Presentation is loading. Please wait.

Presentation is loading. Please wait.

Chapter 17 “Water and Aqueous Systems” Chemistry Golden Valley High School Stephen L. Cotton.

Similar presentations

Presentation on theme: "Chapter 17 “Water and Aqueous Systems” Chemistry Golden Valley High School Stephen L. Cotton."— Presentation transcript:

1 Chapter 17 “Water and Aqueous Systems” Chemistry Golden Valley High School Stephen L. Cotton

2 Section 17.1 Water and it’s Properties l OBJECTIVES: –Explain the high surface tension and low vapor pressure of water in terms of the structure of the water molecule and hydrogen bonding.

3 Section 17.1 Water and it’s Properties l OBJECTIVES: –Describe the structure of ice.

4 The Water Molecule l Water is a simple triatomic molecule. l Each O-H bond is highly polar, because of the high electronegativity of the oxygen (N, O, F, and Cl have high values) l bond angle of water = 105 o l due to the bent shape, the O-H bond polarities do not cancel. This means water as a whole is polar.

5 The Water Molecule δ is the lowercase Greek symbol delta δ-δ- δ+δ+ δ+δ+ H H O δ- means a partial negative charge δ+ means a partial positive charge Thus, water has a partial negative end (0xygen) and a partial positive end (Hydrogen) – and it is called “polar” because of these areas of difference

6 The Water Molecule l Water’s bent shape and ability to hydrogen bond gives it many special properties! l Water molecules are attracted to one another by dipole interactions – p. 446 l This hydrogen bonding gives water: a) its high surface tension, b) its low vapor pressure.

7 High Surface Tension l liquid water acts like it has a “skin” –glass of water bulges over the top l Water forms round drops –spray water on greasy surface l All because water hydrogen bonds. l Fig. 15.4, p.447 – how does this insect walk on the water?

8 Surface Tension l One water molecule can hydrogen bond to another because of this electrostatic attraction. l Also, hydrogen bonding occurs with other molecules surrounding them on all sides. H H O ++ ++ -- H H O ++ -- ++

9 Surface Tension l A water molecule in the middle of a solution is pulled in all directions.

10 Surface Tension l Not true at the surface. l Only pulled down and to each side. l Holds the molecules at the surface together tightly. l This causes surface tension.

11 Surface Tension l Water drops are rounded, because all molecules on the edge are pulled to the middle- not outward to the air!

12 Surface Tension l Glass has polar molecules. l Glass can also hydrogen bond. l This attracts the water molecules. l Some of the pull is up a cylinder.

13 Meniscus l Water curves up along the side of glass. l This makes the meniscus, as in a graduated cylinder l Plastics are non- wetting; no attraction

14 Meniscus In Glass In Plastic

15 Surface tension l All liquids have surface tension –water is just higher than most others l How can we decrease surface tension? –Use a surfactant - surface active agent –Also called a wetting agent, like detergent or soap –Interferes with hydrogen bonding

16 Low vapor pressure Hydrogen bonding also explains water’s unusually low vapor pressure. –Holds water molecules together, so they do not escape easily –This is a good thing, because lakes and oceans would evaporate very quickly due to their large surface area!

17 Ice l Most liquids contract (get smaller) as they are cooled. –They get more dense. l When they change to solid, they are more dense than the liquid. l Solid metals sink in their own liquid metal. –But, ice floats in water. l Why?

18 Ice l Water becomes more dense as it cools, until it reaches about 4ºC. –Then it becomes less dense by starting to expand (Want your water pipes to freeze in the winter?) l As the molecules slow down, they arrange themselves into honeycomb shaped crystals. l These are held together by hydrogen bonds. Fig. 15.5, p.449

19 H H O H H O H H O H H O H H O H H O H H O H H O H H O H H O H H O H H O Liquid – random shaped arrangement Solid – honeycomb shaped arrangement

20 Ice l 10% lower density than water. l Water freezes from the top down. –The layer of ice on a pond acts as an insulator for water below l Why is ice less dense than liquid water? –The structure of ice is a regular open framework of water molecules arranged like a honeycomb.

21 Ice l A considerable amount of energy is required to return water in the solid state to the liquid –The heat absorbed when 1 g of water changes from solid to liquid is 334 J. –This is the same amount of energy needed to raise the temperature of 1 g of liquid water from 0 o C to 80 o C!

22 Section 17.2 Homogeneous Aqueous Solutions l OBJECTIVES: –Distinguish between a solvent and a solute.

23 Section 17.2 Homogeneous Aqueous Solutions l OBJECTIVES: –Describe what happens in the solution process.

24 Section 17.2 Homogeneous Aqueous Solutions l OBJECTIVES: –Explain why all ionic compounds are electrolytes.

25 Section 17.2 Homogeneous Aqueous Solutions l OBJECTIVES: –Demonstrate how the formula for a hydrate is written.

26 Solvents and Solutes l Solution - a homogenous mixture, that is mixed molecule by molecule. l Solvent - the dissolving medium l Solute -the dissolved particles l Aqueous solution- a solution with water as the solvent. l Particle size less than 1 nm; cannot be separated by filtration – Fig. 15.6, p.450

27 1. Solute A solute is the dissolved substance in a solution. A solvent is the dissolving medium in a solution. 2. Solvent Salt in salt waterSugar in soda drinks Carbon dioxide in soda drinks Water in salt waterWater in soda Parts of a Solution:

28 Solvents

29 Concentrated vs. Dilute

30 Aqueous Solutions l Water dissolves ionic compounds and polar covalent molecules best. l The rule is: “like dissolves like” l Polar dissolves polar. l Nonpolar dissolves nonpolar. l Oil is nonpolar. –Oil and water don’t mix. l Salt is ionic- makes salt water.

