chapter 10 - h. chem

the kinetic-molecular theory of matter

• based on the idea that particles of matter are always in motion

• used to explain the properties of solids, liquids, and gases in terms of the energy of particles and the forces that act between them

• theory of gases

  • an ideal gas is a hypothetical gas that perfectly fits all the assumptions of the kinetic-molecular theory

  • based on the following five assumptions

    • gases consist of large numbers of tiny particles that are far apart relative to their size

      • these particles (normally molecules and atoms) occupy a volume that is about 1000 times greater than the volume occupied by an equal number of particles in the liquid or solid state

      • these particles of gases are much farther apart than molecules of liquids and solids

      • most of the volume occupied by a gas is empty space = gases have lower density than liquids and solids & gases are easily compressed

    • collisions between gas particles and between particles and container walls are elastic collisions

      • elastic collisions = collision in which there is no net loss of total kinetic energy

      • kinetic energy is transferred between two particles during collisions but the total kinetic energy of the two particles remains the same as long as the Kelvin temperature is constant

    • gas particles are in continuous, rapid, random motion; they have kinetic energy (energy of motion)

      • move in all directions

      • their ke overcomes the attractive forces between them (except !! near the temperature at which the gas condenses and becomes a liquid)

    • no  forces of attraction between gas particles

      • their collisions = no sticking but bouncing apart

    • temperature of a gas depends on the average kinetic energy of the particles of the gas

• KINETIC ENERGY EQUATION:

  • m = mass of the particle (kg)

    • because all particles of a specific gas have the same mass; their ke's are dependent only on their speeds

  • v is its speed 

  • BIG IDEA: the avg. speeds and kinetic energies of gas particles increase with an increase in temp and decrease with a decrease in temp (all related to Kelvin)

  • all gases at the same temp have the same ke 

    • = same temp but lighter gas particles (higher average speed) while heavier gas particles (lower average speed)

• kmt and nature of gases

  • only applies to ideal gases that don't actually exist

    • but many gases behave nearly identical if pressure is not very high and the temp is not too low

  • expansion

    • gases do not have definite shape or volume = they completely fill any container in which they are enclosed and take its shape

    • gas particles move rapidly in all directions (assumption 3) without significant attraction between them (assumption 4)

  • fluidity

    • attractive forces are insignificant (assumption 4) = gas particles glide past one another (ability to flow like liquids)

    • liquids and gases flow = both called fluids

  • low density

    • 1/1000 the density of the same substance in the liquid or solid state

    • bc particles are much farther apart (assumption 1)

  • compressionability 

    • the gas particles (which are naturally further apart) are pushed clower together

    • the volume of the given sample of gas is greatly decreased

    • why pressurized containers can hold more than 100 times the number of gas particles than nonpressurized containers could

  • diffusion and effusion

    • diffusion = such spontaneous mixing of the particles of two substances caused by their random movement 

      • result of third assumption (random and continuous motion)

      • readily diffuse and mix together bc of the rapid motion of the molecules and the empty space between them

    • effusion = process by which gas particles pass through a tiny opening

      • rates of effusion of diff. gases = proportional to velocities of their particles

      • BIG IDEA: molecules of low mass effuse faster than molecules of high mass

• deviations of real gases from ideal behaviors

  • real gas = gas that does not behave completely according to the assumptions of the kinetic-molecular theory

    • at very high pressures and low temps = gas will be closer together and their kinetic energy will be insufficient to overcome the attractive forces

  • kmt holds truer for games whose particles have little attraction for each other 

    • noble gases (helium and neon) show ideal gas behavior over a WIDE RANGE of temps. and pressures

      • particles are monatomic (nonpolar) 

    • particles of gases (nitrogen and hydrogen) are nonpolar diatomic molecules

      • the behavior of these approximate that of the ideal gas under CERTAIN conditions

      • BIG IDEA: the more polar the molecules of a gas are, the greater the attractive forces between them and the more the gas will deviate from ideal gas behavior

liquids

• least common state of matter bc a substance can exist in the liquid state only within a relatively narrow range of temperatures and pressures

• properties of liquids and kmt

  • definite volume and indefinite shape (of container)

  • motion and arrangements

    • particles in a liquid are in constant motion but they are closer tg than they are in a gas = attractive forces between particles in a liquid are more effective than those between particles in a gas

      • attraction is caused by…

        • intermolecular forces (dipole/dipole forces, london dispersion forces, and hydrogen bonding)

        • some molecules at the surface of a liquid can have neough kinetic energy to overcome these forces and enter the gas state

