Thiolates are excellent nucleophiles and react with many electrophiles 2) S-alkylation: Thiols are readily alkylated to give thioethers: RSH + R'Br + (base) → RSR' + HBr 3) Oxidation of thiols to disulfides a) Reaction of analkylthiol(RSH)with bromine or iodine gives a disulfide: (R-S-S-R) The thiol is oxidized in the process and the halogen is reduced Disulfides are named by naming the R groups attached to the sulfur atoms followed by the word disulfide. b) Oxidation by more powerful reagents such as hydrogen peroxide yields sulfonic acids (RSO2OH). R-SH + 3H2O2 → RSO2OH + 3H2O sulfonic acid 4) Acidity The proton of the thiol is much more acidic than the hydroxylic proton of alcohols. E.g. Butanethiol has a pKa of 10.5 vs 15 for butanol. Thiophenol has a pKa of 6 vs 10 for phenol. The nucleophilic thiol(ate) is much more reactive than the hydroxyl of alcohols BIOLOGICAL IMPORTANCE 1) Cysteine and cystine: the thiol functional group plays a very important role. When the thiol groups of two cysteine residues are brought near each other in the course of protein folding, an oxidation reaction can generate a cystein unit with a disulfide bond (-S-S-). Disulfide bonds can contribute to a protein's tertiary structure if the cysteines are part of the same peptide chain, or contribute to the quaternery structure of multi-unit proteins by forming fairly strong covalent bonds between different peptide chains. Sulfhydryl groups in the active site of an enzyme can form noncovalent bonds with the enzyme's substrate, contributing to catalytic activity. Active site cysteine residues are the functional unit in cystein protease. 2) Cofactors Many cofactors (non-protein-based helper molecules) feature thiols. (The biosynthesis and degradation of fatty acids and related long-chain hydrocarbons is conducted on a scaffold that anchors the growing chain through a thioester derived from the thiol Coenzyme A). The biosynthesis of methane, the principal hydrocarbon on earth, arises from the reaction mediated by Coenzyme M (2- mercaptoethyl sulfonic acid). 3) ) Glutathione The redox agent that mediates the formation and degradation of disulfide bridges in most proteins is glutathione, a versatile coenzyme. In the process, glutathione is converted to its oxidized form glutathione disulfide (GSSG). (Glutathione (GSH) is a tripeptide with a gamma peptide linkage between the amino group of cystein (which is attached by normal peptide linkage to a glycine) and the carboxyl group of the glutamic acid side chain.) • Glutathione is an antioxidant, preventing damage to important cellular components caused by reactive oxygen species such as free radicals and peroxides. • Once oxidized, glutathione can be reduced back by glutathione reductase, using NADPH as an electron donor. 4) Thiols are still used today as one of the main treatments for heavy metal poisoning. By binding to the toxic metals, they help them pass safely out of the body. THIOETHERS (Sulfides) A thioether is similar to an ether except that it contains a sulfur atom in place of the oxygen. The grouping of oxygen and sulfur in the periodic table suggests that the chemical properties of ethers and thioethers are somewhat similar. [R-S-R'] NAMING THIOETHERS • Thioethers are sometimes called sulfides, especially in the older literature and this term remains in use for the names of specific thioethers. (CH3)2S dimethylsulfide • Some thioethers are named by modifying the common name for the corresponding ether. For example, C6H5SCH3 is methyl phenyl sulfide, but is more commonly called thioanisole, since its structure is related to that of anisole, C6H5OCH3 PREPARATION 1) Thio ethers are typically prepared by the alkylation of thiols: R-Br + HS-R' → R-S-R' + HBr 2) Alternatively, thioethers can be synthesized by the addition of a thiol to an alkene: R-CH=CH2 + HS-R' → R-CH2-CH2-S-R' Chemical properties Oxidation a) While, in general, ethers are non- oxidizeable, thioethers can be easily oxidized to the sulfoxides (R-S(=O)-R) S(CH3)2 + H2O2 → O=S(CH3)2 + H2O dimethylsulfoxide Dimethyl sulfoxide (O=S(CH3)2, DMSO) is often used as a polar aprotic solvent. b) Oxidation of a sulfoxide with a peroxyacid yields a sulfone (R-S(=O)2-R • Methionine sulfoxide increases with age in body tissues, which is believed to contribute to ageing. • Side chain of methionone in proteins is readily oxidized to markedly alter the overall properties of the amino acid, thus potentially modifying protein function. Methionine(protein)+ H2O2→ Methionine Sulfoxide (protein)+ H2O Methionine Sulfoxide(protein)+ NADPH+H+→ Methionine(protein)+ NADP++H2O Side chains of methionine in proteins are readily oxidized to markedly alter the overall properties of the amino acids, thus potentially modifying protein function. The sulfur mustards, commonly known as mustard gas, are a class of related cytotoxic and chemical warfare agents with the ability to form large blisters on the exposed skin and in the lungs. - Pure sulfur mustards are colorless, viscous liquids at room temperature. - When used in impure form, such as warfare agents, have an odor resembling mustard plants, garlic, or horseradish, hence the name. - Large-scale production of mustard gas was for the Imperial German Army in 1916. Development of the first chemotherapy drug As early as 1919 it was known that mustard gas was a suppressor of hematopoiesis. In addition, autopsies performed on 75 soldiers who had died of mustard gas during World War I were done by researchers from the University of Pennsylvania who reported decreased counts of white blood cells. This led the American Office of Scientific Research and Development (OSRD) to finance the biology and chemistry departments at Yale University to conduct research on the use of chemical warfare during World War II. As a part of this effort, the group investigated nitrogen mustard as a therapy for Hodgkin's lymphoma and other types of lymphoma and leukemia, and this compound was tried out on its first human patient in December 1942. The results of this study were not published until 1946, when they were declassified. In a parallel track, after the air raid on Bari in December 1943, the doctors of the U.S. Army noted that white blood cell counts were reduced in their patients. Some years after World War II was over, the incident in Bari and the work of the Yale University group with nitrogen mustard converged, and this prompted a search for other similar chemical compounds. Due to its use in previous studies, the nitrogen mustard call

The elements of the nitrogen group: electronic structure and chemical properties. Ammonia. Oxides of nitrogen. Oxoacids of nitrogen and their salts. Oxides of phosphorus. Oxoacids of phosphorus and their salts. Halogenated phosphorus compounds.

Elemental state: N2 inert gas 1.Ammonia (NH3) Colorless, water-soluble gas, unpleasant odour. Aqueous solution is basic (weak base). trig pyramidal w/ dipole character and H bonding. lone pair e- means its bron low base 2. Dinitrogen oxide (N2O, „laughing gas") The only inorganic anesthetic. 3.Nitrogen monoxide (NO) Colorless, odorless gas. 4. Nitrogen dioxide (NO2) Reacts with oxygen at room temperature: NO(g) + O2(g) 2 NO2(g) Mixture of NO and NO2 is very toxic (city smog, acid rain). NO2 is the common anhydride of HNO3 and HNO2. Oxoacids&Salts: HNO3//nitric acid. SA, strong ox agent. nitrates are salts. HNO2// nitrous acid. Nitrites are salts. unstable and WA Phosphorous Oxides: Phosphorus pentoxide (P4O10): White, hygroscopic powder. Reacts vigorously with water (dehydrating agent) Phosphorus trioxide (P4O6) Oxoacids of phosphorus: metaphosphoric acid pyrophosphoric acid orthophosphoric acid *phosphoric acid: aka orthophosphoric acid (H3PO4) Moderately strong or weak acid. In soft drinks. Salts: phosphates. halogenated phosph comps i dont feel like doing this.

