Matter & Energy - PowerPoint Presentation

Matter & Energy

States of Matter

UEQ: In dealing with matter, how are composition, structure, properties and energy related? LEQ: What are the differences between physical and chemical properties? Law of Conservation of Mass Mass cannot be created or destroyed in ordinary chemical and physical changes Therefore, the mass before the reaction will be equal to the mass after the reaction

States of Matter Solid Liquid Gas Holds Shape Fixed Volume Shape of Container Free Surface Fixed Volume Shape of Container Volume of Container heat heat

5 COPPER PHASE - SOLID

6 LIQUID PHASE - COPPER

7 GAS PHASE - COPPER

Some Properties of Solids, Liquids, and Gases Property Solid Liquid Gas Shape Has definite shape Takes the shape of Takes the shape the container of its container Volume Has a definite volume Has a definite volume Fills the volume of the container Arrangement of Fixed, very close Random , close Random , far apart Particles Interactions between Very strong Strong Essentially none particles

9 WHAT MAKES A STATE OF MATTER? The state of matter is determined by the particles that make up the matter and their attraction to each other. solids have a large attraction and are packed tightly together the attraction in liquids is not as great so particles flow gases have no real attraction so particles spread out Matter changes from one state to another when its energy changes (either loses or gains energy). when energy (heat) is added to ice it changes state to a liquid more energy changes the liquid water to water vapor KINETIC THEORY OF MATTER: ALL MATTER IS MADE UP OF PARTICLES IN CONSTANT RANDOM MOTION

10 State of matter Shape Def Shape? (Y/N) Volume Def volume (Y/N) Particle Attraction Particle Arrangement Solid YES YES VERY STRONG CLOSE, VIBRATE IN PLACE Liquid NO YES STRONG CLOSE, SLIDE PAST EACH OTHER Gas NO NO WEAK FREE TO MOVE ABOUT YOU NEED TO KNOW UNDERSTAND THE INFORMATION IN THIS CHART!!!

Properties of Matter Electrical ConductivityHeat Conductivity Density – mass/volume Melting Point Boiling Point Malleability - sheets Ductility - wires

UEQ: In dealing with matter, how are composition, structure, properties and energy related? LEQ: What are the differences between physical and chemical properties? Extensive properties Intensive properties Depends on the amount of matter present Volume, mass, amount of energy in a substance Does not depend on amount of matter present Melting point, boiling point, density, ability to conduct electricity and transfer energy as heat

UEQ: In dealing with matter, how are composition, structure, properties and energy related? LEQ: What are the differences between physical and chemical properties? Physical properties Chemical properties Characteristic observed or measured without changing substance Melting point, boiling point, color, size, ductility, malleable, density, ability to conduct heat or electricity Ability to change from one substance to a new substance Ability to produce a gas, burn, rust Example: Copper w/ moist air

UEQ: In dealing with matter, how are composition, structure, properties and energy related? LEQ: What are the differences between physical and chemical properties? Physical changes Chemical changes Change that does not change the identity or chemical makeup of the substance Cutting, melting, drawing into wire, crushing, temperature and pressure changes Substance changes into new substance b/c chemical bonds have been broken or made Occurs on molecular level Noticed by temperature change, smell/odor, bubbles (gas), rust formation Reactants products

15 CHANGES OF STATE Adding or taking away energy causes substances to change state. Most of these changes in state are familiar. Solid Liquid Gas Melt Evaporate Condense Freeze http://eanes.tx.schoolwebpages.com/education/components/docmgr/default.php?sectiondetailid=6107&fileitem=8730&catfilter=1594

16 CHANGES OF STATE Some changes in state are not so familiar. An important change in state is sublimation. Sublimation is the change of a substance as it goes from a solid directly to a gas. The most common example being dry ice. Dry ice is frozen carbon dioxide gas.

