Classifying Matter

1. Chapter 3: Matter & Energy Classifying Matter

2. Day One And Two Separation of mixtures

3. Separation of Mixtures • Objectives: • Make a mixture out of sand, salt, and iron • Separate the sand, salt, and iron mixture • Determine percent recovered of iron, salt, and sand

4. Separation of Mixtures • You and your partner must submit to me a procedure as to how you are going to separate the salt, sand, and iron mixture before you go into the lab • Materials Used: • ANYTHING YOU WANT! ASK!

5. Day Three Chemical & PhysicalConservation of Matter

6. What Is Matter? • Matter is defined as anything that occupies space and has mass. • Even though it appears to be smooth and continuous, matter is actually composed of a lot of tiny little pieces we call atoms and molecules.

7. Atoms and Molecules • Atoms are the tiny particles that make up all matter. • In most substances, the atoms are joined together in units called molecules. • The atoms are joined in specific geometric arrangements.

8. Classifying Matterby Physical State • Matter can be classified as solid, liquid, or gas based on what properties it exhibits. • Fixed = Property doesn’t change when placed in a container. • Indefinite = Takes the property of the container.

9. 1. Arrangement 2. Movement 3. Volume Solid Liquid Gas

10. Physical and Chemical Properties • Characteristic that is displayed by the substance WITHOUT changing its composition • Examples: • Odor • Boiling Point • Melting Point • Density • Characteristics that is displayed by the substance WITH changing its composition • Examples: • Flammability • Corrosiveness • Acidity • Toxicity PHYSICAL CHEMICAL

11. Physical and Chemical Change • Matter changes its appearance but not its composition • Example: • Phase Changes • Change in appearance • Matter DOES change its composition • Results in a completely NEW substance • Example: • Burning • Heat exchange • Evolution of a gas • Formation of a precipitate Physical Change Chemical Change

12. Day Four Classification Activity

13. Day Five Classifying Matter Separating Mixtures

14. Types of Matter

15. Pure Substances 1. All samples have the same physical and chemical properties. 2. Constant composition = All samples have the same pieces in the same percentages. 3. Homogeneous. 4. Separate into components based on chemical properties. 5. Temperature stays constant while melting or boiling. Mixtures 1. Different samples may show different properties. 2. Variable composition = Samples made with the same pure substances may have different percentages. 3. Homogeneous or heterogeneous. 4. Separate into components based on physical properties. 5. Temperature usually changes while melting or boiling because composition changes. Pure Substances vs. Mixtures

16. Different Physical Property Technique Boiling point Distillation State of matter (solid/liquid/gas) Filtration Adherence to a surface Chromatography Volatility Evaporation Density Centrifugation and decanting Separation of Mixtures • Separate mixtures based on different physical properties of the components. • Physical change.

17. Distillation

18. Filtration

19. Law of Conservation of Mass • Antoine Lavoisier • “Matter is neither created nor destroyed in a chemical reaction.” • The total amount of matter present before a chemical reaction is always the same as the total amount after. • The total mass of all the reactants is equal to the total mass of all the products.

20. Conservation of Mass • Total amount of matter remains constant in a chemical reaction. • 58 grams of butane burns in 208 grams of oxygen to form 176 grams of carbon dioxide and 90 grams of water.

21. Day Six Energy

22. Energy • There are things that do not have mass and volume. • These things fall into a category we call energy. • Energy is anything that has the capacity to do work. • Although chemistry is the study of matter, matter is effected by energy. • It can cause physical and/or chemical changes in matter.

23. Some Forms of Energy • Electrical • Kinetic energy associated with the flow of electrical charge. • Heat or Thermal Energy • Kinetic energy associated with molecular motion. • Light or Radiant Energy • Kinetic energy associated with energy transitions in an atom. • Nuclear • Potential energy in the nucleus of atoms. • Chemical • Potential energy in the attachment of atoms or because of their position.

24. “Losing” Energy • If a process was 100% efficient, we could theoretically get all the energy transformed into a useful form. • Unfortunately we cannot get a 100% efficient process. • The energy “lost” in the process is energy transformed into a form we cannot use.

25. Units of Energy • Calorie (cal) is the amount of energy needed to raise one gram of water by 1 °C. • kcal = energy needed to raise 1000 g of water 1 °C. • food calories = kcals.

26. Exothermic Processes • When a change results in the release of energy it is called an exothermic process. • The excess energy is released into the surrounding materials, adding energy to them. • Often the surrounding materials get hotter from the energy released by the reaction.

