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MIDTERM POST LABORATORY DISCUSSION

MIDTERM POST LABORATORY DISCUSSION. Adapted from ppt created by Andrea D. Leonard University of Louisiana at Lafayette. POLARITY. Properties of Organic Compounds Polarity. A covalent bond is nonpolar when two atoms of identical or similar electronegativity are bonded.

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MIDTERM POST LABORATORY DISCUSSION

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  1. MIDTERM POST LABORATORY DISCUSSION Adapted from ppt created by Andrea D. Leonard University of Louisiana at Lafayette

  2. POLARITY

  3. Properties of Organic CompoundsPolarity • A covalent bond is nonpolar when two atoms of • identical or similar electronegativity are bonded. • Thus, C–C and C–H bonds are nonpolar bonds. • A covalent bond is polar when atoms of different • electronegativity are bonded. • Thus, bonds between carbon and N, O, and the • halogens are polar bonds.

  4. Properties of Organic CompoundsPolarity • If a single bond is polar, • the molecule is polar because it contains a net dipole. • Hydrocarbons contain • only nonpolar C–C and • C–H bonds, so they are • nonpolar molecules.

  5. Properties of Organic CompoundsPolarity • If the individual polar • bonds (dipoles) cancel • in a molecule, the • molecule is nonpolar. • If the individual bond • dipolesdo not cancel, the • molecule is polar.

  6. INTERMOLECULAR FORCES

  7. Intermolecular Forces

  8. BOILING POINT AND MELTING POINT

  9. Intermolecular ForcesBoiling Point and Melting Point • The boiling point is the temperature at which a • liquid is converted to the gas phase. • The melting point is the temperature at which a • solid is converted to the liquid phase. • The stronger the intermolecular forces, the higher • the boiling point and melting point.

  10. Intermolecular ForcesBoiling Point and Melting Point

  11. Intermolecular ForcesBoiling Point and Melting Point • Both propane and butane have London • dispersion forces and nonpolar bonds. • In this case, the larger molecule will have stronger • attractive forces.

  12. The Liquid StateVapor Pressure • Evaporation is the conversion of liquids into the • gas phase. • Evaporation is endothermic—it absorbs heat from • the surroundings. • Condensation is the conversion of gases into the • liquid phase. • Condensation is exothermic—it gives off heat to • the surroundings.

  13. The Liquid StateVapor Pressure • Vapor pressure is the pressure exerted by gas • molecules in equilibrium with the liquid phase. • Vapor pressure increases with increasing • temperature.

  14. The Liquid StateVapor Pressure • The stronger the intermolecular forces, the lower • the vapor pressure at a given temperature.

  15. The Liquid StateViscosity and Surface Tension Viscosity is a measure of a fluid’s resistance to flow freely. • A viscous liquid feels “thick.” • Compounds with strong intermolecular forces • tend to be more viscous than compounds with weaker forces. • Substances composed of large molecules tend to be more viscous, too, because large molecules do not slide past each other as freely.

  16. The Liquid StateViscosity and Surface Tension Surface tension is a measure of the resistance of a liquid to spread out. • Interior molecules in a • liquid are surrounded by • intermolecular forces on • all sides. • Surface molecules only • experience intermolecular • forces from the sides and • from below.

  17. The Liquid StateViscosity and Surface Tension • The stronger the intermolecular forces, the stronger • the surface molecules are pulled down toward the • interior of a liquid and the higher the surface tension. • Water has a very high surface tension because • of its strong intermolecular hydrogen bonding. • Small objects can seem to “float” on the surface • of water. • These normally heavy objects and are held up by • the surface tension only.

  18. The Solid StateTypes of Solids • Solids can be either crystalline or amorphous. • A crystalline solid has a regular arrangement of • particles—atoms, molecules, or ions—with a repeating structure. • An amorphous solid has no regular arrangement of its closely packed particles. • There are four different types of crystalline solids—ionic, molecular, network, and metallic.

