Final Exam Notes - Review Flashcards

Chemical bonding and structure Models

• Bond making releases energy ( exothermic reaction)

• Bond breaking requires energy.

• Intra-molecular bonds are bonds between atoms within a molecule:

  • Ionic bonds = between ions (e.g., NaCl → Na^+ Cl^-, where Na is 1+ and Cl is 7-).

  • Covalent bonds = share an electron pair (e.g., H. + .H → H:H or H_2).

  • Dative covalent bonds = one atom donates an electron pair into a covalent bond (e.g., H_2O + H^+ → [H:O:H]^+H Hydronium ion).

• Inter-molecular forces are bonds between different molecules; these forces are much weaker:

  • Van der Waals forces = non-polar molecules, permanent electrostatic forces.

  • Dipole-Dipole interactions = caused by the presence of polar bonds.

  • Hydrogen bonding = a very strong dipole-dipole bonding caused when hydrogen binds with an electronegative atom.

Covalent Bond

• A covalent bond is a type of intra-molecular bond where atoms share an electron pair (e.g., H. + .H → H:H or H_2).

• Covalent bonds can be sigma (\sigma), pi (\pi), or delta (\delta).

  • Sigma (\sigma) = found in single bonds; formed by head-to-head overlap of atomic orbitals (2 s-orbitals, s+p-orbital, or p+p-orbital).

  • Pi (\pi) = found in double bonds; weaker than \sigma bonds; found in unsaturated compounds; sharing 2 P orbitals.

  • Delta (\delta) = found in quadruple bonds; delta orbital formed when 2 d orbitals overlap.

• Properties of covalent bonds:

  • Low melting and boiling points.

  • Poor electrical conductors.

  • Mostly insoluble in water.

  • Can be gases, liquids, or solids.

  • Bond energy, polarization, bond length, electronegativity, bond order, bond angle (the angle between a pair of bonds with an atom in common).

Valence Bond Theory

• Valence bond theory states that the overlap between 2 atomic orbitals forms a covalent bond (which can be dative or covalent).

Hybridization of Atomic Orbitals

• Atoms hold electrons in orbitals: 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6, …

• Hybridization is the phenomenon of promoting electrons up an orbital to allow it to form a bond (e.g., sp3 hybridization).

• Types of hybridization:

  • sp = one electron from s-orbital to p-orbital.

  • sp2 = one electron from s-orbital to p-orbital which already has 2 electrons.

  • sp3 = one electron from s-orbital to p-orbital which already has 3 electrons.

• The atomic orbitals have appropriate shapes and steric direction.

• Overlap of orbitals leads to covalent bond formation.

• Only 2 atoms can participate in a covalent bond; covalent bonds have characteristic saturability.

• Molecules of the compounds can have single or multiple (double or triple) bonds.

Coordinate Covalent Bond (Dative Bond)

• One atom donates an electron pair into a covalent bond (e.g., H_2O + H^+ → [H:O:H]^+H Hydronium ion).

• It is an intra-molecular bond.

• In a dative bond, the tail of the arrow indicates the source of the pair of electrons: C: + D → C:D (or C → D).

• In order for a dative bond to form, the acceptor atom must be electron deficient.

• All bonds have the same length and same average bond enthalpy.

• Example: Formation of Ammonium ion (NH4^+) from Ammonia (NH3) and Hydrogen ion (H^+).

Delocalized Covalent Bond

• Electrons belonging to certain molecules, not attracted to certain molecules or bonds, are said to be delocalized.

• These can be imagined as clouds of electrons.

• In some molecules, electrons delocalize, thus stabilizing the molecule.

• Examples: cyclic and acyclic conjugated systems, lipids, carotene (10 conjugated bonds), lycopene (11 conjugated bonds).

• The number of conjugated bonds & light with less energy + longer wavelength may be absorbed.

• Conjugation = the overlap of 2 p-orbitals across a sigma bond.

Polar and Non-Polar Covalent Bonds

• Non-polar = a covalent bond between the same atoms; atoms attract to an equal degree, thus sharing electron pairs (e.g., H2, O2, N2, Cl2, F_2).

• Polar covalent bond = formed between different atoms with close electronegativities; shared electron pair is towards the atom of greater electronegativity (e.g., H-Cl). Cl has greater electronegativity than H atom.

