CHEM 10 CHP 3

1. CHEM 10 CHP 3 MATTER & ENERGY

2. Matter • Anything that occupies space and has mass • Composed ultimately of atoms

3. Structure Determines Properties • the atoms or molecules have different structures in solids, liquid and gases, leading to different properties

4. Classifying Physical States of Matter: Solid: it has a definite shape and volume Liquid: it has a definite volume, but no shape; takes on shape of container Gas: it has no definite shape and no definite volume (fills shape of container, but mostly empty space)

5. TYPES OF SOLIDS Crystalline: regular repeated patterns Salt, sugar, quartz Amorphous: no patterns Plastic, gel, glass

6. Substances and Mixtures Substance - matter with a definite , fixed composition Also known as pure substances Examples - elements or compounds Mixture– two or more substances that mix, homogeneous or heterogeneous Examples – coffee or cereal See Figure 3.8: memorize!

7. Classifying Matter

8. Substances and Mixtures: Definitions Homogenous - uniform appearance - has same properties throughout, like coffee Heterogeneous - contains two or more physically distinct phases, like ice cubes in water or oil on top of water Phase - homogeneous part of system - separated by boundaries System - body of matter under consideration

9. Elements Fundamental/elementary substances that cannot be broken down by chemical means into simpler substances Atom is the smallest particle of an element 91 elements occur naturally in nature

10. Compounds Made of two or more elements with some type of bonding arrangement Examples: water H2O ethanol CH3CH2OH sugar (sucrose) C12H22O11 table salt NaCl

11. ELEMENTS & COMPOUNDS Element names and symbols: - Elements are made of atoms or diatomic molecules - Compounds are made of atoms from two different elements, which form molecules or formula units

12. ELEMENTS & COMPOUNDS Examples: Au is symbol for atom of gold, pure element Learn all the seven diatomic gases. H2 is symbol for a diatomic molecule for pure element hydrogen H2O consists of a molecule that has two hydrogen atoms and 1 oxygen atom

13. Properties of a Substance A property is a characteristic of a substance Each substance has a set of properties that are characteristic of that substance and give it a unique identity

14. Physical Properties The inherent characteristics of a substance that are determined without changing its composition. Examples: taste color physical state melting point boiling point

15. (from another textbook)

16. Chemical Properties Describe the ability of a substance to form new substances, either by reaction with other substances or by decomposition.

17. Chemical Properties of Chlorine It will not burn in oxygen. It will support the combustion of certain other substances. It can be used as a bleaching agent. It can be used as a water disinfectant. It can combine with sodium to form sodium chloride.

18. Physical Changes Changes in physical properties (such as size, shape, and density) or changes in the state of matter without an accompanying change in composition. Examples: tearing of paper change of ice into water change of water into steam heating platinum wire No new substances are formed.

19. Chemical Changes In a chemical change new substances are formed that have different properties and composition from the original material. (Same as chemical properties)

20. Formation of Copper(II) Oxide Heating a copper wire in a Bunsen burner causes the copper to lose its original appearance and become a black material. The formation of copper(II) oxide from copper and oxygen is a chemical change. The copper (II) oxide is a new substance with properties that are different from copper. Heating a copper wire in a Bunsen burner causes the copper to lose its original appearance and become a black material. The black material is a new substance called copper(II) oxide. Copper is 100% copper by mass. Copper (II) oxide is: 79.94% copper by mass 20.1% oxygen by mass.

21. Formation of Copper(II) Oxide Neither Cu nor O2 contains Cu2+ or O2- A chemical change has occurred. Copper(II) oxide is made up of Cu2+ and O2- 4.2

22. Decomposition of Water The composition and physical appearance of hydrogen and oxygen are different from water. The hydrogen explodes with a pop upon the addition of a burning splint. The oxygen causes the flame of a burning splint to intensify. Water is decomposed into hydrogen and oxygen by passing electricity through it. They are both colorless gases. But the burning splint is extinguished when placed into the water sample.

23. products reactant yields CHEMICAL EQUATIONS Water decomposes into hydrogen and oxygen when electrolyzed.

24. 2H2O 2H2 O2 reactant yields products Water decomposes into hydrogen and oxygen when electrolyzed.

25. heat product yield reactants Copper plus oxygen yields copper(II) oxide.

26. 2Cu O2 2CuO product yield reactants Copper plus oxygen yields copper(II) oxide. heat

28. 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

29. 46.0 g 32.1 g mass reactants = sodium + sulfur  sodium sulfide 78.1 g → 78.1 g reactant 78.1 g product mass products

30. ENERGY!!!! • Energy is the capacity to do work • Two main forms of energy: • Potential • Nonpotential

31. Potential Energy Energy that an object possesses due to its relative position. Stored energy: positional, chemical, etc.