  19. The Solid StateCrystalline Solids • An ionic solid is composed • of oppositely charged ions • (NaCl). • A molecular solid is • composed of individual • molecules arranged • regularly (H2O).

  20. The Solid StateCrystalline Solids • A network solid is composed • of a vast number of atoms • covalently bonded together • (SiO2). • A metallic solid is a lattice • of metal cations • surrounded by a cloud of • e− that move freely (Cu).

  21. The Solid StateAmorphous Solids Amorphous solids have no regular arrangement of their particles. • They can be formed when liquids cool too quickly • for regular crystal formation. • Very large covalent molecules tend to form • amorphous solids, because they can become folded and intertwined. • Examples include rubber, glass, and plastic.

  22. Energy and Phase ChangesConverting a Solid to a Liquid liquid water solid water The amount of energy needed to melt 1 gram of a substance is called its heat of fusion.

  23. Energy and Phase ChangesConverting a Liquid to a Gas gaseous water liquid water The amount of energy needed to vaporize 1 gram of a substance is called its heat of vaporization.

  24. Energy and Phase ChangesConverting a Solid to a Gas gaseous CO2 solid CO2

  25. SOLUBILITY

  26. SolutionsIntroduction A solution is a homogeneous mixture that contains small particles. Liquid solutions are transparent. A colloid is a homogeneous mixture with larger particles, often having an opaque appearance. Solutions consist of two parts: • The solute is the substance present in a lesser amount. • The solvent is the substance present in a larger amount. An aqueous solution has water as the solvent.

  27. SolutionsIntroduction Three different types of solutions: a solution of gases (O2, CO2, and N2) an aqueous solution of NaCl (a solid in a liquid) Hg(l) dissolved in Ag(s) (a liquid in a solid)

  28. SolutionsIntroduction • A substance that conducts • an electric current in water • is called an electrolyte. • A substance that does • not conduct an electric • current in water is • called a nonelectrolyte. NaCl(aq) dissociates into Na+(aq) and Cl−(aq) H2O2 does not dissociate

  29. SolubilityGeneral Features Solubility is the amount of solute that dissolves in a given amount of solvent. • It is usually reported in grams of solute per 100 mL • of solution (g/100 mL). • A saturated solution contains the maximum • number of grams of solute that can dissolve. • An unsaturated solution contains less than the • maximum number of grams of solute that can dissolve.

  30. SolubilityBasic Principles Solubility can be summed up as “like dissolves like.” • Most ionic and polar covalent compounds are • soluble in water, a polar solvent.

  31. SolubilityBasic Principles • Small neutral molecules with O or N atoms that can hydrogen bond to water are water soluble. Ethanol can hydrogen bond to water.

  32. SolubilityBasic Principles • Nonpolar compounds are soluble in nonpolar • solvents (i.e., like dissolves like). • Octane (C8H18) dissolves in CCl4 because both are • nonpolar liquids that exhibit only London dispersion • forces. octane + CCl4 octane CCl4

  33. Properties of Organic CompoundsSolubility • The rule of solubility is “like dissolves like.” • Most organic compounds are soluble in organic • solvents. • Hydrocarbons and other nonpolar organic • compounds are insoluble in water. • Polar organic compounds are water soluble only • if they are small and contain a N or O atom that • can hydrogen bond with water.

  34. Properties of Organic CompoundsSolubility CH3CH2CH2CH2CH2CH3 hexane CH3CH2—OH ethanol • small nonpolar molecule • no O or N present • H2O insoluble • organic solvent soluble • small polar molecule • O atom present • H2O soluble • organic solvent soluble

  35. Properties of Organic CompoundsSolubility cholesterol • very large molecule • O atom present • too many nonpolar C—C and C—H bonds • H2O insoluble • organic solvent soluble

  36. SolubilityBasic Principles • When solvation releases more energy than that • required to separate particles, the overall process • is exothermic (heat is released). • When the separation of particles requires more • energy than is released during solvation, the • process is endothermic (heat is absorbed).