Acidity and basicity of aqueous solutions. Reaction of ions with water: hydration, solvation, hydrolysis. Acid-base titration; indicators.

Hydration: the interaction of ions with water molecules. Solvation: the interaction of ions with the molecules of a solvent. Hydrolysis: the reaction of an ion with water. strong acids or bases react with water, so hydrolyze. HA + H2O acid H3O+ + A− conjugate base Strong acid shifted to the rightWeak acid shifted to the left The conjugate pair a strong acid/base is a weak base/acid, the conjugate pair of a weak acid/base is a strong base/acid Titration is a procedure for determining the amount of a substance by adding a measured volume of a solution with known concentration of another substance until the reaction of the substances is just complete. (Volumetric analysis.) Acid-base titration: a titration that is based on an acid-base (neutralization) reaction. Gives an acid/ base titration curve plots pH of solution vs the amount of acid/base added. Has equivalence point where stoichiometric amount of reactant has been added Acid-base indicators are weak acids or bases having differently colored acidic and basic forms ie Phenol Red yellow at 6.5, red at 8pH used to show change in pH of titrations

Hydrogen and its compounds: water, hydrogen peroxide, ammonia, hydrogen sulfide, hydrides. Alkali metals: electronic structure and chemical properties. Oxides, hydroxides. Salts of alkali metals.

Hydrogen is the must abundant element in universe, colorless odorless flammable gas. Forms covalent compounds usually, sometimes ionic (HYDRIDES with 1a and 2a group metals). Isotopes include protium, deuterium, tritium. Water: essential for life. max density at 4*C. A hydrophylic solvent Hydrogen Peroxide: colorless liquid, spont decomp into water and oxygen gas. Strong oxidiser.used as antiseptic. Ammonia: Hydrogen Sulfide: smells bad, slighty soluble Hydrides: Alkali Metals: Oxides: three types: ionic 9metal) oxides that are basic in water. molecular oxides (nonmetal) which are acidic in water and mostly soluble in water. covalent network oxides which are solid, hard, h2o insoluble. Hydroxides: Salts of Alakli metals: These elements are located in group 1 of the periodic table They are very reactive elements Because they are so reactive they do not occur freely in nature These elements have only one electron in their outer shell Theyre ready to lose their single electron in ionic bonding with other elements Like all metals they are malleable, ductile, and good conductors of electricity These metals tend to be softer then most other metals Cesium and Francium are the most reactive Alkali Metals Exposure to water can result in an explosion They have lower densities than other metals. They have one loosely bound electron. They have low ionization energies and low electronegativities.

intermolecular forces. Hydrogen bonding. Criteria of H-bond formation. Hydrogen bonding in organic compounds. Van der Waals forces: dipole-dipole and London forces.

IMFs: Attractive forces between molecules. Boiling and melting points are determined by the intermolecular forces and the molecular mass. Energy: 0-40 kJ/mol Types: 1. Hydrogen bonding (10-40 kJ/mol) =Electrostatic attraction between a partially positively charged hydrogen atom and a lone pair of electron of a partially negatively charged atom (N, O or F). --- Hbond is Organic Compounds: Alcohol molecules can form hydrogen bonds with one another, but ether molecules cannot. Ether molecules can form hydrogen bonds with water molecules. 2. Van der Waals forces (0-10 kJ/mol) a, dipole-dipole forces: A polar molecule has a permanent dipole. Polar molecules tend to allign themselves so that the positive end of one molecule is near the negative end of another. Dipole-dipole force: the electrostatic attraction between opp charged poles of the molecules. ie HCl b, London (dispersion) force: Exist between any two molecules. instantaneous dipoles. ex. I2 London force: the weak electrostatic attraction between the oppositely charged poles of the instantaneous dipoles

Structure and nomenclature of alkenes. Constitutional and cis-trans isomerism. Chemical reactions of carbon-carbon double bonds: electrophilic addition. Markovnikov's rule. Polymerization. Alkynes, acetylene and its reactions.

IUPAC rules: Nomenclature of unsaturated hydrocarbons 1. Ending: -ene for C=C and -yne for C≡C. 2. Select the longest chain that includes the multiple bond. 3. Number the chain from the end nearest the multiple bond. 4. If both a double and a triple bond is present, the double bond receives the lowest numbers Structure: - sp2 hybridized carbon atom - trigonal planar arrangement - bond angle ~ 120° Consitutional and Stereoismoerism (Described in previous FC) Chemical Reactions: a. Addition of Halogens: No catalyst required. Diatomic halogen attacks double bond to form two C-Halogen bonds b. Addition of water Alkene+ H2O---(ACID CATALYST)--> form C-H and C-OH bond (Alcohol formed) c. Addition of Acid Note that the Hydrogen of the acid ( HCl, HF, HBr, H2SO4) is electropositive. H and A bond to One Carbon each. d. Addition of Hydrogen C=C + H2 --> C-C with two new CH bonds. A CATALYST NEEDED. (Pd) ------------- MARK RULE: Used in addtn of unsymm reagent to unsymm alkenes The addition is regiospecific: it gives only one product. Markovnikov's rule: the electropositive part of the reagent (H) bonds to the carbon of the double bond that has the greater number of H atoms attached to it... Ie nucelophile goes to more substituted carbon not one more H.. the other H will go to less sub. -------------- Electrophilic Mech.. described in previous FC but essentially an electrophile attacks C=C and plucks off H, leaving an unstable carbocation intermediate.. intermediate reacts with Nucleophile to regain stability aka 2 e- -------------- Polymerisation: [Polymer]: large molecule built up from small repeating units. [Monomer]: small molecule from which the repeating units are derived. [Polymerization]: the process converting a monomer to polymer. occurs via free radical mechanism ----------- Alkynes: sp hybridized carbon atom, linear arrangement, bond angle = 180° (triple bond) two pi one sigma Addition occurs more slowly than with alkenes. a, Halogenation b, Hydration c, Addition of acids d, Hydrogenation Acetylene (systematic name: ethyne) is the chemical compound with the formula C2H2. It is a hydrocarbon and the simplest alkyne. Reactions: A hydrogen atom on a triply bonded carbon is weakly acidic therefore, when added to Na+, a REDOX reaction will occur.. cool! More bonds between two C's leads to more acidity aka willingness to give up H. Makes sense. In its Hydration reaction, Hg2+ forms a complex with the triple bond and activates it (catalyst). Alcohol intermediate, followed by rearrangement to C=O double bond and C-CH3 bond instead.

The elements of the carbon group, electronic structure. Carbon-carbon bonds: sp3, sp2 and sp hybridization of the carbon atom. Chemical properties. Oxides of carbon, carbonic acid and its salts. Silicates.

