1. Chapter 2 Energy and Matter

2. Energy Defined as… Ability to do work or to produce heat Examples: –Sunlight –Power Bar –Car

3. ENERGY Three main categories are: 1.Radiant ( example: Sunlight… discuss later) 2.Kinetic Energy of Motion 3.Potential Energy of Position

4. Law of Conservation of Energy Energy is neither created nor destroyed. Can be converted / change forms.

5. Measuring Energy Joule (named for James Joule, physicist)  SI unit  4.184 J = 1 cal [note use as conversion factor] 1 cal –Amount of heat energy needed to raise 1g of water 1 o C Calorie = food calorie  1 Cal = 1000 cal

6. Practice Convert 5.15 J to cal and Cal: Convert 300.0 J to cal Convert 433.2 cal to Joules Convert 2500.0 Cal to Joules

7. Energy Crisis With your table partner discuss what is meant by Energy Crisis… Note examples… What are possible solutions to the crisis? Are there present day solutions?

9. 2-2 Temperature Scale of ‘hotness’ Amount of heat (energy, molecular movement) in something Scales –Fahrenheit –Celsius –Kelvin

10. Heat Flow In what direction does heat flow? –From hot things to cold things »OR –From cold things to hot things

11. Heat Flow Endothermic –Heat flows ‘in’ Exothermic –Heat flows ‘out’ What was the heat flow for our COLD PACK LAB

12. Conversions o C to o F o F = 1.8( o C) + 32 o F to o C o C = 0.55 ( o F - 32) o C to K; K to o C K = o C + 273 C = K - 273

14. Calorimetry Molar heat capacity (C molar ) Specific heat capacity (C p ) –Heat required to raise 1 g by 1 K Equation: q = m · C p · ΔT

15. Calorimetry What is the heat required to raise 400.00 g of water by 34.50 o C? What is the heat lost when 200.00 g of iron changes from 115.50 o C to 22.00 o C?

16. Calorimetry Large beds of rocks are used in some solar heated homes to store heat. Assume that the specific heat of rock is 0.082 J/g-K. Calculate the quantity of heat absorbed by 50.0 kg of rocks if their temperature increases by 12.0 o C

17. Flashback… Law of conservation of energy (First law of thermodynamics): –Energy is neither created nor destroyed –Energy is conserved –Energy is transformed

18. Exit Problem What is the function of the coffee cups? Determine the final temperature of water if 525.00 g of water at 57.80 o C RELEASES 6500.00 J of energy. The specific heat of water is 4.184 J/g o C SHOW YOUR WORK AND BOX YOUR ANSWER!!!

19. Calorimetry You have heated a 55.00 g piece of iron, C p = 0.385 J/g-K, to 200.00 o C. You then put the iron into water in a calorimeter. There are 300.00 g of water at 22.00 o C. What is the final temperature of the mixture?

20. Calorimetry In the calorimetry lab you will be mixing an acid and a base and studying the temperature changes. You mix 35.00 mL each of 1 M HCl and 1 M NaOH in a calorimeter. The temperature increases from 21.0 to 27.5 o C. What is the enthalpy change for the reaction in kJ/mol HCl?

22. Matter 10/28/15 Defined: –Anything that has mass and occupies volume

23. States of Matter Solid –Compact –Dense –Atoms held tight –Don’t move –Distinct volume –Don’t fill a container –Discreet shape Liquid –Compact but flows –Less dense than solid –Atoms more loosely held –Atoms/molecules move –Distinct volume –Doesn’t fill container –Takes shape of container

24. States of Matter Gas –Not compact –Not very dense –Diffuses –No distinct volume –Fills and takes shape of container –Flows Plasma –Extremely energetic –Flows –Found in sun, not on earth

25. Properties Physical Properties –Characteristics that can be observed without changing the identity of a substance Chemical Properties –Ability to undergo a change in identity

26. Changes in Matter Physical Changes –Changes in form Folding Tearing Melting Crushing

27. Changes in Matter Chemical Changes –Changes in Identity Burning Rusting Decomposing

28. Changes in Matter

30. Law of Conservation of Mass Mass can be neither created nor destroyed in chemical reactions. The total mass of the products is the same as the total mass of reactants.

32. Pure Substances Element –A–A substance that cannot be separated into simpler substances

33. Elements ~111 presently known elements (92 naturally occurring) Building blocks of all substances At room temperature: –2 liquid –11 gases –All others solid Distribution of elements in galaxies, earth’s crust, seawater and air, and human bodies

34. Elements Names of the elements –Greek –Latin –German –Properties of elements –Scientist who discovered it –Location where discovered

35. Elements Arranged in the Periodic Table (inside front cover) Symbols –One or two letters Usually part of name Some symbols are Latin/Greek name

36. Elements Classification –Metal –Nonmetal –Metalloid See Table 3.5 (page 54)

37. Elements Metals: –U–Usually solid at room temperature –G–Good conductors of heat and electricity –H–High luster –D–Ductile –M–Malleable –H–High melting point; high density –U–Usually don’t combine with each other –R–Readily combine with nonmetals

38. Nonmetals: –Solids (C, P, S, Se, I); Liquid (Br); Gases (all others) –Poor conductors of heat and electricity; no luster –Low melting point; low density –Will combine with each other (CO 2 ) –Will combine with metals or metalloids –Some found uncombined in nature (noble gases)

39. Elements Metalloids –Have properties of both metals and nonmetals –Some used for semiconductors in electronics

40. Pure Substances II Compound –Two or more elements combined through a chemical reaction –Different properties than elements which compose it

41. Compounds Two or more elements chemically combined New properties Definite proportions Can be chemically separated Molecular or Ionic

42. Compounds Molecular –Held together with covalent bonds –Molecule: smallest uncharged individual unit of a compound –Water is an example

43. Compounds Ionic –Ion: positively or negatively charged atom or group of atoms Cation – positive Anion – negative –Held together by ionic bond – attraction between positive and negative charges

44. Compounds Diatomic Molecules –Always only 2 atoms –7 naturally occurring Hydrogen, oxygen, nitrogen, flourine, chlorine, bromine, iodine H 2, O 2, N 2, F 2, Cl 2, Br 2, I 2,

46. Substances and Mixtures Pure Substance: a particular kind of matter with a definite, fixed composition –Elements (copper, gold, oxygen) –Compounds (sugar, salt, water) Mixture: a blend of two or more pure substances –Not chemically combined

47. Matter Pure substances (homogeneous composition) Mixtures of two or more substances ElementsCompounds Solutions (homogeneous composition – one phase) Heterogeneous mixtures (two or more phases) Figure 3.2 (page 48)

48. Types of Mixtures Heterogeneous mixtures –Visibly different parts –Chocolate chip cookies; granite –Two or more phases (usually)

49. Types of Mixtures Homogeneous mixtures –Different parts not visible (uniform throughout) –One phase –Seawater; air

50. Separating Mixtures Do NOT cause chemical changes Heterogeneous Mixtures –Filtration

51. Separating Mixtures Homogeneous Mixtures –Distillation

52. Separating Mixtures Homogeneous Mixtures –Chromatography

53. Separating Mixtures Homogeneous Mixtures –Crystallization