31 The Solution Process l Called “solvation”. l Water breaks the + and - charged pieces apart and surrounds them. l Fig. 17.7, p. 451 l In some ionic compounds, the attraction between ions is greater than the attraction exerted by water –Barium sulfate and calcium carbonate

32 How Ionic solids dissolve in water H H O H H O H H O H H O H H O H H O H H O H H O H H O These ions have been surrounded by water, and are now dissolved! These ions have been pulled away from the main crystal structure by water’s polarity.

33 l Solids will dissolve if the attractive force of the water molecules is stronger than the attractive force of the crystal. l If not, the solids are insoluble. l Water doesn’t dissolve nonpolar molecules (like oil) because the water molecules can’t hold onto them. l The water molecules hold onto other water molecules, and separate from the nonpolar molecules. l Nonpolars? No repulsion between them

34 Electrolytes and Nonelectrolytes l Electrolytes- compounds that conduct an electric current in aqueous solution, or in the molten state –all ionic compounds are electrolytes because they dissociate into ions (they are also called “salts”) l barium sulfate- will conduct when molten, but is insoluble in water!

35 Electrolytes and Nonelectrolytes l Do not conduct? = Nonelectrolytes. –Most are molecular materials, because they do not have ions l Not all electrolytes conduct to the same degree –there are weak electrolytes, and strong electrolytes –depends on: degree of ionization

36 The ammeter measures the flow of electrons (current) through the circuit. If the ammeter measures a current, and the bulb glows, then the solution conducts. If the ammeter fails to measure a current, and the bulb does not glow, the solution is non- conducting. Electrolytes vs. Nonelectrolytes

37 Electrolytes and Nonelectrolytes l Strong electrolytes exist as nearly 100 % ions l Weak electrolytes have only a fraction of the solute that exists as ions l How do you know if it is strong or weak? Refer to the rules on the handout sheet.

38 Electrolyte Summary l Substances that conduct electricity when dissolved in water, or molten. l Must have charged particles that can move. l Ionic compounds break into charged ions: NaCl  Na 1+ and Cl 1- l These ions can conduct electricity.

39 l Nonelectrolytes do not conduct electricity when dissolved in water or molten l Polar covalent molecules such as methanol (CH 3 OH) don’t fall apart into ions when they dissolve. l Weak electrolytes don’t fall completely apart into ions. l Strong electrolytes do ionize completely.

40 Water of Hydration (or Water of Crystallization) l Water molecules are chemically bonded to solid salt molecules (not in solution) l These compounds have fixed amounts of water. l The water can be driven off by heating: l CuSO 4. 5H 2 O CuSO 4 + 5H 2 O l Called copper(II)sulfate pentahydrate. - heat + heat

41 Hydrates l Table 15.2, p.455 list some familiar hydrates l Since heat can drive off the water, the forces holding it are weak l If a hydrate has a vapor pressure higher than that of water vapor in air, the hydrate will effloresce by losing the water of hydration

42 Hydrates l Some hydrates that have a low vapor pressure remove water from the air to form higher hydrates- these are called hygroscopic –used as drying agents, or dessicants –packaged with products to absorb moisture

43 - Page 456

44 Hydrates l Some compounds are so hygroscopic, they become wet when exposed to normally moist air - called deliquescent –remove sufficient water to dissolve completely and form solutions: Fig , page 457

45 Section 17.4 Heterogeneous Aqueous Systems l OBJECTIVES: –Distinguish between a suspension and a solution.

46 Section 17.4 Heterogeneous Aqueous Systems l OBJECTIVES: –Identify the distinguishing characteristics of a colloid.

47 Mixtures that are NOT Solutions l Suspensions: mixtures that slowly settle upon standing. –Particles of a suspension are greater in diameter than 100 nm. –Can be separated by filtering (p.459) l Colloids: heterogeneous mixtures with particles between the size of suspensions and true solutions (1-100 nm)

48 Mixtures that are NOT Solutions l The colloid particles are the dispersed phase, and are spread throughout the dispersion medium l The first colloids were glues. Others include mixtures such as gelatin, paint, aerosol sprays, and smoke l Table 17.5, p.491 list some common colloidal systems and examples

49 Mixtures that are NOT Solutions l Many colloids are cloudy or milky in appearance when concentrated, but almost clear when dilute –do not settle out –can not be filtered out l Colloids exhibit the Tyndall effect- the scattering of visible light in all directions. –suspensions also show Tyndall effect

50 The Tyndall Effect Colloids and suspensions scatter light, making a beam visible, due to their large particle size. Solutions do not scatter light, because of their small particle size. Which glass contains a colloid? solution colloid

51 - Page 461 Note that you can easily see the “sunbeam”, probably due to the presence of fog in the forest

52 Mixtures that are NOT Solutions l Flashes of light are seen when colloids are studied under a microscope- light is reflecting- called Brownian motion to describe the chaotic movement of the particles. Observed by Robert Brown. l Table 17.6, p.492 summarizes the properties of solutions, colloids, and suspensions

53 Mixtures that are NOT Solutions l Emulsions- dispersions of liquid in liquid (2 immiscible liquids + an emulsifier) –an emulsifying agent is essential for maintaining stability; it has one polar end, and the other end is nonpolar –oil and water not soluble; but with soap or detergent added, they will be. l Oil + vinegar in dressing – are they soluble? –What makes up Mayonnaise? Margarine?


Download ppt "Chapter 17 “Water and Aqueous Systems” Chemistry Golden Valley High School Stephen L. Cotton."

Similar presentations

Ads by Google