• The degree of iconicity may range from 0 to 1.

• Degree of ionization = the proportion of neutral particles that are ionized.

Ionic Bond

• Ionic bond results from the transfer of electrons from a metal to a non-metal.

• Arises from elements with low electronegativity reacting with elements with high electronegativity.

• Example = table salt (sodium chloride); sodium gives up its one outer shell electron completely to chlorine, which needs one electron to fill its shell.

• Ionic compounds are made of ions and are called salts or crystals.

• Properties of ionic compounds:

  • High melting and boiling points.

  • In aqueous solution, they conduct electricity.

  • Hard solid at 22°C.

  • Non-conductors of electricity in solid state.

  • Good conductors in liquid state or dissolved water.

• Cations = donation of electrons; Anions = acceptance of electrons.

• Held together by electrostatic attraction of opposite charges.

Basic Characteristics of Molecules

• Geometry: Conjugated structure = flat; Incomplete = not flat. Example: Conjugated structures are flat. Systems where conjugation is not complete will not be flat (these exceptions are rare).

• Bond length gives some idea of the degree of conjugation; the greater the conjugation, the more stable the structure.

• Bond angles: It is an angle formed between 3 atoms across at least 2 bonds. (e.g., Beryllium chloride, Boron trichloride, Water H_2O).

• Bond angles depend on the type of hybridization.

• Bond length: This is the average distance between nuclei of 2 bonded atoms in a molecule; measured in nanometers. The bond length between 2 equal atoms will vary depending on the molecular environment of atoms.

• Bond energy: Is the measure of bond strength in a chemical bond. Energy needed to break bond or energy liberated when bond is formed.

Dipole Moment

• Occurs when there is a separation in charge.

• Determines the direction of electron shifting in bond.

• Arises from difference in electronegativity; the greater the difference, the greater the dipole movement.

• Dipole moment is a vector quantity which determines the direction of shifting of electron in a bond. Its direction is towards the more electronegative atom or from negative to positive charge.

• \mu = l . q (q = size of charge, l = length of dipole = to distance between the centers of positive + negative charges).

Non-Covalent Interactions

• Non-covalent interaction differs from covalent bonds as it does not involve the sharing of electrons. It involves more dispersed variations of electromagnetic interactions between molecules.

• Examples of non-covalent interactions are: Hydrogen bonds, ionic bonds, Van der Waals forces, hydrophobic bonds.

• Intramolecular forces are bonds between atoms within a molecule such as: Ionic bonds, covalent bonds, dative covalent bonds.

• Intermolecular forces are bonds between different molecules; much weaker than intramolecular forces. Examples: forces:

  • Van der Waals forces = non-polar molecules, permanent electrostatic.

  • Dipole-Dipole interactions = caused by presence of polar bonds.

  • Hydrogen bonds = very strong dipole-dipole bonding.

Hydrogen Bond

• Hydrogen bond is a special case of dipole-dipole attraction when H bonds with an electronegative atom (EQW) leading to a highly covalent polar bond.

• Weaker than covalent or ionic bonds, but stronger than van der Waals.

• Occurs in organic molecules (e.g., DNA) and inorganic molecules (e.g., water).

• H bonds hold the 2 complimentary DNA strands together; forms between nitrogenous bases.

• Also occur in folding of peptide chain and determine secondary, tertiary + quaternary protein structure.

• Also responsible for solvent properties of water.

• H bonding has a considerable effect on physical properties such as melting point boiling point vaporization and heat of fusion.

Ionic Bond

• Chemical bonding that requires electrostatic attraction between 2 oppositely charged ions.

• Complete transfer of electrons from metal to non-metal.

• In solid form arranged as crystal lattice.

• Important in binding of ligands to receptor sites in some enzymatic reactions.

Van der Waals Forces

• Attractive forces between polar molecules.

  • Permanent dipole-permanent dipole: weak electrostatic attractive forces between permanent dipoles; bind substances to active and receptor sites.

  • Dipole-induced dipole interactions: Polar molecule causes an inductive dipole into a non-polar molecule; 2 molecules interact by orientation (e.g., interaction of water and bromine).

  • Induced dipole-induced dipole (London dispersion forces): Found in non-polar molecules; dependent on number of electrons in atom; LDF increase in with & in polarizability which is dependent on number of electrons found in atom. Transient dipoles induce dipoles in neighboring molecules; have weak electrostatic forces; increasing the molecule size increases number of LDFs.