Includes Carbon: NONMETAL and Si a metalloid. E structure ends in 2p2 or 3p2 w/ 4 ve-. Hybrid states: 4 ve so can be sp3, sp2 and sp hybridized (form 4 bonds round itself) Allotropes: diamond, graphite and fullerenes. Carbon// combustible and reducing agent. Oxides: 1. Carbon monoxide (CO) Odorless, colorless, very toxic (binds to hemoglobin). Combustible 2. Carbon dioxide (CO2) Colorless, odorless, toxic gas. CO2 is soluble in water. The aqueous solution is acidic. 3. Carbonic Acid (H2CO3). CO2 is the anhydride of H2CO3. carbonates are its salts. 4. Hydrogen cyanide (HCN) Volatile liquid, almond odour. toxic: CN− ion is a very good complexing agent Salts: cyanides. Silicon// Elemental state: covalent network solid. Diamond-like structure, but semiconductor. 1. Silicon dioxide (SiO2, quartz, sand) Covalently bonded (SiO4) structural units. Usage: special glasses 2.Silicic acids Orthosilicic acid (H4SiO4), metasilicic acid (H2SiO3) Sodium metasilicate (Na2SiO3, water glass). 3. Silicons: Synthetic compounds (chains or rings) containing Si-O bonds. Usage: heart valve implants, medical tubing. -> Organosilicons: Si atoms and alkyl groups

The elements of the oxygen group. Electronic structure and chemical properties. Oxides. Acidic anhydrides and basic anhydrides. Peroxides and superoxides. Hydrogen sulfide and other sulfides. Oxides of sulfur. Oxoacids of sulfur and their salts.

Includes O and S. Electron structure of O is 1s2 2s2 2p4. S is 1s2 2s2 2p6 3s2 3p4. both one of six organogenic elements. O: Low solubility in water. Reacts with almost all elements via redox rxns. ox state usually -2. −1 in peroxides (e.g. H2O2) −1/2 in superoxides (e.g. KO2) allotrope is ozone. Compounds include water and h2o2 oxides include: -ionic (metal oxides) some insoluble in aq. they are basic anhydrides -molecular oxides (nonmetal oxides) most soluble and are acidic anhydrides. -covalent network oxides: solid hard insoluble. S: Elemental state: S8 molecules. Yellow, crystalline material. Low melting point (115°C) Insoluble in water (nonpolar), soluble in nonpolar solvents. Reacts with metals and nonmetals at high temperature to form Ionic sulfides S compounds: Hydrogen sulfide (H2S) Toxic gas having rotten egg odour. Slightly soluble in water. Aqueous solution is weak acid. Reducing agent. salts: sulfides, insoluble!!!! 1.Sulfur dioxide (SO2) Colorless, toxic gas. Soluble in water. Reacts with water 2.Sulfurous acid (H2SO3) Weak acid. Easily decomposes: H2SO3 Salts: sulfites. 3.Sulfur trioxide (SO3) Colorless liquid. Exothermic run with h2o. 3. sulfuric acid h2so4 SA. exothermic hydration. hygroscopic: removes water from organic compounds. concentrated, strong oxidizing agent. salts are sulfates!!!

Common-ion effect. Buffers. Application of buffers. Henderson-Hasselbalch equation. Types of buffers. Buffer capacity.

Le Châtelier's Principle states that if an equilibrium gets out of balance, the reaction will shift to restore the balance. If a common ion is added to a weak acid or weak base equilibrium, then the equilibrium will shift towards the reactants, in this case the weak acid or base buffers:A buffer is a solution that has the ability to resist changes when a limited amount of acid or base is added to it. • either a weak acid and its conjugate base or a weak base and its conjugate acid. A buffer solution contains both an acid species and a base species in equilibrium application of buffers: H-H equation: pH=pka+log[A-]/[HA] pH=pkb+log[B+]/[OH-] Types of Buffers: Buffer capacity is the number of moles of a strong monoprotic acid or base added to 1L of a buffer solution causing 1 unit change in the pH. Buffer capacity depends on: • the amount of acid and conjugate base (or base and conjugate acid) Having more of these, the buffer capacity is greater. • the ratio of amounts of acid and conjugate base (or base and conjugate acid) This ratio should be close to 1 (between 1:10 and 10:1), otherwise the buffer capacity is too low. ACIDIC BUFFERS: HA- H+ + A- adding acid to it: equilib shifts left. tiny decrease in pH. Adding base to it(OH ions) equilib shifted(LeChat) to right and tiny increase in pH. BASIC BUFFERS: Adding base tiny decrease in pH(common ion). Adding an acid tiny decrease in pH (LeChat) pH > 7.4 alkalosis pH < 7.4 acidosis Hydrogencarbonate/ CO2 buffer prevents serious changes in blood.

Complex compounds. Formation and stability of complexes. Coordination number. Chelate complexes. Biological importance of chelates.

Ligand and central ion/atom bonded covalently (dative) ligand has a nonbonding e- pair, central atom had empty orbitals in valence shell Coordinate covalent bond: the ligand donates, the central ion or atom accepts the electron pair. Coordination number: the total number of bonds the central ion or atom forms with ligands. Monodentate ligand: „one-toothed" ligand, a ligand that is bonded to the central ion or atom through one bond. Polydentate ligand („has many teeth"): a ligand that can bond to the central ion or atom with two or more bonds A chelate is a chemical compound composed of a metal ion and a chelating agent. A chelating agent is a substance whose molecules can form several bonds to a single metal ion. In other words, a chelating agent is a multidentate ligand. An example of a simple chelating agent is ethylenediamine. Chelates play important roles in oxygen transport and in photosynthesis. Furthermore, many biological catalysts (enzymes) are chelates. Chelate: a STABLE complex formed by polydentate ligands.

Liquid state. Changes of state, phase changes related to pressure and temperature. Phase diagrams. Intermolecular forces. Association of molecules. Solid state. Types of solids; crystalline solids. The crystalline lattice and crystal systems. Sublimation and lyophilization

Liquids are continuously vaporising. Vapor pressure: the partial pressure of the vapor over the liquid measured at a given temperature. The vapor pressure increases with the temperature. Phase Diagram: graph that shows conditions (temp v pressure)under which the different states of a substance are stable. shows melting pt curve, boiling pt curve, sublimation. A is trile point where three phases coexist in equillib, C is critical point above which liquid state doesnt exist Solid State: units are close and fixed, nearly incompressible w/ well defined shape, volume. Can be crystalline or amorphous Crystalline: Well-defined, ordered structure. True melting point. Crystal lattice: geometric arrangement of lattice points in a crystal. E.g. diamond, table salt, dry ice. Amphorous Disordered structure. No true melting point. Gradual phase change. E.g. glass, polymers. TYPES OF CRYSTALLINE SOLIDS: -Molecular//molecule,atoms. Intermolecular forces ie h2o, co2 Low MP, SOFT, Noncond -metallic ie Fe, Ag, Na. MP varies, hardness varies, CONDUCTING -ionic: ie NaCl. HIGH MP, hard, with solid NC and liquid conducting -covalent network: ie diamond, made of atoms with covalent bonding. Very high MP, very hard, noncond Sublimation: solid to gas Examples: I2, solid CO2 Lyophilization aka freeze drying: a dehydration process at low T and p. Steps: 1. Freezing the substance. 2. Removing the solvent at low p (in vacuum): sublimation Purpose: preserve material.

Metathesis reactions: precipitation, gas formation and neutralization. Solubility of electrolytes, solubility product constant (Ksp). Monatomic and polyatomic ions, naming of these ions and their salts.

