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Introductory Chemistry , 3 rd Edition Nivaldo Tro

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  1. Introductory Chemistry, 3rd EditionNivaldo Tro Chapter 3 Matter and Energy Roy Kennedy Massachusetts Bay Community College Wellesley Hills, MA 2009, Prentice Hall

  2. Are matter & energy related • Matter is any particle with mass and volume • Energy is simply matter that is moving • 0 Kelvin is defined as the temperature when matter does not moving • So temperature is related to moving mass • Therefore: temperature and mass are related to energy • That’s why any chemistry or physics equation with energy must relate mass and temperature. Tro's "Introductory Chemistry", Chapter 3

  3. Around you • Everything you can see, touch, smell or taste in your room is made of matter. Tro's "Introductory Chemistry", Chapter 3

  4. What Is Matter? • Matter is anything with mass. • Typically, we think of tiny little pieces of mass as atoms and molecules because those 117 elements behave Newtonian. There are over 200 smaller particles that behave Quantunian. Tro's "Introductory Chemistry", Chapter 3

  5. Energy: it’s just Mass and Velocity • 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.

  6. 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. Tro's "Introductory Chemistry", Chapter 3

  7. Any matter can exist in one of 3 States • Solid • Liquid • Gas Tro's "Introductory Chemistry", Chapter 3

  8. Structure Determines Properties • The atoms or molecules have different structures in solids, liquids, and gases − leading to different properties.

  9. Solids • The particles in a solid are packed close together and are fixed in position. • Although they may vibrate. • The close packing of the particles results in solids being incompressible. • The inability of the particles to move around results in solids retaining their shape and volume when placed in a new container and prevents the particles from flowing. Tro's "Introductory Chemistry", Chapter 3

  10. Solids, Continued • Some solids have their particles arranged in an orderly geometric pattern—we call these crystalline solids. • Salt and diamonds. • Other solids have particles that do not show a regular geometric pattern over a long range—we call these amorphous solids. • Plastic and glass. Tro's "Introductory Chemistry", Chapter 3

  11. Liquids • The particles in a liquid are closely packed, but they have some ability to move around. • The close packing results in liquids being incompressible. • The ability of the particles to move allows liquids to take the shape of their container and to flow. However, they don’t have enough freedom to escape and expand to fill the container. Tro's "Introductory Chemistry", Chapter 3

  12. Gases • In the gas state, the particles have complete freedom from each other. • The particles are constantly flying around, bumping into each other and the container. • In the gas state, there is a lot of empty space between the particles. • On average. Tro's "Introductory Chemistry", Chapter 3

  13. Gases, Continued • Because there is a lot of empty space, the particles can be squeezed closer together. Therefore, gases are compressible. • Because the particles are not held in close contact and are moving freely, gases expand to fill and take the shape of their container, and will flow. Tro's "Introductory Chemistry", Chapter 3

  14. Matter Pure Substance Mixture Constant Composition Variable Composition Homogeneous Matter: is it pure or impure Heterogeneous • Pure Substance = All samples are made of the same pieces in the same percentages. • Salt • Mixtures= Different samples may have the same pieces in different percentages. • Salt water Tro's "Introductory Chemistry", Chapter 3

  15. Mixtures Heterogeneous Homogeneous 1. Made of multiple substances, whose presence can be seen. 2. Portions of a sample have different composition and properties. 1. Made of multiple substances, but appears to be one substance. 2. All portions of a sample have the same composition and properties. Tro's "Introductory Chemistry", Chapter 3

  16. Matter Summary

  17. Matter has Properties • Physical Properties are the characteristics of matter that can be changed without changing its composition. • Characteristics that are directly observable. • Chemical Properties are the characteristics that determine how the composition of matter changes as a result of contact with other matter or the influence of energy. • Characteristics that describe the behavior of matter. Tro's "Introductory Chemistry", Chapter 3

