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Gen-Chem110-Ch01

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The Foundations of Chemistry

The Foundations of Chemistry

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  • Chem110

Gen-Chem110-Ch01 Gen-Chem110-Ch01 Presentation Transcript

  • Chapter 1 The Foundations of Chemistry HW27Qs: 3, 4, 6, 10, 14,18, 20, 26, 28, 30, 32, 34, 36, 38, 40, 44, 46, 50, 54, 56, 58, 60, 62, 64, 66, 68,70 Dr. Shiunchin C. Wang
  • 1.1 Matter and Energy CHEMISTRY MATTER What is Chemistry? Chemistry is the science that describe matter regarding with its composition, structure, properties. matter is anything has mass and occupies space . Mass Occupies Space + Mass is a measurement of the quantity of matter in a sample of any materials. Mass is different from weight. Mass does not vary in different place, but weight does. Weight : a measurement of the gravitational attraction of the earth for the body. Volume: a 3-D space. Subdiscipline of chemistry: Analytical, Biological, Biophysical, Inorganic, Organic, Material, Physical, Polymer, Nanotechnology etc.
  • 1.1 Matter and Energy (Cont.)
    • Matter posses certain ability called “energy”.
    • (def) ENERGY : the capacity to do work or transfer hear .
    Types of Energy: Electrical energy (electrical power), thermal energy (heat), nuclear energy (energy release from nucleus of atom, E=mc 2 ), & chemical energy (associate with atoms).
  • Natural Laws
    • Scientific (natural) law
      • A general statement based the observed behavior of matter to which no exceptions are known.
    • Law of Conservation : neither created nor destroyed .
    • Law of Conservation of Mass
    • Law of Conservation of Energy
    • Law of Conservation of Mass-Energy
      • Einstein’s Relativity
      • E=mc 2
    • Law of Conservation of Matter : no observable change in the quantity of matter during “a chemical reaction” or “physical change”.
    • E.g.. Magnesium burns in oxygen to from magnesium oxide.
    • 2 Mg(s) + O 2 (g)  2MgO(s)
    • mass of 2 Mg + mass of 1 O 2 = mass of 2 MgO
  • 1.1 Matter and Energy (Cont.) Law of Conservation of Energy : Energy cannot be created or destroyed in a chemical reaction or in a physical change. It can only be converted from one form to another . Relationship between matter and energy by Albert Einstein in 1940: E= mc 2 Exothermic reaction : convert chemical energy into heat energy. Endothermic reaction : heat, light, and electrical energy can be converted into chemical energy
  • 1.1 Endothermic vs. Exothermic
    • Exo thermic reaction : reactions that release energy (chemical energy) in the form of heat (thermal energy).
    Endo thermic reaction : reactions that absorb energy from surrounding (heat, light, electrical energy) to from chemical energy. H 2 O(s) + 6.02 kJ  H 2 O(l) solid ice liquid water
  • 1.2 Chemistry – A molecular View of Matter
  • Non-Mixtures: Elements Atom: smallest particle of an element. Atoms consists 3 fundamental particles: electrons, protons, neutrons.
  • Non-Mixtures: Compounds
    • Compounds
      • substances composed of two or more elements in a definite ratio by mass, C 2 H 5 OH
      • can be decomposed into the constituent elements
        • Water is a compound that can be decomposed into simpler substances – hydrogen and oxygen
    • Molecule: the smallest particle of an element or compound that can have a stable independent existence.
    • 7 diatomic molecules: H 2 , N 2 , O 2 , F 2 , Cl 2 , Br 2 , I 2 .
    • Polyatomic: P 4 (tetra-phosphorus) and S 8 .
  • (p.8) Ex. 1.1 distinguish different matters
    • Atom : smallest particle of an element
    • Molecule : a single stable atom of an element or compound consist of more than one atom.
    • Element contains a single kind of atom, periodic table atoms in the stable form.
    • Compound contains atoms of 2 or more different elements.