Metathesis: rxns in which two compounds react to form two new compounds, with no changes in oxidation number. Reactions in which the ions of two compounds exchange partners. Precipitation: If a reaction taking place in aqueous solution produces an insoluble product, that product will precipitate, forming a solid. Precipitation often happens with double replacement reactions. Gas Formation: Any reaction that produces a product in the gas phase is a gas forming reaction Whenever a metathesis reaction results in the products H2CO3, H2SO3, or H2S, a gas will be formed. The compound H2S is itself a gas and you need to recognize it as such. The other two products, H2CO3 and H2SO3 decompose to form gases. Neutralization: acid and base react to neutralize one another to make water and a salt. (pH balances) Solubility of electrolytes: Electrolyte = A solute that produces ions in solution. Electrolytic Solution = A solution that conducts electricity. When an electrolyte dissolves in water (or another solvent) it forms an electrolytic solution Electrolytes can be further classified into strong and weak electrolytes. Strong electrolyte = A solute that dissolves almost completely to ions (i.e. soluble inorganic salts, strong acids & bases). Weak electrolyte = A solute that only partially dissolves into ions (i.e. weak acids and bases).weak electrolyte dissolves in water to give an equilibrium between a molecular substance and a small amount of ions Ksp: the equilibrium constant for a solid substance dissolving in an aqueous solution. It represents the level at which a solute dissolves in solution. A more a substance dissolves, the higher the Ksp value it has. Example: aA(s)⇌cC(aq)+dD(aq) To solve for the Ksp it is necessary to take the molarities or concentrations of the products (cC and dD) and multiply them. If there are coefficients in front of any of the products, it is necessary to raise the product to that coefficient power(and also multiply the concentration by that coefficient). Monatomic and Polyatomic ions: Naming these Ions& Salts: name a few.

Homogenous and heterogeneous systems. Homogenous systems: mixtures of gases, solutions and alloys. Microheterogeneous systems, colloidal state, colloidal size. Properties of colloids. Types of colloids: hydrophobic and hydrophilic, dispersed and association colloids (micelles).

Mixture(three types): a material that can be separated by physical means into two or more substances. HOMOGENOUS SYSTEMS: uniform whole aka solutions -mixture of gases: -Solutions: -Alloys: HETEROGENOUS SYSTEMS: ie oil+water. have physically distinct parts Microheterogeneous mixtures = colloids Colloidal state:dispersion of particles of one substance (dispersed phase) throughout another substance(continuous phase). ie milk, mayo, fog Colloid Properties:A colloid has the ability to scatter light. The scattering of light by colloidal-size particles is known as the Tyndall effect. Colloid types: according to the nature of interaction between dispersed phase and continuous phase 1.hydrophobic-unstable. in h2o as cont phase is solvent hating. The dispersed phase aggregates into larger particles and form a heterogeneous mixture. (Coagulation) 2. hydrophilic-in h2o as cont phase is solvent loving. strong attractions btwn phases results in stable colloid. Can also be classified based on PARTICLES of dispersed phase: 1. Macromolecules- dispersed phase consists of macromolecules. The size of the macromolecules is in the colloidal range. ie dna, protein 2.association colloids- Amphiphilic molecules associate and form colloidal-sized particles, or micelles. colloid in which the dispersed phase consists of micelles. 3. dispersed colloids- a gas, liquid or solid microphase is dispersed in a continuous phase. ie smoke, fog. can be further classified ased on continuous phase.

The electronic structure and chemical properties of d-transition metals. Oxidation states of d- transition metals. Most important representatives: e.g. Mn, Fe, Cu, Pt and Zn.

Outermost shell electron configuration: from (n-1)d1ns2 to (n-1)d10ns2 High density and melting points. Oxidation states: can be different. Different color for each of them. //Manganese Valence-shell: 3d54s2 Oxidation states: +2, +3, +4, +6, +7 -Manganese dioxide (MnO2) Brown precipitate. -Potassium permanganate (KMnO4) Strong oxidizing agent. In a strongly acidic solution //Fe Valence-shell: 3d64s2 Oxidation states: +2, +3 Rust: FeO(OH) Biological importance: essential trace element //Cu Oxidation states: +1 and +2 (Cu2O and CuO) Essential trace element. Very good conductor (2nd). Copper sulfate (CuSO4) Blue, water soluble substance. Analytical reagent, effective antifungal agent //Pt Pt is [Xe] 6s1 4f14 5d9. also, the configuration for Pt2+ ion is [Xe] 4f14 5d8 //Zn Oxidation state: +2 Essential trace element - Stabilizes the conformation of certain proteins Zinc sulfate (ZnSO4) Antiseptic agent.

Polar covalent bond. Electronegativity. Dipole moment. Geometrical arrangement of molecules.

PoCo bond:Use H2O as an example: In a polar covalent bond, the electrons shared by the atoms spend a greater amount of time, on the average, closer to the Oxygen nucleus than the Hydrogen nucleus. This is because of the geometry of the molecule and the great electronegativity difference between the Hydrogen atom and the Oxygen atom. EN: intrinsic ability of an atom to attract the shared electrons in a covalent bond F highest at 4. Metals on left side of periodic table attract electrons weakly, lower EN Halogens and other reactive nonmetals on right side of periodic table attract electrons strongly, higher electronegativities. Diff in EN gives dipole moment, nonzero dipole moment ((μ): the quantitative measure of charge separation in a molecule originated from unequal distribution of electrons between atoms in the bond;net molecular polarity) gives poco bonds dipole moment == magnitude of charge (Q) multiplied by the distance (r) between the charges; units are D (Debye) = 3.36 x 1030 C.m Geometric Arrangement of Molecules: In molecules there are 2 types of electron: bonding pairs and non-bonding or lone pairs. The combinations of these determine the shape of the molecule. Single bonds have a big impact on shape, double bonds have little effect. The outer pairs of electrons around a covalently bonded atom minimize repulsions between them by moving as far apart as possible. 2 ep w/ 2bonding linear 3 ep w/ 3bonding trig plan. 2bonding1non is bent 4 ep (sp3) tetrahedral, trig pyramidal, bent (2bonding 2 nonbonding) 5bonding: sp3d, trigonal bipyramidal 6 bonding: octahedral sp3d2 In a molecule with two or more polar bonds, each bond has a dipole moment contribution = bond dipole Net dipole moment = vector sum of its bond dipoles. Nonzero dipole moment: polar covalent bond(s) and an appropriate geometry. symm=nonpolar, nonsym=polar ie h2o

Polarization in organic compounds; polar covalent bond, inductive effect, functional groups having inductive effect. Conjugative effect: position of electron pairs. Butadiene: delocalization of double bonds; mesomeric structures. Resonance theory. Groups having conjugative effect.

