Chapter 2 - Matter and Change

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This presentation covers Chapter 2 on the general topic of Matter and Change. Substances or mixtures? Chemical or physical changes?

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Chapter 2 - Matter and Change

  1. 1. Chapter 2 – Matter & Change By Kendon Smith – Columbia Central HS – Brooklyn, MI
  2. 2. Chapter 2 – Matter and ChangeI. Properties of Matter 1. Matter is defined as anything that has mass and volume. a. Mass: the amount of matter an object contains. b. Volume: the amount of space an object occupies.
  3. 3. 2. Describing Mattera. Extensive Property: property that depends on the amount of matter in a sampleb. Intensive Property: property that depends on the type of matter in a sampleExamples: Identify the following properties of a sample as extensive or intensive.Mass: ___ Color: ___ Flammability: ___Volume: ___ Density: ___ Melting Point: ___
  4. 4. 3. Identifying Substances a. Substance: matter that has uniform and definite composition. - Uniform means it is the same throughout the sample. - Definite composition means chemically the same all over the planet!Examples: Water? Sand? Copper? Salt?
  5. 5. b. Physical Property – a quality or condition of a substance that can be measured or observed without changing a substance’s composition.Examples:color, solubility, odor, hardness, density, melting point, boiling point, physical state, mass, volume
  6. 6. 4. States of MatterA. Solids – form of matter with definite shape and definite volume - Particles are packed tightly togetherB. Liquids – form of matter with indefinite shape and definite volume - Liquids take the shape of their container - Particles in liquids are almost incompressible - Particles in liquids usually expand slightly when heated. - Particles in a liquid are close together, but can flow past each other
  7. 7. 4. States of MatterC. Gases – form of matter with indefinite shape and indefinite volume - Particles in gases very far apart compared to liquids and solids. - Particles in gases are easily compressed, pushing closer together. * Vapors – Gaseous substances that are liquids or solids at room temperature.
  8. 8. States of Matter
  9. 9. 5. Physical changes – alter a material without changing its chemical composition - Physical changes can be classified as reversible or irreversible a. Reversible changes: melt, freeze, boil, condense b. Irreversible changes: break, split, grind, cut, crush
  10. 10. II. Mixtures 1. Mixture – a physical blend of two or more substances. a. Homogeneous mixture: - completely uniform composition, - parts are not readily distinguished Examples – brine, gatorade, jello, creamy peanut butter
  11. 11. II. Mixtures 1. Mixture – a physical blend of two or more substances. b. Heterogeneous mixture: - not uniform composition - parts are readily distinguished Examples – sand and water, a salad, jello w/ fruit, chunky peanut butter
  12. 12. 2. Solution – a homogeneous mixture - may be solid, liquid, or gas Examples – brine, sugar water, air, stainless steel3. Phase – any part of a sample that with uniform composition and properties How many phases? Brine = ___ Oil/Vinegar = ___ Air = ___ Sand/Water = ___
  13. 13. 4. Separating Mixtures: Mixtures can be separated using physical means.a. Decant, or pour off, a liquid layer. Ex. Oil/Water Mixb. Filtration: Separates solids from liquids with their dissolved particles.c. Distillation – Boil a liquid to become a vapor and condense it back.d. Magnet – Pull out magnetic particles from a mixture.e. Sorting - Physically separating by appearance Ex. size, color, shape
  14. 14. III. Elements and Compounds1. Distinguishing elements and compoundsa. Elements – the simplest form of matter with a unique set of properties - Elements are the basic building blocks of all other substances. - Cannot be separated into simpler substances by chemical means. Examples: hydrogen, oxygen, carbon
  15. 15. III. Elements and Compoundsb. Compounds – two or more elements chemically combined in a fixed proportion - Compounds can only be separated into simpler substances by chemical means, but not by physical means, like mixtures. Examples: water (H2O) sucrose (C12H22O11) carbon dioxide (CO2) sugar water???
  16. 16. III. Elements and Compoundsc. Chemical change: produces matter with a different composition - Bonds are broken to chemically separate elements. - New bonds are formed to chemically join elements.Example 1: sucrose (table sugar) HEAT C12H22O11 12 C + 11 H2O
  17. 17. III. Elements and Compoundsc. Chemical change: produces matter with a different composition - Bonds are broken to chemically separate elements. - New bonds are formed to chemically join elements.Example 2: water (electrolysis) ELECTRICITY 2 H2O 2 H2 + O2
  18. 18. III. Elements and Compounds2. Properties of Compoundsa. Properties of compounds can be quite different from those of their component elements. - Each compound is a new, pure substance with unique properties.