  18. H2O Physical verses H2O Chemical Chapter One

  19. Tro's "Introductory Chemistry", Chapter 3

  20. Some Physical Properties of Iron • Iron is a silvery solid at room temperature with a metallic taste and smooth texture. • Iron melts at 1538 °C and boils at 4428 °C. • Iron’s density is 7.87 g/cm3. • Iron can be magnetized. • Iron conducts electricity, but not as well as most other common metals. • Iron’s ductility and thermal conductivity are about average for a metal. • It requires 0.45 J of heat energy to raise the temperature of one gram of iron by 1°C. Tro's "Introductory Chemistry", Chapter 3

  21. Tro's "Introductory Chemistry", Chapter 3

  22. Some Chemical Properties of Iron • Iron is easily oxidized in moist air to form rust. • When iron is added to hydrochloric acid, it produces a solution of ferric chloride and hydrogen gas. • Iron is more reactive than silver, but less reactive than magnesium. Tro's "Introductory Chemistry", Chapter 3

  23. Quiz: is it a Physical or Chemical Property • Salt is a white, granular solid = physical. • Salt melts at 801 °C = physical. • Salt is stable at room temperature, it does not decompose = chemical. • 36 g of salt will dissolve in 100 g of water = physical. • When a clear, colorless solution of silver nitrate is added to a salt solution, a white solid forms = chemical. Tro's "Introductory Chemistry", Chapter 3

  24. Matter has Properties, Matter can also go through Changes • Changes that alter the state or appearance of the matter without altering the composition are called physical changes. • Changes that alter the composition of the matter are called chemical changes. • During the chemical change, the atoms that are present rearrange into new molecules, but all of the original atoms are still present. Tro's "Introductory Chemistry", Chapter 3

  25. Is it a Physical or Chemical Change? • A physical change results in a different form of the same substance. • The kinds of molecules don’t change. • A chemical change results in one or more completely new substances. • Also called chemical reactions. • The new substances have different molecules than the original substances. • You will observe different physical properties because the new substances have their own physical properties. Tro's "Introductory Chemistry", Chapter 3

  26. Phase Changes ArePhysical Changes • Boiling = liquid to gas. • Melting = solid to liquid. • Subliming = solid to gas. • Freezing = liquid to solid. • Condensing = gas to liquid. • Deposition = gas to solid. • State changes require heating or cooling the substance. • Evaporation is not a simple phase change, it is a solution process. Tro's "Introductory Chemistry", Chapter 3

  27. Quiz: is it a Physical or Chemical change • Evaporation of rubbing alcohol = physical. • Sugar turning black when heated = chemical. • An egg splitting open and spilling out = physical. • Sugar fermenting into alcohol = chemical. • Bubbles escaping from soda = physical. • Bubbles that form when hydrogen peroxide is mixed with blood = chemical. Tro's "Introductory Chemistry", Chapter 3

  28. 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. Tro's "Introductory Chemistry", Chapter 3

  29. Distillation: different boiling points Tro's "Introductory Chemistry", Chapter 3

  30. Filtration: different solubility's

  31. Summary • Moving Matter has Energy. Motion is related to temperature. All energy formulas are relations between mass and temperature • Matter has 3 states • Matter has properties • Matter can change States/Properties/Change are all related to temperature and how much you have Tro's "Introductory Chemistry", Chapter 3

  32. 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. • butane + oxygen  carbon dioxide + water 58 grams + 208 grams  176 grams + 90 grams 266 grams = 266 grams Tro's "Introductory Chemistry", Chapter 3

  33. Law of Conservation of Energy • “Energy can neither be created nor destroyed.” • The total amount of energy in the universe is constant. There is no process that can increase or decrease that amount. • Note: neither Mass nor Energy are ever destroyed Tro's "Introductory Chemistry", Chapter 3

  34. Energy • The Fundamental Principle of the Universe is Energy • From the Greeks to Newton to Quantum Mechanics Energy is known as the capacity to do work and is simply calculated by knowing the mass and velocity of a particle. • The harder you swing an ax the faster you can fall a tree. • Guess what happens when you walk into a wall .005 mph or 500 mph Tro's "Introductory Chemistry", Chapter 3

  35. Energy: it’s just Mass and Velocity • 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.

  36. To get Energy (electrical, thermal, light, nuclear, chemical) • You take slow moving particles and make them move faster As slow moving water falls, gravity pulls it faster. The water falls on top of a turbine, which moves a coil in a magnet to generate electricity.

