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Structure Determines the Properties of Liquids and Solids PowerPoint Presentation
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Structure Determines the Properties of Liquids and Solids

Structure Determines the Properties of Liquids and Solids

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Structure Determines the Properties of Liquids and Solids

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  1. Structure Determines the Properties of Liquids and Solids • The atoms or molecules have different structures in solids, liquids and gases which leads to different properties.

  2. The Structures of Solids, Liquids and Gases

  3. Gases • In the gas state, the particles have complete freedom from each other. • The particles are constantly flying around, bumping into each other and the container. • In the gas state, there is a lot of empty space between the gas particles. • Because there is a lot of empty space, the particles can be squeezed close together: gases are compressible. • Because the particles are not held in close contact and are moving freely, gases expand to fill and take the shape of their container.

  4. Liquids • The particles in a liquid are closely packed, but they have some ability to move around. • The close packing results in liquids being incompressible. • The ability of the liquid particles to move allows liquids to take the shape of their container and to flow. However, they don’t have enough freedom to escape and expand to fill the container as gases do.

  5. Properties of Liquids Viscosity • Some liquids flow more easily than others which means there is less of an attraction between the molecules. • The resistance of a liquid to flow is called viscosity. -The greater the attractive forces between the molecules, the larger the viscosity. Surface Tension • Liquids tend to minimize their surface causing the surface of the liquid to resist penetration – a phenomenon called surface tension. • Stronger attractive forces between liquid molecules results in a larger surface tension.

  6. There are two different Forces of Attraction within a Liquid • Cohesive Forces are forces that try to hold the liquid molecules to each other and are the forces responsible for surface tension. • Adhesive Forces are forces that bind a substance to a surface such as evidenced by • capillary action: a liquid rising in a tube • Meniscus: a liquid rising up the sides of a graduated cylinder

  7. Attractive Forces and Properties • Like dissolves Like • miscible = liquids that do not separate, no matter what the proportions, are miscible. • Polar molecules dissolve in polar solvents: • water, alcohol, CH2Cl2 • molecules with O or N higher solubility in H2O due to H-bonding with H2O. • Nonpolar molecules dissolve in nonpolar solvents: • ligroin (hexane), toluene, CCl4. • If a molecule has both polar & nonpolar parts, then there will be hydrophilic - hydrophobic competition.

  8. Immiscible Liquids When liquid pentane, a nonpolar substance, is mixed with water, a polar substance, the two liquids separate because they are more attracted to their own kind of molecule than to each other. The one that is less dense floats on top of the more dense liquid.

  9. - - - - - - - - - - - - - - + - - + - - - - - - - - + - + + - - + - + + - - - - - - - - - - - - - - + + - - - - - - - - Why are molecules attracted to each other? • Intermolecular attractions are due to attractive forces between opposite charges • + ion to - ion • + end of polar molecule to - end of polar molecule • H-bonding especially strong • larger charge = stronger attraction • Even nonpolar molecules can have temporary induced dipoles: • also known as London Forces or Induced Dipoles • caused by electrons on one molecule distorting the electron cloud on another • all molecules have dispersion forces

  10. Types of Intermolecular Forces

  11. + - + - _ + + _ + _ + _ Attractive Forces Dispersion Forces – all molecules Dipole-to-Dipole Forces – polar molecules + - + -

  12. Strength of the Dispersion Force • Depends on how easily the electrons can move, or be polarized. • The more electrons and the farther they are from the nuclei, the larger the dipole that can be induced. • The strength of the dispersion force gets larger with larger molecules. • Because of the kinds of atoms that are bonded together and their relative positions in the molecule, some molecules have a permanent dipole. • All polar molecules have a permanent dipole with the more electronegative atom(s) drawing the electrons away from the less electronegative atom(s).

  13. Dipole-to-Dipole Attraction • Polar molecules have a permanent dipole • a positive (+) end and a negative (-) end. • the + end of one molecule will be attracted to the – end of another.

  14. Hydrogen Bonding • Molecules that have HF, OH or NH groups have particularly strong intermolecular attractions • unusually high melting and boiling points • unusually high solubility in water

  15. H-Bonds vs. Chemical Bonds • Hydrogen bonds are not chemical bonds! • Hydrogen bonds are attractive forces between molecules. • Chemical bonds are attractive forces that make molecules: covalent bonds are between nonmetals with nonmetals and ionic bonds are between metals and nonmetals.

  16. Attractive Forces & Properties

  17. How much Energy is needed to overcome these forces? Which are stronger?

  18. Liquids Evaporate: a change in state from a liquid to a gasLiquids Condense: A change in state from a gas to a liquid • Over time, liquids evaporate – the molecules of the liquid mix with and dissolve in the air. • Evaporation occurs at the liquid surface. • The molecules on the surface experience a smaller net attractive force than molecules inside the liquid. • All of the surface molecules do not escape at once, only the ones with enough kinetic energy to overcome the attractions holding them in the liquid state can escape into the atmosphere as gas molecules. • liquids that evaporate quickly are called volatile liquids, while those that do not are called nonvolatile • Gases can lose kinetic energy and condense • into liquids.

