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State Of Matter

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  • 1. STATE OF MATTER
      • MATTER
      • Anything that occupies space and has mass.
      • Include things we can see and touch (water, trees) as well as things we can’t see (air).
      • THREE STATES OF MATTER
      • All substances/matter can exist in 3 states : solid, liquid and gas.
      • Solid - molecules are held close together in an orderly fashion with little freedom of motion.
      • Liquid - molecules are close together but are not held so rigidly in position and can move past one another.
      • Gas - molecules are separated by distances that are large compared with the size of the molecules.
  • 2.
      • The three states of matter are interconvertable .
    • The physical properties of a substance depends on the state of the substance. When a substance undergoes a change in state, many of its physical properties change.
  • 3. THE GASEOUS STATE
    • Under certain condition of pressure & temp., most substances can exist in any one of the three state of matter.
    • E.g. water - solid :ice ; liquid : water ; gaseous : steam / water vapor
    • In gases - molecular motion is totally random, forces of interaction between molecules are so small, each molecules moves freely and essentially independently of other molecules.
    • SUBSTANCES THAT EXIST AS GASES
    • Under normal condition of pressure & temp. (1 atm , 25 o C), elements that exist as gases are:
    • The noble gases (Group 8A elements) : He, Ne, Ar, Kr, Xe, Rn - monatomic species.
    • Hydrogen, nitrogen, oxygen, fluorine, and chlorine - exist as diatomic molecules : H 2 , N 2 , O 2 , F 2 , Cl 2 .
    • Allotrope of oxygen, Ozone (O 3 ) .
  • 4.
    • Compounds that exist as gases are HCl, CO, CO 2 , NH 3 , N 2 O, NO, NO 2 , SO 2 , H 2 S, HCN, CH 4 .
    • Only O 2 essential for our survival.
    • H 2 S, HCN - deadly poisons.
    • CO, NO 2 , SO 2 , O 3 - less toxic.
    • He, Ne, Ar - chemically inert.
    • Most gases are colorless except F 2 , Cl 2 , NO 2
    • General Properties of gas ;
    • Compressible
    • Have low density (about 2 kg per m 3 )
    • Diffuses quickly (mix thoroughly)
    • Fills up a container uniformly
    • Exert pressure uniformly on all sides of a container independently of the height or depth .
  • 5. PRESSURE OF A GAS
    • Gas exert pressure on any surface with which they come into contact - gas molecules are constantly in motion and collide with the surface.
    • Instrument to measure atmospheric pressure : barometer
    • - Long tube filled with mercury (Hg) is inverted into a dish of mercury.
    • -Atm. pressure pushing on the surface of the Hg
    • in the tube is proportional to atmospheric pressure.
    • - Height of Hg (h)  atm pressure.
    h Atmospheric pressure
  • 6.
    • A second type of barometer, a manometer has two arms, one opened to the atmosphere and one closed or connected to a container filled with gas. The pressure exerted by the atmosphere or by gas in a container is proportional to the difference in the mercury levels (h).
  • 7.
    • The std atmospheric pressure (1 atm) is equal to the pressure that supports a column of mercury exactly 760 mm (76 cm) high at 0 o C at sea level.
    • 1 atm = 760 mmHg (mmHg represents pressure exerted by a column of mercury 1 mm high)
    • SI Units - Pascals, Pa : Pressure =force/ area
    • 1 Pa = 1 N/m 2
    • 1 atm =101,325 Pa =1.01325 x10 5 Pa =1.01325 x10 2 kPa.
  • 8. THE GAS LAWS
    • Important generalizations regarding the macroscopic behavior of gaseous
    • substances.
    • 1. The Pressure -Volume Relationship : Boyle’s Law
    • -studied by Robert Boyle in 17th century.
    • Volume of a fixed amount of gas maintained at constant temperature is
    • inversely proportional to the gas pressure
    • V  1/P,  : proportional to
    • or V = k 1 x 1/P k 1 : proportionality constant
    • PV = k 1
    • For a given sample of gas under two diff sets of conditions at constant
    • temp. : P 1 V 1 = k 1 = P 2 V 2
  • 9. P (atm) 1/V 0.6 0.3 2 4 P versus V graph at constant temp.: Volume of gas doubles as the pressure is halved. P (atm) V (L) P versus 1/V
  • 10. The Temperature-Volume Relationship : Charles’ Law
    • At constant pressure, the volume of a gas sample expands when heated and contracts when cooled.
    • Study on the temp. – vol. relationship at various pressure showed that :
      • At any given pressure, the plot of vol. vs temp. yields a straight line.
      • Extending the line to zero vol., the intercept on temp. axis is -273.15 o C (absolute temperature)
    V T ( o C) -273.15 P1 P2 P3 P4
  • 11.
    • Lord Kelvin identified the temp. -273.15 o C as theoretically the lowest attainable temp ., called absolute zero.
      • absolute temp. scale , now called Kelvin temp. scale.
      • Absolute zero : 0 K = -273.15 o C
      • 273.15 K = 0 o C
      • 373.15 K = 100 o C
      • Relationship between o C and K:
      • T (K) = t( o C) +273.15 o C
  • 12.
    • Charles’ Law states that : the vol. of a fixed amounts of a gas maintained at constant pressure is directly proportional to the absolute temp. of the gas.
