States of Matter

# States of Matter

## States of Matter

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##### Presentation Transcript

1. States of Matter Solids, Liquids, Gases, Plasma, and Phase Changes

2. Is there any movement within the structure of a solid?

3. The general properties of solids reflect the orderly arrangement of their particles and the fixed locations of their particles. A Model for Solids • In most solids, the atoms, ions, or molecules are packed tightly together. • Solids are dense and not easy to compress. • Because the particles in solids tend to vibrate about fixed points, solids do not flow.

4. The shape of a crystal reflects the arrangement of the particles within the solid. Crystal Structure and Unit Cells • In sodium chloride, sodium ions and chloride ions are closely packed in a regular array. • The ions vibrate about fixed points in the crystal.

5. A Model for Solids When you heat a solid, its particles vibrate more rapidly as their kinetic energy increases. • The melting point (mp) is the temperature at which a solid changes into a liquid.

6. A Model for Solids When you heat a solid, its particles vibrate more rapidly as their kinetic energy increases. • The melting point (mp) is the temperature at which a solid changes into a liquid. • At this temperature, the disruptive vibrations of the particles are strong enough to overcome the attractions that hold them in fixed positions.

7. A Model for Solids • In general, ionic solids have high melting points because relatively strong forces hold them together. • Freezing Point: The temperature at which a liquid turns into a solid. (Reverse of melting point) • Sodium chloride, an ionic compound, has a rather high melting point of 801°C.

8. A Model for Solids • In general, ionic solids have high melting points because relatively strong forces hold them together. • Sodium chloride, an ionic compound, has a rather high melting point of 801°C. • Molecular solids have relatively low melting points. • Hydrogen chloride, a molecular compound, melts at –112°C.

9. A Model for Solids • Different forces hold substances together. • Ionic compounds are held together by the attraction of anions to cations. • Covalent compounds are held together by several different intermolecular forces, the forces that hold covalent compounds together • Hydrogen Bonds (positive and negative ends attract) • Dipole – Dipole forces (positive and negative ends attract) • Dispersion forces (temporary positive and negative ends form that attract)

10. Crystal Structure and Unit Cells Some substances can exist in more than one form. • Diamond is one crystalline form of carbon. • A different form of carbon is graphite. • In 1985, a third crystalline form of carbon was discovered. This form is called buckminsterfullerene.

11. Crystal Structure and Unit Cells In diamond, each carbon atom in the interior of the diamond is strongly bonded to four others. The array is rigid and compact. In graphite, the carbon atoms are linked in widely spaced layers of hexagonal arrays. In buckminster-fullerene, 60 carbon atoms form a hollow sphere. The carbons are arranged in penta-gons and hexagons.

12. Crystal Structure and Unit Cells The physical properties of diamond, graphite, and fullerenes are quite different. • Diamond has a high density and is very hard. • Graphite has a relatively low density and is soft and slippery. • The hollow cages in fullerenes give them strength and rigidity.

13. Crystal Structure and Unit Cells • Allotropes Diamond, graphite, and fullerenes are crystalline allotropes of carbon. • Allotropes are two or more different molecular forms of the same element in the same physical state.

14. Crystal Structure and Unit Cells • Allotropes Only a few elements have allotropes. • In addition to carbon, these include phosphorus, sulfur, oxygen (O2 and O3), boron, and antimony.

15. Crystal Structure and Unit Cells • Non-Crystalline Solids Not all solids are crystalline in form; some solids are amorphous. • An amorphous solid lacks an ordered internal structure. • Rubber, plastic, and asphalt are amorphous solids. • Their atoms are randomly arranged.

16. Crystal Structure and Unit Cells • Non-Crystalline Solids Other examples of amorphous solids are glasses. • A glass is a transparent fusion product of inorganic substances that have cooled to a rigid state without crystallizing. • Glasses are sometimes called supercooled liquids. • The irregular internal structures are intermediate between those of a crystalline solid and those of a free-flowing liquid.

