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Chapter 13 – States of Matter

Chapter 13 – States of Matter

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Chapter 13 – States of Matter

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  1. Chapter 13 – States of Matter Chem 311 Period 3

  2. Kinetic Theory: All matter consists of tiny particles in CONSTANT MOTION PARTICLES IN GAS are small, hard sphere w/ insignificant volume Particles are relatively far apart (compared to those of solids / liquids) Particles are NOT attracted or repulsed by each other Motion of one particle of gas INDEPENDENT of motion of other particles • Kinetic Energy: the energy an object has because of its motion • THUS  Motion of particles of Gas: • Rapid, constant, random (unaffected by each • other) • As a result, gases can spread out into space w/o limit (any shape) •  travel in straight paths until they collide with • another particle or object - they change direction only when they rebound/collide with another particle or object. • All collisions between particles in a gas are perfectly elastic • Elastic collision: Kinetic energy is transferred WITHOUT LOSS from one particle to another – total kinetic energy remains CONSTANT

  3. Measuring Air Pressure • A Barometer is a device used to measure atmospheric pressure and can be read the same way as a thermometer. • The height of the mercury in a barometer depends on weather and altitude. The higher the altitude, the lower the pressure. • The atmospheric pressure pushes the mercury from the reservoir into the glass tube. The pressure at sea level is 760mm, or 29.2 inches. The tip of the tube is a vacuum- an empty space with no pressure or particles.

  4. GasPressure • Gas Pressure is the result of simultaneous collisions of billions of rapidly moving gas particles against an object. The more collisions, the higher the pressure. • The collision / movement of these gas particles can be explained with the Kinetic theory (which states that matter consists of tiny particles in constant motion). • The Pascal (Pa) is the SI unit of pressure, another unit of measurement is millimeters of mercury (mm Hg) and standard atmospheres (atm).

  5. Kinetic Theory + Distribution Curve • Increase in kinetic energy= increase in temperature • Higher temperature has wider range of kinetic energies • Increase in avg kinetic energy causes a rise of the temperature

  6. Kelvin • At absolute zero (0 K) particles have no kinetic energy/no motion. There is no lower temperature than absolute zero Kelvin. • Kelvin temperature is directly proportional to the avg kinetic energy of a substance • For example, particles of helium at 200K have 2x the avg kinetic energy as helium at 100K

  7. A Model for Liquids Both the particles in gases and the particles in liquids have kinetic energy. This energy allows the particles in gases and liquids to flow past each other. Substances that can flow are known as fluids. Since fluids flow the way they do, they conform to the shape of their containers.

  8. Pt. II Unlike the particles of gases, inter-molecular attractions reduce the amount of space between the particles in liquids. Since the particles of liquids are close together, liquids have a definite volume. Liquids are consequently much more dense than gases Changing the amount of pressure on a liquid hardly effects its volume. Since liquids and solids are mostly incondensable, they are known as condensed states of matter.

  9. 5) Vaporization with kinetic theory • Vaporization=liquid converted to a vapor or gas • Kinetic Theory- all matter consists tiny particles in constant motion • Most molecules do not have high enough kinetic energy (movement of particles) to overcome attractive force & vaporize

  10. 5) Vaporization with kinetic theory • Molecules above a certain point of high kinetic energy (rapid motion) can vaporize • Molecules evaporate from surface of liquid to a gas

  11. 5)Manometer • Measures equilibrium vapor pressure when mercury levels in tubes fluctuate • Liquid added, pressure increases-moving mercury in U-Tube • Difference in levels of mercury = vapor pressure in mm of Hg.

