UNIT 6 PHYSICAL BEHAVIOR OF MATTER I CLASSIFICATION
UNIT 6: PHYSICAL BEHAVIOR OF MATTER
I. CLASSIFICATION OF MATTER Substance - Definite Composition (Homogenous) Physically Separable Mixture of Substances
I. CLASSIFICATION OF MATTER Element Compound (Fe, K, Ca, Two or more Ne) different elements Chemicallybonded Separable Individual atoms Ionic (Metal and Nonmetal) Molecular (covalent bond) (Nonmetals)
CHECKS FOR UNDERSTANDING 1. A) B) C) D) A compound differs from an element in that a compound Is homogeneous Has a definite composition Has a definite melting point Can be decomposed by a chemical reaction 2. Which of the following substances cannot be separated by chemical change? A) Nitrogen (g) B) Sodium chloride (s) C) Carbon dioxide (g) D) Magnesium Sulfate (aq)
I. CLASSIFICATION OF MATTER Homogenous Uniform throughout (air, tap water, solutions) Heterogeneous Nonuniform; distinct phases
CHECK FOR UNDERSTANDING 1. A pure substance that is composed only of identical atoms is classified as a A) A compound B) An element C) A heterogeneous mixture D) A homogeneous mixture 3. A) B) C) D) A heterogeneous material may be An element A compound A mixture Pure substance
II. SEPARATING MATTER • Certain types of matter can be separated using various methods. • Monatomic Elements - _______ be CANNOT decomposed (broken apart) using PHYSICAL CHEMICAL _______ or _______ means. • Diatomic Elements and Compounds (ie – O 2 CHEMICAL MEANS and H 2 O) – can be decomposed using _________ only
II. SEPARATING MATTER Mixtures – can be separated using MEANS PHYSICAL __________ Separation by particle size Filtration – Evaporation – Separation by boiling point Separation by polarity Chromatography – Separation by boiling point Distillation –
CHECK FOR UNDERSTANDING 1. Which of the substances could be decomposed by a chemical change? A) sodium B) aluminum C) magnesium D) ammonia 2. A sample of a material is passed through a filter paper. A white deposit remains on the paper, and It is a heterogeneous a clear liquid passesmixture through. The clear liquid is then evaporated, leaving a white residue. What can you determine about the nature of the sample? In a mixture the elements are not bonded with each are other and of can bedifferences physically between separated. 3. What some the a In a compound elements and can mixture of ironthe and oxygen are andbonded compound
THINK ABOUT THIS What happens to the spacing and speed of particles at each of the phases? SOLID LIQUID GAS
III. FORMS OF MECHANICAL ENERGY Kinetic Energy of movement (similar to temperature) (how fast atoms are moving) Potential Energy Stored energy (energy of position) More spread out (gas) = High PE Closer together (solid) = Low PE
IV. HEATING AND COOLING CURVES (ANIMATION) Heating Curve: ______ - Energy is being ENDOTHERMIC ABSORBED ____ l g s l liquid gas s g solid Sublimation (video)- Solid changes directly to a gas Heating Curve Animation
IV. HEATING AND COOLING CURVES A B Kinetic ↑ Energy Potentia Conl Energy stant Phase solid B C C D Constant ↑ ↑ Constant s l liquid melting D E Constant ↑ l g boiling E F ↑ Constant gas
CHECK FOR UNDERSTANDING 1. A substance begins to a melt. What happens to the potential and kinetic energy? PE increase, KE stays the same 2. The temperature of a substance refers to what type of energy? Kinetic energy 3. How does the speed and space of water molecules compare when in a liquid phase to a gas phase Molecules move faster and more spread out in gas phase
