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Matter and Energy

Aim: What is Matter? Do Now: What is the difference between a heterogeneous and homogeneous mixture? Homework: Castle Learning Due Monday. Matter and Energy. Unit IV. I. Chemistry :. Definition : the branch of science that deals with the study of matter, energy, and their interaction.

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Matter and Energy

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  1. Aim: What is Matter? Do Now: What is the difference between a heterogeneous and homogeneous mixture?Homework: Castle Learning Due Monday

  2. Matter and Energy Unit IV

  3. I. Chemistry: • Definition: the branch of science that deals with the study of matter, energy, and their interaction. a)Matter: defined as anything that has massand occupies volume. • Mass – is the measurement of the amount of matter in an object OR is the quantity of matter an object has. • The fundamental unit for mass is the gram (g). • The instrument used in chemistry to measure the mass of an object is the triple beam balance or the digital scale.

  4. Figure 1: Triple Beam Balance From: http://www.kirkwood.k12.mo.us/parent_student/khs/BartinJ/sci%20skills%20book/using_a_3beam_bal.jpg

  5. Volume: defined as the amount of space an object occupies. • The fundamental unit for volume is the liter (l). Other units used to measure the volume of substances include the cubic centimeter (cm3) or the (cc). • The instrument used in chemistry to measure the volume of a substance is the graduated cylinder. Fig 2: Graduated Cylinder showing the meniscus From: http://www.electrickiva.com/quiz/exam03/meniscus.gif

  6. b) The States of Matter: • Gases, liquids and solids are all made up of microscopic particles, but the behaviors of these particles differ in the three phases. 1. SOLID: state of matter that has BOTH A DEFINITE VOLUME and a DEFINITE SHAPE.  Fig 3: Arrangement of Particles in a Solid From: http://www.ul.ie/~walshem/fyp/solid.gif

  7. 2. LIQUID – state of matter that has DEFINITE VOLUME but NO DEFINITE SHAPE.  A key property of a liquid is that they FLOW and can be POURED. Fig. 4: Arrangement of Particles in a Liquid From:http://www.ul.ie/~walshem/fyp/states%20of%20matter.htm#liquid

  8. 3. GAS – state of matter that has NO DEFINITE VOLUME and NO DEFINITE SHAPE.  • A Gas ALWAYS TAKES BOTH THE VOLUME AND THE SHAPE OF ANY CONTAINER INTO WHICH IT IS PLACED.  If a gas is NOT in a container, it will spread out as far as it can. Fig. 5: Arrangement of Particles in a Gas From :http://www.ul.ie/~walshem/fyp/gas.gif

  9. Ex: #1 Under the same conditions of temperature and pressure, a liquid differs from a gas because the particles of the liquid: a) are in constant straight-line motion b) take the shape of the container they occupy c) have no regular arrangement d) have stronger forces of attraction between them

  10. II. The Classification of Matter • a) Substances: Elements and Compounds • 1. ELEMENTS ARE PURE (homogeneous) SUBSTANCES THAT CANNOT BE BROKEN DOWN (decomposed) CHEMICALLY INTO SIMPLER KINDS OF MATTER. • More than 100 elements have been identified, though Fewer than 30 are Important in Living Things.  • All of the Elements are arranged on a Chart known as THE PERIODIC TABLE.  • Among the information provided in The Periodic Table are the ATOMIC NUMBER, THE CHEMICAL SYMBOL, AND THE ATOMIC MASS FOR EACH ELEMENT.