Biological Meaning of Non-Covalent Interactions

• Non-covalent bonds determine the shape of many biological molecule.

• It also stabilizes the complexes that contain 2 or more molecules.

• 4 main types in biological systems: H-bonding, ionic, vanderwaal, hydrophobic.

Electronic Effects in Organic Molecules

• Include: Inductive effect, Electrometer effect, Mesmeric effect etc.

• These show differences in electron behavior when elements other than carbon and Hydrogen participate in molecular bonds Inductive and mesmeric effects.

• Inductive effect = Permanent effect where Sigma electrons are displaced towards more electronegative species. It is a weak effect; it can be transmitted through a chain of carbon.

  • I- effect = It is shown by electron attracting or electron withdrawing species (e.g., A ← C → X).

  • +I effect = shown by electron donating species.

• Mesmeric effect = : It involves complete transfer of Telecoms or lone pair of electron towards more electronegative Atom - similar to resonance effect.

Resonance Structure

• Chemical phenomenon observed in the characteristic compounds having double bonds in organic compounds.

• Benzene → 2 resonance structure; real structure when combined, the C-C bonds are neither single or double bonds but are somewhat intermediate (between.

Reaction Types According to the Structural Change in Reagent Molecules

• Synthesis reaction: 2 or more elements combine to form more complex structure.

• Decomposition reaction: Chemical species break down to form simpler element burnings.

• Combustion: organic molecule → CO2 + H2O

• Exothermic reaction.

Reactions of Addition, Elimination, Substitution, Rearrangement and Oxidation-Reduction Reactions

• Addition = reactions in which 2 molecules react to form a single Product.

• Elimination = Reaction in which a small molecule is removed from adjacent carbon atoms resulting in formation of multiple bonds between them.

• Substitution = replacement, atom or functional group in molecule replaced by different atoms g-functional group.

• Rearrangement = form an isomer, migration of atoms or groups of atoms from one position to another within a molecule.

• Oxidation: dehydrogenation (-2H), oxygenation (+O); Loss of electron.

• Reduction: hydrogenation (+2H) gain q electron.

Types of Reagents

• Radicals = very reactive species having unpaired electrons.

• Nucleophiles = -reagents containing an atom having unshaved I love pair of electrons; Nvcleophe = electron rich and seeks electron deficient sites (nucleus Loving).

• Electrophiles = -Positively charged or neutral species, deficient g- electrons + accept apair g- electrons, Electron Loving.

Activated Complex and Transition State

• Activated complex = a high energy state that a reaction goes through in order to change reactants to products.

• Transition state = -short - lived configuration of atoms at a local energy maximum in a reaction energy diagram; has partial bonds + extremely short lifetime; cannot be isolated.

Complex Chemical Reactions

• Competitive = 2 or more starting materials compete with other starting materials to give side products in addition to main products.

• Successive.

• Reversible = Where conversions of reactants to products and products to reactants occur simultaneously.

• Chain reactions = series of chemical reactions where products of one reaction contributes to reactants of another reaction Allowing reaction to continue with minimal of no outside influence.

Heat of Chemical Reactions

• Heat of chemical reaction = the amount of heat that must be added or removed during chemical reaction in order to keep all substances at the same temp.

• Endothermic = positive heat.

• Exothermic = negative heat.

Isomerism

• Isomer = same molecular formula but different chemical structure.

• Structural Isomer = same molecular formula but different structural formula.

• Optical Isomer = isomers that cannot be super imposed on each other, have to have a chiral centre.

• Constitutional (chain + positional Isomers).

• Chain = same molecular formula but different structures g- carbon chain; Number of possible isomers increase with longer hydrocarbons.

• Positional = - same molecular formula but different positions of side chain group.

• stereoisomerism.

• Same molecular formula, same structural formula , but different arrangement of atoms or group of atoms in space.

• 2 subtypes = -① Elz isomer Geometrical ↳ compound must have G-C double covalent bond ↳ Each carbon in c=c bond must be bonded to 2 deft group -② Optical Isomer ↳ compound has same molecular formula, same structural formula but different position of atoms or groups g- atoms in space ↳ only possible on compounds with chiral carbon ↳ optical isomers also called enantiomers ↳ Always drawn in tetrahedral shape ↳ To distinguish optical isomer you can pass plane -polarized light from their solutions ↳ Tetaric acid was og optical isomer.

