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POLAR BONDS AND MOLECULES NOTES 16.3. Covalent Bonds  bond in which two atoms share a pair of electrons. 1. Single bond = 1 shared pair of electron 2.

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Presentation on theme: "POLAR BONDS AND MOLECULES NOTES 16.3. Covalent Bonds  bond in which two atoms share a pair of electrons. 1. Single bond = 1 shared pair of electron 2."— Presentation transcript:


2 Covalent Bonds  bond in which two atoms share a pair of electrons. 1. Single bond = 1 shared pair of electron 2. Double bonds = 2 shared pairs 3. Triple bonds = 3 shared pairs

3 Bond Polarity: 1. bonding pairs of electrons in covalent bonds are pulled between the nuclei of the atoms sharing the electrons. 2. When bonds are pulled equally the bond is a nonpolar covalent (occurs between like atoms.)

4 3. When a covalent bond occurs between different atoms then electrons are shared unequally which is a polar bond.

5 a. This leads to partial charges on the atoms (  - or  + ) b. The polarity of a bond can be shown with an arrow pointing to the negative side. c. Table 16.4 p. 462

6 Polar Molecules - when one end of a molecule is slightly negative and other is slightly positive. - Water is polar because the way the bonds cause the molecule to bend. You have partial charges on two sides of molecule.

7 - CO 2 is not polar because the double bonds keep the molecule linear and the charges cancel.

8 Attractions  molecules are attracted to each other through a variety of forces.

9 1. These forces are responsible for determining whether a compound is a gas, liquid or solid. 2. Van der Waals forces consist of two possible types.

10 a. Dispersion  caused by the movement of electrons. Increases as the # of electrons increase. b. Dipole Interactions  occurs when polar molecules are attracted to one another.

11 1. Hydrogen bonds occur when H already in a polar compound bond with a partial negative of another molecule. - extremely important in determining the properties of water and biological molecules such as proteins.

12 Intermolecular Attractions  The physical properties of a compound depend on the type of bonding it displays.


14 Water Molecule: 1. It is a triatomic molecule with two polar covalent bonds (H – O). 2. It has a bent shape leading to a partially + (δ+) and – (δ) ends to the molecule.

15 3. Because of the polarity the molecule will form Hydrogen bonds with other water molecules.

16 Water Molecule -- Polarity - Water is polar because the way the bonds cause the molecule to bend. You have partial charges on two sides of molecule.

17  H – bonding gives H 2 O many of its properties. 1) high surface tension 2) low vapor pressure 3) high specific heat 4) high boiling point

18 Surface Properties: 1. Water molecules experience an uneven attraction the molecules are hydrogen- bonded on only one side of the drop. The molecules pull toward the body of the liquid.

19 2. This pull is called surface tension. 3. A liquid that has strong intermolecular forces has high surface tension. 4. You can decrease surface tension by adding a surfactant, an agent such as soap that interferes with the hydrogen bonds.

20 Specific Heat Capacity  water has a high specific heat capacity. Ability to absorb heat without changing temperatures.


22 Evaporation and Condensation: 1. The hydrogen bonds of water helps hold the molecules together, and therefore requiring a high amount of energy to break the bonds to turn to a vapor.

23 2. The less hydrogen bonding the easier to vaporize. 3. The reverse of evaporation is condensation, when water condenses it releases energy (heat). 4. Temperatures in the tropics would be much higher if water didn’t absorb heat.

24 5. Temperatures in the polar regions would be much lower if water vapor did not release heat when condensing out of the air.

25 Ice: A typical liquid cools, it contracts slightly. Its density increases because its volume decreases. The solid would sink because its density is higher than the liquid.

26  As water cools it acts like a typical liquid, until it reaches 4 o C then the density begins to decrease. As Ice forms at 0 o C the volume expands and it has lower density than the surrounding water.


28 Solvents and Solutes 1. Chemically pure water never exists in nature, because water dissolves so many substances (Universal solvent). 2. Water samples containing dissolved substances are called aqueous solutions.

29 In a solution, the dissolving medium is the solvent. In a solution, the particles dissolved are the solutes. 3. Solutions are homogeneous mixtures.

30 4. Solvents and solutes can be solids, liquids, and gases. 5. Ionic compounds and Polar covalent molecules dissolve most readily in water, but nonpolar covalent do not.

31 The Solution Process: 1. As you place a solute in a solvent the particles begin to collide with one another. The solvent attracts the solute particles until substance is dissolved.

32 2. In some ionic compounds, the solvent can’t break the ionic bonds and the salt doesn’t dissolve. 3. Polar solvents dissolve ionic and polar molecules, nonpolar solvents dissolve nonpolar compounds.

33 Raises boiling point – Salt in water Lowers freezing point – Salt on road Solute added to solution:

34 Electrolytes & Nonelectrolytes 1. Compounds that conduct an electric current in aqueous solution or molten state are electrolytes. - All ionic compounds are electrolytes.

35  Compounds that do not conduct electric current are nonelectrolytes. - Most molecular compounds and compounds of carbon are nonelectrolytes.