  37. To get Energy (electrical, thermal, light, nuclear, chemical) • You take slow moving particles and make them move faster

  38. To get Energy (electrical, thermal, light, nuclear, chemical) • You take slow moving particles and make them move faster Binding energy is simply the amount of energy (and mass) released, when free nucleons join to form a nucleus; a gluon is released or absorbed Einstein's mass-energy equivalence formula E = mc² can be used to compute the binding energy

  39. Kinds of EnergyKinetic and Potential • Potential energy is energy that is stored; slow moving • Water flows because gravity pulls it downstream. • However, the dam won’t allow it to move, so it has to store that energy. • Kinetic energy is energy of motion, or energy that is being transferred from one object to another; fast moving. • When the water flows over the dam, some of its potential energy is converted to kinetic energy of motion. Tro's "Introductory Chemistry", Chapter 3

  40. There’s No Such Thing as a Free Ride When atoms contact each other, frictions is produced. You will often notice friction as sound or heat. So instead of useful energy, “anti-energy” friction slows your car down. Tro's "Introductory Chemistry", Chapter 3

  41. 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. Tro's "Introductory Chemistry", Chapter 3

  42. Energy Use 2008 Tro's "Introductory Chemistry", Chapter 3

  43. Ex 3.5, A candy bar has 225 Cal, convert to Joules • Write down the Given quantity and its unit. Given: 225 Cal 3 sig figs • Write down the quantity you want to Find and unit. Find: ? J • Write down the appropriate Conversion Factors. Conversion Factors: 1 Cal = 1000 cal 1 cal = 4.184 J • Write a Solution Map. Solution Map: Cal cal J • Follow the solution map to Solve the problem. Solution: • Significant figures and round. Round: 225 Cal = 9.41 x 105 J 3 significant figures Units and magnitude are correct. • Check. Check:

  44. Chemical Potential Energy • The amount of energy stored in a material is its chemical potential energy. • The stored energy arises mainly from • the attachments between atoms in the molecules • the attractive forces between molecules. Tro's "Introductory Chemistry", Chapter 3

  45. Exothermic Processes • When a change results in the release of energy it is called an exothermic process. • An exothermic chemical reaction occurs when the reactants have more chemical potential energy than the products. • 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. Tro's "Introductory Chemistry", Chapter 3

  46. Surroundings reaction Reactants Amount of energy released Potential energy Products An Exothermic Reaction Tro's "Introductory Chemistry", Chapter 3

  47. Endothermic Processes • When a change requires the absorption of energy it is called an endothermic process. • An endothermic chemical reaction occurs when the products have more chemical potential energy than the reactants. • 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. Tro's "Introductory Chemistry", Chapter 3

  48. Surroundings reaction Products Amount of energy absorbed Potential energy Reactants An Endothermic Reaction Tro's "Introductory Chemistry", Chapter 3

  49. Temperature Scales 100°C 373 K 212°F 671 R Boiling point water 298 K 75°F 534 R Room temp 25°C 0°C 273 K 32°F 459 R Melting point ice -38.9°C 234.1 K -38°F 421 R Boiling point mercury -183°C 90 K -297°F 162 R Boiling point oxygen BP helium -269°C 4 K -452°F 7 R -273°C 0 K -459 °F 0 R Absolute zero Celsius Kelvin Fahrenheit Rankine

  50. 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. Tro's "Introductory Chemistry", Chapter 3