Configurational Isomerism

• Configuration = spatial arrangement of atoms around a chiral center in the molecule.

• Configurational isomers = 2 molecules with same constitution but different configuration.

Enantiomerism of Compounds with Chiral Centers

• Enantiomers rotate plane of polarized light to right orleft In opn.ca isomers.

• In a pair of optically active enantiomers, each enantiomer will rotate the plane of polarized light in equal and opposite directions. Therefore referred to as e) andy enantiomers depending on direction of Observed direction.

Fischer Projections and Wedge - ana -dash Structural Formulas

• Fischer projection = A 2D representation of a 3D organic molecule by projection All bonds depicted as horizontal or vertical lines it ↳ A horizontal bonds to chiral carbon represents wedge c = ☐ ↳ Dotted lines can bre represented by vertical bonds (I :-).

Sigma -diastereomerism g- Molecules with More Than Chiral Centers

• Epimers - more +"" " atoms \ - Epimers are diaotereome.rs that have more than one chiral center but differ from each other in absolute configuration at only 1 chiral center - The molecular structure of glucose and mannose differs in configuration of C2 atom ↳ Therefore are a-epimers.

• Pi -diastereomerismcgeomekc isomerism -2 chiral carbons with double bond called Di -dices tomes.

• Diastereomers have more than 1 chiral carbon.

• If 2 Chiral carbons have double bond to each other then compound called pi -diastercomes.

Conformational Isomerism

• Rotation isomerism.

• Conformation =Different spatial orientation of atoms in molecule that result from rotations or twisting about single bonds.

• Many conformers at carbon-carbon single bond Newman projection formulae visualizes conformations of a c-C Sigma bond from front to back - Front carbon represented by a dot - Back carbon represented as a circle

• Basic types of conformations: Eclipsed conformation; Staggered conformation; sawhorse; Newnan projection; Chair conformer / free of angle strain; Boat conformer.

Types of Dispersion

• True solutions: homogenous mixture; Particle size less than 1nm.

• Colloidal micro solution f ↳ heterogenous mixture; Particle size lessthan inn.< loonm 1- 100hm; easily visible.

• Suspensions: heterogenous mixture; Particle size < 100hm; easily visible.

Molecular and Ionic Solutions

• A solution = a homogenous mixture composed of 2 or more substances

• When a substance is dissolved in a liquid solvent ↳ liquid - solvent ↳ substance -solute.

• In molecular solutions bonds are not broken as they are in non-molecular solutions.

• Ionic solution = is a solution containing ions ↳ Formed by dissolving ionic compounds in solvents typically water.

Mechanism of Dissolution

• Dissolution = solute un gas, liquid or solid phase dissolve in solvent to produce solution solubility and concentration g- solutions /given solute.

Concentration

• The amount of solute that is dissolved in a given solvent or solution.

Temperature Dependence of Solubility

• The solubility of a solute in a solvent is dependent on temp.

• Solid solutions = increased solubility with A temp = T solubility; Barium Nitrate = TT so does solubility in NaOH = Tf. :. Sob Nacl = fairly independent g- T.

Solutions of Electrolytes

• Strong electrolytes: All ionic substances that dissolve; completely break into ions in water; strong acids, bases + soluble salts.

• Weak electrolytes: Molecular substances that dissolve; Partially break into ions in water; weak acids +bases.

• Substances that dissolve in water either electrolytes or non electrolytes; Non electrolytes = form non-conducting solutions chis solve as molecule); Electrolytes = form conducting solutions Clisson as ions.

• Oswald> dissolution Law: If all acids contain same active ion , chemical affinity must correspond to number of these active ions in solution.

Colloid Dispersed Systems

• What are colloids? ↳ mixture of substances; when a substance divided into minute particles in a second substance Cdispersion medium -particle size = I -100 nm.

• colloid system = Mixture of substances which are so fine grained that constituents cannot be seen under a microscope To produce system dispersed phase must be insoluble in dispersing medium.

Classification of Colloids

• Lyophilic sons = solvent loving - If dispersing medium is water the colloidal particles hydrate particles may carry electrical charge - In acid = particles (+) charge . - in alkaline= particles f) Charge - Reversible sols. E.g / Arabic gum, Gelatin.