  11. More than 90 Percent of the Mass of living things is composed of JUST FOUR ELEMENTS:  OXYGEN, O, CARBON, C, HYDROGEN, H, AND NITROGEN, N. • Each Element has different Chemical Symbol which consist of One or Two Letters. Figure 12: Atomic Mass, Number, and Chemical Symbol

  12. 2. Compounds: Substances that consist of two or more elements that are combined chemically by bonds. • The elements in a compound can only be separated chemically by breaking the bonds that hold them together. • Compounds have fixed ratios of their components. Water will always have 2 hydrogen atoms for every one oxygen atom. • Compounds are homogeneous – one cannot distinguish between the components of the compound. If given a sample of water, you could not determine what is hydrogen and what is oxygen. Examples of Compounds water carbon dioxide glucose

  13. Two or more substances that are combined physically; the components can be easily separated. b) Mixtures • In homogeneous mixtures, the substances are completely mixed. This means that you cannot see the individual components. The mixture appears to be only one substance. • In heterogeneous mixtures, the substances are not completely mixed. This means that you can see the individual components. The mixture appears to be only two different substances in the same container. Figure 14: Heterogeneous Mixture: SOIL Figure 13: Homogeneous Mixture: Salt Water

  14. III. Energy: a) Definition: The capacity for doing work. • b) Energy has several forms including: • potential energy- the energy stored in the chemical bonds that exist between particles of matter. • kinetic energy – the energy of motion. Temperature is a direct measure of the average kinetic energy of particles. • c) Law of Conservation of Mass and Energy: • This scientific law states that neither mass nor energy can be created or destroyed. They can only be converted from one form into another. • Units: Joule (J) or Kilojoules (KJ), calorie (cal) or Kilocalorie (Kcal).

  15. d) Types of Energy • Light Energy: the energy associated with light waves and other forms of electromagnetic radiation. • Electrical Energy: the energy associated with electrical current (flow of electrons). • Chemical energy: the energy associated with chemical changes (the breaking and reformation of chemical bonds). • Heat Energy: the energy associated with the temperature of substances. Heat energy is the least useful form of energy. • Mechanical Energy: the energy associated with doing work. • Atomic and Nuclear Energy: the energy associated with changes in the mass of atoms and the energy that binds atoms together.

  16. Ex: #5 A compound differs from a mixture in that a compound always has a a) homogeneous composition b) maximum of two components c) minimum of three components d) heterogeneous composition Ex: #6 Which substance cannot be decomposed into simpler substances? a) ammonia (NH3) b) aluminum (Al) c) methane (CH4) d) methanol (CH3OH)

  17. e) Energy and Chemical Change • All chemical reactions require energy to occur. • Chemical bonds are forces of attraction that hold atoms, elements, compounds, and molecules together. • Chemical bonds store energy. • For a reaction to occur, a chemical must absorb enough energy to break the bonds that hold it’s atoms together. • The amount of energy that must be absorbed by a chemical to begin a reaction is called the Activation Energy (Ea).

  18. IV. Free Energy and Spontaneity a) Spontaneous Reactions: chemical reactions that occur without the addition of an outside source of energy. b) Non-spontaneous Reactions : chemical reactions that require an external energy source to occur.

  19. V. Types of Reactions • Exothermic Reactions • Energy releasing processes, ones that "generate" energy, are termed exothermic reactions. Figure 6 Ex: The burning (oxidation or combustion) of gasoline.

  20. Figure 7: Exothermic Reaction

  21. b) Endothermic Reactions: • Reactions that require energy to initiate the reaction are known as endothermic reactions. Figure 8: Endothermic Reaction Ex: The melting of ice.

  22. NOTE: • Enthalpy (H) indicates whether a reaction is exothermic or endothermic. • If DH = (+), the reaction is endothermic and energy is absorbed. • If DH =(-), the reaction is exothermic and energy is released.

  23. Note: • All natural processes tend to proceed in such a direction that the disorder or randomness of the universe increases. • Endergonic = Endothermic • Exergonic = Exothermic

  24. Label the following diagram and indicate what each arrow stands for. D B E A F C

  25. VI. Measuring Energy a) Temperature • Temperature is a measure of the average kinetic energy of matter. • The greater the average kinetic energy, the greater the velocity of the particles of matter, the greater the temperature (and vice versa). • The instrument used to measure the temperature (average kinetic energy) of matter is the thermometer.