IV. HEATING AND COOLING CURVES Cooling Curve: ______ - Energy. RELEASED is being EXOTHERMIC ____ gas g l liquid l s g s solid Deposition - Gas changes directly to a solid
IV. HEATING AND COOLING CURVES A B Kinetic ↓ Energy Potentia Conl Energy stant Phase gas B C C D Constant ↓ ↓ Constant g l condensing D E Constant ↓ liquid l s Freezing E F ↓ Constant solid
CHECK FOR UNDERSTANDING 1. As a substance condenses, what happens to its potential and kinetic energy? PE decreases, KE stays the same 2. What phase is a substance in when it has its highest kinetic energy? gas 3. How does the speed and space of water molecules compare when in a liquid phase to a solid phase Molecules move slower and are closer together in solid phase
V. TEMPERATURE VS. HEAT Amount of energy Average kinetic energy transferred from one of its particles (how substance to fast they’re moving) another Joules (J) or Celsius (o. C) or Kelvin (K) Calories (cal) (K = o. C + 273) 1 cal = 4. 18 J T q
V. TEMPERATURE HEAT Temperature Scales (See Ref. VS. Tabs. ): K = o. C + 273 K = Kelvin Celsius o. C = degrees Convert: 200 degrees Celsius to Kelvin o o K = C + 273 K = 200 C + 273 = 473 K Law of Conservation of Energy: Energy (heat) cannot be created or destroyed. Energy (heat) can be TRANSFERRED. Heat Transfer: HEAT ALWAYS MOVES FROM WARMER OBJECTS TO COLDER OBJECTS
VI. MEASUREMENT OF HEAT ENERGY • The amount of heat given off or absorbed in a reaction can be calculated using the following equation: (See Ref. T Tabs. on Table ______ ) q = m c ∆T • Specific Heat: q = heat (J) m = mass (g) c = specific heat (J/g*o. C) ∆T = change in temperature (o. C) The amount of heat it takes to raise the temperature of 1 g of a substance 1 o. C 4. 18 _____ (Found B on Table • Specific Heat for water: _______ in Ref. Specific heat. Tabs) for. J/g*K concrete is 0. 84 J/(g. K) – why concrete is much hotter than water on a sunny
CHECK FOR UNDERSTANDING 1. You wake up in the morning and your barefoot touches the ceramic floor and it feels cold. Explain which way heat is being transferred. Heat moves from your body (warm) to the floor (cold) 2. You are cooking pasta in a boiling metal pot of water. You grab the metal handles with your bare hands (ouch!). Explain which way heat is being transferred. Heat moves from metal handles (warm) to your hands (cold). 3. Why do you feel cold after you get out of a hot shower. (link)
VI. MEASUREMENT OF HEAT ENERGY Example: How many joules are absorbed when 50. 0 g of water are hater from 30. 2 o. C to 58. 6 o. C? m = 50. 0 g Ti = 30. 2 o. C Tf = 58. 6 o. C q=? q = m c ∆T q = (50. 0 g) (4. 18 J/go. C ) (58. 6 o. C – 30. 2 q = 5935. 6 J 5940 J
VI. MEASUREMENT OF HEAT ENERGY Example: How many joules of heat energy are released when 50. 0 g of water are cooled from 70. 0 o. C to 60. 0 o. C? q = m c ∆T m = 50. 0 g q = (50. 0 g) (4. 18 J/go. C ) (60. 0 o. C Ti = 30. 2 o. C Tf = 58. 6 o. C – 70. 0 o. C) q = - 2090 J ( - means heat is q=? released) Example: 50. 0 g of water goes from 289. 6 K to 309. 6 K. A) Is heat energy released or absorbed? B) Calculate the amount energy. m = 50. 0 g q = m c ∆T Ti = 289. 6 K o. C ) (36. 6 o. C q = (50. 0 g) (4. 18 J/g o 16. 6 C – 16. 6 o. C ) Tf = 309. 6 K q = 4180 J o 36. 6 C
VII. HEAT OF FUSION • Heat of Fusion: Amount of heat absorbed (endothermic) to change a substance from s to l at its melting point 334 J/g B on • Heat of Fusion for water: ______ (Found Table ____ in Ref. Tabs) • Equation: (Found on. TTable _______ in Ref. Tabs) q = m. Hf