36 3. Some very polar molecular compounds are nonelectrolytes in pure state, but electrolytes in an aqueous state. 4. You can have strong or weak electrolytes. Depends on how well the solute dissolves into ions.

37 Water of Hydration 1. Water molecules are an integral part of crystal structure; this is called water of hydration. Also, called a hydrate. 2. Effloresce  the ability to lose the water hydration.

38 3. Hygroscropic  the ability to remove water from the air. a. These are used as drying agents (desiccants). Ex. Silica gel b. Agents that became wet from solutions from H2O in air when exposed to air are deliquescent.


40 Suspensions  mixtures from which particles settle out of solution upon standing.  Differs from solution because component parts are much larger.

41 Colloids  contain particles that are intermediate in size between suspensions and solutions. 1. The properties of colloids differ from both suspensions and solutions.

42 2. Colloids are cloudy in appearance when concentrated but clear to almost clear when diluted. 3. Particles do not settle of a mixture. 4. Colloids exhibit the Tyndall Effect, the scattering of visible light.

43  Emulsions  dispersions of liquids in liquids. An emulsifying agent is essential for the formation of an emulsion.  (Ex. Mayo → vinegar, oil, and egg)


45 Solution Formation 1. Solutions are homogeneous mixtures and can be solids, liquids or gases.

46 2. Factors that affect how fast a substance dissolves. a. agitation b. temperature c. surface area  the smaller the particle, the faster it dissolves.

47 Solubility  the amount that dissolves in a given quantity of a solvent at a given temperature to produce a saturated solution.  Particles can move from solid to a solvated state and back to a solid again.

48  This is a saturated solution (contains the maximum amount of solvent.)  A solution that contains less solute than a saturated solution is unsaturated.

49 2. Two liquids are said to be miscible if they dissolve in each other. (ex. Water & ethanol) 3. Liquids that are insoluble in each other are immiscible (ex. Oil & Vinegar)

50 Factors Affecting Solubility 1. Temperature a. for most solids as temperature increases solubility increases. b. For most gases as temperature decreases solubility increases.

51 2. Pressure a. Gas solubility increases as the partial pressure of gas above the solution increases (Ex. Carbonated drinks  contain dissolved CO2 in H2O) and decreases as pressure decreases.

52 3. Supersaturated solution  contains more solute than it should theoretically continue to hold. a. Crystallization of the solution can occur by adding a small crystal (seed crystal).


54 Molarity 1. Concentration is a measure of the amount of solute that is dissolved in a given quantity of solvent. a. dilute solution  low concentration b. concentrated  high concentration

55 2. Molarity (M)  the number of moles of a solute dissolved per liter of solution. a. also known as molar concentration and read as “molar”. b. To calculate the molarity of any solution - calculate the number of moles in 1 L of the solution.

56 Molarity = moles of solute Liters of solution

57 3. Making Dilution: a. by adding solvent to a solution you can lower its molarity. 1) moles of solute do not change.

58 4. Percent Solutions a. If both solute and solvent are liquids, a convenient way to make a solution is to measure volumes. 1)If 20 mL of pure alcohol is diluted with water to a total volume and 100 mL the final solution is 20% alcohol by volume.

59 Percent volume = Volume of Solute x 100% Solution Volume


61 Decrease in Vapor Pressure 1. Properties of solutions differ from those of the pure solvent. 2. Properties that depend on the number of particles dissolved in a given mass of solvent are called colligative properties.

62 3. 3 Important colligative properties of solutions: a. vapor pressure lowering b. boiling point elevation c. freezing point depression

63 4. A solution that contains a nonvolatile solute always has a lower vapor pressure than the solvent. (Volatile  easily vaporized.)

64 Boiling Point Elevation: 1. By adding a nonvolatile solute would increase the boiling point. 2. Attractive forces occur between the solvent and solute therefore you need more energy to overcome these forces.

65 Freezing Point Depression: 1. When a substance freezes the particles of the solid take on an orderly pattern. The presence of a solute disrupts this. Therefore more energy must be withdrawn. 2. The more solute you add the lower the freezing point. (Ex. Salt and water)

66 Percent by Mass

67  A solute in solution is the number of grams of solute dissolved in 100 g of solution. Percent by Mass = __Mass of solute__x 100% Mass of solute + mass of solvent

68 Example:  % by mass = 10 g NaOH___ x 100% 10 g NaOH + 90 g H2O = 10 % NaOH

69 Molarity (M)  Number of moles of solute in one liter of solution.  Molarity = # of moles of solute # of liters of solution

70 Example:  If 0.500 moles of NaOH is dissolved in 1.00 L of solution, what is the molarity that is produced?  M = 0.500 mole NaOH 1.00 L = 0.500 M NaOH

71  Moles of solute = M1 x V1 = M2 x V2

72 Molality (m)  Concentration of a solution expressed in moles of solute per kilogram of solvent.  Molality = # of moles solute__ Mass of solvent (kg)

73 Example:  If 8.50 g of ammonia, which is one – half mole of ammonia, is dissolved in exactly 1 kg of water, what molality is produced?  m = 0.500 mole NH3 1 kg H2O = 0.500 m NH3

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