Polarity Ranking of the Functional Groups: (most polar first) Amide > Acid > Alcohol > Ketone ~ Aldehyde > Amine > Ester > Ether > Alkane In general, the presence of an oxygen is more polar than a nitrogen because oxygen is more electronegative than nitrogen. The combination of carbons and hydrogens as in hydrocarbons or in the hydrocarbon portion of a molecule with a functional group is always NON-POLAR. Inductive Effect: e- dist on SIGMA bond Inductive effect: The effect on electron density in one portion of a molecule due to electron-withdrawing or electron-donating groups elsewhere in the molecule. Inductive effect (I) Electron distribution on a σ bond. There is the Positive inductive effect (+I) (CENTRAL ATOM MORE EN) Groups with positive inductive effect (+I) (Electron donating groups) Alkyl (saturated hydrocarbon groups): e.g. -CH3 Negatively charged groups: e.g. COO− and the Negative inductive effect (−I) (CENTRAL ATOM LESS EN) Groups with negative inductive effect (−I) (Electron withrawing or releasing groups) -F, -Cl, -Br, -I -OH, -NH2, -COOH, -COOR, -CN, -NO2, -COH -Ar Positively charged groups: e.g. NH3+ ------ CONJUGATIVE EFFECT: e- disttrbution on a PI bond 1. Positive conjugative or mesomeric effect (+K): They release (donate) electrons to the rest of the molecule by delocalization -F, -Cl, -Br, -I -OH, -OR -NH2, -NR2 2. Negative conjugative or mesomeric effect (−K) They withdraw electrons from the rest of the molecule by delocalization. COOH, -COOR, -CN, -NO2, -COH ------ BUTADIENE The π electrons are not localized between the carbon atoms of the double bonds, but rather delocalized over four atoms. C+ stabilized by resonance: that explains why, when you add smthg like Br+ to a conjugated diene you'll be mroe 1,4 product than 1,2

Oxidation and reduction. Standard potential, strength of oxidizing and reducing agents involving metals and nonmetals. Reactions of metals with metal ions, water, diluted and concentrated acids, bases, reactions of halogens with halide ions.

Reduction is gaining electrons Oxidation is gaining ox state, or losing electrons. OILRIG..... Standard reduction potential: [voltage associated with a reduction reaction at an electrode when all solutes are 1 M and all gases are at 1 atm and the temperature is 298 K] -more positive E° value, greater the tendency to be reduced(CATHODE GETS REDUCED).. More negative will be oxidised (ANODE) TO Get EMF of cell: cat- anode ((anode is lower always) Strength of oxidizing and reducing agents involving metals and nonmetals //Metal + diluted acid The metal is oxidized: H+ of acid is reduced: 2H+ +2e− →H2 Spontaneous if E°metal < 0.0 V //Metal + water The metals is oxidized − Water is reduced: H2O + 2 e− → H2 + 2 OH− Spontaneous if E°metal < -0.83 V Metals react with cc. oxidizing acids above 0V Metals react with diluted acids from 0 to -0.83V Metals react with diluted acids and water below -0.83V //Metals with water: Na reacts readily, Al reacts only after removal of the protecting oxide layer //Metals with diluted acids: ZINC and IRON react readily Electropositive metals react only with concentrated, oxidizing acids like HNO3 //Metals + bases Aluminum has a negative standard potential and is amphoteric: reacts with both acids and bases. //Metals + metal ions: a) Cu2+ +Zn→Cu+Zn2+ Spontaneous, copper is oxidizing agent (coppers reduced) and zinc is reducing agent (zincs oxidized) //Halogens + halide ions: a) chlorine + bromide: Spontaneous, because ε°(Cl−/Cl2/Pt) = 1.36 V and ε°( Br−/Br2/Pt) = 1.07 V, so chlorine is oxidizing agent (chlorine is being reduced) and bromine must be the reducing agent (bromine is oxidized.) b) Cl− + Br2 → no reaction tho

Types of organic chemical reactions. Substitution, addition and elimination reactions. Homolysis and heterolysis, radical and ionic mechanisms. Nucleophilic and electrophilic reagents. Examples for electrophilic addition, substitution (AE and SE) and nucleophilic addition, substitution (AN and SN).

Substitution: ----------- Substitution: The most general form for the reaction may be given as Nuc: + R-LG → R-Nuc + LG: The electron pair (:) from the nucleophile(Nuc) attacks the substrate (R-LG) forming a new bond, while the leaving group (LG) departs with an electron pair. The principal product in this case is R-Nuc. The nucleophile may be electrically neutral or negatively charged, whereas the substrate is typically neutral or positively charged. The two main mechanisms are the SN1 reaction and the SN2 reaction. S stands for chemical substitution, N stands for nucleophilic, and the number represents the kinetic order of the reaction.[4] SN1 forms carbocation(Very electrophilic) before nu: attack. In the SN2 reaction, the addition of the nucleophile and the elimination of leaving group take place simultaneously. SN2 occurs where the central carbon atom is easily accessible to the nucleophile. By contrast the SN1 reaction involves two steps. SN1 reactions tend to be important when the central carbon atom of the substrate is surrounded by bulky groups, both because such groups interfere sterically with the SN2 reaction (discussed above) and because a highly substituted carbon forms a stable carbocation. Addition: Here's the basic pattern: break a C-C multiple bond (also called a π bond) and form two new single bonds ("σ-bonds" to carbon. Note that the reaction occurs only at the carbons that are a part of a multiple bond - nothing else on the molecule is affected. ------------ a. Addition of Halogens: No catalyst required. Diatomic halogen attacks double bond to form two C-Halogen bonds b. Addition of water Alkene+ H2O---(ACID CATALYST)--> form C-H and C-OH bond (Alcohol formed) c. Addition of Acid Note that the Hydrogen of the acid ( HCl, HF, HBr, H2SO4) is electropositive. H and A bond to One Carbon each. d. Addition of Hydrogen C=C + H2 --> C-C with two new CH bonds. A CATALYST NEEDED. (Pd) Addition to unsymmetric Alkenes: regiospecific: gives only one product. Markovnikov's rule: the electropositive part of the reagent (H) bonds to the carbon of the double bond that has the greater number of H atoms attached to it. SO BASICALLY IF A RXN FOLLOWS THIS, HALOGEN WILL ATTACH TO Carbon WITH MORE CARBS BONDED TO IT:) Elimination: ---------- two substituents are removed from a molecule in either a one or two-step mechanism.[2] The one-step mechanism is known as the E2 reaction, and the two-step mechanism is known as the E1 reaction. Essentially, a C-C pi bond forms. Two single bonds broken [C-H bond and an adjacent to C-(atom) single bond] --------------------------- Homo/Heterolysis: (Breaking single bonds) Homolytic fission is where each atom of the bond keeps an electron each resulting in species called free radicals. Radicals are important intermediates in organic chemistry and we will talk about them later. As the bond breaks to give two similar species each keeping an electron this form of bond breaking is called Homolytic Fission. one of the atoms carry a negative charge after bond cleavage indicating that it has both the electrons of the bond and the other has no electrons at all. Hence it is electron deficient thus positively charged. As the electrons are not divided equally after bond cleavage this is called Heterolytic Fission Way to read all of that! Now you get a fun fact.. Alkynes: sp hybridized carbon atom - linear arrangement - bond angle = 180° Radical and Ionic Mechanisms: -------------------------- Free Radical Polymerization Mechanism: The process converting a monomer to polymer. Initiation, propogation, and termination factors. ^^ IDK what is ionic mech.^^ ------------ Nucleophilic and electrophilic reagents. Electrophilic reagents: chemical species which get e- Although this definition embraces all oxidizing agents and all Lewis acids, electrophilic reagents are ordinarily thought of as cationic species, such as H+, NO2+, Br+, or SO3 (or carriers of these species such as HCl, CH3COONO2, or Br2), which can form stable covalent bonds with carbon atoms. Electrophilic reagents frequently are positively charged ions (cations). An electrophile that accepts an electron pair at hydrogen is called a Brønsted acid, or just "acid". Nucleophilic reagents are the opposite of electrophilic reagents. Nucleophilic reagents GIVE UP e- -frequently negatively charged ions (anions). I⊖ > Br⊖ >Cl > F⊖; HS⊖ > HO⊖; PH3 > NH3. Thus, other things being equal, larger atoms are better nucleophiles. FACTORS FOR GOOD NUC: 1.Charge!!! (more negative better) 2. Electronegativity!! (More EN better) 3. Solvent (Polar Protic, NU decreases down the pt) (Polar Apro, Nu increases going up pt) 4. Steric hindrance!!!!: less bulky=better nut Typical nucleophilic reagents are hydroxide ion (OH−), halide ions (F−, Cl−, Br−, and I−), cyanide ion (CN−), ammonia (NH3), amines, alkoxide ions (such as CH3O−), and mercaptide ions (such as C6H5S− EXAMPLES for: electrophilic addition: C=C (ETHENE) + H+ ->carbocation->C+ reacts with Nuc-> C-C-Nu substitution (AE and SE) halogen addtn: Ethene+Cl2-> (C-CL)2H4 and nucleophilic addition, substitution (AN and SN).