VII. HEAT OF FUSION Example: How many joules are required to melt 255 g of ice at 0. 00 o. C? m = 255 g Hf = 334 J/g q=? q = m. Hf q = (255 g) (334 J/g) q = 85, 170 J 85, 200 J Example: What is the total number of joules of heat needed to change 150 g of ice to water at 0. 00 o. C? q = 50, 100 J 5. 0 x 10 4 J or 50. k. J
VIII. HEAT OF VAPORIZATION • Heat of Vaporization: Amount of heat absorbed (endothermic) to change a substance from l to g at its boiling point • Heat of Vaporization for 2260 water: J/g ______B(Found on Table ____ in Ref. Tabs) • Equation: (Found on. TTable _______ in Ref. Tabs) q = m. Hv
VIII. HEAT OF VAPORIZATION Example: How many joules of energy are required to vaporize 423 g water at 100 o. C and 1 atm? q = m. Hv m = 423 g q = (423 g) (2260 J/g) Hv = 2260 J/g q=? q = 955, 980 J 956, 000 J Example: What is the total number of joules required to completely boil 125 g of water at 100 o. C at 1 atmosphere? q = 282, 500 J 283, 000 J
IX. CALORIMETRY • Used to: - Measure the amount of heat given off in a reaction. - Use q = m c ∆T to find the amount of heat lost or gained in a sample
UNIT 6: PART B BEHAVIOR OF GASES
X. ENDOTHERMIC AND • Energy is either absorbed or released in EXOTHERMIC (REVISITED) chemical reactions • Remember: - Breaking bonds is __________ ENDOTHERMIC - Heat is ______ the reaction from the ENTERING surroundings - Ex) heat + Br 2 Br + Br EXOTHERMIC - Creating bonds is _________ EXITING - Heat is ______ the reaction from the surroundings - Ex) N+ N N 2+ energy
X. ENDOTHERMIC AND • Where does the heat come from (or go to)? EXOTHERMIC (REVISITED) ___________ The surroundings • For exothermic reactions, heat (energy) leaves the reaction and moves into the _____________. Therefore, warmer makingsurroundings the surrounding temperature _________ • Exothermic Chemical Equation: Reactant(s) Product(s) + HEAT
X. ENDOTHERMIC AND (REVISITED) • For. EXOTHERMIC endothermic reactions, heat (energy) leaves the surroundings and moves Into the reaction _____________. Therefore, colder making the surrounding temperature _________ • Endothermic. Reactant(s) Chemical Equation: + HEAT Product(s) ENDOTHERMIC REACTION
X. ENDOTHERMIC AND EXOTHERMIC (REVISITED) DISORDER Entropy: Degree of __________ in a system Entropy Increases, if Substance goes from a s l g 1) _________________________ More moles of gas are produced 2) _________________________ Large molecule breaks up 3) _________________________
X. ENDOTHERMIC AND • Enthalpy: It is a measure of the heat released or absorbed in a EXOTHERMIC (REVISITED) reaction. If heat is released then the reaction is an exothermic reaction, heat will be on the products side and the ΔH will be negative. If heat is absorbed then the reaction is an endothermic reaction, heat will appear on the reactants side and the ΔH will be positive. (see Table ____ on Reference Tables). I Examples: Tell whether each of the following reactions are endothermic or exothermic (you may have to use table I). Then tell whether the temperature of the surroundings increases or decreases as a result. Entropy Endo/Exothermic ↑ _______Surroundings Exo (-∆H) warmer ___________C 6 H 12 O 6 + 6 O 2 6 CO 2 + 6 H 2 O +
