Matter: Properties and Changes

1. Matter: Properties and Changes Chapter 3 Chapter 15.1, 15.3 Chapter 12.4

2. Chapter Objectives • Distinguish between physical and chemical changes • Define and classify matter by composition • Define properties of liquids, solids, and gasses • Identify observable characteristics of a chemical reaction • Explain and apply the fundamental law of conservation of mass

3. Vocabulary – Ch. 3.1 – 3.2 • Physical property • Extensive property • Intensive property • Chemical property • States of matter • Solid • Liquid • Gas • Vapor • Physical Change • Chemical Change • Law of Conservation of Mass • Phase Change

4. Vocabulary – Chapter 15.1, 15.3 • Energy • Heat • Joule • Specific Heat • Specific Heat Equation (q = mC∆T) • Heat of Vaporization (∆Hvap) • Heat of Fusion (∆Hfus) • Heating Curve

5. Vocabulary – Ch. 12.1, 12.4 • Pressure • Barometer • Atmosphere • Melting Point • Vaporization • Evaporation • Vapor Pressure • Boiling Point • Sublimation • Freezing point • Condensation • Deposition • Phase diagram • Triple Point

6. Ch. 3.1 - What is Matter?? • Matter as defined from Ch. 1 is: • Anything that takes up space and has mass • Mass – measure of the amount of matter in an object.

7. What is a Pure Substance? • Matter that has a uniform and unchanging composition • Table salt is ALWAYS Sodium Chloride • Water is always made of 2 Hydrogen atoms and 1 Oxygen atom • Is seawater a substance?

8. Properties of Matter • Chemistry is the study of matter. • Matter is classified according to its physical and chemical properties. • Physical Properties: Can be observed without changing the identity of substance.

9. Some Physical Properties • State at Room temp (liquid, solid or gas) • Melting point • Boiling Point • Viscosity (resistance to flow) • Density • Color • Odor • Brittle/ductile • Electrical/thermal conductivity

10. Physical Properties • Intensive Properties – Independent of amount of material. They property is the same no matter how much is there. • Example: Density • Extensive Properties – Dependent of the amount of material present. • Example: Length, mass, volume

11. Properties of Matter • Chemical Properties: the ability to combine or change into other substances. • Examples: flammability, oxidation, rotting

12. States of Matter • State of Matter: Its physical form. • There are three physical states: Solid: • Definite shape • Definite volume • Closely packed particles

13. States of Matter Liquid: • particles move past each other (flow) • definite volume • takes the shape of its container (indefinite)

14. States of Matter Gas: • flows • takes the shape of its container (indefinite shape) • Fills the container completely. (indefinite volume) • Note: A vapor refers to a gaseous state of a substance that is a solid or liquid at room temperature.

15. Physical or Chemical Property? • Bending of aluminum • Salt dissolving in water • Magnesium burning in air • Baking soda is a white powder • Fluorine is a highly reactive element

16. Ch. 3.2 - Changes in Matter • Physical changes are those which alter the substance without altering its composition. • Change of phase  one physical state to another Melting of ice - composition unchanged, i.e. ice is water in solid form (H2O) They generally require energy, the ability to absorb or release heat (or work).

17. Phase Changes (Ch. 12.4) • What are the phase changes of water? 1. Melting – changing of a solid to a liquid (heat of fusion = ∆Hf) 2. Vaporization – changing from a liquid to a gas (heat of vaporization = ∆Hvap) 3. Sublimation – Changing from a solid to a gas (heat of sublimation = ∆Hsub) What do these processes have in common?

18. Phase Changes • Phase changes in the opposite direction have names too. • liquid to a solid: • gas to a liquid: • gas to a solid: What do these have in common? Answer:

19. James Joule 1818-1889 UNITS OF ENERGY 1 calorie = heat required to raise temp. of 1.00 g of H2O by 1.0 oC. 1000 cal = 1 kilocalorie = 1 kcal 1 kcal = 1 Calorie (a food “calorie”) But we use the unit called the JOULE 1 cal = 4.184 joules (exactly)

20. Heats of Fusion & Vaporization • Heat of Fusion (∆Hfus) – The amount of heat (in joules) needed to melt 1 g of substance. • For ice: 334 J/g • q (heat) = ∆Hfus*m (m= mass of ice/water) • Heat of Vaporization (∆Hvap) – The amount of heat (in joules) needed to vaporize 1 g of substance • For water: 2260 J/g • q (heat) = ∆Hvap*m (m= mass of water/steam)

21. Example Problems • How much heat does it take to melt 20.5 g of ice at 0⁰C? • q = 334 J/g * 20.5 = 6850 J (6.85 kJ) • How much heat is released when 50.0 g of steam at 100 ⁰C condenses to water at 100 ⁰C? • q = - 2260 J/g * 50.0 g = -113,000 J (-113 kJ)

22. Specific Heat Capacity • Specific Heat Capacity – amount of heat (q) required to raise the temperature of one gram of a substance by 1 degree. • C = J (energy gained or lost) mass (g) * Temp Change(⁰C)

23. Heat Capacity Values Substance Spec. Heat (J/g•⁰C) Water 4.184 Ethylene glycol 2.39 Al uminum 0.897 glass 0.84

24. Calculating Heat Gained or lost • The heat, q, gained or lost by a substance can be calculated by knowing the mass of the object, the temperature change, and the heat capacity. • q = mC∆T

25. Calculations involving Heat • Example 1: A 5.00 g piece of aluminum is heated from 25.0⁰Cto 99.5⁰C. How many joules of heat did it absorb? • q = m * C * ∆T = 5.00 g * 0.897 J/g*⁰C * 74.5⁰C = 334 J

