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Warm-up 9/7/10. New Unit – Matter and Energy Warm-up #1: Describe an apple. Objectives. Students will apply their understanding of physical and chemical properties and changes. (CS 4.12.1*, CLG 4.1.1*). AGENDA . Discuss warm-up Notes on Matter Matter Video Clip Wrap-up.

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Warm-up 9/7/10


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    Presentation Transcript
    1. Warm-up 9/7/10 • New Unit – Matter and Energy • Warm-up #1: Describe an apple.

    2. Objectives • Students will apply their understanding of physical and chemical properties and changes. (CS 4.12.1*, CLG 4.1.1*)

    3. AGENDA • Discuss warm-up • Notes on Matter • Matter Video Clip • Wrap-up

    4. Chemistry Unit Two Matter and Energy

    5. Matter • Matter • anything that has a mass and takes up space. • Law of Conservation of Mass/Matter • Matter can be described using ..

    6. PROPERTIES Types of Properties

    7. Characteristics of Matter • Physical Properties • Examples: • Color,smell, taste, hardness, density, texture, melting/boiling/freezing points, magnetic attraction, solubility, electrical conductivity, temperature, state or phase

    8. Two Types of Physical Properties • Extensive • Intensive • Depends on the type of matter  NOT size of sample • examples: color, melting point, specific heat, density, appearance, etc…

    9. Characteristics of Matter • Chemical Properties • Examples: • Oxidation, Corrosion, Hydrolysis, Combustion, Flammability, Reaction to Acid or Base.

    10. Two Types of Changes • Physical Change • . • Does not change the chemical composition of the matter!!

    11. Characteristics of Matter • Chemical Change • forms a new substance that has different physical and chemical properties than the original substance. • Also known as a .

    12. World of Chemistry • http://www.learner.org/resources/series61.html?pop=yes&pid=797

    13. Kinetic Theory • All matter is made of tiny particles in . • Potential Energy (PE) • energy due to the position or condition • at the atomic level: • the distance between the particles • = lower PE = higher PE • Kinetic Energy (KE) • energy due to motion • =higher KE = lower KE

    14. Phases of Matter

    15. Phases of Matter

    16. Plasma • extraordinary state of matter • consists of high energy particles • electrons are stripped from their nuclei • examples: *Most Abundant State of Matter in the Universe!*

    17. Phase Changes Changes of State • energy (heat) to a substance causes phase changes • The potential energy of the particles is • During a phase change, temperature

    18. Phase Changes • Melting • Δ (adding energy) • Freezing • Δ (removing energy) • Melting point & freezing point of a substance occur temperature.

    19. Phase Changes • Boiling • Δ (adding energy) • Evaporation • Δ (adding energy) • Condensation • Δ (removing energy) • Difference between boiling & evaporation: • a specific temp. below the surface • any temp. at the surface

    20. Phase Changes • Deposition • Δ (removing energy) • Examples: • Sublimation • Δ (adding energy) • Examples:

    21. Phase Change Graphs (T vs t) Liquid Melting Solid AB -heat Δ KE -move faster -temp.  -solid BC -heat Δ PE -get farther apart -temp. stay same -melting CD -heat Δ KE -move faster -temp.  -liquid

    22. Phase Change Graph (T vs t) Gas Boiling DE -heat Δ PE -get farther apart -temp. stay same -boiling EF -heat Δ KE -move faster -temp.  -gas

    23. Phase Change Graph (T vs t) A C B E D F

    24. Phase Change Graph (T vs t) A B C D E F

    25. Phase Change Graph (T vs t) Boiling Freezing Melting What is the boiling point? What is the melting point? What is the freezing point?

    26. Phase Change Graph (T vs t) If melting & freezing points occur at the same temperature, how do you know which change is occurring?

    27. Phase Change Graph (T vs t) What is this substance? - How do you know? -

    28. Heat Calculations • Heat (q) • Energy transferred from an object at a higher temperature to an object at a lower temperature. (heat lost = -heat gained) • q = mcT • q=mHfus • q=mHvap

    29. Heat Calculations • A 10.0g sample of iron at 50.4oC is cooled to 25.0oC in 50.0g of water. Calculate the amount of heat lost by the iron. ciron= 0.449 J/goC *How much heat is gained by the water? • A 2.1g ice cube at –8.0oC melts completely and warms to 12.5oC. How much heat was required? Hfus ice = 334 J/g cice = 2.03 J/goC cwater = 4.18J/goC

    30. Classification of Matter Matter

    31. Matter • Pure Substances • Mixtures • Element

    32. Matter (cont’d) • Compound • Homogeneous Mixtures • Heterogeneous Mixture

    33. Classifications of Mixtures • Solutions • Colloids • Suspension

    34. Solutions • SOLUTION • ELECTROLYTE

    35. Types of Solutions • Gas-Gas • Liquid-Gas • Gas-Liquid • Liquid-Liquid • Solid-Liquid • Solid-Solid

    36. Characteristics of Solutions • Homogeneous Mixture • Soluble- • Insoluble • Immiscible • Miscible

    37. Solvation • When a solid solute is placed in a solvent, the solvent particles completely surround the surface of the solid solute. • If attractive forces between the solute particles and the solvent are greater than the attractive forces holding the the solute particles together, the solvent particles pull the solute particles apart and surround them.

    38. - + + + - + + + - - - + - + + - + - + + - + - - + - + + + - + - - + + + + Process of Solvation H2O H = O = - - + - + NaCl Na = Cl = + -

    39. Water- Universal Solvent • molecule • Dipoles allow

    40. Solvation Factors • Agitation • Increasing surface area of solute • Increasing temperature of solvent

    41. Heat of Solution • Endothermic- • Exothermic-

    42. Solubility • Refers to • g / 100 g • vs vs

    43. Concentration • How much solute is dissolved in a specific amount of solvent

    44. Molarity • / • Calculate the molarity of 1.60 L of a solution containing 1.55 g of dissolved KBr? • How many grams of CaCl2 would be dissolved in 1.0 L of a 0.10M solution of CaCl2?

    45. Diluting Solutions • M1V1 = M2V2 • What volume of a 3.00 M KI stock solution would you use to make ).300 L of a 1.25 MKI solution?

    46. Colligative properties of Solns • Physical properties of solutions that are affected by the number of particles but not the identity of dissolved solute particles