Force Of Attraction Between Gas Particles example essay topic

Chapter 13- States of Matter Gases Kinetic Molecular Theory. Model to describe properties of gases. Kinetic = move. Theory describes behavior in terms of particles in motion. Model makes assumptions about size, motion, and energy of gas particles Particle size. Gases consist of small particles that are separated from one another by empty space.

Because gas particles are far apart, there are no significant attractive or repulsive forces among them Particle Motion. Gas particles are in constant random motion collide with other particles or with the walls of their container. Have elastic collisions-no kinetic energy is lost. Kinetic energy can be transferred between colliding particles, but the total kinetic energy of the 2 particles don't change Particle energy. 2 factors determine the kinetic energy of a particle. Mass, velocity KE = 1/2 mv 2.

In a sample of a dingle gas all particles have the same mass but all particles do not have the same velocity. Because of this particles do not have the same kinetic energy. Kinetic energy and temperature are related. Temperature- a measure of the average kinetic energy of the particles of matter. At a given temperature, all gases have the same average kinetic energy Explaining the behavior of Gases Low Density. Theory states- great deal of space exists between gas particles.

The rest are fewer chlorine molecules than gold atoms in the same volume Cl = solid Au = gold Diffusion & Effusion Theory states-no force of attraction between gas particles. Gas particles can move easily past each other. The space into which a gas flows is already occupied by another gas. This random motion causes gases to mix until they are evenly distributed. Diffusion- Movement of one material through another. Move from high concentration to low concentration.

The rate of diffusion depends on the mass of the particles. For lighter particles to have the same average kinetic energy as heavier particles, they must have a greater velocity (KE = 1/2 mv 2. Effusion is process related to diffusion. Graham's Law of Effusion- States that the rate of effusion for a gas is inversely proportional to the square root of the molar mass (GM). This law applies to rates of diffusion. Using this law you can setup a proportion to compare the diffusion rates for 2 gases Formula: Gas pressure.

Pressure-defined as force per unit area. Gas particles exert pressure when they collide with the walls of their container Atmospheric pressure. Because the particles in air move in every direction, they exert pressure in all directions Measuring Air pressure. Barometer-instrument used to measure atmospheric pressure.

Height of the Hg in the tube is determined by 2 forces. Gravity pushing down on Hg; Humidity / temperature causes air pressure to change. An increase in air pressure causes Hg to rise, a decrease it falls. Manometer-instrument used to measure gas pressure in a closed container. Flask in connected to a U-tube containing Hg. Before the gas is released Hg is at the same height in arms.

After the gas is released the height in the 2 arms are no longer equal. If gas pressure atmospheric pressure the liquid on the left decreases and the liquid on the left increases Pgas = Patm+Phg. If gas pressure atmospheric pressure the liquid on left increases and liquid on right decreases Pgas = Patm-PHg Units of Pressure Pascal (Pa) -Si unit of pressure 1 Pa = 1 N / M 2 Other units: mmHg, torr, atmosphere (atm), psi, kilo pascal (kPa l) 1.00 atm = 760 torr = 101.3 kPa = 14.7 psi (at 0 C) Dalton's Law of Partial Pressure. Law states that the total pressure of a mixture of gases is equal to the sum of the pressures of all the gases in the mixture. Partial pressure depends on: Number of moles of gas, size of the container, temperature of mixture.

It doesn't depend on the identity of the gas Patm = p 1 + p 2 + p 3... When gases are made in a lab they are made over water, so the gas is saturated with water. Pgas = Total-Water. Water pressure will be given to you Practice problems on pg 392 4-6. Water pressure will be given to you.

Forces of attraction. Intramolecular force-Force within atoms, ions, molecules (bonding). Ex: dispersion, H-bonds. Intermolecular force-force between atoms, ions, or molecules. Ex: dispersion, H-bonds Dispersion force. Weak forces that result from temporary shift of electrons in the electron cloud.

Occurs when 2 non polar molecules are in close contact or collide. Electrons in the cloud repel each other causing the electrons to have a higher density in one region of each cloud. Forms temporary dipole; When the dipoles are close together a weak dispersion force exists; Occurs in identical non polar molecules Dipole-Dipole forces. Attractions between oppositely charged regions of polar molecules. Contain permanent dipoles.