• Dispersed particles have charge (+) or (-). E.gl metal sols (gold, platinum).

Kinetics + Optical Properties of Colloids

• Collisions of collard with dispersion medium molecub passes kinetic energy to colloid; KirlCS_ striking from all sides; zig zag motion a conoid parties; Ultramicroscope examines colloid dispersion.

• continuous collisions w/ colloid particles + molecules of dispersion medium; PASSES KINETC ENERGY to colloid particles by striking it from all sides; Results in ZIGZAG motion of colloid particles known as Brownian movement.

• Determined by: -size of colloid particles; -wavelength of light beam ↳ Some colloids have characteristic colours by absorbing specific wavelengths of light Eg/Asos = yellow Fe@H)z= Brick colour.

• If beam of light passes through colloid so. → light is reflected (scatters) → by colloid particles → Path g- light observed. In true solutions -Path of light cannot be seen. : colloid particles reflect light ↳ can view path of light True sd = no path of light viewed.

Stability + Coagulation a- Colloid Solutions

• Important factor of stability = equal charge of particles stability of bounds ↳ hepubion of same -charge colloid particles; tragulatim ↳ process of forming pnnecipit.

• Stability: remove chase of ate of conoids of particles by heating ↳ ① removing charge of particles ②Precipitation by adding by heating = sticking = sedimentation another colloid particle with ② Precipitation by t another opos.ie charge colloid Wl opposite charged particles ③+ soluble ions - -precipitate ③ t soluble ions = precipitate- ④Mutual neutralization ④ mutual neutralization of charged of particles particles = partial attraction . coagulate precipitate.

Biological Meaning

Abundant in environment. -clouds -milk -blood, In bimolecular + physiological interactions that sustain life. -blood -lymph -cellular protoplasm / colloid Socs., Soap + detergent, Medicine + cosmetics

Couigative Properties

• particles have TSA; C. particles not simple molecules =Are aggregations of molecules ↳ colloidal Sd . -No particles v. small compound to true Sor., Relative to in vapour pressure; Elevation in boiling point; Depression in freezing point; to osmotic pressure.

Biological Significance

• Abundant in nature: -Clouds -blood - milk.Bimolecular + physiological interactions that sustain life, -blood -lymph -cellular protoplasm, Plant sap, Medicine + make - up, soap + detergent.
  VP of Liquid = pressure exerted from vapour above liquid surface at equilibrium -dependent on - tempt nature of liquid:

• Diffusion: = net movement of molecules from an area of high concentration to an area of low concentration -due to random movement g- police.

• Osmosis = Spontaneous net movement of molecules through a selectively semi - permeable membrane from a region a- T water potential to the region of t water potential.

• Osmotic pressure = measure of tendency of a solution to take in a pure solvent by osmosis.

Chemical Kinetics + Thermodynamics

Chemical kinetics:

• study of rate of reactions. If reactions free energ is in and negative = Fast reaction. catalyst increases rate of reaction. catalyst provide an alternate pathway from reactant to product chemical reaction rate. Reaction rate i -The rates of change in concentrations or amounts of reactants / products per unit.

• Change on concentration measured in polls Thermodynamics + kinetics affect rate cnaiieinia.cm?cEnfe9-u.anb /prom" . Rate = -DCI = T AT 1\¥ ↳ -change in concentrationÑreactant1timÑ ↳ change in concentration of product /time concentration dependence on the reaction rate. As reaction goes on , rate to - At start of reaction : -A reaction rate -due to t concentration of reactants - As time goes on , reactants react = form product - As reaction progresses there are less successful collisions between reactants - After reaction taken place : - rate of reaction = constant -As all reactants have formed products

• Kinetic Equation Rate Law : V = KIA] " (B) b k = rate constant at a fixed temperature , theT U = t reaction rate At B = molar concentrations of species At B a + D= stoichiometric coefficients of balanced equation.

• Determining Step + Temperature dependence on rate of reaction Increase in temp =P in rate of reaction Arrhenius equation - Arrhenius equation give quantalive basis of relationship between activation energy and rate at which reaction proceeds Explains role T with Tlemp t - rateto with Tfa Activation Equation i teneras → Average kinetic rate constant-K = Ae -Ea/Rt energy