Phase Diagram

Energy Changes Accompanying Phase Changes Solid Liquid Gas Melting Freezing Deposition Condensation Vaporization Sublimation Energy of system Brown, LeMay, Bursten, Chemistry 2000, page 405

solid liquid gas Heat added Temperature ( o C) A B C D E Heating Curve for Water 0 100 LeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World , 1996, page 487

solid liquid gas vaporization condensation melting freezing Heat added Temperature ( o C) A B C D E Heating Curve for Water 0 100 LeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World , 1996, page 487

21 KEY POINTS: A, C, and E show temperature is changing As heat is added temperature changes just as expected. The state of matter does not change. The substance is gaining energy but does not have enough to change to a new state. The equation Q = mC Δ T is used alongside this graph. Q = heat, measured in Joules (J) m= mass, measured in grams (g) C= specific heat, measured in J/g- o C Δ T = change in temperature, measured in Celsius ( o C) We can measure each of these variables given enough information

22 A NOTE ON SPECIFIC HEAT Substance Melting point Boiling point Bolognium 20 °C 100 °C Unobtainium 40 °C 140 °C Foosium 70 °C 140 °C Description: One paper cup is filled with water; the other is empty. When the cups are placed above a lit Bunsen burner, the empty cup burns, whereas the water-filled one doesn't. This is because the water absorbs most of the heat of the flame due to its high specific heat. Specific Heat Capacity -  The ability of materials to "hold" a certain amount of thermal energy at a given temperature.   Consider a day at your favorite Carolina beach. On a typical summer day the sand is so hot it will burn your feet, yet the water feels much cooler and quite comfortable. The same holds true if you consider the parking lot. Each material experiences the same summer sun, but each heats up differently. Later that night the sand and parking lot would feel quite cool to the feet, and the water temperature will not have changed. Each material “holds” a certain amount of specific heat, and each will gain and lose that heat at different rates. The sand loses its heat really fast (low specific heat) the water loses heat slowly (high specific heat). The reverse is true, water must gain a lot of heat to change temperature, sand needs only a little to warm up quickly. Q = m x C x Δ T Not as bad as it looks…read on

23 Q = m x C x Δ T Determine the energy needed (in Joules) when 55.6 grams of water at 43.2 °C is heated to 78.1 °C. Q = ? m = 55.6 g Δ T = T f – T i (final temp – initial temp) 78.1 o – 43.2 o = 34.9 o C C = specific heat of water 4.18J (found on your reference table) Q = 55.6g x 4.18J x 34.9 o C = 8111 J of energy needed

24 TRY ANOTHER PROBLEM Determine the energy required when cooling 456.2 grams of water at 89.2 °C to a final temperature of 5.9 °C. What is mass? What is the change in temperature? What is the specific heat of water? (remember your reference tables Answer: Q = 456.2g x 4.18J x – 84.3 o C = -160753 J 456.2 g 5.9 – 89.2 = -83.3 o C 4.18J http://www.sciencegeek.net/Shockwave/SpecificHeat.htm Click on the link below to work some sample problems.

25 Heat is added but temperature is NOT changing. The flat lines at B and D represent where the substance is going through a change of state. At B it is melting/freezing, at D is boiling/condensing. The temperature does not change when the state of matter changes because the added energy is used to break the attraction holding the particles. NOTE AS MORE ENERGY/HEAT IS ADDED THE PARTICLES GET FURTHER APART. When no temperature change occurs the equations q=m H f or q = mH v are used to make calculations. H f and H v will be found on your reference sheets. B and D show an unchanging temperature RETURN TO THE GRAPH

26 REVIEW A solid has a ____________ shape and takes up a _____________ amount of space. A ______________ has no definite shape and no definite volume. A liquid has a ____________ shape and _____________ volume . During a change in state, as energy is added, temperature (does or does not ) change. The change of state when a substance goes from a solid directly to a gas is called _____________.

27 Use the heating curve below to answer the following questions. What is the melting point of the substance? What is the boiling point of the substance? Which letter represents heating of the solid? Which letter represents heating of the vapor? Which letter represents melting of the solid? Which letter represents boiling of the liquid?