  19. Evaporation and Condensation will reach equilibrium in a closed container! When water is just added to the flask and the flask is capped, all of the water molecules are in the liquid state. Shortly, the water starts to evaporate. Initially, the speed of evaporation is much faster than the speed of condensation. Eventually, the condensation and evaporation reach the same speed, i.e. they are occurring at the same rate. The air in the flask is now saturated with water vapor. As long as the conditions don’t change, the partial pressure exerted by the vapor is constant and is called the vapor pressure and is dependent on temperature and intermolecular attractions.

  20. Factors that Effect the Rate of Evaporation • Increasing the surface area increases the rate of evaporation. • Increasing temperature increases the rate of evaporation. • Weaker attractive forces between liquid molecules results in a faster rate of evaporation. • As the higher energy molecules from the liquid escape, the total kinetic energy of the liquid decreases which cools the liquid. • The remaining molecules redistribute their energies, generating more high energy molecules. • The result is that the liquid continues to evaporate until there is no more liquid.

  21. Why do we feel cold after a shower? Evaporation

  22. What happens when a liquid is Boiled? • In an open container, as a liquid is heated the average kinetic energy of the molecules increases. When they have enough kinetic energy to overcome atmospheric pressure and the intermolecular forces holding them in the liquid state they escape into the atmosphere as gas molecules. • the rate of evaporation increases • Eventually, the temperature is high enough for molecules in the interior of the liquid to escape – a phenomenon called boiling. • The temperature at which the vapor pressure of the liquid is the same as the atmospheric pressure is called the boiling point. • The boiling point is dependent on what the atmospheric pressure is • the temperature of boiling water on the top of a mountain will be lower than that for boiling water at sea level.

  23. Temperature and Boiling • As a liquid is heated, the temperature increases until it reaches the boiling point. • Once the liquid starts to boil, the temperature remains the same until all of the liquid turns into a gas. • Two phases exist during boiling: liquid and gas. • The same is true for melting. Two phases, solid and liquid, exist until all of the solid has turned into a liquid.

  24. Change in state of water Increasing E Decreasing E

  25. Phase Changes are Physical Changes • Boiling = liquid to gas • Melting = solid to liquid • Sublimation = solid to gas • Condensing = gas to liquid • Freezing = liquid to solid • Deposition = gas to solid • state changes require heating or cooling the substance

  26. Solids • the particles in a solid are packed close together and are fixed in position • though they are vibrating • the close packing of the particles results in solids being incompressible • the inability of the particles to move around results in solids retaining their shape and volume when placed in a new container; and prevents the particles from flowing. • Some solids have their particles arranged in an orderly geometric pattern – we call these crystalline solids: • salt and diamonds • Other solids have particles that do not show a regular geometric pattern over a long range – we call these amorphous solids: • plastic and glass

  27. Why is Sugar a Solid ButWater is a Liquid? • The chemical state in which a material exists depends on the attraction between molecules and their ability to overcome this attraction. • The attractive forces between ions or molecules depends on their structure: • the attractions are electrostatic • depend on shape, polarity, etc. • The ability of the molecules to overcome the attraction depends on the amount of kinetic energy they possess: • Gases > Liquids > Solids

  28. Types of Crystalline Solids • Ionic Solids: consists of oppositely charged ions packed together. These are very stable with high melting points. Examples: NaCl; CsF; BaO • Molecular solids: contain molecules as their fundamental particles. These tend to melt at low temperatures because the intermolecular forces between the molecules are relatively week. Examples: ice; white phosphorous (P4 molecules); sulfur (S8 molecules) • Atomic Solids: consists of atoms as the fundamental particles. Depending on the way atoms interact, these solids vary greatly in their properties. Examples: Argon; Gold; Diamond.

  29. Atomic Crystalline Solids: Metallic Bonds • Atomic solids are solids whose composite units are individual atoms. • The solids are held together by either covalent bonds (diamonds, have very high melting points), dispersion forces (Xenon, have low melting points) or metallic bonds (Iron, positively charged ions surrounded by electrons and have variable melting points). • The metal atoms release some of their electrons to be shared by all the other metal atoms in the crystal. • The metallic bond is the attraction of the metal cations for the mobile electrons • often described as islands of cations in a sea of electrons.

  30. Metallic Bonding: the model of metallic bonding can be used to explain the properties of metals • The luster, malleability, ductility, electrical and thermal conductivity of metals are all related to the mobility of the electrons in the solid. • The strength of the metallic bond varies, depending on the charge and size of the cations – so the melting points of metals vary.

  31. Alloys • Alloy: a substance that contains a mixture of elements and has metallic properties. Two common types: • Substitutional alloy: some of the host metal atoms of similar sizes. Example: brass, 2/3 copper 1/3 zinc. • Interstitial alloy: some interstices (holes) among the closely packed metal atoms are occupied by atoms much smaller than the host atoms. Example: steel, contains carbon atoms in the “holes” of an iron crystal. Varying the ratio of carbon from less than 0.2% to 1.5% change the steel from soft to very hard dependent on the intended use.