Ch 1 and 2 ppt


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  • STOP
  • Handout Chemistry Reference Table
  • 2 Na, 2 S, 3 O 1 Mg, 2 N, 6 O 10 Fe, 15 O
  • A bathtub filled with lukewarm water contains more thermal energy than a teaspoon of boiling water Teaspoon has a higher temperature
  • K = 33 + 273 = 306 K F = (1.8x33) = 59.4 ◦ F
  • Ch 1 and 2 ppt

    1. 1. Science, Chemistry and You
    2. 2. Chemistry • Definition – study of the composition and properties of matter and the energy transformations accompanying changes in the structure of matter
    3. 3. Major Branches of Chemistry • Inorganic Chemistry – Study of all the elements other than Carbon • Organic Chemistry – Study of compounds containing carbon • Biochemistry – study of chemical processes in living things • Nuclear Chemistry – study of radioactivity, the nucleus and the changes that the nucleus undergoes
    4. 4. Aristotle Early Greek Theories • 400 B.C. - Democritus thought matter could not be divided indefinitely. • 350 B.C - Aristotle modified an earlier theory that matter was made of four “elements”: earth, fire, water, air. Democritus • Aristotle was wrong. However, his theory persisted for 2000 years. • This led to the idea of atoms in a void.
    5. 5. The Rise of Modern Chemistry • The Greek idea of the 4 basic elements was not disputed until the mid 1600s • Robert Boyle proposed that elements are substances that cannot be chemically decomposed into simpler substances. Earth, air, fire and water could not be called elements • In 1774 Joseph Priestly discovered a gas in which substances burned easily, Antoine Lavoisier named the gas Oxygen Boyle Priestly
    6. 6. John Dalton • 1800 -Dalton proposed a modern atomic model based on experimentation not on pure reason. • All matter is made of atoms. • Atoms of an element are identical. • Each element has different atoms. • Atoms of different elements combine in constant ratios to form compounds. • Atoms are rearranged in reactions. • His ideas account for the law of conservation of mass (atoms are neither created nor destroyed) and the law of constant composition (elements combine in fixed ratios).
    7. 7. Reaction of the Day Table sugar + sulfuric acid  Carbon + H20 H2SO4 C12H22011 (s)  12 C (s) + 11 H2O (g)
    8. 8. Ch 2 - Matter Matter – anything that takes up space and has mass
    9. 9. Chemical and Physical Properties of Matter Physical properties – color, shape, texture, odor, taste, electrical conductivity, and density density – how closely packed the molecules are malleable – substances that can be easily hammered into shapes ductility – substances that can be stretched into wires conductivity – substances that can transfer heat or electricity Chemical properties – describe how matter acts in the presence of other materials
    10. 10. What is each picture modeling? Density, malleability, ductility, conductivity
    11. 11. Physical or Chemical Change
    12. 12. Physical vs. Chemical Change Physical Change • Atoms do not rearrange • Only physical properties change. Chemical properties do not change. • Physical changes are generally easy to reverse. • No energy is produced by the substance. Chemical Change • Atoms are rearranged into different molecules • Both physical and chemical properties are changed • Changes are not reversible without another reaction • Energy is often produced ( fire or heat, for example)
    13. 13. Identify each of the following as a Physical or Chemical Change. Put a P next to Physical Changes and a C next to Chemical Changes 1. A piece of wood burns to form ash. 2. Water evaporates into steam. 3. A piece of cork is cut in half. 4. A bicycle chain rusts. 5. Food is digested in the stomach. 6. Water is absorbed by a paper towel. 7. Hydrochloric Acid reacts with zinc. 8. A piece of an apple rots on the ground. 9. A tire is inflated with air. 10. A plant turns sunlight, CO2, and water into sugar and oxygen. 11. Sugar dissolves in water. 12. Eggs turn into an omelette. 13. Milk sours. 14. A popsicle melts. 15. Turning brownie mix into brownies.