What is Matter? Anything with mass and volume (takes up space)Matter can be broken into 2 categories: 1. Pure substances 2. Mixtures

Pure Substances Fixed, definite composition /uniform/unchanging Pure substances can be broken down further into two more categories: elements and compounds

Pure Substances Element – a substance that cannot be separated or broken down into simpler substances composed of identical atoms EX : copper wire, aluminum foil Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Pure Substances Compound composed of 2 or more elements in a fixed ratio Must break bonds to get elemental form properties differ from those of individual elements EX : table salt ( NaCl ) Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Pure Substances Law of Definite Composition A given compound always contains the same, fixed ratio of elements. Law of Multiple Proportions Elements can combine in different ratios to form different compounds. Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Pure Substances For example… Two different compounds, each has a definite composition. Carbon, C Oxygen, O Carbon monoxide, CO Carbon, C Oxygen, O Oxygen, O Carbon dioxide, CO 2 Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Mixtures Variable combination of two or more pure substances . Heterogeneous Homogeneous Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Mixtures Homogeneous Uniformly mixed (even) very small particles no Tyndall effect Also called a solution Tyndall Effect particles don’t settle EX : rubbing alcohol Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

UEQ: In dealing with matter, how are composition, structure, properties and energy related? LEQ: How is matter classified? solutions Solute – substance being dissolved (sugar) Solvent – substance doing the dissolving (water) RULE OF THUMB!!! USUALLY YOU WILL HAVE MORE SOLVENT THAN SOLUTE!!!What is solubility?Amount of solute that can be dissolvedVaries with temperatureBased on a saturated solution Concentration – high C – high solute/solvent, low C – low solute/solvent

WHICH IS WHICH 1 st ROUND Identify the solute and solvent in each of the following solutions. 1.0 g of sugar dissolved in 100 g of water. 50 mL of water mixed with 20 mL isopropyl alcohol A tincture of Iodine is prepared with 0.10 g I 2 and 10.0 mL of ethyl alcohol. 40 % ethanol a rubbing alcohol. Sugar is the smaller quantity that is dissolving. Therefore it the solute and water is the solvent. Since both water and isopropyl alcohol are liquids, the one smaller volume, is the solute and water is the solvent. Iodine is the solute and ethyl alcohol is the solvent. Ethanol is the solvent and water is the solute.

WHICH IS WHICH 2 nd ROUND Identify the solute and solvent in each of the following solutions. 10 g NaCl and 100 g of water. 50 mL ethanol and 1 0 mL H 2 O 2.0 L oxygen and 8.0 L nitrogen. 100 g silver and 40 g mercury. 100 mL H 2 O and 5.0 g sugar

KINDS OF SOLUTION Dilute small amount of solute dispersed in the solvent Concentrated large amount of solute is dissolved in the solvent

Types of Solutions Based on Solute Concentration… Hypotonic (lower solute concentration) to the solution Hypertonic (higher solute concentration) to the solution

Types of Solutions Based on Solute Concentration… Isotonic solutions are equal in their solute concentrations. We say that they are isotonic to each other.

VOCABULARY Immiscible Two liquids that are insoluble (Oil & Vinegar) Miscible Two liquids that are soluble in each other (Alcohol & Water) SOLVENT WATER KEROSENE ACETONE ALCOHOL Oil IMMISCIBLE MISCIBLE MISCIBLE S/MISCIBLE

Solubility Solids are more soluble at... high temperatures. Gases are more soluble at... low temperatures & high pressures (Henry’s Law). EX : nitrogen narcosis, the “bends,” soda

Solubility SATURATED SOLUTION no more solute dissolves UNSATURATED SOLUTION more solute dissolves SUPERSATURATED SOLUTION becomes unstable, crystals form increasing concentration