    atom molecule element compound Krypton, Kr V V V X Ethane, C 2 H 6 X V X V Nitrogen, N 2 X V V X Aspirin, C 9 H 8 O 4 X V X V Sulfur dioxide, SO 2 X V X V Mole of copper , Cu = 6.022 x 10 23 particles X X V X
  • 1-4 Chemical and Physical Properties
    • Matters have:
    • Physical Properties : can be observed by color, density, hardens etc. e.g., ice  water  steam.
    • Chemical Properties : change in composition, e.g., 2Mg + O 2  2MgO
    • Extensive Properties : The Volume or Mass depend on amount
    • Intensive (intrinsic) Properties : all chemical properties are intensive properties.
    T (melting) = T (freezing) T (evaporation) = T (condensation) T (deposition) = T (sublimation)
  • 1-7 Measurements in Chemistry
    • (p.22) Scientific Notation : a standard form for writing numbers as:
    • n x 10 m
    • where m = z (integer, … -3, -2, -1, 0, 1, 2, 3…),
    • and 1  |n| < 10.
    Example: 1 mole of gold = 197 gram of gold 1 mole of gold = 6.022 X 10 23 atoms of gold = 602,200,000,000,000,000,000,000 Mass of gold = 3.27 X 10 -27 gram = 0.000 000 000 000 000 000 000 327 gram
  • Significant Figures (p 23) Determination of 123,000 significant figures: (1) 1.23 x 10 5 has 3 significant figures, (2) 1.230 x 10 5 has 4 significant figures. Special case: 24 0 000 has 3 significant figures. Therefore, it is best to use “scientific notation” to represent the significant figures. In this case will be 2.40 X 10 5 . Note: Measurement of volumetric cylinder by reading the bottom of the meniscus . Significant Figures = Exact numbers + 1 estimate number
  • 1-10 The Unit Factor Method
    • 1 mile = 5280 ft || 1 yd = 3 ft || 1 ft = 12 in || 1 in = 2.54 cm
    Q 1.47 mi = ? in 1 L = 1000 mL = 1 dm 3 || 1 mL = 1 cm 3 = 1 cc 1 gal = 4 qt = 8pt || 1 qt = 0.964 L
  • 1-12 Density & Specific Gravity
    • Mass (m) = density (d) X Volume (V)
    • Water Density = 1.000 g/cm 3 = 1.000 g/mL (at 3.89  C – 25  C)
    Q. What is the density of 47.3 mL ethyl alcohol (= ethanol) with mass 37.32 g? A. d = m/V = 37.32 g/ 47.3mL = 0.789 g/mL Specific Gravity (SG.): is the ratio of its density to the density of water
    • Q. Table salt density is 2.16 g/mL at 20  C
    • SG of table salt = (2.16 g/mL)/ (1.00 g/mL) = 2.16
    • Ex1.17 (p.34) Battery Acid = 40.0% sulfuric acid, H 2 SO 4 + 60.0% water by mass (m/m). Its specific gravity is 1.31. Calculate the mass of pure H 2 SO 4 in 100.0 mL of battery.
    • Ans: Specific gravity = 1.31  so, density = 1.31 g/mL (1.31 g soln in 1 mL soln)
  • 1.13 Heat and Temperature
    • Temperature: measures the intensity of heat.
    • 0  C (Celsius)= 32  F (Fahrenheit) = 273 K (Kelvin)
    • 100  C= 212  F = 373 K
    • Q. 400 K = ?  F
    • (400K – 273.15) (1C/1K) = 127 C
    • 127C X (1.8F/1C) + 32F = 261 F
  • 1.14 Heat Transfer and the Measurement of Heat
    • Specific Heat of a substance (J/g  C) : the amount of heat required to raise the temperature of one gram of the substane on degree Celsius with no change in phase .
    • Heat Capacity (J/  C): the amount of heat required to raise its temperature 1  C.
    Heat Energy: Q = m c  T Q (joule, J) = Heat Energy; where 1calorie = 4.184 joule m = mass (gram) c = Specific heat = J/g  C; C(ice) = 2.09 J/g  C C (water) = 4.18 J/g  C C (steam) = 2.03 J/g  C  T = change of temperature Heat Capacity = C = Q/∆T
    • (Ex. 1.20) How much heat, in joules, is required to raise the temperature of 205 g of water from 21.2  C to 92.4  C?
    • Ans Q = m C  T
    • = (205 g)(4.18J/g  C)(92.4-21.2  C) = 6.02 X 10 4 J
    (Ex 1.21) A 588-gram chunk or iron is heated to 97.5  C. Then it is immersed in 247 grams of water originally at 20.7  C. When thermal equilibrium has been reached, the water and iron are both at 36.2  C. Calculate the specific heat of iron. Ans T final = 36.2  C [(588 g) (c) (97.5 – 36.2  C)] iron = [(247 g) (4.18J/g  C) (36.2 - 20.7  C)] water specific heat of iron, c = 0.444 J/g  C Thermal Equilibrium Heat gained = eat lost (m c  T) gained = (m c  T) lost