Thermodynamics. State functions (state properties). Exothermic and endothermic reactions. Internal energy. The first law of thermodynamics. Heat of reaction and enthalpy. Hess's law.

Thermodynamics= study of transfer of energy from one form to another, the inter-relation between heat, work and internal energy of a system State Functions property of a system that depends only on the current state of the system, not on the way in which the system acquired that state (independent of path). A state function describes the equilibrium state of a system. For example, internal energy, enthalpy, and entropy Enthaply: its important to have a way of expressing heat change of the system during a constant pressure reaction, a new STATE FUNCTION is defined, ENTHALPY = H. • enthalpy (heatcontent) of reaction equals the heat of reaction at constant pressure (qp) Bond enthalpy = average enthalpy change for breaking a particular bond in a mole of gaseous substance Exothermic: heat released from reaction Endothermic: heat abs in run • Endothermic ∆H > 0 • Exothermic ∆H < 0 q > 0 heat transferred INTO a system q < 0 heat transferred OUT of a system defined as enthalpy. Change in enthalpy (∆H) is the heat gained or lost by the system when a process occurs under constant pressure delta H is a state function and extensive property The total amount of heat that is evolved or absorbed by a system in a chemical reaction (qp) • The amount of heat required to return a system to the given temperature at the completion of the reaction HESS'S LAW: • The heat exchange is independent of the path. A reaction can be carried out in one step or any number of steps • The net enthalpy change is the sum of all enthalpies.

Electrolytes, electrolytic dissociation, degree of dissociation. Strong and weak electrolytes. Equilibrium in electrolytes. Self-ionization of water. Ion product of water. pH and pOH. Acid and base ionization equilibrium, dissociation constants (Ka and Kb).

electrolyte is a compound that ionizes when dissolved in a solvent. strong electrolytes diss completely. ie NaCl. weaker ones (slightly soluble or insoluble salts) do not.Ksp is the product solubility constant The conjugate pair a strong acid/base is a weak base/acid, the conjugate pair of a weak acid/base is a strong base/acid Kc is equillib constant from product cxns over reactant coxs (no coefficients!)Ksp is the solubility product constant ie Ksp(AgCl) = [Ag+][Cl−] H2o self io: Pure water has a very small conductivity resulting from self-ionization (or autoionization), a reaction in which two like molecules react to give ions. Kw= [h3o][oh]=1x10^-14 neutral solution [oh]=[h+]. acidic has more H+ vive. so acidic, [H]>1.0*10^-7 pH=-log[H+] pH+pOH=14 pH under 7 means acidic. wA or wB only dissociate to a small extent at equilibrium. Bigger Ka means stronger acid (ka= [H][A]/[HA]). smaller pKa means stronger acid. polyprotic weak acids: lose several protons per molecule but each molecule has a lower Ka.

Entropy and disorder. The second and the third laws of thermodynamics. Entropy change for a reaction. Spontaneous reactions. Free energy. Free energy and equilibrium.

entropy is measure of randomness/ disorder. Suniv is constantly increasing. 2nd law of thermodynamics: The entropy of the universe increases in any spontaneous process. ∆Suniv must increase during a spontaneous process, even if ∆Ssyst decreases. Third Law: The entropy of a pure crystalline substance at T=0 (absolute zero) is zero. Decrease in temperature ⇒ decrease in thermal energy ⇒ decrease in translational, vibrational and rotational energy ⇒ decrease in entropy In spontaneous reactions: ∆H<0, ∆S>0 ∆G < 0 ⇒ spontaneous in forward direction (exergonic process) ∆G > 0 ⇒ non-spontaneous in forward direction spontaneous in reverse direction (endergonic process) GIBBS: criterion for spontaneity prediction that focuses only on the system. DEFINED as G=H-TdeltaS. An extensive state function. so High entropy with large negative deltaH favors high entropy

Classification of organic compounds based on their functional groups. Structure, nomenclature and structural isomerism of alkanes (paraffins). Conformation. Conformers of ethane and butane. Naming of alkyl groups. Characteristic chemical reaction: substitution, free radical mechanism.

hydroxyl (-oh)//alcohol carbonyl (CHO)// Aldehydes --CO// Ketone carboxyl//carboxylic acid amino(NH2)//Amines phosphate(OPO3)//organic phosphates sulfyhydryl (SH)// thiols lUPAC rules: Nomenclature of unsaturated hydrocarbons 1. Ending: -ene for C=C and -yne for C≡C. 2. Select the longest chain that includes the multiple bond. 3. Number the chain from the end nearest the multiple bond. 4. If both a double and a triple bond is present, the double bond receives the lowest numbers. ----- Alkenes:carbon-carbon double bond. Alkynes: C-C triple bond Structure, nomenclature and structural isomerism of alkanes (paraffins) Isomerism: can be constitutional or stereoisomers. isomers are different compounds that have the same molecular formula. When the group of atoms that make up the molecules of different isomers are bonded together in fundamentally different ways, we refer to such compounds as constitutional isomers. For example, in the case of the C4H8 hydrocarbons, most of the isomers are constitutional. Isomers that differ only in the spatial orientation of their component atoms are called stereoisomers. Stereoisomers always require that an additional nomenclature prefix be added to the IUPAC name in order to indicate their spatial orientation, for example, cis (Latin, meaning on this side) and trans (Latin, meaning across) In n-alkanes, no carbon is bonded to more than two other carbons, giving rise to a linear chain. When a carbon is bonded to more than two other carbons, a branch is formed. The smallest branched alkane is isobutane. Alkenes have a general molecular formula C n H 2n and alkynes have a general...molecular formula of C n H (2n - 2) - sp2 hybridized carbon atom - trigonal planar arrangement - bond angle ~ 120° Alkenes: Rotation about a double bond is restricted. Alkenes exhibit one form of stereoisomerism. To understand how alkenes can form stereoisomers, recall that the C=C double bond consists of a σ bond between the atoms and a Π bond that lies above and below the plane of the molecule. The strength of the Π bond depends directly on the degree of physical overlap between adjacent p-orbitals. This implies that it is impossible to rotate about the double bond without breaking the Π bond completely. Conformers of ethane and butane: Shown with Newman Projections BUTANE: Syn (Eclipsed). Dihedral Angle= 0, 360 Gauche (staggered): 60 Eclipsed: 120 Anti (Staggered): 180 Eclipsed: 240 Gauche (Staggered): 300 → gauche means methyls are neighbors. Eclipsed (SYN)->Staggered(GAUCHE)->eclipsed->Staggered(ANTI)->Eclipsed->Staggered(GAUCHE) ANTI is least energy, Syn is most energy Ethane: Staggered and Eclipsed conformations only. Staggered conformations are lower energy and more commpnly found. Naming of alkyl groups. Single chain of carbons, Methyl, ethyl, propyl(Isopropyl), butyl (isobutyl, sec-butyl, tert-butyl..increasing number of carbons bonded to central C), pentyl, hexyl, octyl Characteristic chemical reaction: Substitution: The most general form for the reaction may be given as Nuc: + R-LG → R-Nuc + LG: The electron pair (:) from the nucleophile(Nuc) attacks the substrate (R-LG) forming a new bond, while the leaving group (LG) departs with an electron pair. The principal product in this case is R-Nuc. The nucleophile may be electrically neutral or negatively charged, whereas the substrate is typically neutral or positively charged. The two main mechanisms are the SN1 reaction and the SN2 reaction. S stands for chemical substitution, N stands for nucleophilic, and the number represents the kinetic order of the reaction.[4] In the SN2 reaction, the addition of the nucleophile and the elimination of leaving group take place simultaneously. SN2 occurs where the central carbon atom is easily accessible to the nucleophile. By contrast the SN1 reaction involves two steps. SN1 reactions tend to be important when the central carbon atom of the substrate is surrounded by bulky groups, both because such groups interfere sterically with the SN2 reaction (discussed above) and because a highly substituted carbon forms a stable carbocation. free radical mechanism: 1. Initiation: Splitting or homolysis of a chlorine molecule to form two chlorine atoms, initiated by ultraviolet radiation or sunlight. A chlorine atom has an unpaired electron and acts as a free radical. 2. chain propagation (two steps): a hydrogen atom is pulled off from methane leaving a 1˚ methyl radical. The methyl radical then pulls a Cl· from Cl2. This results in the desired product plus another chlorine radical. This radical will then go on to take part in another propagation reaction causing a chain reaction. If there is sufficient chlorine, other products such as CH2Cl2 may be formed. 3. chain termination: recombination of two free radicals: The last possibility in the termination step will result in an impurity in the final mixture; notably this results in an organic molecule with a longer carbon chain than the reactants.