X. ENDOTHERMIC AND EXOTHERMIC (REVISITED) Entropy Endo/Exothermic _______Surroundings 1. __________C 6 H 12 O 6 (s) + 6 O 2(g) 6 CO 2 (g) + 6 H 2 O (l)+ heat ______ 2. _________ 2 H 2 O + 484 k. J 2 H 2 (g) + O 2 (g) exo(-∆H) warmer ↑ _____ endo(+∆H) 3. __________ N 2 (g) + 3 H 2 (g) 2 NH 3 colder ↑ (g) warmer ↓ ____exo (-∆H) exo(-∆H) ↑ 4. ___________ Na. OH(s) Na+ (aq) warmer + OH- (aq) warmer SKIP _____exo(-∆H) ↑ 5. ___________ C 2 H 5 OH(l) + 3 O 2 (g) warmer 2 CO 2 (g) colder endo(+∆H) + 3 H _____ ↓ 2 O(l) 6. ____________2 KCl. O 3(s) 2 KCl(s) + O 2(g) + heat ____ 7. _______H+(aq) + C 2 H 3 O 2(aq) + heat HC 2 H 3 O 2(l) ______
XI. KINETIC MOLECULAR THEORY Ideal Gas: A set of gases that • DO NOT ACTUALLY EXIST a ________________, but represents standard that we can use to predict the behavior of gases. • Kinetic Molecular Theory of Ideal Gases tell how ideal gases behave: Particles are small and spread far apart 1. Travel in constant, random, straight line motion Can not 2. happen Have no attractive forces (IMF’s) in 3. real-world Don’t lose energy when they collide, but they can 4. transfer energy in the collision Move faster when hotter HIGH 5. • However, Real gases behave ideally under LOW
XII. VAPOR PRESSURE • Vapor Pressure: (VP) Pressure a liquid “feels” pushing it to evaporate (turn to gas) • ______ VP = _____________ HIGH Evaporates easily Example: Ethanol has higher VP than H O, so it evaporates more easily 2 • Factors for Vapor Pressure Strength of intermolecular force: molecules held together by dipole 1. dipole (polar) have lower VP than Van der Waals (nonpolar) (H 2 O does not evaporate as fast as methane (CH 4)) 2. Temperature/Pressure: increase temp. , increase VP
• Vapor of Four Liquids (See Table _____) XII. Pressure VAPOR PRESSURE H • Questions: k. Pa (101. 3 k. Pa = 1 atm = 760 mm. Hg = 760 torr. ) 1. Measured What isin: the vapor pressure of propanone at 35 o. C? _______ 48 k. Pa 2. What temperature does water boil at if the pressure is 70 90 degrees C k. Pa? ______ 117 degrees C 3. What is the normal boiling point of ethanoic acid? _____ Occurs because solids have very weak IMF (usually Van de Waals). Sublimation: They go directly from s g and have HIGH VP
CHECK FOR UNDERSTANDING 1. How could you change the boiling point temperature of a substance without adding anything to the substance? Change your altitude (higher altitude (lower pressure), lower boiling point) 2. You spill a glass of water on the floor. How could you get the water to evaporate faster, without using a mop or something to soak it up? Increase the temperature of the room. Spread out the puddle (increase surface area).
XIII. EVAPORATION VS. BOILING POINT Converting l g at the surface of a liquid that is NOT BOILING (leaving clothing out to dry) 1. Substance (strong IMF, slower evaporation) When the vapor pressure of liquid = air pressure Temperature when a liquid boils at standard pressure 2. Temperature (high temp. , faster evaporation) (101. 3 k. Pa or 1 atm = sea level) 3. Surface area (increase, faster evaporation) Water = 100 o. C Water Boiling at 70. 0 o. C
CHECK FORCO, UNDERSTANDING You move to Denver, where the altitude is 5, 183 feet above sea level (mile high city – approximately). 1. What would this high altitude due to the spacing of air particles? The air pressure would be lower, so the particles would be more spread out. 2. Based on your answer to question 1, what would this do to the vapor pressure of water? The vapor pressure would be less (less pressure pushing on the water – approximately 90 k. Pa in Denver) 3. Based on your answer to question 2, what would this do to the boiling point temperature of water in Boiling temperature would be lower (95 degrees Celsius) Denver?