    • more ordered than gases because of the stronger intermolecular forces and the lower mobility of the liquid particles

      • the particles are not boudn tg in fixed positions but instead…. MOVE ABOUT CONSTANTLY

        • explains why they and gases are called fluids

          • fluid = substance that can flow and therefore take the shape of its container

          • most liquids naturally flow downhill because of gravity but soem liquids can flow in toher directions

            • liquid He can flow uphill if near absolute zero (theoretical temperature)

  • relativley high density

    • the higher density is a result of the close arrangement of liquid particles

    • most substances are only slightly less dense (about 10%) in a liquid state than in a solid state

    • water is one of the only substances that becomes less dense when it solidifies

    • at the same temp and pressure, different liquids can differ greatly in density

  • relative incompressibility

    • when liquid water at 20 degrees celsius is compressed by a pressure of 1000 atm = its volume decreases by only 4%

      • typical of all liquids and is similar to the behavior of solids

    • liquids are much less compressible than gases because liquid particles are more closely packed together

    • can transmit pressure equally in all directions like gases

  • ability to diffuse

    • any liquid gradually diffuses throughout another liquid in which it can dissolve

      • the constant, random motion of particles causes diffusion in liquids as it does in gases

    • diffusion in liquids is much slower than in gases bc the particles are so close together plus the attractive forces between teh particles of a liquid slow this movement down

    • BIG IDEA: as the temp of a liquid increases, diffusion occurs more rapidly

      • bc the average kinetic energy and thereby the average speed of the particles has increased

  • surface tension !!

    • = a force that tends to pull adjacent parts of a liquid's surface together, thereby decreasing surface area to the smallest possible size

    • results from the attractive forces between particles of a liquid

      • BIG IDEA: the higher the force of attraction, the higher the surface tension

    • water has a higher surface tension because the hydrogen bonds water molecules can form with each other only occur with the water molecules BELOW; there is nothing to bond to above = water molecules are drawn together and toward the body of the liquid, creating a high surface tension

    • capillary action = the attraction of the surface of a liquid to the surface of a solid is related to the surface tension

  • evaporation and boiling

    • the process by which a liquid or solid changes to a gas = vaporization

    • the process by which particles escape from the surface of a nonboiling liquid and enter the gas state = evaporation

      • occurs bc particles of a liquid have different kinetic energies

        • BIG IDEA: particles with higher-than-average energies move faster

        • some surface particles with higher-than-average energies can overcome the intermolecular forces that bind them to the liquid

    • boiling is the change of a liquid to bubbles of vapor that appear throughout the liquid

  • formation of solids

    • when a liquid is cooled, the average energy of its particles decreases

    • if the energy is low enough, attractive forces pull the particles into an even more orderly arrangement, therefore, a solid

    • freezing/solidification: the physical change of a liquid to a solid by removal of energy as heat

    • all liquids freeze but the temperatures vary

solids

• properties of solids and kmt

  • particles of solids are more closely packed than those of a liquid or gas

  • intermolecular forces between particles are therefore much more effective in solids

    • interparticle attractions (dipole-dipole attractions, london dispersion forces, and hydrogen bonding) exert stronger effects in solids than in the corresponding liquids or gases

    • attractive forces tend to hold the particles of a solid in rleatively fixed positions with only vibrational movement around fixed points

      • bc the motions of these particles are restricted like this, solids are more ordered than liquids and are much more ordered than gases

  • two types of solids:

    • crystalline solids

      • they consist of crystals

        • crystal is a substance in which the particles are arranged in an orderly, geometric, repeating pattern

    • amorphous solids

      • one in which the particles are arranged randomly 

  • definite shape and volume

    • solids maintain a definite shape without a container

      • crystalline solids are geometrically regular

      • even the fragments of a shattered crystalline solid have distinct geometric shapes that reflect their internal structures

      • amorphous solids maintain a definite shape but do not have the distinct geometric shapes of crystalline solids

    • the volume of a solid only changes slightly with a change in temp. or pressure

      • solids have definite volume bc of how tightly packed their particles are = very little empty space into which the particles can be compressed

      • crystalline solids generally do not follow bc their particles are held in relatively fixed positions

  • definite melting point

    • melting - the physical change of a solid to a liquid by te addition of energy as heat

    • melting point - the tmep at which a solid becomes a liquid 

      • at this temp, the ke of the particles within the solid overcome the attractive forces holding them together 

      • the particles can then break out of their positions in crystalline solids, which have definite melting points