27. Endothermic Processes • When a change requires the absorption of energy it is called an endothermic process. • The required energy is absorbed from the surrounding materials, taking energy from them. • Often the surrounding materials get colder due to the energy being removed by the reaction.

28. Temperature Scales • Fahrenheit scale, °F. • Used in the U.S. • Celsius scale, °C. • Used in all other countries. • A Celsius degree is 1.8 times larger than a Fahrenheit degree. • Kelvin scale, K. • Absolute scale.

29. Temperature Scales • The Fahrenheit temperature scale used as its two reference points the freezing point of concentrated saltwater (0 °F) and average body temperature (96 °F). • More accurate measure now sets average body temperature at 98.6 °F. • Room temperature is about 72 °F.

30. Temperature Scales, Continued • The Celsius temperature scale used as its two reference points the freezing point of distilled water (0 °C) and boiling point of distilled water (100 °C). • More reproducible standards. • Most commonly used in science. • Room temperature is about 22 °C.

31. Fahrenheit vs. Celsius • A Celsius degree is 1.8 times larger than a Fahrenheit degree. • The standard used for 0° on the Fahrenheit scale is a lower temperature than the standard used for 0° on the Celsius scale.

32. The Kelvin Temperature Scale • Both the Celsius and Fahrenheit scales have negative numbers. • Yet, real physical things are always positive amounts! • The Kelvin scale is an absolute scale, meaning it measures the actual temperature of an object. • 0 K is called absolute zero. It is too cold for matter to exist because all molecular motion would stop. • 0 K = -273 °C = -459 °F. • Absolute zero is a theoretical value obtained by following patterns mathematically.

33. Kelvin vs. Celsius • The size of a “degree” on the Kelvin scale is the same as on the Celsius scale. • Although technically, we don’t call the divisions on the Kelvin scale degrees; we call them kelvins! • That makes 1 K 1.8 times larger than 1 °F. • The 0 standard on the Kelvin scale is a much lower temperature than on the Celsius scale. • When converting between kelvins and °C, remember that the kelvin temperature is always the larger number and always positive!

34. Day Seven Heat Capacity Specific Heat

35. Change in Heat • Example • Window in the winter time • Energy always flows in the same direction → • When does the energy flow stop?

36. Heat Capacity • Heat capacity is the amount of heat a substance must absorb to raise its temperature by 1 °C. • cal/°C or J/°C. • Metals have low heat capacities; insulators have high heat capacities. • Specific heat = heat capacity of 1 gram of the substance. • cal/g°C or J/g°C. • Water’s specific heat = 4.184 J/g°C for liquid. • Or 1.000 cal/g°C. • It is less for ice and steam.

37. Heat Capacity • Heat capacity is the amount of heat a substance must absorb to raise its temperature by 1 °C. • cal/°C or J/°C. • Metals have low heat capacities; insulators have high heat capacities. • Specific heat = heat capacity of 1 gram of the substance. • cal/g°C or J/g°C. • Water’s specific heat = 4.184 J/g°C for liquid. • Or 1.000 cal/g°C. • It is less for ice and steam.

38. Specific Heat Capacities

39. Heat Gain or Loss by an Object • The amount of heat energy gained or lost by an object depends on 3 factors: how much material there is, what the material is, and how much the temperature changed.

40. Hg has a specific heat of 0.139 J/g⁰C. How much heat is required to raise the temperature of a 22.80 grams sample from 16.2⁰C to 32.5⁰C?

41. How many joules of heat are required to raise the temperature of 200 grams of water from 20.0⁰C to 50.0 ⁰C?

42. Day Eight Calorimetry

43. All food has energy, so how can we measure it? • Energy remember is a transfer of heat • Some food (Bugles for instance) we can burn and it will continue to burn on its own until it uses up all the energy in the food. • If we can measure the heat it gives off we can calculate the energy.

44. Measure energy in food • If its giving off heat then we can measure the temperature change in the surrounding air • However, the energy would dissipate very quickly and it would not be a good way to get the temperature change • We use calorimeters!!!

47. What is Heat Capacity? • heat capacity is for objects whose size is predetermined and we can factor out the mass. • Units for Heat Capacity is (J/ ⁰C)

48. A bomb calorimeter was filled with propane which was then ignited. This reaction released 104,000 J of energy. Initially, the temperature of the calorimeter was 25⁰C, after the reaction the temperature was measured at 47.5 ⁰C. What is the heat capacity of this calorimeter?

49. Day Nine Work Day

50. Day Ten Bugle Lab