    14. 14. Demonstration of the day Vinegar + baking soda Acetic acid + sodium bicarbonate  carbon dioxide + water + sodium acetate Heterogeneous mixture containing, solid, liquid and gas phases
    15. 15. The Division of Matter Two major categories: 1) pure substances - consists of only one type of matter, which cannot be separated into other kinds of matter by any physical processes. Ex: Olive oil 2) mixtures – material that can be separated by physical means into two or more pure substances. Ex: Oil and vinegar salad dressing
    16. 16. Two Types of Mixtures • Heterogeneous – a mixture in which the substances are not uniformly mixed Ex: oil & vinegar dressing, granite has quartz & mica • Homogeneous – a substance in which the particles are uniformly mixed Ex: dough & air
    17. 17. Elements and Their Symbols Element - pure substance that cannot be broken down into simpler substances
    18. 18. Elements and Their Symbols • Atoms – smallest particles that maintain the physical and chemical characteristics of an element • Monoatomic elements – elements that do not naturally combine or bond together. Ex: Ne, He, Ar • Diatomic elements - elements that bond into two-atom units. Ex: O2, H2 • Polyatomic elements – elements composed of multi-atom units. Ex: S8
    19. 19. Elements and Their Symbols Symbol – letter given to represent the name of each element Hydrogen Oxygen Calcium Magnesium Manganese Sodium
    20. 20. Compounds and Their Formulas • Compounds are made up of atoms from two or more different elements, chemically bonded together • Formulas tell the type and number of atoms that are present in compounds Common Compounds and Their Formulas Compound Formula Atoms Ammonia NH3 1 nitrogen, 3 hydrogen Rust Fe2O3 2 iron, 3 oxygen Salt NaCl 1 sodium, 1 chlorine Sucrose C12H22O11 12 carbon, 22 hydrogen, 11 oxygen
    21. 21. Sample Problems How many atoms of each element are present in each of the following groups? a.Na2S2O3 b.Mg(NO3)2 c. 5 Fe2O3
    22. 22. Molecule • The smallest independent units of compounds • Consist of two or more atoms that are chemically bonded together • Ex: H20, NH3, H2SO4 • Homework: Read pgs 21-28 Section Review Questions 2A, pg 29, #1-3
    23. 23. Tuesday September 14, 2010 • Go over homework problems
    24. 24. 2B Energy in Matter • Every chemical reaction either releases or absorbs energy • Exothermic reactions – release energy (get hot) Ex: lighting a match • Endothermic reactions – absorb energy (get cold) Ex: ice pack
    25. 25. Energy – the ability to do work • There are many forms of energy • Chemistry is concerned with the relationship among chemical, thermal, electrical and nuclear energy
    26. 26. Energy Conservation • Thermodynamics – the study of energy flow • First Law of Thermodynamics or Law of Conservation of Mass-Energy –matter and energy can neither be created nor destroyed, simply changed from one form to another • Second Law of Thermodynamics – during any energy transformation, some energy goes to an unusable form
    27. 27. Energy Conservation • Entropy – randomness or disorder of a system • There is a tendency for all natural processes to increase in entropy (disorder)
    28. 28. Heat, Energy & Temperature • Kinetic Energy – energy of motion All matter contains particles that are moving • Thermal Energy – sum of all the kinetic energy of an object • Temperature measures the average kinetic energy of all the particles in a sample • Heat – thermal energy that is transferred from one object to another • Amount of heat transferred between objects is determined by the temperature difference between them and the mass of the hotter object
    29. 29. Which contains more thermal energy? A teaspoon of boiling water or a bathtub full of lukewarm water Which has a higher temp?
    30. 30. The Measurement of Energy • Joule – standard unit of measurement for energy • BTU – English unit of measurement for thermal energy, the amount of heat required to raise one pound of water by one degree Fahrenheit • Calorie – amount of energy required to raise the temperature of one gram of water one degree Celsius • 1 cal = 4.184 J
    31. 31. Temperature Scales Celsius scale – freezing point of water is 0◦ C boiling point of water is 100◦ C Kelvin scale – uses absolute zero (point at which molecules no longer move) as the zero point freezing point of water is 273 K boiling point of water is 373 K Fahrenheit scale – freezing point of water is 32◦ F boiling point of water is 212◦ F
    32. 32. Conversion between scales K = ◦ C + 273 ◦ C = K - 273 ◦ F = (1.8 x ◦ C) ◦ C = (◦ F-32)/1.8 Sample Problem: The weatherman announces that the high for the day is expected to be 33◦ C What is this temperature on the Kelvin scale and the Fahrenheit scale?
    33. 33. Phase Changes of Matter • Condensation –gas to liquid • Vaporization – liquid to gas • Freezing – liquid to solid • Melting –solid to liquid • Sublimation – solid to gas • Deposition – gas to solid
    34. 34. Tuesday Homework Read pgs 29 – 39 Section Review Questions 2B Pg 36, questions 1 - 4
    35. 35. Wednesday • Do Review Questions pg 40 & 41
    36. 36. Thursday Go Over Review
    37. 37. Friday • Test Ch 1&2