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Matter And Energy

Matter And Energy. Chemistry Ch 3. The Nature of Matter. Chemists are interested in the nature of matter and how this is related to its atoms and molecules. Gold. Mercury. Chemistry & Matter. We can explore the MACROSCOPIC world — what we can see —

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Matter And Energy

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  1. Matter And Energy Chemistry Ch 3

  2. The Nature of Matter Chemists are interested in the nature of matter and how this is related to its atoms and molecules. Gold Mercury

  3. Chemistry & Matter • We can explore the MACROSCOPIC world — what we can see — • to understand the PARTICULATE worlds we cannot see. • We write SYMBOLS to describe these worlds.

  4. A Chemist’s View of Water Macroscopic H2O (gas, liquid, solid) Symbolic Particulate

  5. A Chemist’s View Macroscopic 2 H2(g) + O2 (g) --> 2 H2O(g) Particulate Symbolic

  6. Kinetic Nature of Matter Matter consists of atoms and molecules in motion.

  7. STATES OF MATTER • SOLIDS— have rigid shape, fixed volume. External shape can reflect the atomic and molecular arrangement. • Reasonably well understood. • LIQUIDS— have no fixed shape and may not fill a container completely. • Not well understood. • GASES— expand to fill their container. • Good theoretical understanding.

  8. OTHER STATES OF MATTER • PLASMA— an electrically charged gas; Example: the sun or any other star • BOSE-EINSTEIN CONDENSATE— a condensate that forms near absolute zero that has superconductive properties; Example: supercooled Rb gas

  9. Physical Properties What are some physical properties? • color • melting and boiling point • odor

  10. Physical properties of matter are categorized as either Intensive or Extensive: • Intensive - Properties that do not depend on the amount of the matter present. • Color • Odor • Luster - How shiny a substance is. • Malleability - The ability of a substance to be beaten into thin sheets. • Ductility - The ability of a substance to be drawn into thin wires. • Conductivity - The ability of a substance to allow the flow of energy or electricity. • Hardness - How easily a substance can be scratched. • Melting/Freezing Point - The temperature at which the solid and liquid phases of a substance are in equilibrium at atmospheric pressure. • Boiling Point - The temperature at which the vapor pressure of a liquid is equal to the pressure on the liquid (generally atmospheric pressure). • Density - The mass of a substance divided by its volume

  11. Physical properties of matter are categorized as either Intensive or Extensive: • Extensive - Properties that do depend on the amount of matter present. • Mass - A measurement of the amount of matter in a object (grams). • Weight - A measurement of the gravitational force of attraction of the earth acting on an object. • Volume - A measurement of the amount of space a substance occupies. • Length

  12. Physical Changes • can be observed without changing the identity of the substance Some physical changes would be • boiling of a liquid • melting of a solid • dissolving a solid in a liquid to give a homogeneous mixture — a SOLUTION.

  13. Chemical Properties and Chemical Change • Burning hydrogen (H2) in oxygen (O2) gives H2O. • Chemical change or chemical reaction — transformation of one or more atoms or molecules into one or more different molecules.

  14. Sure Signs of a Chemical Change • Heat • Light • Gas Produced (not from boiling!) • Precipitate – a solid formed by mixing two liquids together

  15. Physical vs. Chemical • physical • chemical • physical • physical • chemical • Examples: • melting point • flammable • density • magnetic • tarnishes in air

  16. Physical vs. Chemical • chemical • physical • chemical • physical • physical • Examples: • rusting iron • dissolving in water • burning a log • melting ice • grinding spices

  17. PURE SUBSTANCE MIXTURE yes no yes no Is the composition uniform? Can it be chemically decomposed? Colloids Suspensions Matter Flowchart MATTER yes no Can it be physically separated? Homogeneous Mixture (solution) Heterogeneous Mixture Compound Element

  18. Types of Mixtures • Variable combination of 2 or more pure substances. Heterogeneous –visibly separate phases Homogeneous – Same throughout

  19. Alloys

  20. Law of Definite Proportions • Joseph Proust (France 1799) • A given compound always contains elements in a certain proportion by mass. (Constant composition).

  21. Atoms combine in whole number ratios, so their proportion by mass will always be the same. • Example: H2O is always made up of 2 atoms of H and one atom of O. The mass ratio of O to H in water is always 16:2 or 8:1. Percent Mass of Compounds • Percent composition consists of the mass percent of each element in a compound: Percent by mass=

  22. Example: KCl • KCl always contains one atom of K for every one atom of Cl • In KCl, potassium and chlorine always have a ratio of “39.09 to 35.45” or “1.1 to 1” by mass.

  23. Law of Multiple Proportions (John Dalton) • When the same two elements combine to form more than one compound: the ratios of the mass of one element in the first compound to its mass in the second compound, (as it combines with the same mass of the other element), can always be expressed as ratios of small whole numbers( ex: 1:3 or 2:5).

  24. Example of Law of Multiple Proportions • Carbon combines with oxygen to form CO and CO2 . 12.01 16.00 2:1 12.01 32.00

  25. Practice Problem 1 In the carbon compounds ethane (C2H6) and ethene (C2H4), what is the lowest whole number ratio of H atoms that react with the same number of C atoms? 24.02 6.06 6:4 or 3:2 24.02 4.04

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