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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

Transcript

  • 1. Science, Chemistry and You
  • 2. Chemistry • Definition – study of the composition and properties of matter and the energy transformations accompanying changes in the structure of matter
  • 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. 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. 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. 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. Reaction of the Day Table sugar + sulfuric acid  Carbon + H20 H2SO4 C12H22011 (s)  12 C (s) + 11 H2O (g)
  • 8. Ch 2 - Matter Matter – anything that takes up space and has mass
  • 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. What is each picture modeling? Density, malleability, ductility, conductivity
  • 11. Physical or Chemical Change
  • 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. 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. 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. 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. 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. Elements and Their Symbols Element - pure substance that cannot be broken down into simpler substances
  • 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. Elements and Their Symbols Symbol – letter given to represent the name of each element Hydrogen Oxygen Calcium Magnesium Manganese Sodium
  • 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. 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. 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. Tuesday September 14, 2010 • Go over homework problems
  • 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. 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. 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. Energy Conservation • Entropy – randomness or disorder of a system • There is a tendency for all natural processes to increase in entropy (disorder)
  • 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. Which contains more thermal energy? A teaspoon of boiling water or a bathtub full of lukewarm water Which has a higher temp?
  • 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. 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. 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. 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. Tuesday Homework Read pgs 29 – 39 Section Review Questions 2B Pg 36, questions 1 - 4
  • 35. Wednesday • Do Review Questions pg 40 & 41
  • 36. Thursday Go Over Review
  • 37. Friday • Test Ch 1&2