Chemistry Helpers
Matter and Energy

 

Physical Behavior of Matter
 
Energy
    Energy is defined as capacity to do work
     Work is done when a force is used to change the position or motion of matter.
     Potential Energy: energy due to position
     Kinetic Energy: energy of motion
     Ways that energy can be expresses (forms of energy)
S - Sound
A – Atomic (nuclear)
C - Chemical
H - Heat
E - Electrical
M - Mechanical
L - Light
     Law of Conservation of Energy:
Energy can not be created or destroyed, only transferred (it can only change form). Example: Battery – Chemical ® electrical ® sound in radio
     Exothermic: Release of heat energy
     Endothermic: Absorption of heat energy
 
 
 
 
 
Measurement of Energy:
     Thermometer: a device used to measure the temperature of a system
     Temperature is the average kinetic energy of a substance.
NOTE: a) Heat is NOT in the definition
                                           b) Average is of all the particles in the sample.
                  
Thermometry:
Celsius and Kelvin are the two scientific temperature scales.
     Boiling Point: temperature at which liquid becomes a gas
     Freezing Point: temperature at which liquid becomes solid
 
For water at Standard Pressure (1 atm or 101.3kPa)
           Boiling Point is 100°C or 373 K
          Freezing Point is 0°C or 273 K
 
To convert form one temperature scale to the other:
°C to K °C + 273 = K
K to °K – 273 = °C
 
Absolute Zero: The theoretical temperature at which all molecular motion ceases. (0K or -273°C)
 
          STP stands for: Standard temperature (273 K) and pressure (1 atm)
                                          This information can be found on Table A
 
 
Calorimetry: Process of measuring heat exchange during a reaction
     Calorimeter: device that uses changes in temperature of water to measure heat exchange during a reaction.
     Calorie (cal): a unit of heat energy – amount of heat needed to raise the temperature of 1g of water 1°C.
     Joule (J): a unit of heat energy – 4.18 are needed to raise the temperature of 1g of water 1°C
     Specific Heat (C): amount of heat energy needed to raise the temperature of 1g of any substance by 1°C.
 
For Water: Specific Heat = 4.18J/g K or 1 cal/g K
FORMULA: q = mCDT    Table T
                             Where:       q = heat
                                                m = mass
                             C = specific heat
                             T2-T1 = change in temperature
 
Example:
1.    How many joules of heat are absorbed by 45g of water when the temperature of that water is raised from 293K to 323K?
Step 1: q = mC (T2-T1)
Step 2: q = ?
            m = 45g
        C = 4.18 J/gK
           T2-T1 = 30K
Step 3: q = (45g) (4.18J/gK) (30K)
Step 4: q = 5.8 J
 
2.    A reaction chamber inside a calorimeter contains 245g of water at 23°C. After the reaction, the temperature of the water is at 28°C. How many joules were released during the reaction?
Step 1: q = mC(T2-T1)
Step 2: q = ?
           m = 245g
       C = 4.18 J/gK
          T2-T1 = 5K
Step 3: q = (245g) (4.18J/gK) (5K)
Step 4: q = 5120.5 J
 
 
Matter
     Matter is anything that has mass and takes up space
     Homogeneous matter: uniform throughout – elements/compounds (substances) or solutions (mixtures).
     Heterogeneous matter: not uniform throughout - *most mixtures (ex: salad, soil, chex mix)
 
Matter can be classified into two main groups:
1. Pure Substances              2. Mixtures 2 or more substances physically combined
          - elements/compounds               - homogeneous/heterogeneous
      Elements cannot be chemically          Homogeneous: solutions (aq)
                   decomposed                                            (ex: salt water, air) 
      Compounds: 2 or more elements       Heterogeneous: you can
                  chemically combined                pick out the parts (ex. Chex mix)
 
     Phases of Matter:
     Gas: substances with no shape, no volume. *Takes the shape or volume of the container.
     Liquid: substances with a definite volume but no definite shape.
     Solid: substances with definite shape and volume.
The phase in which a substance exists is dependent upon the bonding between the particles. The stronger the bond, the more likely the substance can exist in the solid phase.
 
     Matter can change from one form to another. As this occurs, energy also changes.
*As one proceeds from ice to water to water vapor, there is a(n) increase   in kinetic energy.
* The changes of phase are not chemical, they are physical changes.
 