Classification and preparation of organic halogen compounds. Nomenclature of alkyl halides. Mechanisms of nucleophilic substitution. Elimination reactions. Aromatic halogenated compounds and their reactivity.

simplest organochlorine compound: chloromethane, aka methyl chloride (CH 3 Cl). Others include bromomethane (CH 3 Br), chloroform (CHCl 3 ), and carbon tetrachloride (CCl 4 ) Organohalogens can be made in various ways: a. Direct halogenation of hydrocarbons with chlorine gives organochlorines; with bromine, organobromines. b. Alcohols can be converted into organohalogens by reaction with hydrogen halides. c. Aromatic organohalogens such as chlorobenzene are synthesized by treatment of benzene with halogen and a Lewis acid catalyst such as aluminum chloride. d. Organohalogen compounds are also produced by adding halogen or hydrogen halide to alkenes and alkynes. e. Organoiodines and organofluorines are prepared by displacement reactions

States of matter. The gaseous state. State functions. Kinetic theory of gases. Maxwell's distribution of molecular speed. Gas laws. The ideal gas law. Real gases.

solid: rigid, low compressibility, w/ fixed shape and V liquid: fluid, low compressibility, fixed V deltaShape gas: fluid, high compressibility, fits in most containers KINETIC THEORY OF GASES: The kinetic theory of ideal gases is based on five postulates 1. The size of gas molecules is negligible. 2. Gas molecules move randomly in all directions at various speeds. 3. The intermolecular forces are negligible. 4. Collisions are elastic. (The total kinetic energy remains constant.) 5. The average kinetic energy is proportional to the absolute temperature. -State Functions: -Kinetic Thoery of Gases: -Maxwell's Distribution of molecular speed: The speeds of molecules vary over a range of values. Maxwell showed theoretically how molecular speeds are distributed. Molecular speeds vary widely, but most are close to the average speed. As the temperature increases, the average speed increases. -Gas Laws: All gases behave quite simply with respect to pressure, temperature, volume and molar amount. 1. Boyles Law: pV = constant at a fixed T for a given amount of gas 2. Charles's law: V/T = constant at a fixed p for a given amount of gas 3.Avogadro's law: V/n = constant at a fixed T and p All gases behave quite simply with respect to pressure, temperature, volume and molar amount -Ideal Gas Law: pv=nrt -Real Gases: At high pressures, the gases deviate noticeably from the ideal gas law. - The volume of molecules becomes important. The space through which they can move differs from V. - Intermolecular forces become significant. The actual pressure is less than predicted. Van der Waals created equation for real gases

Types of solutions. The solution process. Factors determining solubility. Saturated solutions. Solubility and the molecular structure. Concentration of solutions, ways of expressing concentration. Solubility product constant. Solubility of gases. Effects of temperature and pressure on solubility. Partition coefficient. Osmosis and its biological importance.

solution is solvent plus solute. types include gas(gas in gas), liquid(gas in liquid, liquid in liquid, solid in liquid), solid solutions(liquid in solid, solid oin solid) -A saturated solution contains the maximum amount of a solute that will dissolve in a given solvent at a specific temperature. -An unsaturated solution contains less solute than the solvent has the capacity to dissolve at a specific temperature. -A supersaturated solution contains more solute than is present in a saturated solution at a specific temperature. A dynamic equilibrium - ions continually exchange between the solid and solution form. Solubility: the maximum amount of solute that can dissolve in a specific amount of solvent (usually in 100 g) at a given temperature. Solubility is the composition of saturated solution SOLUTION: LIKE DISSOLVES LIKE: Polar solutes are soluble in polar solvents. Nonpolar solutes are soluble in nonpolar solvents.Ionic compounds are more soluble in polar solvents. Disorder is favorable (2nd law of thermodynamics). For solids: Solubility ↑ as temperature ↑ - usually. If ∆Hsoln > 0 (endothermic) solubility ↑ If ∆Hsoln < 0 (exothermic) solubility ↓ For gases: Solubility ↓ temperature ↑ - always. Kinetic energy plays a primary role Entropy is also a factor Formation of a solution increases entropy Water breaks the + and - charged pieces apart and surrounds them ⇒ hydration (solvation). Hydration: the process of surrounding solute molecules by water. Solvation: the process of surrounding solute molecules by solvent particles. COMPOSITION OF SOLUTIONS: M: c=mol solute/litre solution Mass%: (mass sol/mass solution*100%) Vol%: (vol sol/vol soltn*100%) Mass/Vol%: mass sol/vol solution m(molality): mol solute/mass solvent (kg) mole fraction: mol solute/totalmoles solution

Halogens: electronic structure and chemical properties. Hydrogen halides and their salts, halogen oxoacids and their salts. The noble gases. Electronic structure of noble gases.