  19. 19. a. Properties of compounds can be quite different from those of their component elements.Example 1: Water – made of hydrogen and oxygenProperties: Hydrogen: Extremely flammable gas Oxygen: Gas that supports combustion Water: Non-flammable liquid!
  20. 20. a. Properties of compounds can be quite different from those of their component elements.Example 2: Table Salt – made of sodium and chlorine Soft, silvery metal thatProperties: Sodium: reacts violently with water Poisonous, yellowish gas Chlorine: white, crystalline solid Sodium chloride: that we must eat!
  21. 21. III. Elements and Compounds3. Distinguishing Substances and Mixtures a. Fixed composition – the material is a substance b. Variable composition – the material is a mixture KEY CONCEPT! * Homogeneous mixtures look like substances because both appear to be made of one kind of matter.
  22. 22. Matter Substance Can be Mixture (definite (variable composition) physically composition) separatedElement Compound Homogeneous Heterogeneous (uniform; called (not uniform; a solution) distinct phases) Can be chemically separated
  23. 23. Matter Substance Mixture (definite (variable lo composition) ok composition) s lik eElement Compound Homogeneous Heterogeneous (uniform; called (not uniform; a solution) distinct phases)
  24. 24. 4. Symbols and Formulas a. Chemical symbol – one or two letters which represent each element - First letter of the symbol is always capitalized, while the second letter is always lower case. - Most symbols are derived from the Latin names for the elements.
  25. 25. Name Latin Symbol Sodium Natrium NaPotassium Kalium KAntimony Stibium Sb Gold Aurum Au Silver Argentum Ag Iron Ferrum Fe Lead Plumbum PbMercury Hydrargyrum HgTungsten Wolfran W Copper Cuprum Cu
  26. 26. b. Chemical formulas – use symbols to show the relative proportions of elements in a compound. - Because compounds have fixed compositions, the chemical formula for a compound is always the same. - Subscripts are written next to the symbol and give us the proportion of each element in the compound. i. The number one is never written as a subscript, it is assumed! ii. If a subscript is written outside parenthesis, it multiplies the proportions of all the elements inside the parenthesis.
  27. 27. Examples: What are the proportions of elements in each of the following compounds?1. Table sugar (sucrose) = 12 carbons, 22 hydrogens, 11 oxygens2. C2H5OH (ethanol) = 2 carbons, 6 hydrogens, 1 oxygen3. Al(OH)3 = 1 aluminum, 3 oxygens, 3 hydrogens4. (NH4)2CO3 = 2 nitrogens, 8 hydrogens, 1 carbon, 3 oxygens5. Co3(PO4)2 = 3 cobalts, 2 phosphorus, 8 oxygen
  28. 28. IV. Chemical Reactions1. Chemical Changes: always involve a change in chemical compositiona. The ability of a substance to undergo chemical change is called a chemical propertyb. Words that signify chemical change: burn, rot, rust, decompose, ferment, explode, corrodec. Chemical properties can only be observed while a substance is undergoing a chemical change, during what is called a chemical reaction
  29. 29. IV. Chemical Reactions2. Chemical reaction – one or more substances change into new substancesa. Reactant – a substance present at the start of the reaction (BEFORE!)b. Product – a substance produced during a reaction (AFTER!) REACTANTS PRODUCTS
  30. 30. 3. Some basics rules for writing chemical reactions:a. Reactants written on the left and products on the right with an arrow between.b. Reactants are separated from each other with a + symbol. Same for products.c. The arrows always points toward the products and can be read as “yields” or “produces” or “changes into”.
  31. 31. Example 1 - Oxidation of Iron: The reaction between iron and oxygen yields iron oxide. Iron + Oxygen Iron oxideExample 2: During combustion, ethanol reacts with oxygen and produces carbon dioxide and water. ethanol + oxygen carbon dioxide + water
  32. 32. 4. Recognizing Chemical Changes a. Clues that a chemical reaction has taken place: 1. Energy is absorbed or released - hot or cold! Examples: Burning wood – gives off heat Cook food – absorbs heat *Note: There are energy changes during changes of state, which are only physical changes, so energy change alone is not sufficient!
  33. 33. a. Clues that a chemical reaction has taken place:2. Change in color = new chemical has been formed!3. Change in odor = smells are caused by chemicals!4. Production of gas = you will see bubbles!5. Production of a precipitate = a solid formed from liquid mixtures.
  34. 34. 5. Conservation of Massa. Law of Conservation of Mass: mass cannot be created or destroyedb. In a chemical reaction, the mass of all the reactants must equal the mass of all the products.
  35. 35. Examples:10g of ice melt = 10g of water (phys. change)2.5 kg of wood burn & produce 0.4 kg of ash?- 2.1 kg of water vapor, carbon dioxide, and other gases released to air1.5 grams of magnesium ribbon burns and the ashes weigh 2.1 grams- combined with oxygen in the air to produce magnesium oxide

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