Solubility Table LeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World , 1996, page 517 0 10 20 30 40 50 60 70 80 90 100 Solubility vs. Temperature for Solids Solubility (grams of solute/100 g H 2 O) KI KCl 20 10 30 40 50 60 70 80 90 110 120 130 140 100 NaNO 3 KNO 3 HCl NH 4 Cl NH 3 NaCl KClO 3 SO 2 shows the dependence of solubility on temperature gases solids

Temp. ( o C) Solubility (g/100 g H 2 O) KNO 3 (s) KCl (s) HCl (g) SOLUBILITY CURVE Solubility  how much solute dissolves in a given amt. of solvent at a given temp. unsaturated: solution could hold more solute; below line saturated : solution has “just right” amt. of solute; on line supersaturated : solution has “too much” solute dissolved in it; above the line

T o Sol. T o Sol. Solids dissolved in liquids Gases dissolved in liquids As T o , solubility As T o , solubility

Electrolytes Timberlake, Chemistry 7 th Edition, page 290 Electrolytes - solutions that carry an electric current NaCl( aq ) Na + + Cl - HF( aq ) H + + F - strong electrolyte weak electrolyte nonelectrolyte

Electrolyte Imbalances Electrolyte Normal range (mmol / L) Excess Defiency Sodium Na + 135 - 145 Hypernatremia (increased urine excretion; excess water loss) Hyponatremia (dehydration; diabetes-related low blood pH; vomiting, diarrhea) Potassium K + 3.5 – 5.0 Hyperkalemia (renal failure, low blood pH) Hypokalemia (gastointestinal conditions) Hydrogen carbonate HCO 3 - 24 - 30 Hypercapina (high blood pH; hypoventilation) Hypocapnia (low blood pH; hyper-ventilation; dehydration) Chloride Cl - 100 - 106 Hyperchloremia (anemia, heart conditions, dehydration) Hypochloremia (acute infections; burns; hypoventilation)

Mixtures Heterogeneous Substances can be identified Not uniformly mixed Visible particles settle Also classified as suspensions Examples: granite, wood, blood

Mixtures Suspension heterogeneous large particles Tyndall effect – scattering of light by particles in a mixture particles settle EX : fresh-squeezed lemonade Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Mixtures Colloid Heterogeneous Looks like a solution b/c particles do not settle, but acts like a suspension b/c particles do not dissolve medium-sized particles Tyndall effect particles don’t settle EX : milk, fog, clay Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

MATTER Can it be physically separated? Homogeneous Mixture (solution) Heterogeneous Mixture Compound Element MIXTURE PURE SUBSTANCE yes no Can it be chemically decomposed? no yes Is the composition uniform? no yes Colloids Suspensions Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Matter Flowchart Examples: graphite pepper sugar (sucrose) paint soda Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem element hetero. mixture compound solution homo. mixture hetero. mixture

Mixtures Examples: mayonnaise muddy water fog saltwater Italian salad dressing Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem colloid suspension colloid solution suspension

Compounds vs. Mixtures Compounds have properties that are uniquely different from the elements from which they are made. A formula can always be written for a compound e.g. NaCl  Na + Cl 2 Mixtures retain their individual properties. e.g. Salt water is salty and wet

Mixture vs. Compound Mixture Fixed Composition Bonds between components Can ONLY be separated by chemical means Variable Composition No bonds between components Can be separated by physical means Alike Different Contain two or moreelements Can be separated intoelements Involve substances Compound Different Topic Topic

Methods of Separating Mixtures Magnet Filter Decant EvaporationCentrifuge Chromatography Distillation

Gases Unit Symbol Definition/ relationship Pascal Pa SI pressure unit Millimeters of mercury mm Hg Pressure that supports 1mm mercury column in a barometer torr torr 1 torr = 1 mm Hg Atmosphere atm Average atmospheric pressure at sea level and 0 C. 1 atm = 760 mm Hg = 760 torr = 1.01325 X 10^5 Pa Pounds per square inch psi 1 psi = 6.89286 x 10^3 Pa 1 atm = 14.700 psi