  26. 1. Thermometer: • A thin glass, capillary tube that contains a fluid (mercury (Hg)) that when heated will expand. • When the fluid expands, it rises. This correlates to an increase in temperature. 2. Temperature Scales: • There are several scales that are used to represent the temperature of substances. They include: Fahrenheit, Celsius, and Kelvin. • Scientists most frequently use the Celsius and Kelvin scales.

  27. 212° 100 ° 373 Boiling Point 32 ° 0 ° 273 Freezing Point 0 Absolute Zero Fahrenheit Celsius Kelvin Figure 9: Temperature Scales and the Boiling/Freezing Points of Water at Standard Conditions

  28. Ex: #2 As ice cools from 273 K to 263 K, the average kinetic energy of its molecules will: a) decrease b) increase c) remain the same

  29. Conversion Formulas: C = Celsius K = Kelvin 1. C = K – 273 2. K = C + 273 Ex: #3 What Kelvin temperature is equal to 25°C? a) 248 K b) 298 K c) 100 K d) 200 K

  30. b) Heat Energy • The device that is used to measure the amount of heat energy that a substance contains is called a calorimeter. Figure 10: Calorimeter

  31. Substances store energy in their chemical bonds. • When the chemical bonds are broken, heat energy is released. • The heat energy that is released by the substance should theoretically be equally to the chemical energy stored in the chemical bonds of that substance. • A calorimeter indirectly determines the amount of heat energy stored by a substance by measuring the change in temperature for a known quantity of water which immerses that substance.

  32. 1. Determination of the Heat Energy Transfer • To determine the amount of heat energy gained/lost by a substance, several variables must be known. These include: Q = heat energy lost/gained (Joules) m = mass (grams) c = specific heat capacity (J/g. C o) D T = change in Temp Cwater = 4.2 J/g. C o Reference Table B Q = mc D T Reference Table T Note: Specific heat capacity indicates the ease at which a substance absorbs or releases heat energy. It is a constant value for any given substance.

  33. Figure 11: Calorimeter

  34. Ex: #4 What is the total number of joules of heat energy absorbed by 15 grams of water when it is heated from 30°C to 40°C? a) 10 b) 63 c) 150 d) 630

  35. 2. Determination of the Amount of Energy to Required for a Phase Change a) Heat of Fusion – the amount of energy required to convert a solid to a liquid at its melting point and standard pressure. Q = heat energy lost/gained (Joules) m = mass (grams) Hf = Heat of Fusion (for water - reference table B) Q = mHf 334 J/g a) Heat of Vaporization – the amount of energy required to convert a liquid to a gas at its boiling point and standard pressure. Q = mHv Q = heat energy lost/gained (Joules) m = mass (grams) Hv = Heat of Vaporization (for water - reference table B) 2260 J/g

  36. Calculate the following thermochemistry problem. 1. Find the initial temperature of a 120g sample of water that has reached a final temperature of 90ºC by absorbing 2520J of energy. NOTE: Cwater = 4.2 J/g. C o • A 10 g sample of water at -20º C is constantly heated till it reaches a final temperature of 110º C. Find the amount of heat energy needed for this change to occur given the following constants: • C water vapor = 1.996J/g. C° • C water = 4.2 J/g. C° • C ice = 2.108 J/g. C° • Hf =334 J/g • Hv = 2260 J/g

  37. VII. Heating/Cooling Curves • Heating/Cooling curves shows the change in kinetic and potential energies of substances. In addition, the curves indicates the points of phase change. a) Kinetic Energy: the energy associated with the velocity of the particles of matter. Temperature is a direct measure of the average kinetic energy of matter. As temperature increases, so does the average kinetic energy of matter (and vice versa). • Note: During phase changes, the KE of matter remains the same. b) Potential Energy: the stored energy found within substances. The potential energy of a substance increases as it is converted to a phase of matter having greater entropy (and vice versa). • Note: When KE changes, PE remains the same.