     The phase change Diagram: (Heating and Cooling Curves)
1.    The Heating Curve
The following is the uniform heating of a substance from a temperature above its boiling point. T1 = Freezing/Melting Point, T2 = Boiling/Condensation Point.
 
a.     Energy changes along the curve:
If there is a change in the temperature there is a change in kinetic because there is a change in the average motion of the particles.
During a phase change there is no change in temperature, therefore no change in kinetic energy. Instead, energy goes to breaking bonds so it is a potential energy change.
 
(1)Energy Changes:
Along AB    DKE, no DPE
            BC    no DKE, DPE
            CD    DKE, no DPE
            DE    no DKE, DPE
            EF     DKE, no DPE
 
2.    A cooling curve would be the opposite:
 
 
 
 
Properties of Gases, Liquids and Solids:
     Gases (g): Transparent, compressible, expand without limit, have no shape/volume. **Take the shape and volume of their container.
 
1.    Gases exert pressure: 
a.     STP: defined as standard temperature and pressure
*Found on Table A
*Pressure can also be 760 torr or 760 mm Hg
 
2.    The Kinetic Molecular Theory of Gases: (The Ideal Gas Laws)
a)    Ideal gases are those whose gas particles (molecules):
                                                                    i.            travel in random, constant, straight line motion.
*At absolute zero, all kinetic energy ceases, so all particles stop moving.
                                                                  ii.            are separated by great distances relative to the size of the molecule so that the volume of the actual molecule is considered negligible.
                                                                iii.            have no real attractive forces between them.
                                                               iv.            have collisions that may result in the transfer of energy between particles {but remember overall energy is conserved}
 
Real gases do have volumes and do exhibit attractive forces between their particles or we would have no atmosphere!
 
 
b)    For real gases to behave like ideal gases:
Temperature must be high
Pressure must be low
 
c)    The two most common real gases that behave most like ideal gases are H2 and He because they are the smallest and least dense.
 
3.    Graham’s Law of Diffusion: the lighter / less dense gas travels faster / further than heavier / more dense gas particles under the same conditions of temperature and pressure. *Lighter particles diffuse faster.
 
4.    Dalton’s Law of Partial Pressure: the total pressure on a gaseous system is the sum total of all its parts.
Ptotal = P1 + P2 + P3 + …….
Example: The total pressure on a system contains gas X and gas Y is 3atm. If gas X has a partial pressure of 2.5atm, what is the pressure exerted by gas Y?
Step 1: Ptotal = P1 + P2 + P3 + …….
Step 2: Ptotal = 3atm
 
                                           PX = 2.5atm
Step 3: 3atm = PY + 2.5atm
Step 4: PY = 0.5atm
 
5.    Avogadro’s Hypothesis: under the same conditions of temperature and pressure, 2 equal volumes of two different gases will have the same number of particles regardless of their masses. At STP, this number of particles is Avogadro’s number, which is 6.02 x 1023
 
6.    The Combined Gas Law: The pressure a gas exerts on its surroundings is related to its temperature and volume such that:
FORMULA: P1V1 = P2V2
                      T1        T2
 
*note*         Temperature must be in KELVIN!!! Pressure 1 and 2 and Volume 1 and 2 must be in the same units.
 
Example 1:
A gas has a volume of 1400mL at 20K and 101.3kPa. What is the new volume if the temperature changes to 40K and pressure changes to 50.65kPa?
Step 1: P1V1 = P2V2
                T1        T2
Step 2: P1 = 101.3kPa             P2 = 50.65kPa
             V1 = 1400mL               V2 = ?
             T1 = 20K                      T2 = 40K
Step 3: (101.3kPa) (1400mL)   = (50.65kPa) (x)
                20K                     40K
Step 4: V2 = 5600mL
 
Example 2:
What will the volume of a sample of gas be if 42mL of the gas at 720torr and 27°C changes to 240torr and -123°C?
Step 1: P1V1 = P2V2
                     T1           T2              
 
Step 2: P1 = 720 torr     P2 = 240 torr                        
                  V1 = 42mL         V2 = ?
           T1 = 27°C           T2 = -123°C
 