Kinetic –Molecular Theory of gases Ideal Gas: hypothetical gas that perfectly fits all the assumptions of the kinetic molecular theory Five Assumptions: Gases consist of objects with a defined mass and zero volume. The gas particles travel randomly in straight-line motion where their movement can be described by the fundamental laws of mechanics All collisions involving gas particles are elastic; the kinetic energy of the system is conserved even though the kinetic energy among the particles is redistributed; continuous, rapid, random motion. The gas particles do not interact with each other or with the walls of any container The gas phase system will have an average kinetic energy that is proportional to temperature; temperature depends on the kinetic energy of the particles.

Dalton’s Law of Partial Pressures The total pressure of a gas mixture is the sum of the partial pressures of the component gases Pt = P1 + P2 +P3 + … Example: Scuba Diving Standard conditions for a gas are 0° C and 1 atm.

Gas Laws Boyle’s Law P1V1=P2V2 Charle’s law V1/T1 = V2/T2 Guy- Lussac’s Law P1/T1 = P2/T2 Combined Gas Law P1V1/T1 = P2V2/T2

Examples Boyle’s Law A sample of oxygen gas has a volume of 150.0 mL when its pressure is 0.947 atm. What will the volume of the gas be at a pressure of 0.987 atm if the temperature remains constant? A gas occupies 12.3 liters at a pressure of 40.0 mm Hg. What is the volume when the pressure is increased to 60.0 mm Hg? If a gas at 25.0 °C occupies 3.60 liters at a pressure of 1.00 atm , what will be its volume at a pressure of 2.50 atm ? A balloon filled with helium gas has a volume of 500 mL at a pressure of 1 atm. The balloon is released and reaches an altitude of 6.5 km, where the pressure is 0.5 atm. If the temperature has remained the same, what volume does the gas occupy at this height?

Examples Charles’s Law A sample of neon gas occupies a volume of 752 mL at 25 C. What volume will the gas occupy at 50 C if the pressure remains constant? A sample of neon gas has a volume of 752 mL at 25 C. What will the volume at 100 C be if the pressure is constant? A sample of nitrogen gas is contained in a piston with a freely moving cylinder. At 0.0 C, the volume of the gas is 375mL. To what temperature must the gas be heated to occupy a volume of 500mL?

Examples Calculate the decrease in temperature when 2.00 L at 20.0 °C is compressed to 1.00 L. 600.0 mL of air is at 20.0 °C. What is the volume at 60.0 °C? A gas occupies 900.0 mL at a temperature of 27.0 °C. What is the volume at 132.0 °C? At 225.0 °C a gas has a volume of 400.0 mL. What is the volume of this gas at 127.0 °C? At 210.0 °C a gas has a volume of 8.00 L. What is the volume of this gas at -23.0 °C?

Examples The gas in a container is at a pressure of 3.00atm at 25 C. Directions on the container warn the user not to keep it in a place where the temperature exceeds 52 C. What would the gas pressure in the container be at 52 C? 3/298 = x/325, A 600 mL sample of nitrogen is heated from 27 °C to 77 °C at constant pressure. What is the final volume?

Gay-Lussac’s Law Determine the pressure change when a constant volume of gas at 1.00 atm is heated from 20.0 °C to 30.0 °C. A gas has a pressure of 0.370 atm at 50.0 °C. What is the pressure at standard temperature? A gas has a pressure of 699.0 mm Hg at 40.0 °C. What is the temperature at standard pressure?