      • amorphous solids, however, have no definite mp

        • have the ability to flow over a arnage of temps

        • classified as supercooled liquids

          • which are substances that retain certain liquid properties even at temps at which they appear to be solid

        • these properties exist bc the particles in amorphous solids are arranged randomly but unlike the particles in a true liquid amorphous solids’ particles are not constantly changing their positions

  • high density and incompressibility

    • have high density because of how close their particles are packed together

    • generally less compressible than liquids and are sometimes considered incompressible

      • some solids (wood and cork) may seem compressible but they are not

        • they have pores that are filled with air

          • when subjected to intense pressure, the pores compressed, not the solid matter in the wood or cork itself

  • low rate of diffusion

    • incredibly low (zinc plate and copper plates will diffuse very little if put together for a long time)

• crystallien solids

  • exist as single crystals or as groups of crystals fused together

    • crystal structure

      • the total three-dimensional arrangement of particles of a crystal

      • lattice - arrangement of particles in the crystal can be represented by a coordinate system

        • contains many unit cells packed together

      • unit cell - the smallest portion of a crystal lattice that shows the 3d pattern of the entire lattice

    • bdining forces in crystals

      • ionic crystals

        • consist of postive and negative ions arrangd in a regular pattern

          • these are strongly bonded together

            • gives them certain properties:

              • hard 

              • brittle

              • high melting points

              • good insulators

        • monatomic or polyatomic

        • ionic crystals form when g1/2 metals combine with g16/17 nonmetals or nonmetallic polyatomic ions

      • covalent network crystals

        • each atom is covalently bonded to its neartes neighboring atoms

          • covalent bonding extends throughout a network that includes a very large number of atoms

          •  in the subscript = indicates that the component within the parentheses extends indefinitely

        • properties:

          • very had

          • brittle

          • high melting points

          • nonconductors/semiconductors

      • metallic crystals

        • consists of metal cations surrounded by a sea of delocalized valence electrons

        • the electrons come from the metal atoms and belong to the crystal as a whole

        • freedom of these delocalized electrons to move explains the high electric conductivity of metals

      • covalent molecular crystals

        • covalently bonded moelcuels held together by intermolecular forces

          • nonpolar molecules - weak london dispersion forces between moleucles

          • polar molecules  - molecules held together by dispersion forces, by stronger dipole-dipole forces and maybe even hydrogen bonding

        • forces that hold polar and nonpolar molecules together in the structure are much weaker than the covalent chemical bonds between the atoms within each molecule

          • = low mps

          • easily vaporized

          • relatively soft

          • good insulators

  • amorphous solids

    • amorphous = without shape

    • not arranged in a regular pattern

changes of state

• phase - any part of a system that has uniform composiotn and properties

• condensation - the process by which a gas changes to a liquid

• vapor - a gas in contact with its liquid or solid phase 

• the rate of condesnation equals the rate of evaporation

• equilibrium - a dynamic condition in which two opposing changes occur at equal rates in a closed system

  • equilibrium vapor pressure of a liquid

    • the pressure exerted by a vapor in equilibrium with its corresponding liquid at a given temperature

    • increases in evp is proportional to the the increase in temp 

      • increasing the temp = increases avg ke in the molecules = increases the number of molecules that have enough energy to escape from the liquid phase into the vapor phase = increase evaporation = increases the number of molecules in the vapor phase = increases evp

    • every liquid has a specific evp at a give tmep bc of characteristic forces of attraction between particles

    • low percent evaporation = low evp

    • volatile lqiudis - liquids that evaporate readily

      • relatively weak forces of attraction between their particles

    • nonvolatile iquids do NOT evaporate readily & have strong attractive forces between their particles

• boiling

  • evp can be used to explain boiling

    • the conversion of a liquid to a vapor within the liquid as well as its surface

  • temp increases = evp increases

  • boiling point - is the temperature at which the evp of the liquid equals the atmospheric pressure

    • lower atmospheric pressure - lower bp

    • at the boiling point = all of the energy absorbed is used to evaporate the lqiiuid and the temp remains constant as long as the pressure does not change

      • if the pressure above the lqiuid increases - temp of the liquid will rise until the vp equals the new pressure and the liquid boils once again

• energy and boiling

  • energy must be added continuously in order to keep a liquid boiling

    • stops boiling almost immediately after it is removed from a heat source

    • temp of a boiling liquid and its vapor are the same temp

    • the temp at the boiling point remains constant despite the continuous addition of energy

      • added energy is used to overcome the attractive forces between molecules of the liquid during the l-g change and is stored in the vapor as potential energy

