Kinetic-Molecular Theory Flashcards
Vapor
Gas produced from a liquid or solid under normal conditions (vapor pressure must be equal to that of the atmosphere)
Ideal gas
Imaginary gas that perfectly fits all of the assumptions of the kinetic-molecular theory. A gas that behaves independently of its environment (mostly noble gases and diatomic molecules)
Assumptions
A. Gases consist of a large number of tiny particles with much space in between
B. Collisions occur as elastic collisions, no net loss of kinetic energy
C. Gases are in constant, rapid, random motion
D. No attractive or repulsive forces between gases
E. Average kinetic energy is constant and depends on temperature
Johann van Helmont (1662)
Used “gas” to describe the most energetic phase of matter (derived from the Greek “chaos”, original matter on Earth)
Expansion
Allows gas to take the shape of the container
Pressure
Force per unit area
Barometer
Instrument to measure atmospheric pressure
Manometer
Instrument to measure pressure produced in a lab reaction or tank
Density
Measure of mass per volume (for gases g/l)
Diffusion of gases
Random spreading of gas in uniform dimensions
Effusion
Controlled escape of gas through a hole or pore
Condensation temperature
The point where gas molecules form a liquid
Real gases
Don’t behave completely according to the kinetic-molecular theory; intermolecular forces provide interactions that may change gas properties or behaviors
Van der Waal forces
General terms for weak intermolecular attractions between molecules
London dispersion forces
Caused by motion of electrons, increases as # of electrons increases. Can occur in any molecule to induce temporary polar behavior.
Dipole-dipole interactions
Electrostatic attractions based on distinct polarity of molecules
Hydrogen bonding
Hydrogen bonded to a highly electronegative atom is attracted to an unshared pair of electrons in a nearby molecule
Definite volume
Won’t expand, will take the shape of the container (liquid)
Incompressible (liquid)
Liquids equalize pressure, particles are close together
Fluidity
Ability to flow and be poured
Diffusion of liquid
Particles have motion and randomly spread, but more slowly than gases
Cohesion
Force of attraction between molecules of the same substance
Adhesion
Force of attraction between different substances
Surface tension
Force that pulls adjacent parts of a liquid’s surface together resulting in the smallest surface area possible
Capillary action
Attraction of the surface of a liquid to the surface of a solid; tends to move a liquid against gravity
Definite shape
Exists in a fixed form
Definite volume
All surfaces are free and defined
Incompressible (solid)
Structure can’t be altered without a phase change
Amorphous
Without form, that which appears to have inconsistent or random particle arrangement
Crystal
Homogeneous substance bounded by plane surfaces making definite angles, giving a geometric form
Crystal lattice
Pattern of pints in a crystal
Crystallography
Science of classifying crystal shape with X-ray diffraction
Unit cell
Smallest portion of a crystal lattice that shows the pattern of the lattice
Ionic crystal
Atoms held together by electrostatic forces, hard, high melting point (ex: NaCl)
Covalent crystal
Covalent bonds, like one big molecule, very high melting point (ex: diamond)
Metallic crystal
Atoms of metal sit on lattice points and outer electrons float around the lattice, very dense, high melting point (ex: copper)
Molecular crystal
Contain recognizable molecules in their structure, held together by van der Waal forces and hydrogen bonding, soft, low melting point (ex: rock candy/sucrose)
Isometric (cubic)
3 axes at right angles are equal lengths
Iso = same
Hexagonal
3 equal axes intersect at 60 degree angles, a vertical axis of a different length is at a right angle to these axes (note: four axes)
Tetragonal
3 axes at right angles, only two are at equal lengths
Trigonal
3 equal axes at non right angles
Orthorhombic
3 unequal axes at right angles
Monoclinic
3 unequal axes, one non right angle
Triclinic
3 unequal axes, three non right angles
Supercooled liquids
Substances that retain certain liquid properties even at temperatures at which they appear to be solid
Gas liquification
Michael Faraday 1823, English chemist; cooling and compressing gases at the same time would yield a liquid; a critical point for pressure and temperature. Faraday liquified Cl, CO2, HSO2, and HBr
Volatile liquids
Liquids that easily evaporate due to weak intermolecular forces
Molar heat of vaporization
Amount of heat energy needed to melt one mole of solid at its melting point
Molar heat of fusion
Amount of heat energy needed to melt one mole of solid at its melting point
Closed system
No substances are added or lost, but energy changes freely occur
Physical equilibrium
State in which 2 opposing physical changes occur at equal rates in the same system
Equilibrium vapor pressure
Pressure exerted by a vapor in equilibrium with its liquid
Dynamic equilibrium
Forward and backward reaction reached a point where there is no further change in the quantities of substances, but the reactions are still taking place
Henri Le Châtelier’s Principle 1884
A. If a system is subjected to stress, the equilibrium will change in order to relieve the stress
B. Equation will shift based on concentration, partial pressure of gas (molecule # is important), and temperature; consider if the reaction is endothermic (H is positive) or exothermic (H is negative)
C. Catalysts don’t change equilibrium, they only speed up the rate at which dynamic equilibrium is met
Phase
Part of a system that has uniform properties and composition; phase changes are physical changes based on heat content
Triple point
The temperature and pressure conditions at which the solid, liquid, and vapor of the substance can coexist at equilibrium
Critical point
Indicates a critical temperature and pressure
Critical temperature
Temperature above which the substance cannot exist as a liquid
Critical pressure
Lowest pressure at which the substance can exist as a liquid at the critical temperature
Supercritical fluid
A material that can be either liquid or gas, used above the critical point where gases and liquids can coexist. This fluid has gaseous properties of penetrating substances, and liquid properties of dissolving materials into their components
Density of water
Liquid: 1 g/ml, less dense as ice (.92g/ml), hence it floats on water
High specific heat index of water
Absorbs and releases a lot of energy when getting hot or cold; helps regulate air temperature
Water pH
7
Percentage of water in life
70-90% of all living matter
Hydrate
Crystallized substance that contains a definite # of water molecules in the structure (if water lost, anhydrous)
Hydrolysis
Chemical reaction in which water molecules are split and the parts are added to the products
Effervescence
Rapid evolution of a gas when a dry substance reacts with water
Efflorescence
The loss of some or all of the water of crystallization when exposed to a warm or dry environment
Deliquescence
Absorbing water by a substance from the air, making it moist or resulting in a solution