Combined Gas law A gas has a volume of 800.0 mL at minus 23.00 °C and 300.0 torr . What would the volume of the gas be at 227.0 °C and 600.0 torr of pressure? 500.0 liters of a gas are prepared at 700.0 mm Hg and 200.0 °C. The gas is placed into a tank under high pressure. When the tank cools to 20.0 °C, the pressure of the gas is 30.0 atm. What is the volume of the gas? What is the final volume of a 400.0 mL gas sample that is subjected to a temperature change from 22.0 °C to 30.0 °C and a pressure change from 760.0 mm Hg to 360.0 mm Hg? What is the volume of gas at 2.00 atm and 200.0 K if its original volume was 300.0 L at 0.250 atm and 400.0 K.? A helium-filled balloon has a volume of 50 L at 25 C and 1.08 atm. What volume will it have at 0.855 atm and 10 C?

Avogadro’s Law Equal volumes of gases at the same temperature and pressure contain equal number of molecules V = kn 1 mol = 22.4 L STP = standard temperature and pressure temp = 0 C or 273 K Pressure = 1 atm Examples: What volume does 0.0685 mol of gas occupy at STP? What quantity of gas, in moles, is contained in 2.21 L STP?

Ideal Gas Law PV = nRT n= moles of substance R = ideal gas constant, 0.0821 L* atm /(mol*K) Mathematical relationship among pressure, volume, temperature and the number of moles of a gas Examples: What is the pressure in atmospheres exerted by a 0.500 mol sample of nitrogen gas in a 10.0 L container at 298 K? What pressure, in atms , is exerted by 0.325 mol of hydrogen gas in a 4.08 L container at 35 C?

Diffusion and Effusion Graham’s Law of Effusion States that the rates of effusion of gases at the same temperature and pressure are inversely proportional to the square roots of their molar masses Rate of effusion of A = Mb Rate of effusion of B Ma Rate of diffusion or effusion depends on the velocities of gas molecules Effusion – process where gas molecules under pressure pass through a tiny opening. Ex: punctured tire Diffusion – mixing of two gases by random molecular motion. Ex: ammonia, vinegar, or perfume

Energy

Copyright © 2007 Pearson Benjamin Cummings. All rights reserved. (a) Radiant energy (b) Thermal energy (c) Chemical energy (d) Nuclear energy (e) Electrical energy

Types of Energy Potenial Kinetic Chemical Electromagnetic Electrical Heat Mechanical Nuclear

The energy something possesses due to its motion, depending on mass and velocity. Potential energy Energy in Energy out kinetic energy kinetic energy

Energy Kinetic Energy – energy of motion KE = ½ m v 2 Potential Energy – stored energy Batteries (chemical potential energy) Spring in a watch (mechanical potential energy) Water trapped above a dam (gravitational potential energy) mass velocity (speed) B A C

Exothermic Reaction Reactants  Products + Energy 10 energy = 8 energy + 2 energy Reactants Products - D H Energy Energy of reactants Energy of products Reaction Progress

Endothermic Reaction Energy + Reactants  Products + D H Endothermic Reaction progress Energy Reactants Products Activation Energy

Effect of Catalyst on Reaction Rate reactants products Energy activation energy for catalyzed reaction Reaction Progress No catalyst Catalyst lowers the activation energy for the reaction. What is a catalyst? What does it do during a chemical reaction?

Hess’s Law Calculate the enthalpy of formation of carbon dioxide from its elements. C(g) + 2O(g)  CO 2 (g) Use the following data: 2O(g)  O 2 (g) D H = - 250 kJC(s)  C(g) DH = +720 kJCO2 (g)  C(s) + O2(g) DH = +390 kJ Smith, Smoot, Himes, pg 141 2O(g)  O 2(g) DH = - 250 kJ C(g) + 2O(g)  CO2(g) D H = -1360 kJ C(g)  C(s) DH = - 720 kJ C(s) + O 2(g)  CO2(g) DH = - 390 kJ

Fission vs. Fusion Fuse small atoms 2H 2 He NO Radioactive waste Very High Temperatures ~5,000,000 oC(SUN) Split large atomsU-235 Radioactive waste (long half-life) NuclearPowerPlants Alike Different Create Large Amountsof EnergyE = mc2 Transmutation of ElementsOccurs Change Nucleusof Atoms Fusion Different Topic Topic Fission