symbol electron configuration fluorine F [He]2s22p5 chlorine Cl [Ne]3s23p5 bromine Br [Ar]3d104s2 4p5 iodine I [Kr]4d105s2 5p5 astatine At [Xe]4f14 5d106s2 6p5 The elements of Group 17, the Halogens, are a very similar set of non-metals. They all exist as diatomic molecules, X2, and oxidise metals to form halides. The halogen oxides are acidic, and the hydrides are covalent. Fluorine is the most electronegative element of all. Generally, electronegativity and oxidising ability decrease on descending the Group. The result of this decreasing electronegativity is increased covalent character in the compounds, so that AlF3 is ionic whereas AlCl3 is covalent. The melting points, boiling points, atomic radii and ionic radii all increase on descending the Group. The most characteristic chemical feature of the halogens is their ability to oxidise. Fluorine has the strongest oxidising ability, so other elements which combine with fluorine have their highest possible oxidation number. No fluorine oxides as F is more electronegative than O. Chlorine, bromine and iodine each form several oxides which are thermally unstable, such as chlorine dioxide ClO2. The only fluorine oxoacid, HOF, is unstable at room temperature, but there are many oxoacids of the other halogens. The best known salts of these are; hypochlorite, chlorate(I) CIO-, chlorite, chlorate(III) ClO2-, hypochlorate, chlorate(V) CIO3-, perchlorate, chlorate(VII) ClO4- . Perchlorate is strongest acid, rest decrease in strength w/ fewer O. These are all powerful oxidising agents. Hydrogen halides (HF, HCl, HBr, HI) Aqueous solutions are acidic. Acidity order: HF << HCl < HBr < HI weak acid: HF. Rest r strong acidz Valence shell: ns2np6 (He: ns2) Closed, octet. Monatomic, inert gases. Extremely low reactivity. colorless, odorless, low MP BP

Reaction kinetics. Elementary reactions. Molecularity of a reaction. Reaction rate. Rate law, rate constant. Reaction order. Activation energy. Activated complex. Types of complex chemical reactions. Catalysts, homogenous and heterogeneous catalysis. Enzymes.

the study of the relative rates of chemical reactions. Reactions are often done at constant volume and the quantity of reactant or product is represented as concentration. Elementary RXN: a single step rxn w/ a single transition state, no intermediates The molecularity of a reaction refers to the number of reactant particles involved in the reaction. As there can only be discrete number of particles, the molecularity must be an integer. Molecularity can be described as either unimolecular, bimolecular, or termolecular Reaction Rate given by the nature of reactants concentration of reactants Temperature Catalysts The relationship between reactant concentration and reaction rate is called the rate equation or rate law. Rate Law: experimentally determined characteristic of the reactions. Differential Rate Law: expresses how rate depends on concentration. Integrated Rate Law: expresses how concentration depends on time. Reaction rate ∝ k * concentration k: rate constant (depends on temperature, but independent of the concentration) v=k[A]^x[B]^y. x and y are rxn order. ALWAYS defined in terms of reactant cxns, not product cxns. Increasing the temperature usually has a dramatic effect on reaction rates. "rule of thumb" that a 10°C rise in temperature doubles the rate of a reaction. 1st order: units of k are s^-1. Second order is M^-1time^-1, 3rd is MTime^-1 activation energy (Ea):the energy required to go from reactants to transition state. Species between the reactants and products is called the transition state (or activated complex), and is the highest potential energy point of the reaction. Homo same state. duh. Enzymes are catalysts that drive reaction rates forward. Lower AE. should be noted that speeds up both forward and reverse reactions and cannot increase the final equilibrium yield, but it gets to the final equilibrium state faster. - does not influences ∆H and equilibrium constant ENZYMES act on certain molecules, called substrates (reactants) while leaving the rest of system unaffected. • homogeneous. • An enzyme is typically a protein molecules that contains one or more active sites where interactions with substrates take place

The periodic table. Explanation of periodic properties: relationship to electron configuration. Periodic classification of the elements. Atomic and ionic radius.

the table provides concise and fundamental information not only about every individual element, but also about general trends across all the elements. Each square shows atomic number (protons), atomic mass and element symbol. The Atomic number increases from the top left to the bottom right. It ascends sequentially across each period. Nonmetals: Non-Metal: found in groups 13-18 of the periodic table They are not able to conduct electricity or heat well These elements are very brittle compared to metals These elements are not ductile (can not be rolled into wires or pounded into sheets) They occur as two states of matter: gases (ex. Oxygen) and solids (ex. Carbon) They do not have any metallic luster, nor do they reflect light They usually exist as molecules in their elemental form. They are generally gases at room temperature. They generally form negative ions. Alkali Metal: These elements are located in group 1 of the periodic table They are very reactive elements Because they are so reactive they do not occur freely in nature These elements have only one electron in their outer shell They are ready to loss their single electron in ionic bonding with other elements Like all metals they are malleable, ductile, and good conductors of electricity These metals tend to be softer then most other metals Cesium and Francium are the most reactive Alkali Metals Exposure to water can result in an explosion They have lower densities than other metals. They have one loosely bound electron. They have low ionization energies and low electronegativities. Alkaline Earth: These elements are found in the second group of the periodic table They are metallic They have a oxidation number of +2 Having an oxidation number of +2 makes these elements very reactive They are not found free in nature because of their reactivity They are present in the earth's crust but not in their basic form. They have high boiling and melting points. They have low density, electron affinity, and electronegativity. They react easily with halogens and water. They are softer and stronger than other metals (except the alkali metals). Metalloid: These elements are located on the border between metals and non-metals These elements have properties of metal and non-metals Some of these elements are semi-conductors (they have a ability to carry an electrical charge under certain conditions) They can be shiny or dull. Their shape is easily changed. They typically conduct heat better and electricity better than nonmetals, but not as well as metals. Halogen: These elements are located in group 17 of the periodic table There are only 5 Halogens They are non-metallic elements "Halogen" means "salt-former" and compounds that contain halogens are called "salts" They have 7 electron in their outer shell They have an oxidation number of -1 At room temperature Halogens exist in all three states of matter Transition Metals: These elements are found in group 3 through 12 of the periodic table There are 38 elements in this group Like all metals they are ductile, malleable, and conduct electricity, and heat Their oxidation states are variable This group contains iron, cobalt, and nickel, the only elements known to produce and electric field They have low ionization energies. They have positive oxidation states. They have high boiling and melting points.

aromatic compounds. Structure of benzene. Stability of aromatic compounds. Hückel's rule. Naphthalene, anthracene and phenanthrene. Derivatives of benzene: toluene, xylenes, ethylbenzene, vinylbenzene (styrene). Aryl and aralkyl groups. 3,4-benzpyrene and carcinogens.

~Benzene~ •Benzene (C6H6): Simplest aromatic hydrocarbon (or arene) •Four degrees of unsaturation( a highly unsaturated hydrocarbon) • Whereas unsaturated hydrocarbons such as alkenes, alkynes and dienes readily undergo addition reactions, benzene does not... ie c6h6--Br2--> NR! but C2H4--Br2-->addition product Observations: -toxic but w/ added methyl (TOULENE) it becomes nontoxic (water soluble) - reacts slowly with Br2 to form bromobenzene: substitution product. - Addition products are not observed. C-C bond length: double shorter duh -Reasonance hybrid! How to tell if Aromatic:(HUCKNEL'S RULE): A planar cyclic molecule with alternating double and single bonds is aromatic if it has (4n+2) π electrons (n is a small integer) Benzene: n=1 ⇒ (4n+2)=6 electrons Benzene is stable and the electrons are delocalized.. ie benz 3 double bonds, six pi e-... if n=2, then 10 pi e- ---if there are 4n pi e-, compound is antiaromatic and behave likeordinary alkenes Nomenclature: alphabetically important, overall lowest numbers *if bonds are: 1,2: ortho (o). 1,3: meta(m). 1,4 para(p)... or you can just say like 1,4-dibromobenzene not p-dibromobenzene.