1. Classification of Matter Chemistry Apps Chapter 9 Mr. Gilbertson

2. Composition of Matter Substances (Pure) – cannot be reduced to simpler substances by physical means. Elements – all atoms in a sample are identical. Simplest form of matter Cannot be subdivided by physical means

3. Elements 118 recognized elements 90 natural elements 28 synthetic elements (man-made)

4. Compounds  Made up of two or more elements chemically combined to form new substances.  Have properties which are significantly different from the properties of the elements from which they are formed.  Made up of particles which are either molecules or ion pairs.

5. Mixtures Heterogeneous – non-uniform composition consisting of two or more distinct phases. Physical combination of two or more substances which can be separated by physical means. Each part of the mixture maintains its distinct properties. Homogeneous – uniform composition appearing to be in a single phase. Solutions are the most common examples. One substance (solute) is completely dissolved in another (solvent).

6. Solution Homogeneous mixtures in which one substance it thoroughly distributed in another (single phase at equilibrium) May exist in any phase (solid, liquid, gas) Kool aid, Steel, air, and seawater, are all good examples.

7. Solution Two parts: Solvent – the part that is in greatest quantity or the part in which the other is dissolved. (the water in kool-aid) Solute – the substance being dissolved or distributed (the sugar and powder in kool-aid)

8. Important Terminology System – the total environment of the reaction under study (must be defined) Phase – A region within the system displaying uniform properties Interface – the boundary between two adjacent phases Organic – containing carbon or being produced as a result of life processes Inorganic – not organic

9. Heterogeneous Mixture

10. Classification of Matter

11. Special Mixtures Colloids – a heterogeneous mixture that does not settle out. Exhibits Tyndall effect and Brownian Motion. Particle size is between those of solutions and suspensions. Cannot be filtered out. Suspensions – a heterogeneous mixture that has very large particles which always settle out. Exhibits Tyndall Effect. Can be filtered out.

12. Tyndall Effect Property of all colloids and suspensions which causes them to scatter light as a result of the reflection of very large particles.

13. Describing Matter

14. Physical Properties Any characteristic of a material that you can observe without changing the substances that make up the material. Color, shape, size, density, melting point, etc. Often can be measured directly. All substances have physical properties which distinguish it from other substances.

15. Physical Properties Extensive – depend on the amount of substance present (mass, dimensions, weight, etc.) Intensive – independent of the amount of substance (Density, color, hardness, luster, etc.)

16. Physical Properties

17. Properties used to identify Mixtures are often separated using physical means. The differences between the properties of the substance in a mixture are used to separate the parts. Magnetic properties, solubility, particle size, boiling point, color, are all properties that might be useful for separation.

18. Separation using physical properties Panning for Gold Distillation

19. Separation of matter magnetism A refinery

20. Water purification using Colloids

21. Physical Changes Any change which does not involve a change in composition (the identity of the material does not change). Examples: cutting wood, breaking glass, tearing paper, melting, vaporizing, etc. Changes in size, shape, or state, or arrangement of parts are considered to be physical changes.

22. Chemical Changes Always involve changes in composition, one substance changes to another and its properties change. Evidences: Gases given off, color changes, formation of a precipitate, energy changes (heat or light) are all evidences of a chemical change. (Rule of thumb) May also absorb energy in the form of heat and light.

23. Energy Changes in Chemical Reactions Exothermic reactions – Give off energy to the surroundings (heat up) Endothermic reactions – absorb energy from their environment (cool down) Activation Energy – the energy that is required to cause a reaction to begin.

24. Video about Chemical and Physical changes Physical and Chemical Changes Physical and Chemical Changes

25. Chemical Properties A characteristic of a substance which indicates whether it can undergo a certain chemical change. Describes how a substance will react in the presence of other substances. Cannot be observed without changing the composition of the substances. Flammability, corrosive, toxic, acidic, etc.

26. Flammability

27. Electrolysis Used to separate elements in a compound by passing an electric current through a solution.

28. Chemical Properties

29. Conservation of Mass First proposed by Antoine Lavoisier. The mass of products in a chemical reaction must equal the mass of the reactants. Matter cannot be created or destroyed during a chemical change. Requires a closed system.

30. Antoine Lavoisier Proposed Law of Conservation of Mass

31. Chemical Change

32. Mass is conserved in chemical reactions Sometimes seems to be untrue, but if it is contained in a closed system so that no new matter can enter or exit the reaction vessel. Mass of all products must equal mass of reactants.

33. Silver Ag Properties: The melting point of silver is 961.93°C, boiling point is 2212°C, specific gravity is 10.50 (20°C), with a valence of 1 or 2. Pure silver has a brilliant white metallic luster. Silver is slightly harder than gold. It is very ductile and malleable, exceeded in these properties by gold and palladium. Pure silver has the highest electrical and thermal conductivity of all metals. Silver possesses the lowest contact resistance of all metals. Silver is stable in pure air and water, although it tarnishes upon exposure to ozone, hydrogen sulfide, or air containing sulfur.

34. Heat and Temperature Heat is the transfer of energy from a body at a higher temperature to a body at a lower temperature. Temperature is a measure of the average kinetic energy in the particles of a system. Often confused because the flow of heat will result in a temperature change.

35. Calorimeter Used to measure heat transfer from one body to another Insulated to restrict heat flow to the Environment.

36. Amount of heat required to raise the temperature of one gram of a substance 1 Celsius degree (1 Kelvin). ^Q =mc^T ^Q – heat change (Joules) m – mass (g) C – specific heat (J/g o C) ^T – change in temp ( o C or K) Specific Heat

37. Conservation of energy and specific heat. ^Q loss = ^Q gained Heat lost by sample must equal heat gained by the rest of the system in a calorimeter. So….. mc^T unknown = mc^T water + mc^T cup mc^T water + mc^T cup C unknown = --------------------------------------- m^T unknown Where :^T = T f -T i calorimeter democalorimeter demo

38. THE END