26. Calculations involving Heat • Example 2: 10.2 g of cooking oil at 25.0 ⁰C is placed in a pan and 3.34 kJ of heat is required to raise the temperature to 196.4 ⁰C. What is the specific heat of the oil? • q = m*C*∆T • C = q/(m ∆T) • C = 3340 J/(10.2 g * (196.4-25.0) ⁰C) • C = 1.91 J/g* ⁰C

27. Calculations involving Heat • Important Points! • q (heat) is a positive quantity. The sign (+ or -) refers to whether the system you’re looking gained it (+) or lost it (-). • From the previous example, the oil would lose 3340 J of heat upon cooling back to 25.0 ⁰C. (-3340 J heat lost) • Specific heat capacity is like a bucket. It is a measure of how much energy an object absorbs before the temperature changes.

28. Heating Curve for Water Note that T is constant as ice melts

29. Heating/Cooling Curve for Water

30. +333 J/g +2260 J/g Heat & Changes of State What quantity of heat is required to melt 500. g of ice (at 0oC) and heat the water to steam at 100oC? Heat of fusion of ice = 333 J/g Specific heat of water = 4.2 J/g•K Heat of vaporization = 2260 J/g

31. Heat & Changes of State What quantity of energy as heat is required to melt 500. g of ice (at 0⁰C) and heat the water to steam at 100 oC? 1. To melt ice at 0⁰C q = (500. g)(333 J/g) = 1.67 x 105 J 2. To raise water from 0 oC to 100oC q = (500. g)(4.2 J/g•K)(100 - 0)K = 2.1 x 105 J 3. To vaporize water at 100oC q = (500. g)(2260 J/g) = 1.13 x 106 J 4. Total energy = 1.51 x 106 J = 1510 kJ

32. Practice problem If we add 6050 J of heat to 54.2g of ice at -10.0⁰C, whatwillit be at theend? Whattemperaturewillit be? Thespecificheat of ice is 2.03 J/g*⁰C.

33. Pressure • Pressureis the force acting on an object per unit area: • Gravity exerts a force on the earth’s atmosphere • A column of air 1 m2 in cross section exerts a force of about 105 N (101,300 N/m2). • 1 Pascal (Pa) = 1 N/m2 . So, 101,300 N/m2 = 101,300 Pa or 101.3 kPa. • Since we are at the surface of the earth, we ‘feel’ 1 atmosphere of pressure.

34. Barometer Vacuum • A barometer measures atmospheric pressure • The pressure of the atmosphere at sea level will hold a column of mercury 760 mm Hg. • 1 atm = 760 mm Hg 760 mm Hg 1 atm Pressure

35. Units of pressure • 1 atmosphere (atm) = 760 mm Hg = 760 torr • 1 atm = 101,300 Pascals = 101.3 kPa • Can make conversion factors from these. • What is 724 mm Hg in atm ? • What is 724 mm Hg in kPa

36. Phase Changes Vapor pressure is the pressure exerted by a vapor over a liquid. The vapor pressure increases with increasing temperature. This is why water evaporates even though it’s not 212˚F.

37. Phase Changes However, when the vapor pressure of the water is the same as the atmospheric pressure the water is … boiling.

38. Phase Diagram A phase diagram is a graph of pressure vs temperature that shows in which phase a substance exists under different conditions of T & P.

39. Solid Liquid Gas Phase Diagram for water Boiling Pressure 1 Atm Melting Condensation Freezing Sublimation Deposition Temperature

40. Ch. 3.2 - Chemical Properties and Changes • Chemical Properties – the ability to combine or change into other substances. • Example: Water won’t react with aluminum, but reacts with sodium and potassium (violently) and iron (slowly but surely). • A chemical property always relates to a chemical change – the changing of one or more substances into other substances. • A chemical change is also known as a chemical reaction.

41. Signs of a Chemical Change • Color change • Odor • Gas • Formation of a precipitate (solid) • Heat

42. Chemical Properties and Changes • Chemical changes are a rearrangement of atoms in the substance. • Chemical changes follow the Law of Conservation of Mass. • This means that when there is a chemical change, matter is neither created or destroyed, just changed in form. Massreactants = Massproducts

43. Conservation of Mass Example • Reactants are the substances you start with. • Started with sugar and sulfuric acid • alcohol and air (O2) • Products are the new substances that are made. - Carbon, water vapor, and heat - water & carbon dioxide

44. Conservation of Mass Example • Example: 10.00 g of red mercury (II) oxide powder is placed in an open flask. It is heated until it has fully converted to liquid mercury and oxygen gas. • Written as: • Mercury oxide  mercury + oxygen • The liquid mercury has a mass of 9.26 grams • The mass of oxygen formed in this reaction is 10.0 g (total) – 9.26 g (mercury) = 0.74 g oxygen

45. Practice • 57.48 g of sodium reacts with chlorine gas to form 146.10 g of sodium chloride. • Sodium + Chlorine  sodium chloride • How much chlorine gas was used in the reaction?

46. Elements • An element is a pure substance that cannot be broken down into simpler substances • Each element has a symbol and they are arranged in a periodic table

47. Elements • Each element has it’s own symbol in the periodic table. • Either a single capital letter: H is for hydrogen • Or two letters, the first is capital the second is ALWAYS small: He is for helium. • Co is different than CO.

49. Practice Name the element or symbol. • C • Calcium • Cl • Iodine • K • Mercury

50. Compounds • Compounds are pure substances that are combinations of elements in a fixed ratio. • Example: Table salt is a combination of sodium and chlorine! NaCl (1:1 ratio)