  37. SolubilityIonic Compounds—Additional Principles General Rules for the Solubility of Ionic Compounds Rule[1] A compound is soluble if it contains one of the following cations: • Group 1A cations: Li+, Na+, K+, Rb+, Cs+ • Ammonium, NH4+

  38. SolubilityIonic Compounds—Additional Principles General Rules for the Solubility of Ionic Compounds Rule[2] A compound is soluble if it contains one of the following anions: • Halide: Cl−, Br−, I−, except for salts with • Ag+, Hg22+, and Pb2+ • Nitrate, NO3− • Acetate, CH3CO2− • Sulfate, SO42−, except for salts with Ba2+, • Hg22+, and Pb2+

  39. ALCOHOLS

  40. Structure and Properties of Alcohols • An alcohol contains an O atom with a bent shape • like H2O, with a bond angle of 109.5o. • Alcohols have two polar bonds, C—O and O—H, • with a bent shape, therefore it has a net dipole.

  41. Structure and Properties of Alcohols • Alcohols have an H atom bonded to an O atom, • making them capable of intermolecular hydrogen • bonding. • All of these properties give alcohols much stronger • intermolecular forces than alkanes and alkenes.

  42. Structure and Properties of Alcohols • Therefore, alcohols have higher boiling and melting • points than hydrocarbons of comparable size and • shape. stronger intermolecular forces higher boiling and melting point

  43. Structure and Properties of Alcohols • Alcohols are soluble in organic solvents. • Low molecular weight alcohols (6 C’s or less) • are soluble in water. • Higher molecular weight alcohols (6 C’s or more) • are not soluble in water. 2 C’s in chain water soluble 8 C’s in chain water insoluble

  44. ACIDS AND BASES

  45. Introduction to Acids and Bases The earliest definition was given by Arrhenius: • An acid contains a hydrogen atom and dissolves • in water to form a hydrogen ion, H+. HCl(g) H+(aq) + Cl−(aq) acid • A base contains hydroxide and dissolves in water • to form −OH. NaOH(s) Na+(aq) + −OH(aq) base

  46. Introduction to Acids and Bases • The Arrhenius definition correctly predicts the • behavior of many acids and bases. • However, this definition is limited and sometimes • inaccurate. • For example, H+does not exist in water. Instead, it • reacts with water to form the hydronium ion, H3O+. H+(aq) + H2O(l) H3O+(aq) hydronium ion: actually present in aqueous solution hydrogen ion: does not really exist in solution

  47. Introduction to Acids and Bases The Brønsted–Lowry definition is more widely used: • A Brønsted–Lowry acid is a proton (H+) donor. • A Brønsted–Lowry base is a proton (H+) acceptor. This proton is donated. HCl(g) + H2O(l) H3O+(aq) + Cl−(aq) • HCl is a Brønsted–Lowry acid because it donates • a proton to the solvent water. • H2O is a Brønsted–Lowry base because it accepts • a proton from HCl.

  48. Introduction to Acids and BasesBrønsted–Lowry Acids • A Brønsted–Lowry acid must contain a hydrogen • atom. • Common Brønsted–Lowry acids (HA): HCl hydrochloric acid H2SO4 sulfuric acid H O HBr hydrobromic acid acidic H atom H C C H O H HNO3 nitric acid acetic acid

  49. Introduction to Acids and BasesBrønsted–Lowry Acids • A monoprotic acid contains one acidic proton. HCl • A diprotic acid contains two acidic protons. H2SO4 • A triprotic acid contains three acidic protons. H3PO4 • A Brønsted–Lowry acid may be neutral or it may • carry a net positive or negative charge. HCl, H3O+, HSO4−

  50. Introduction to Acids and BasesBrønsted–Lowry Bases • A Brønsted–Lowry base is a proton acceptor, • so it must be able to form a bond to a proton. • A base must contain a lone pair of electrons that • can be used to form a new bond to the proton. This e− pair forms a new bond to a H from H2O. + H + H2O(l) H N H H N H + −OH(aq) H H Brønsted–Lowry base

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