17. What is the difference between an amorphous solid and a crystalline solid?

18. What is the difference between an amorphous solid and a crystalline solid? Particles in a crystalline solid are arranged in an orderly, repeating pattern or lattice. Particles in an amorphous solid are arranged randomly.

19. Kinetic Theory • The word kinetic refers to motion. • The energy an object has because of its motion is called kinetic energy. • According to the kinetic theory, all matter consists of tiny particles that are in constant motion. • The particles in a gas are usually molecules or atoms.

20. Kinetic Energy and Temperature • As a substance is heated, its particles absorb energy, some of which is stored within the particles. • This stored portion of the energy, or potential energy, does not raise the temperature of the substance. • The remaining absorbed energy does speed up the particles, that is, increases their kinetic energy. • This increase in kinetic energy results in an increase in temperature.

21. Kinetic Energy and Temperature • Average Kinetic Energy The particles in any collection of atoms or molecules at a given temperature have a wide range of kinetic energies. • Most have kinetic energies somewhere in the middle of this range. • We use average kinetic energy when discussing the kinetic energy of a collection of particles in a substance.

22. Kinetic Energy and Temperature • Average Kinetic Energy At any given temperature, the particles of all substances, regardless of physical state, have the same average kinetic energy.

23. InterpretGraphs • The figure below shows the distribution of kinetic energies of water molecules at two different temperatures. • The green curve shows the distribution of kinetic energy in cold water. • The purple curve shows the distribution of kinetic energy in hot water.

24. Kinetic Energy and Temperature • Average Kinetic Energy The average kinetic energy of the particles in a substance is directly related to the substance’s temperature. • An increase in the average kinetic energy of the particles causes the temperature of a substance to rise. • As a substance cools, the particles tend to move more slowly, and their average kinetic energy decreases.

25. Kinetic Energy and Temperature • Average Kinetic Energy Absolute zero (0 K, or –273.15oC) is the temperature at which the motion of particles theoretically ceases. • No temperature can be lower than absolute zero. • Absolute zero has never been produced in the laboratory. • A near-zero temperature of about 0.000 000 000 1 K, which is 0.1 nanokelvin, has been achieved.

26. Kinetic Energy and Temperature • Average Kinetic Energy The coldest temperatures recorded outside the laboratory are from space. • Astronomers used a radio telescope to measure the temperature of the boomerang nebula. • At about 1 K, it is the coldest known region of space.

27. The Kelvin temperature of a substance is directly proportional to the average kinetic energy of the particles of the substance. Kinetic Energy and Temperature • Average Kinetic Energy and Kelvin Temperature

28. What is the result of increasing the temperature of a gas sample? A. Adecrease in the average kinetic energy of the sample B. No effect on the sample C. An increase in the average kinetic energy of the sample D. Theparticles slow down.

29. What is the result of increasing the temperature of a gas sample? A. Adecrease in the average kinetic energy of the sample B. No effect on the sample C. An increase in the average kinetic energy of the sample D. Theparticles slow down.

30. The particles in a gas are considered to be small, hard spheres with an insignificant volume. Kinetic Theory and a Model for Gases • The kinetic theory as it applies to gases includes the following fundamental assumptions about gases. • Within a gas, the particles are relatively far apart compared with the distance between particles in a liquid or solid. • Between the particles, there is empty space. • No attractive or repulsive forces exist between the particles.

31. The motion of particles in a gas is rapid, constant, and random. Bromine molecule Kinetic Theory and a Model for Gases • The kinetic theory as it applies to gases includes the following fundamental assumptions about gases. • Gases fill their containers regardless of the shape and volume of the containers. • An uncontained gas can spread out into space without limit.

32. All collisions between particles in a gas are perfectly elastic. Kinetic Theory and a Model for Gases • The kinetic theory as it applies to gases includes the following fundamental assumptions about gases. • During an elastic collision, kinetic energy is transferred without loss from one particle to another. • The total kinetic energy remains constant.

33. Gas Pressure Gas pressure results from the force exerted by a gas per unit surface area of an object. • Moving bodies exert a force when they collide with other bodies.

