Chapter 10: States of Matter

Kinetic-Molecular Theory

Kinetic-molecular theory- based on the idea that particles of matter are always in motion
- used to account for the behavior of atoms and molecules that make up matter
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The states of matter (solid, liquid, gas) are determined by a battle between the kinetic energy of the substance and the strength of the intermolecular attractive forces...
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GASES
KINETIC MOLECULAR THEORY OF GAS

ideal gas- a hypothetical gas that perfectly fits all the assumptions of the kinetic-molecular theory (5 assumptions)
1. Gases consist of large numbers of tiny particles that are far apart relative to their size
2. Collisions between gas particles and between particles are container walls are elastic collisions*
3. Gas particles are in continuous, rapid, random motion. They therefore possess kinetic energy- energy in motion
4. There are no forces of attraction between gas particles (they just bounce apart)
5. The temperature of a gas depends on the average kinetic energy of the particles of the gas
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*elastic collisions- one in which there is no net loss of total kinetic energy
*Kinetic Energy= 1/2mass*speed squared
(all particles of a specific gas have same mass and the kinetic energy therefore depends only on speed, so at the same temperature lighter particles have faster speeds)

The state of matter is determined by a battle between the kinetic energy of the substance and the strength of the intermolecular attractive forces
(PICTUREEE OF MOVEMENT, INTER FORCES AND STRENGHT)
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LIQUIDS
PROPERTIES OF LIQUIDS AND KINETIC MOLECULAR THEORY

liquid- has a definite volume and takes shape of its container
-continually in motion and particles closer together
-attractive forces are more effective cause by the intermolecular forces
- are fluids
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fluid- substance that can flow and therefore take shape of its container
relatively high density- only 10% less then solid form
relatively incompressibility- particles are close together and can transmit pressure equally in all directions
ability to diffuse- also diffuse and mix with other liquids
surface tension- a force that tends to pull adjacent parts of a liquid surface together, thereby decreasing surface area to the smaller possible size (drops are spheres)
capillary actions- attraction of the surface of a liquid to the surface of a solid (causes miniscus)
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EVAPORATION AND BOILING
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vaporization- process by which a liquid or solid changes to a gas
evaporation- process by which particles escape from the surface of a nonboiling liquid and enter the gas state
(evaporation is a form of vaporization)
-these vaporizations are caused by different kinetic energies particles with higher then average energies can overcome, and the intermolecular forces that binds them together in the liquid state are overcome thus entering the gas state
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boiling- change of a liquid to bubbles of vapor that appear throughout the liquid
boiling point- the temperature at which the vapor pressure of a liquid is equal to the atmospheric pressure
formation of a solid- liquids cooled--> average energy decreases--> attractive forces greater--> pull into orderly arrangement--> becomes a solid

freezing/solidification- physical change of a liquid to a solid state by the removal of energy as heat
sublimation- solid to gas phase
melting- solid to liquid phase
condensation- gas to liquid phase
deposition- gas to solid phase
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KINETIC MOLECULAR THEORY AND NATURE OF GASES
-how KMT accounts for physical properties of a gas

Expansion- no definite shape or volume
-when enclosed gas will completely fill any container
-explained by (3rd and 4th assumption KMT of gas)

Fluidity- attractive forces between gas particles are insignificant, they glide past each-other easily
-this is similar to liquids (behave the same) so they are both referred to as fluid

Low Density- 1/1000 density of liquid or solid state of the substance and must further apart

Compressability- can easily compress and decrease volume of a gas


DIFFUSION AND EFFUSION- gases mixing uniformly with air and substances
diffusion- such spontaneous mixing of the particles of two substances caused by their random motion
effusion- process by which gas particles pass through a tiny opening (rate of effusion is directly proportional to velocity so low masses effuse faster)