  38. Figure 15: The Heating Curve of Water

  39. Explanation of Figure 15 Point State Kinetic Energy Potential Energy A Solid Increases RTS B (s) to (l) melting RTS Increases C Liquid Increases RTS D (l ) to (g) boiling RTS Increases E Gas Increases RTS RTS = remains the same

  40. c) Phase Changes • Substances must either gain or lose energy for a change of phase to occur. Increase in Enthalpy (and Entropy)Process Solid to liquid Melting Liquid to gas Boiling (vaporization) Solid to gas Sublimation Decrease in Enthalpy (and Entropy)Process Gas to Liquid Condensation Liquid to solid Freezing Gas to solid Deposition

  41. d) Heats of Fusion and Vaporization: • Heat of Fusion: the amount of energy needed to be gained by a substance to convert it from a solid to a liquid. • This is an intensive property. • Example: Water - Hf = 334 J/g • Heat of Vaporization: the amount of energy needed to be gained by a substance to convert it from a liquid to a gas. • This is an intensive property. • Example: Water - Hv = 2260 J/g

  42. e) Heats of Fusion/Vaporization: Calculations Q = mHf Q = mHv Q = heat energy in joules m = mass in grams Hf = Heat of Fusion Hv = Heat of Vaporization

  43. Ex: # 6 How many joules of heat are absorbed when 70.0 grams of water is completely vaporized at its boiling point? a) 23, 352 J b) 7, 000 J c) 15, 813 J d) 158, 130 J

  44. Gram Molecular (Formula) Mass 1 mole 22.4 liters 6.02 x 1023 particles VIII. Avogadro’s Hypothesis • Under the same conditions of temperature and pressure, equal volumes of all gases contain the same number of particles. • For example, 1 liter of hydrogen gas will contain the same number of particles as 1 liter of oxygen gas.

  45. IX. The Unique Properties of Gases a) Gases are greatly influenced by changes in temperature and pressure. b) The behavior of gases is due in part to the fact that the atoms/molecules that they are comprised of are greatly dispersed.

  46. c) The Kinetic-Molecular Theory of Gases (Ideal Gases) • 1. Gases are made from molecules that are in constant random motion. • 2. Collisions between gas particles are completely elastic. • 3. The volume of individual gas molecules is insignificant as compare to the overall volume of space that the gas occupies. 4. Individual gas molecules do not have attractive forces for each other. 5. The average kinetic energy of a gas is directly proportional to its temperature.

  47. d) Exceptions to Ideal Gases 1. As opposed to an ideal gas, real gas molecules do have small but significant volumes as well as attractive forces. 2. These deviations become apparent under the conditions of low temperature and high pressure.

  48. e) The Gas Laws • Boyle’s Law: an indirect relationship; states that at a constant temperature the volume of a gas decreases with increasing pressure. • P1V1 = P2V2 KEY: P1 = initial pressure P2 = final pressure V1 = initial volume V2 = final volume Figure 16: Boyle’s Law From: http://chemmovies.unl.edu/chemistry/smallscale/SSGifs/26Fig2.gif

  49. Example #1: Divers get “the bends” if they come up too fast because gas in their blood expands, forming bubbles in their blood. If a diver has 0.05 L of gas in his blood under a pressure of 250 atm, then rises instantaneously to a depth where his blood has a pressure of 50.0 atm, what will the volume of gas in his blood be? Do you think this will harm the diver? Example #2: Part of the reason that conventional explosives cause so much damage is that their detonation produces a strong shock wave that can knock things down. While using explosives to knock down a building, the shock wave can be so strong that 12 liters of gas will reach a pressure of 3.8 x 104 mm Hg. When the shock wave passes and the gas returns to a pressure of 760 mm Hg, what will the volume of that gas be?

  50. 2. Charles Law: a direct relationship; states that at a constant pressure, the volume of a gas will increase with a corresponding increase in the temperature (Kelvin). V2 T2 V1 = T1 KEY V1 = initial volume T2 = final temperature V2 = final volume T1 = initial temperture Figure 17: Charles Law From: http://wine1.sb.fsu.edu/chm1045/notes/Gases/GasLaw/charles.gif

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