• molar enthalpy of vaporization

  • the amount of energy as heat that is needed to vaporize one mole of a liquid at the liquid’s bp at cosntant pressure    

  • mev’s magnitude is a measure of the attraction between particles of the liquid

    • stronger this attraction is = more energy is needed to over come = higher mev

      • = each liquid has its own mev

      • water’s mev’s is really ihgh

        • makes water an effect cooling agent

        • water evaporates from your skin = escaping moelcuels carry a great deal of energy as heat away with them

    • higher temp - greater portion of molecuels have the ke required to escape from the liquid surface and become vapor

• freezing and melting

  • the physical change of a liquid to a solid is = freezing

    • involves a loss of energy in the form of heat by the liquid

      • liquid -> solid and energy

    • change at a constant temperature

    • normal fp is the temp at which the solid and liquid are in equilibrium at 1 atm (760 torr or 101.3 kPa) pressure

      • = energy loss during freezing is a loss of potential energy that was present in the liquid 

      • = significant increase in particle order bc the solid state of a substance is much more ordered than the liquid state (even at the same temp)

  • melting also occurs at a constant temperature

    • as a solid melts it continuously absorbs energy as heat 

      • solid + energy -> liquid 

  • pure crystalline solids have the same fp and mp bc at equilibrium the metling and freezing occur at equal rates

    • solid + energy -><- liquid

• molar enthalpy of fusion

  • the amount of eneryg as heat required to melt on mole of solid at the solid’s mp is the solid molar enthalpy of fusion

  • energy absorbed increaseds the solid’s pe as its particles are pulled apart and the attractive forces holding them tg are overpowered

  • there is a significant decrease in partilce order as the substance make the transformation from solid to liquid

  • the magnitude of the mef depends on the attraction between solid particles

• sublimation and deposition

  • sufficiently low tmeps and pressure conditions = a liquid cannot esit

  • a solid susbtance in these conditions exists in equilibrium with its vapor instead

    • solid + energy -><- vapor

  • sublimation - the change of state from solid directly to gas

  • deposition - the change of state from a gas directly to a solid

phase diagrams

• a grpah of pressure versus temperature that shows the conditions under which the phases of a substance exist 

  • also reveals how the states of a system change with changing temperature or pressure

  • triple point - indicates the temperature and pressure conditions at which the solid/liquid/and vapor of the substance can coexist at equilibrium

  • critical point - indicates the critical temp and critical pressure

    • ctricitcal temperature = the temperature above which the susbtance cannot exist in a liquid state

    • critical pressure - the lowest pressure at which the substance can exist as a liquid at the critical temperature

water

• essential to human life and is incredibly abundant

• structure of water

  • two atoms of hydrogen and one atom of oxygen bonded by polar-covalent bonds

    • bent 109.5, actually 105

      • sp^3 hybridization

  • molecules in solid or liquid water are linked by hydrogen bonding

    • number of linked molecules decreases = temp increases

      • bc increase in ke make hydrogen bond formation difficult

      • usually four-eight molecules per group in liquid water

        • allows for water to exist the way it does at room temp (liquid)

  • empty spaces in the hexagonal arrangement of ice accounts for the low density of ice

    • as ice is heated, the increased energy of the molecules causes them to move and vibrate more vigorously 

      • when the mp is reached, the energy of the molecules is so great tht atht e rigid open structure ice crystals break down and the ice is now liqiuid water

      • the rigid open structure of the ice has broken down, water molecules can crowd together (high density)

      • hydrogen bonds are why the bp of water to be much higher

  • physical properties

    • pure liquid water is…

      • transparent

      • odorless

      • tasteless

      • almost colorless

        • any observable odor or taste is caused by impurities (dissovled minerals, liquids, or gases)

    • freezes and ice melts at 0 degrees celsius

    • 6.0009 kJ/mle is he mef

      • relatively larged compared to that of other solids

    • expands involume as it freezes bc of its molecules forming an open rigid structure

      • density of liquid water > density of frozen water (ice)

        • lower desnity of ice = explains why it floats in liquid water

    • water boils at 100 degrees celsius

      • at this temp the mev is about 40.79 kJ/mol

      • both the bp and the mev of water are quite high compared with those of nonpolar susbtances of comparable molecular mass (ex. methane)

        • because of the strong hydrogen bonding that much be overcome for boiling to occur

        • high mev makes water useful for household steam-heating systems

          • steam sotres lots of energy as heat 

          • when the steam condenses in radiators, the energy is released