DEVIATIONS OF REAL GASES FROM IDEAL GASES
real gas- a gas that does not behave completely according to the assumptions of the kinetic-molecular theory
-noble gases generally ideal gases
-more polar a gas is the less ideal it is
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INTERMOLECULAR FORCES:
1. London Dispersion Forces (LDF)- arise from the random movement of electrons
-appear in all atoms, compounds, etc.
-more electrons equal a stronger london dispersion force
instantaneous/temporary dipole-when a compound/element/etc. is polar for a moment (because from the random movement of electrons they all go to one side for a moment)
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2. Dipole-Dipole Forces- the positive of one polar molecule being attracted to the f- of another
-exist only in polar molecules where permanent dipoles exist
(remember polar molecules have an electronegativity of 0.5-1.9
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Compounds have a higher boiling point over others because it takes more energy to break stronger intermolecular forces apart
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3. Hydrogen Bond- strongest intermolecular force there is
- exists when hydrogen is directly bonded to nitrogen, oxygen, or fluoride (N,O,F)
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Example: Order these compounds from low to high boiling point; H2O, C12H26,C2H6, C4H10, NaCl, C6H6,
-C
2H6, C6H6, C4H10, C12H26, NaCl, H2O
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STRENGTH OF INTERMOLECULAR ATTRACTIVE FORCES

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Pressure: PRESSURE- a force = N(ewton)
area = m(eter)
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vapor pressure- the pressure resulting from an evaporation liquid
volatile- a liquid exhibits a high vapor pressure
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UNITS of PRESSURE: milimeters of mercury (mmHg) Pascal (Pa) kilo pascal (KPa) torr pounds per square inch (psi) and
atmosphere (atm)--> the pressure of the atmosphere at sea level
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CONVERSIONS:
101.3 Kpa= 760 torr= 760 mmHg= 1 atm
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Example: Convert 802.1 mmHg to Kpa
802.1mmHg 101.3 Kpa= 106.9 kpa
760mmHg
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MEASURING PRESSURE
1. barometer- measuring atmospheric pressur
2. manometer- measuring gas pressure (make sure mmHg is the end unit)
- the more the Hg is to the right in the manometer the stronger the Gas (A) is compared to the atmospheric pressure (like the picture below)
- if the Gas has a greater pressure then the atmosphere then you add the difference (11 mmHg in the picture)
- if the Gas has a weaker pressure then the atmosphere then you subtract the difference
Example: If atmospheric pressure is 765 mmHg, what is the pressure of Gas A
765+11=776 mmHg make sure mmHg is always the unit being used in these equations
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Photo_230.jpg Photo_232.jpg this manometer is part of the ex. above
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HEATING CURVES:
Heating Curve- is a graph that shows the temperature and time of a substance as it changes states (every heating curve is different for every substance)
-horizontal lines indicate the substance is undergoing a change of state determined by the temperature which it does so (ex. H2O changes .....................from a solid to liquid at 0ºC so it would show a horizontal line at 0°C)
-should be in Celsius
-can determine heat at any moment on the chart by:
1. To determine the heat during a slope use Q=MC deltaT (heat equals mass times specific heat times change in temperature)
2. To determine the heat during a horizontal stretch use the heat of fusion and heat of vaporization with a stoichometry problem type set up
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Heat of Fusion- energy required to melt 1 mol of a substance
-Hfus for water is 6.02 KJ/mol
Heat of Vaporization- heat required to vaporize 1 mol of a substance
-Hvap for water is 40.7 KJ/mol
Specific Heat for H2O:
ice- 2.06 J/g
water- 4.184 J/g
gas- 1.87 J/g
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Photo_234.jpg
1. heat of the solid is rising
2. energy breaks crystals in ice
3. water temperature is rising
4. energy is warming the water; energy is breaks intermolecular forces holding water together to form gas
5. temperature of the gas is rising
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Example of a Heating Curve problem: How much heat is required to go from -20 degrees celsius to 115 degrees celsius?
1. Q=50(2.06)20= 2060 J
3. Q=50(4.184)100= 20920 J
5. Q=(1.87)15= 1402.5 J
2. 50g (1 mol) (6.02 KJ)= 16.7 KJ times 1000= 16,700 J
18.92g 1 mol
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4. 50g (1 mol) (40.7 KJ)= 112.9 KJ times 1000= 112,900 J
18.92g . 1 mol
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PHASE DIAGRAMS:
Phase Diagram- shows the different phases of a substance when there is pressure on the substance
- different phase diagrams for different substances
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Example Phase Diagram:
Photo_206.jpg

3- The straight sloping line three is on indicates changing- both gas and solid phases exist
1- The Tripe Point- all three phases can exist
2- The Critical Point- Maximum temperature that liquid water can exist, called super critical fluid
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All of this information came from the creditable source of Modern Chemistry by Raymond E. Davis,and the pictures came from creditable examples from a high school chemistry course, thanks for reading!