34. Gas pressure is the result of billions of rapidly moving particles in a gas simultaneously colliding with an object. Gas Pressure • If no particles are present, no collisions can occur. Consequently, there is no pressure. • An empty space with no particles and no pressure is called a vacuum.

35. Gas Pressure Air exerts pressure on Earth because gravity holds the particles in air within Earth’s atmosphere. • The collisions of atoms and molecules in air with objects results in atmospheric pressure.*Observe the effects of atmospheric pressure in the following demo • Atmospheric pressure decreases as you climb a mountain because the density of Earth’s atmosphere decreases as the elevation increases.

36. Vacuum Atmospheric pressure 760 mm Hg (barometric pressure) 253 mm Hg Sea level On top of Mount Everest Gas Pressure A barometer is a device that is used to measure atmospheric pressure. • At sea level, air exerts enough pressure to support a 760-mm column of mercury. • On top of Mount Everest, at 9000 m, the air exerts only enough pressure to support a 253-mm column of mercury.

37. CHEMISTRY&YOU • When weather forecasters state that a low-pressure system is moving into your region, it usually means that a storm is coming. What do you think happens to the column of mercury in a barometer as a storm approaches? Why?

38. CHEMISTRY&YOU • When weather forecasters state that a low-pressure system is moving into your region, it usually means that a storm is coming. What do you think happens to the column of mercury in a barometer as a storm approaches? Why? When a storm approaches, the column of mercury goes down, indicating a decrease in atmospheric pressure.

39. Gas Pressure The SI unit of pressure is the pascal (Pa). • Normal atmospheric pressure is about 100,000 Pa, that is, 100 kilopascals (kPa). • Two older units of pressure are commonly used. • millimeters of mercury (mm Hg) • atmospheres (atm)

40. Gas Pressure One standard atmosphere (atm) is the pressure required to support 760 mm of mercury in a mercury barometer at 25°C. • The numerical relationship among the three units is • 1 atm = 760 mm Hg = 101.3 kPa. • Standard temperature and pressure (STP) are defined as a temperature of 0°C and a pressure of 101.3 kPa, or 1 atm.

41. A Model for Liquids • Substances that can flow are referred to as fluids. • Both liquids and gases can flow. • The ability of gases and liquids to flow allows them to conform to the shape of their containers.

42. A Model for Liquids • Gases and liquids have a key difference between them. • According to kinetic theory, there are no attractions between the particles in a gas. • The particles in a liquid are attracted to each other. • These intermolecular attractions keep the particles in a liquid close together, which is why liquids have a definite volume.

43. A Model for Liquids • Different forces hold substances together. • Ionic compounds are held together by the attraction of anions to cations. • Covalent compounds are held together by several different intermolecular forces, the forces that hold covalent compounds together • Hydrogen Bonds (positive and negative ends attract) • Dipole – Dipole forces (positive and negative ends attract) • London Dispersion forces (temporary positive and negative ends form that attract)

44. A Model for Liquids Class Activity: Try fitting as many drops of water as possible on a penny without any dripping off the penny, and then do the same with drops of acetone. Which could you drop more drops of? Why?

45. A Model for Liquids Class Activity: Try fitting as many drops of water as possible on a penny without any dripping off the penny, and then do the same with drops of acetone. Which could you drop more drops of? Why? Both have surface tension that allows for the addition of more drops, but water has stronger intermolecular forces holding it together so you can add more water.

46. A Model for Liquids • Surface tension: A property of liquids where liquids can resist force on the surface. • Results from the net force of the particles in a liquid pulling down on the surface molecules.

47. The interplay between the disruptive motions of particles in a liquid and the attractions among the particles determines the physical properties of liquids. A Model for Liquids

48. A Model for Liquids • Liquids are much more dense than gases. • Increasing the pressure on a liquid has hardly any effect on its volume. • The same is true for solids. • Liquids and solids are known as condensed states of matter.

49. Describe one way in which liquids and gases are similar, and one way in which they are different.

50. Evaporation • The conversion of a liquid to a gas or vapor is called vaporization. • When this conversion occurs at the surface of a liquid that is not boiling, the process is called evaporation.