Unit 4: Gas Laws Matter and Phase Changes
ENERGY and matter
What is 에너지
A fundamental property of the universe that cannot be easily defined.
Energy No one knows what energy is, only what it does or has the potential to do. It is a property of matter that can sometimes be measured.
the operational definition: The ability to do work*. *When an object moves by a force acting on it. work and energy are interchangeable.
Energy (and work) is measured in joules (J) 1 J = 1 kg m 2 s 2
How can energy affect matter? STATES OF MATTER
Matter Matter is anything that takes up space. Matter mainly exists in three states: Solids Liquids Gases
Kinetic-molecular theory of matter The particles that make up matter are in constant motion (Brownian motion). How much energy the molecules have determines whether the matter is solid, liquid, gas or plasma
The Kinetic Molecular Theory This can be used to explain the properties of solids, liquids and gases in terms of the energy of particles and the forces that act between them.
Energy acts on a system add energy remove energy
Based on the kinetic-molecular theory PROPERTIES OF SOLID, LIQUIDS AND GASES
Solids
two types of solids Crystalline: Consist of crystals Particles are arranged in orderly, repeating, geometric pattern Amorphous: Particles are arranged randomly.
Amorphous Solids Glass is made by cooling molten materials in a way that prevents them from crystallizing. This allows the glass to appear to be a solid yet be transparent. Other examples include plastics and semiconductors (used in electronics).
Properties of solids based on the Kinetic-molecular theory: Solids maintain a definite shape and volume. Crystalline solids have definite melting points. Incompressible Extremely low rate of diffusion.
DEFINITE MELTING POINT melting is the physical change from a solid to a liquid by the addition of heat melting point is the temperature at which a solid becomes a liquid supercooled liquid is a substance that retains certain properties of a liquid even at temperatures where it appears to be a solid (glass, plastics, )
Liquids Liquids are fluids, which means they can flow and take the shape of their container
Properties of liquids based on the Kinetic-molecular theory: liquids have relatively high densities. Surface tension & capillary action. Relatively incompressible Ability to diffuse.
Liquids exhibit surface tension, a force that tends to pull adjacent parts of a liquid s surface together, thereby decreasing the surface area to its smallest possible size. Surface tension resists penetration of objects into a liquid. Liquids have capillary action, the attraction of the surface of a liquid to the surface of a solid.
Gases A gas will expand or contract to fill its container. Molecules of a gas are spread far apart and have greater kinetic energy Almost all the volume of a gas is empty space!
Assumptions of the Kineticmolecular theory of Gases 1. Molecules are very far apart, most of the volume is empty space. 2. No kinetic energy is lost when gas particles collide (elastic collisions) 3. Gas particles are in constant, random motion. 4. There are no forces of attraction between the particles, nor do they repel each other. 5. The temperature of the gas depends on the average kinetic energy of the particles All gases at the same temperature have the same average kinetic energy.
All gases at the same temperature have the same average kinetic energy. Kinetic energy is measured by the following formula: KE = 1/2mv 2 At the same temperature, then, lighter gas particles (such as hydrogen) have higher average speeds than heavier particles (such as oxygen). A proportional relationship!
Physical properties of gases due to the kinetic-molecular theory: Expansion Fluidity Low density Compressibility Diffusion and effusion
Expansion Gases take the shape of any container and expand to fill any space. Fluidity Gas particles glide easily past each other. This allows them to flow. Low density Gas particles are 1/1000 the density the same material is in liquid or solid phase.
Compressibility Gas particles can be pressed close together. Diffusion and effusion Gas particles spread out from an area of high concentration until they are distributed evenly.
Phase Changes (or how energy affects matter) A change from one state of matter (solid, liquid, gas) to another. a Physical change because appearance is changed, not the chemical make-up; it s reversible. Thermal energy is either absorbed (molecules speed up) or released (molecules slow down)
A phase diagram is a graph of pressure versus temperature that shows the conditions under which the phases of a substance exist.
Melting Phase change from a solid to a liquid Molecules speed up, move farther apart, and absorb heat energy
Freezing Phase Change from a liquid to a solid Molecule slow down, move closer together and release heat energy.
Vaporization (Boiling) Phase change from a liquid to gas. It occurs at the boiling point of matter. Molecules speed up, move farther apart, and absorb heat energy.
Evaporation Phase change from a liquid to a gas on the surface of a liquid (occurs at all temperatures) Molecules speed up, move farther apart, and absorb heat energy.
Condensation Phase change from a gas back to a liquid. Molecules slow down, move closer together and release heat energy.
Sublimation Phase change from a solid directly to a gas without becoming liquid first. Molecules speed up, move farther apart, and absorb heat energy.
Crystallization Phase change from a gas to a solid. Molecules slow down, move closer together and release heat energy.
Energy and Kinetic-molecular theory BOILING AND FREEZING
(influenced by pressure)
Boiling is the conversion of a liquid to a vapor (gas). Boiling point of a liquid is the temperature at which the equilibrium vapor pressure of the liquid equals the atmospheric pressure. Energy must be added continuously in order to keep a liquid boiling.
Vapor pressure The pressure exerted by the vapor in equilibrium with a liquid in a closed container at a given temperature is called the liquid s vapor pressure When a liquid vaporizes in a closed container, gas molecules cannot escape. As these gas phase molecules move about, they occasionally collide with the surface of the condensed phase, and in some cases, these collisions will result in the molecules re-entering the condensed phase.
When the rate of condensation becomes equal to the rate of vaporization, neither the amount of the liquid nor the amount of the vapor in the container changes. The vapor in the container is then said to be in equilibrium
Boiling When the vapor pressure increases enough to equal the external atmospheric pressure, the liquid reaches its boiling point. The boiling point is the temperature at which its equilibrium vapor pressure is equal to the pressure exerted on the liquid by the atmosphere.
The boiling points of liquids are the temperatures at which their equilibrium vapor pressures equal the pressure of the surrounding atmosphere. Normal boiling points are those corresponding to a pressure of 1 atm (101.3 kpa.)
Heat of vaporization is the amount of heat needed to vaporize one gram of a liquid at its boiling temperature. The energy used here is the heat of vaporization Heat of vaporization for water is 540 calories/gram.
Cooling via vaporization Molar enthalpy of vaporization: The energy change associated with the vaporization process needed to vaporize one mole of a liquid at the liquid s boiling point. Vaporization is an endothermic process. For example, the vaporization of water at standard temperature is represented by H 2 O(l) H 2 O(g) Hvap = 44.01 kj/mol When water (say on your skin) evaporates, it removes heat.
Melting & Freezing Melting is the physical change of a solid becoming a liquid. Freezing is the physical change of a liquid becoming a solid. Also called solidification or fusion. Freezing point is the temperature at which a liquid becomes a solid and the solid and the liquid are at equilibrium.
Enthalpy of fusion: The amount of heat required to change one mole of a substance from the solid state to the liquid state is the enthalpy of fusion, ΔH fus Fusion (melting) is an endothermic process. Heat of fusion is the amount of heat needed to melt one gram of a solid. For ice, this is 80 calories/gram.
As a solid begins the process of melting, the temperature of the solid stops rising, despite the continual input of heat, and it remains constant until all of the solid is melted. Only after all of the solid has melted will continued heating increase the temperature of the liquid
temperature remains constant while the change of state is in progress. Heat of vaporization for water is 540 calories/gram.
Investigate!