    • Under 2 diff. sets of conditions for a given sample of gas at constant pressure :
          • V 1 / T 1 = k 2 = V 2 / T 2
          • V 1 / T 1 = V 2 / T 2
    • V 1 , V 2 are volumes of the gases at temp. T 1 , T 2 (both in kelvins).
    V  T V = k 2 T or V/T = k 2 k 2 is proportionality constant .
  • 13. The Volume – Amount Relationship : Avogadro,s Law
    • Amedeo Avogadro – complemented the studies of Boyle, Charles and Gay-Lussac. He published a hypothesis that stated :
      • At the same temp. and pressure, equal volumes of different gases contain the same number of molecules (or atoms if the gas is monatomic).
      • The volume of any given gas must be proportional to the number of molecules present;
            • V  n
            • V = k 3 n where n represents the number of moles and
            • k 3 is the proportionality constant.
            • Avogadro’s Law – at constant pressure and temp., the vol. of a gas is
            • directly proportional to the number of moles of the gas present .
  • 14. The Ideal Gas Equation.
    • Boyle,s Law : V  1/P (at constant n and T)
    • Charles’ Law : V  T (at constant n and P)
    • Avogadro’s Law: V  n (at contant P and T)
    • Combine all three expressions :
            • V  nT / P
            • = RnT/P,
            • Or PV = nRT ……………. ideal gas equation
            • R, the proportionality constant is called the gas constant.
    Ideal gas – a hypothetical gas whose pressure- volume-temp. behavior can be completely accounted for by the ideal gas equation. The molecules of an ideal gas do not attract or repel one another, and their vol. is negligible compared with the volume of the container.
  • 15.
    • To apply the ideal gas equation to a real system, we must evaluate the gas constant, R.
    • At 0 o C (273.15K) and 1 atm pressure, many real gas behave like an ideal gas. Exp. show that under these conditions, 1 mole of an ideal gas occupies 22.414L.
    • The conditions 0 o C (273.15K) and 1 atm pressure are called standard temp. and pressure (STP).
        • From PV = nRT
            • R = PV/nT
            • =( 1 atm)(22.414L) = 0.082057 L.atm/K.mol
            • (1 mol)(273.15K)
            • For most calculations, use R=0.0821 L.atm/K.mol and the molar volume of a gas at STP as 22.4L.
  • 16.
        • When pressure. volume, temp. and amount change, a modified form of equation must be employed, which involves initial and final conditions.
        • R = P 1 V 1 / n 1 T 1 , (before change)
        • R = P 2 V 2 / n 2 T 2 , (after change)
        • So,
        • P 1 V 1 / n 1 T 1 = P 2 V 2 / n 2 T 2
        • If n 1 = n 2 as is usually the case because the amount of gas normally does not change,
        • P 1 V 1 / T 1 = P 2 V 2 / T 2
  • 17. Density Calculations
    • From the ideal gas equation ,
    • PV = nRT we can calculate the density of gas.
    • The number of moles of the gas, n, is given by :
    • n = m / M , m = mass of gas (grams) and M is its molar mass
    • Since density is mass per unit volume,
    • n/V = P / RT
    • m/V M = P / RT
    • We can write d = m/V = P M
    • RT
    • Unit of gas density are usually grams per liter (g/L), rather than grams per mL (g/mL) : density of gases is very low at atmospheric condition.
  • 18. Deriving Quantity of Density
    • 1……PV=nRT
    • 2……n=PV/RT
    • 3……n/V=P/RT n= m/ M
    • 4……m/ M V=P/RT d=m/V
    • 5……d/ M = P/RT
    • 6…..d= M P/RT
  • 19. The Molar Mass of a Gaseous Substance
    • Molar Mass of a substance is found by examining its formula and summing the molar masses of its component atoms (if the actual formula of the substance is known).
    • For an unknown gaseous substance, an experiment is needed to determine the density value (or mass and volume data) at a known temp. and pressure.
    • From d = m/V = P M , we get
    • RT
    • M = dRT
    • P
  • 20. Assignment 1
    • A Balloon of volume 0.55mL at sea level at pressure P=1 atm is allowed to rise to a height with P=0.40atm. If the temperature remains constant, find the final volume of the balloon.
    • A 550mL of Fluorine gas, F 2 is heated from 22  C to 87  C at constant pressure.What is its final volume?
    • What is the volume of H 2 evolved when 1.20g of Mg react with excess of HCl at STP.
    • A helium filled balloon has a volume of 6.15m 3 at 14  C and 762mmHg. If its volume expands to 6.37m 3 and pressure falls to 749mmHg. Find its final temperature?
    • A sample of carbon monoxide gas has a volume of 3.20mL at 125  C. Calculate it temperature when its occupy 1.547mL at constant pressure.
  • 21.
      • A gas sample at 25  C and 0.862atm has a density of
      • 2.26g/L Calculate its molar mass.
      • The density of a gas at STP is 1.960g/L.Determine its molar mass.
      • What is the mass of Sulphur dioxide which has a volume of 2.0L at 25  C and 1 atm pressure.?
      • Calculate the relative mass of 3.0g of a gas at 100  C and 0.9atm occupies of 2.0L
      • Find the volume of carbon dioxide formed when 20.0mL of CO reacts with excess oxygen gas at a constant temperatute & pressure.
  • 22. END
    • ****************************