Step 3: (720torr) (42mL)   = (240torr) (x)
              300K                150K
 
Step 4: V2 = 63mL
 
a) Boyle’s Law: As long as temperature remains constant, pressure and volume of a gas will affect each other inversely
 
FORMULA: P1V1 = P2V2
*Note* PV = k (where k is a constant)
 
 
 
Example: Gas X has a volume of 40L at STP. If the temperature remains constant and the pressure is doubled, what is the new volume?
Step 1: P1V1 = P2V2
 
Step 2: P1 = 1atm           P2 = 2atm
             V1 = 40L             V2 = ?
Step 3: (1atm) (40L) = (2atm) (x)
Step 4: V2 = 20L
 
 
b)    Charles Law: As long as the pressure on a given mass of a gas remains constant, the temperature and volume will affect each other directly.
 
FORMULA: V1 = V2
                    T1     T2
 
 
Example: Gas Y has a volume of 500mL at STP. If the pressure remains constant, what is the new volume when the temperature is raised to 373K?
Step 1: V1 = V2
                                           T1     T2
                             Step 2: V1= 500mL      V2 = ?
                                          T1 = 273K        T2 = 373K
                             Step 3:   500mL   =    x
                                            273K         373K    
                             Step 4: V2 = 683.2mL
 
 
     Liquids: no definite shape/ but definite volume with very low compressibility.
1.    Boiling Point: The temperature at which the vapor pressure of a liquid reaches atmospheric pressure; therefore allowing particles to escape as a gas.
*When vapor pressure of a liquid = atmospheric pressure
 
As atmospheric pressure increases one must raise the vapor pressure of the liquid by increasing its temperature.
 
b)    Normal Boiling Point is measured at standard pressure
For water it is:   100°C or 373K
 
2.    Vapor Pressure (for four liquids see Table H)
a)    The pressure exerted by the vapor evaporating off the surface of a liquid.
b)    Each liquid has its own vapor pressure.
c)    As the temperature of the liquid increases, the vapor pressure of that liquid increases.
 
3.    Evaporation: The change of phase from liquid  to gas.
a)    Heat of Vaporization: The amount of heat energy required to vaporize a given mass of a liquid to gas at a constant temperature. This is an endothermic process, energy is being absorbed.
Each substance has its own Heat of Vaporization.
For water at its normal boiling temperature of 100°C and standard pressure the heat of vaporization is 2260 joules per gram. (Table B)
 
4.    Condensation: The change of phase from gas to liquid. This is an exothermic process.
a)    The Heat of Condensation is the direct opposite of the heat of vaporization. The quantity of heat energy is the same as for the heat of vaporization, but instead of being absorbed the heat energy is being released.
 
 
     Solids:  definite shape/definite volume.
          ***Regular Geometric Pattern***
 
1.    Melting (fusion): An endothermic process in which a solid becomes a liquid.
 
For water at Standard Pressure the temperature at which melting occurs is 0°C / 273K.
 
2.    Freezing is the direct opposite of melting, but instead of being an endothermic process where energy is absorbed, it is an exothermic process where energy is released. Water freezes at 0°C / 273K.
 
3.    Heat of Fusion: the amount of heat energy required to change a given mass of solid to liquid at a constant   temperature.
 
Each substance has its own Heat of Fusion.
For water at standard pressure, this quantity of heat is 334 J/gK
Table B
 
4.    Heat of Solidification (crystallization) is the direct opposite of the Heat of Fusion. Since Solidification is an exothermic process, the heat energy is released instead of absorbed.
 
5.    Sublimation: The change of phase from solid to gas, completely skipping the liquid phase.
 
This generally occurs only in solids with high vapor pressures and weak intermolecular forces of attractions.
 
Examples: *paradichlorobenzene (moth balls)
               Dry ice (CO2)
               I2

 



Write Your Questions Below



true or false question gas matter has high thermo energy
January 24, 2012

Kinetic Energy is the energy of motion and temperature measures the average kinetic energy in a system.  As you go up a heating curve you go from a solid to a liquid to a gas as the temerature rises.  Therefore gas would have a greater kinetic energy based upon a single substance (ex water in its three phases of ice, liquid and vapor).

January 24, 2012 -  Replied By Expert

 
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