32. increasing potential energy 50 ft increasing potential energy 20 ft The potential energy of the ball increases with increasing height.

33. Gasoline is a source of chemical potential energy. • The heat released when gasoline burns is associated with a decrease in its chemical potential energy. • The new substances formed by burning have less chemical potential energy than the gasoline and oxygen.

34. Types of Nonpotential Energy Mechanical/kinetic: Energy matter possesses due to its motion. KE = ½ mv2 Chemical: produced in reactions; potential energy in the attachment of atoms or because of their position Electrical: kinetic energy associated with the flow of electrical charge Heat or thermal: q = m * cp *DT Nuclear: potential energy in the nucleus of atoms Radiant or Light: E = hn

35. Moving bodies possess kinetic energy. • A bouncing ball. • The running man.

36. Heat: Quantitative Measurement • A form of energy associated with small particles of matter. Heat • A measure of the intensity of heat, or of how hot or cold a system is. Temperature

37. The SI unit for heat energy is the joule (pronounced “jool”). Another unit is the calorie. (exactly) 4.184 Joules = 1 calorie 4.184 J = 1 cal This amount of heat energy will raise the temperature of 1 gram of water 1oC.

38. 100 g water 30oC 200 g water 30oC 8368 J 4184 J A B 100 g water 20oC 200 g water 20oC Twice as much heat energy is required to raise the temperature of 200 g of water 10oC as compared to 100 g of water. temperature rises 10oC heat beakers

39. The specific heat (capacity), cp, of a substance is the quantity of heat required to change the temperature of 1 g of that substance by 1oC.

40. The units of specific heat in joules are:

41. The units of specific heat in calories are:

43. ( ) specific heatof substance ΔT = heat ( ) mass of substance The relation of mass, specific heat, temperature change (Δt), and quantity of heat lost or gained (commonly called heat transfer) is expressed by the general equation: q = m*cp*DT MEMORIZE THIS!!

44. Calculate the specific heat of a solid in J/goC and in cal/ goC if 1638 J raise the temperature of 125 g of the solid from 25.0oC to 52.6oC. (mass)(specific heat)ΔT = q (m)(cp)ΔT= q q = 1638 J cp = q/mDT m = 125 g ΔT = 52.6oC – 25.0oC = 27.6oC cp = (0.114 cal/goC)

45. A sample of a metal with a mass of 212 g is heated to 125.0oC and then dropped into 375 g of water at 24.0oC. If the final temperature of the water is 34.2oC, what is the specific heat of the metal? When the metal enters the water, it begins to cool, losing heat to the water. At the same time, the temperature of the water rises. This process continues until the temperature of the metal and the temperature of the water are equal, at which point (34.2oC) no net flow of heat occurs.

46. A sample of a metal with a mass of 212 g is heated to 125.0oC and then dropped into 375 g of water at 24.0oC. If the final temperature of the water is 34.2oC, what is the specific heat of the metal? • Calculate the heat gained by the water. • Calculate the final temperature of the metal. • Calculate the specific heat of the metal.

47. heat gained by the water = A sample of a metal with a mass of 212 g is heated to 125.0oC and then dropped into 375 g of water at 24.0oC. If the final temperature of the water is 34.2oC, what is the specific heat of the metal? Heat Gained by the Water temperature rise of the water ΔT= 34.2oC – 24.0oC = 10.2oC

48. heat lost by the metal heat gained by the water = = A sample of a metal with a mass of 212 g is heated to 125.0oC and then dropped into 375 g of water at 24.0oC. If the final temperature of the water is 34.2oC, what is the specific heat of the metal? Heat Lost by the Metal Once the metal is dropped into the water, its temperature will drop until it reaches the same temperature as the water (34.2oC). temperature drop of the metal ΔT = 125.0oC – 34.2oC = 90.8oC

49. The heat lost or gained by the system is given by: (m) (cp) (ΔT) = q cp = A sample of a metal with a mass of 212 g is heated to 125.0oC and then dropped into 375 g of water at 24.0oC. If the final temperature of the water is 34.2oC, what is the specific heat of the metal?

50. Conservation of Energy An energy transformation occurswhenever a chemical change occurs. • If energy is absorbed during a chemical change, the products will have more chemical potential energy than the reactants. • If energy is given off in a chemical change, the products will have less chemical potential energy than the reactants.