THINK ABOUT THIS Expanding Marshmallows 1. Why did this happen to the marshmallow? Decreased the pressure inside and the marshmallow expanded 2. What is the relationship between pressure and volume? As pressure decreases, volume increases 3. What remained constant in this video? temperature
XIV. GAS LAWS A. Pressure and Volume (Boyle’s Law) decreases • As pressureincreases ________, volume temperature ________, and _______ is constant Volume inverse Relationship: Pressure Equation: P 1 V 1 = P 2 V 2 Question: If the pressure on a gas is halved, under constant temperature, what will happen to the volume of the gas? The volume will double
THINK ABOUT THIS You blow up a balloon and it expands each time you exhale into the balloon. 1. What happens to the number of gas particles you increaseseach exhale? put in the balloon 2. What happens to the pressure on the balloon? It increases (more gas particles pushing against the wall of the balloon) 3. What is the relationship between gas particles and pressure More gas particles, more pressure
XIV. GAS LAWS B. Pressure and Number of Gas Particles increases • As number of gas particles ________, increases pressure ________ Pressure Direct Relationship: # of Gas Particles Question: If the number of particles of a gas is tripled, under constant volume, what will happen to the pressure on the gas? The pressure will triple
THINK ABOUT THIS You blow up a balloon inside the school and bring it outside (it is – 10 o. C out). 1. What would you see happen to the balloon as you stand The outside? balloon would get smaller 2. Why would this occur? The cold air would cool the gas inside, bringing the molecules closer together 3. What would you see happen once you brought it back inside? Why? It would expand. Warm air would warm the gas inside, causes the molecules to spread out (volume increases).
XIV. GAS LAWS C. Temperature and Volume (Charles’s Law) increases • As temperature________, volume pressure ________ and _______ is constant Volume Direct Relationship: Temperature Equation: V 1 = V 2 T 1 T 2 Question: If the temperature of a gas is halved, under constant pressure, what will happen to the volume of the gas? The volume will be halved
PRESSURE AND TEMPERATURE OF A GAS On cold days the sensor in my car says that my tires have lower than normal pressure. 1. Why would this happen? The cold air makes the gas inside the tire slow down causing the tire pressure to be less 2. What two properties and being compared? Temperature and pressure 3. What is the relationship between these two properties As temperature decreases, pressure decreases Video Demonstration
XIV. GAS LAWS D. Temperature and Pressure (Gay-Lussac’s Law) • As temperature________, increases pressure________ and _______ is volume constant Direct Pressure Relationship: Equation: P 1 = P 2 T 1 T 2 Temperature Question: If the temperature of a gas is tripled, under constant volume, what will happen to the pressure of the gas? The pressure will triple
XV. COMBINED GAS LAW The relationships among pressure, temperature, and volume can be mathematically represented by an equation known as the combined gas law. • T • Equation (See Table _____) P 1 V 1 = P 2 V 2 T 1 T 2 • Important notes when using the equation: 1. Temperature must be in Kelvin, NOT Celsius 2. For pressure and volume, make sure the units of the initial and final are the same 3. If a variable remains constant, you can take it out of the equation
XV. COMBINED GAS LAW 1. The pressure of a gas at 200. K is increased from 200. k. Pa to 305. k. Pa at a constant 305 volume. What is the new temperature? K 2. If a gas at 8. 00 atm is cooled from 600. K to 150. K in a rigid container, 2. 00 what atm is the final pressure?
XV. COMBINED GAS LAW 3. If I have 2. 90 L of gas at a pressure of 5. 00 atm and a temperature of 320. K, what will be the temperature of the gas if I decrease the volume of the gas to 2. 40 L and decrease the 159 to. K 3. 00 atm? pressure 4. If I initially have a gas at a pressure of 12 atm, a volume of 23 liters, and a temperature of 250 K, and then I raise the pressure to 14 atm and increase 28 L the temperature to 350 K, what is the new volume of the gas?
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