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States of Matter Chapter 5. Ideal Gas Concept A model of the way that particles behave at the microscopic level Involves the relationship between pressure.

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Presentation on theme: "States of Matter Chapter 5. Ideal Gas Concept A model of the way that particles behave at the microscopic level Involves the relationship between pressure."— Presentation transcript:

1 States of Matter Chapter 5

2 Ideal Gas Concept A model of the way that particles behave at the microscopic level Involves the relationship between pressure (P), volume (V), temperature (T) and number of moles of gas (n). Pressure is force per unit area.

3 Atmospheric pressure Atmospheric pressure measured with a barometer has a unit of mmHg (millimeters of mercury) A commonly used unit of measurement is the atmosphere ◦ 1 atm = 760 mmHg ◦ 1 mm Hg = 1 torr ◦ 1 atm =1.01 x 10 5 pascals

4 Kinetic Molecular Theory of Gases Gases are easily compressible. Gases will expand to fill any available volume. Gases have low density. Gases readily diffuse through each other simply because they are in continuous motion. Gases exert pressure on their containers. Gases behave most ideally at low pressures and high temperatures.

5 Boyle’s Law P i V i =P f V f The volume of a gas varies inversely with the pressure exerted by the gas if the number of moles and temperature of gas are held constant

6 Charles’s Law Temperature must be converted to Kelvin before Charles’s law is applied. Watch for the question requesting the “difference” in temperature or volume.

7 Combined Gas Law Often a sample of gas undergoes change involving volume, pressure, and temperature simultaneously.

8 Avogadro’s Law Equal volumes of any ideal gas contain the same number of moles if measured under the same conditions of temperature and pressure.

9 Molar Volume of a Gas The volume occupied by 1 mol of gas At STP (standard temperature and pressure) the molar volume of any gas is 22.4L T = 273 K P = 1 atm

10 Ideal Gas Law Combining Boyle’s, Charles’s and Avogadro’s Law into a single expression pV = nRT P in atmospheres V in liters n in number of moles T in Kelvin

11 Dalton’s Law of Partial Pressures A mixture of gases exerts a pressure that is the sum of the pressures that each gas would exert if it were present alone under the same conditions.

12 The Liquid State Compressibility: ◦ Liquids are practically incompressible ◦ The application of many atmospheres of pressure does not significantly decrease the volume ◦ This is why liquids are used in the brake lines of an automobile and in hydraulics ◦ Instead of compressing the fluid, it transmits it directly to the brake pads, forklift, etc.

13 Viscosity Measure of resistance to flow A function of both attractive forces between molecules and molecular geometry Molecules with complex structures resist sliding smoothly past each other Polar molecules tend to have higher viscosity Glycerol is very viscous, gasoline is not

14 Surface tension Measure of the attractive forces exerted among molecules The net attractive forces on the surface molecules pulls them downward into the body of the liquid Responsible for the spherical shape of drops of liquid Surfactants can be added to liquid to decrease surface tension

15 IRDS Infant respiratory distress syndrome is the developmental insufficiency of surfactant production and structural immaturity in the lungs The surfactant helps prevent collapse of the terminal air-spaces throughout the normal cycle of inhalation and exhalation One of the most commonly used surfactants is Survanta, derived from cow lungs

16 Vapor Pressure of a Liquid Related to evaporation, condensation and boiling point When temperature is too low for a phase change, that’s just the average kinetic energy Evaporation happens when pockets of high-energy molecules possess sufficient energy to escape from the bulk liquid


18 Condensation Evaporation is the conversion of liquid to a gas at a temperature too low to boil. Condensation is the conversion of a gas to the liquid state. Vapor pressure is the pressure exerted by the vapor at equilibrium between rates of evaporation and condensation.

19 Normal Boiling point The boiling point of a liquid is defined as the temperature at which the vapor pressure of the liquid become equal to the atmospheric pressure Normal boiling point is the temperature when vapor pressure of the liquid is equal to 1 atm Therefore at high altitudes (less atmospheric pressure), boiling points are lower. Polar liquids have higher boiling points than nonpolar liquids.

20 Van der waals forces The dipole-dipole interactions of polar molecules decreases vapor pressure and increases boiling point In nonpolar liquids, instantaneous dipoles are created for short periods of time, due to the constant movement of electrons The interaction of these instantaneous dipoles is called London forces. Dipole-dipole and London forces are collectively called Van der Waals forces

21 Hydrogen Bonding When Van der Waals forces are not strong enough to account for high boiling points Molecules where an hydrogen atom is bonded to a small, highly electronegative atom: N, O, or F A large dipole is created

22 More on Hydrogen Bonding Boiling points increase as the electronegativity of the element bonded to hydrogen increases Boiling points from least to greatest: Water > hydrogen fluoride > ammonia > methane

23 The Solid State Solids are virtually incompressible, owing to the small distance between particles. Melting points depend of the strength of attractive forces Polar solids have higher melting points than nonpolar of the same molecular weight

24 Types of solids Crystalline: regular repeating structure Amorphous: having no organized structure

25 Types of Crystalline Solids Ionic: positive and negative ions that electrostatic forces hold together High melting points hard/brittle Covalent: very high melting points Extremely hard Molecular: Van der Waals or hydrogen bonding Frequently volatile, soft, low melting points, poor electrical conductors Metallic: metal atoms, bonds formed by the overlap of orbitals

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