  12. 6) Boiling Point • Boiling occurs – liquid heated to temp where particles have high enough kinetic energy to vaporize • Heating water creates higher level kinetic energy (rapid motion molecules). • Boiling point- temp high enough for vapor pressure of liquid to equal the external pressure pushing on liquid

  13. Solid Organization and Properties • Closely packed ions, atoms, or molecules, well organized • Particles have fixed location and vibrate around fixed points, strong attractions • Solids do not flow • Dense (not easily compressed) • Melting points • Ionic solids- high melting point • Molecular solids- low melting point

  14. Melting Point of Solids • melting point-temperature where solid changes into a liquid • When melted, solid particles have high vibrations that allow them to move out of fixed positions and into liquid form • Equilibrium- temperature where melting point of solid and freezing point of liquid are at same temperature • Melting point is determined by particle organization • Solid melting point exceptions • Cane sugar and wood

  15. Crystal Structures and Systems • Crystal- three dimensional pattern arrangement of solid particles • Crystal shape describes solid particle arrangement • Crystal properties • Has sides or face • Labeled a, b, c • Has angles that intersect at faces • Labeled β, Υ, α • Seven groups of crystals- cubic, tetragonal, rhombic, monoclinic, triclinic, hexagonal, and rhombohedra

  16. Liquid Organization and Properties • Contain kinetic energy • Fluids- particles that flow • Allows liquids to take shape of container • Particles have relatively strong attractions • Keep particles together • Create definite volume • Keeps particles close together • Denser than gases, less dense then solids • Not effected by pressure

  17. Gas Organization and Properties • Particles have kinetic energy (like liquids) • Keeps particles in constant motion • Particles are small hard spheres with insignificant volume • Spacing is high • No attractive forces to hold them together • Particles are in constant motion • Fill container regardless of shape

  18. Solid Liquid Gas Organization and Attraction

  19. Distinguishing Between Crystal and Glass • Properties of Crystal • Particles are arranged in an orderly, 3-D pattern called a crystal lattice • Has sides, or faces • In general, have high melting points (unless molecular) • When shattered, fragments have the same surface angles as the original solid • Properties of Glass • Also called amorphous solid • Crystallization does not occur • The structures of glasses are intermediate between those of crystalline and those of free-flowing liquids • Do not melt at a definite temp. • When shattered, fragments have irregular angles and jagged edges

  20. Crystal vs. Glass Structures • Crystal • The shape of a crystal reflects the arrangement of particles within the solid • The smallest group of particles that retains its geometric shape is the unit cell • The unit cell may be simple cubic, body-centered cubic, or face-centered cubic • Glass • Glasses or, amorphous solids, are transparent fusion products of inorganic substances that have cooled to a rigid state WITHOUT crystallizing

  21. Different Shapes of Crystal Systems

  22. Sublimation • When a solid becomes a gas without passes through the liquid state. • Occurs when a solid has a vapor pressure higher than the pressure at or near room temperature. • Vapor pressure: the measure of a force exerted by a gas above a liquid (true for solids too)

  23. Deposition • When a gas becomes a solid without passes through the liquid state.

  24. Phase diagram-shows the conditions and pressure at which a substance exists as a solid, liquid, and gas • Triple point- the only set of conditions where all three states can exist. • Critical point- beyond this region the physical and chemical properties of water and steam converge to the point where they are identical. Thus, beyond the critical point, we refer to this single phase as a "supercritical fluid". • Boiling, melting, and subliming curves

  25. Plasma: What is it? • Most common state of matter • Electrically charged particles at high energy that collect around electromagnetic fields and form gas-like clouds • It is made of extremely hot ions and electrons in space, but on Earth, it cools into atoms and molecules • The particles are affected by electromagnetic, electric, and magnetic signals but are hardly affected by gravity.

  26. Plasma: What is it? (cont.) • The full range of plasma’s density, temperature, and spatial scales is nearly incomprehensible as it is so wide and varied. • Without sufficient energy, the plasma reverts back to a neutral gas • Energy being thermal, electrical, or light (i.e. ultraviolet) • Ions and electrons move independently in large spaces • Plasmas are still being studied and understood today

  27. Plasma Examples • Flames • neon signs • Nebulae (Clouds in space) • solar wind • Auroras • Galaxies • dense solid state of matter • Stars • Space • Lightning • florescent lights