Neighboring polar molecules orient themselves so that charged regions line up... Forces are usually stronger than dispersion forces. Pg 394-5: pictures of bonds Hydrogen bonds. A dipole-dipole attraction that occurs between molecules containing a hydrogen atom bonded to a small electronegative atom with at least one lone electron pair. Hydrogen must be bonded to Fluorine, Oxygen, Nitrogen. H has a positive charge, O has a negative charge so that one molecule is attracted to the O in another molecule Liquids Kinetic Molecular Model.

Have a constant motion. No foxed position so they take the shape of a container. Force of attraction is greater than in gases so liquids have a fixed volume Density & Compression. At 25 C and 1 atm, liquids are closer than gases because of their greater intermolecular forces holding them together. Like gases, liquids can be compressed, but only by a small volume because the molecules are relatively close together Fluidity.

Ability to flow. Liquids can diffuse through other liquids but at a slower rate than gases because of intermolecular forces. Liquids are less fluid than gases Viscosity. Is the measure of the resistance of a liquid to flow. The attractive force slows their movement. Viscosity is determined by the type of intermolecular forces, shape of the particle, and temperature.

Warm-faster flow, cooler- slow flow larger / longer -slow small / short -fast Surface tension. Particles in the middle of a liquid can be attracted to particles all around them. Particles at the surface have no attractive force from above to balance the attraction for below; so there is a greater force pulling down on particles at the surface. Surface tension- energy required to increase the surface area of a liquid by a given amount. Generally the stronger the attraction between particles the greater the surface tension. Water has a high surface tension because it can form multiple hydrogen bonds.

Surfactants- compound that lower the surface tension of water (break hydrogen bonds) Capillary action. A narrow container; surface of water forms a meniscus 2 forces cause this: . Cohesion- force of attraction between identical molecules. Adhesion- force of attraction between different molecules. The adhesive force between the water and glass is greater than the cohesive force between the water; so water rises along the inner walls of the cylinder Kinetic molecular theory. Are in constant motion.

Have a strong attractive force between particles, this limits the motion, causing more order. Because this solids have a definite shape and volume are much less fluid Density of solids. Particles are packed closely together so they are more dense. Because most solids sink, except water has fewer particles in given volume Crystalline solids. Solid where atoms, ions, or molecules are arranged in a orderly, geometric shape.

Unit cell- smallest arrangement of connected points that can be repeated in 1 direction to form a lattice. Shape of a crystal is determined by the type of unit cell. 7 different crystals on pg 401. Crystalline solids can be classified into 5 groups based on the type of particles they contain.

Pg 403 table 13-4. Amorphous solids- particles are not arranged in a regular repeating pattern. Forms when the motion material cools too quickly and the crystals don't have time to form Phase changes MeltingThsolidThliquid. Solids well absorb heat energy and break the bonds that are holding them together.

Temperature does not change as the solid is melting. Amount of energy needed depends on the strength of that bond. Melting point- temperature at which the liquid phase and solid phase of a given substance can coexist Vaporization liquidThgas. Once all ice melts the addition of energy increases the kinetic energy of the liquid and the temperature increases. Evaporation- when vaporization occurs only at the surface of a liquid. Liquid molecules must absorb energy, break forces of attraction, and enter gas phase ex: sweat.

In a closed container water collects above that liquid and exerts pressure on the surface of the liquid called vapor pressure. Boiling point- temperature at which vapor pressure of a liquid equals the external or atmospheric pressure Sublimation solidThgas. Ex: solid iodine, solid carbon dioxide, mothballs, air freshener. Absorb energy Condensation gasThliquid.

Loses energy, velocity decreases when molecules collide they are more likely to form bonds or increase the force of attraction, and changes to the liquid phase. Reverse vaporization Deposition gasThsolid. Reverse of sublimation energy is released ex: frost & snow Freezing liquidThsolid. Energy is lost, velocity decreases bonds form and hold the molecules in a fixed position.

Reverse of melting. Freezing point- temperature at which a liquid is converted in to a solid. Is equal to the melting point Phase diagrams. Graph of pressure vs. temperature that shown in which phase a substance exists.

Under different conditions of temperature and pressure. 3 curves separate the regions from one another solid, liquid gas. On the curve 2 phases of matter exist. Triple point-temperature at which 6 phase changes can occur and 3 phases of matter can exist.

Critical point- indicates the critical temperature and pressure above which a substance cannot exist as a liquid.