Matter: Properties and Changes

Matter: Properties and Changes
Matter: Properties & Changes (Ch. 3/15.1/15.3/12.4) Matter: Properties and Changes Chapter 3 Chapter 15.1, 15.3 Chapter 12.4 Chapter 3.1 Vocab Physical property Extensive property Intensive property Chemical property States of matter Solid Liquid Gas Vapor Chapter 3.2 Vocab Physical Change Chemical Change Law of Conservation of Mass Phase Change Rev 4

Matter: Properties & Changes (Ch. 3/15.1/15.3/12.4)
Chapter Objectives Distinguish between physical and chemical changes Define and classify matter by composition Define properties of liquids, solids, and gasses Identify observable characteristics of a chemical reaction Explain and apply the fundamental law of conservation of mass Rev 4

Vocabulary – Ch. 3.1 – 3.2 Physical property Extensive property Intensive property Chemical property States of matter Solid Liquid Gas Vapor Physical Change Chemical Change Law of Conservation of Mass Phase Change Rev 4

Vocabulary – Chapter 15.1, 15.3 Energy Heat Joule Specific Heat Specific Heat Equation (q = mC ∆T) Heat of Vaporization (∆Hvap) Heat of Fusion (∆Hfus) Heating Curve Note: We are not going over all of the terms in these chapter sections, just the ones I’ve outlined here. Rev 4

Vocabulary – Ch. 12.1, 12.4 Pressure Barometer Atmosphere Melting Point Vaporization Evaporation Vapor Pressure Boiling Point Sublimation Freezing point Condensation Deposition Phase diagram Triple Point Rev 4

Ch What is Matter?? Matter as defined from Ch. 1 is: Anything that takes up space and has mass Mass – measure of the amount of matter in an object. Rev 4

What is a Pure Substance?
Matter: Properties & Changes (Ch. 3/15.1/15.3/12.4) What is a Pure Substance? Matter that has a uniform and unchanging composition Table salt is ALWAYS Sodium Chloride Water is always made of 2 Hydrogen atoms and 1 Oxygen atom Is seawater a substance? Rev 4

Properties of Matter Chemistry is the study of matter. Matter is classified according to its physical and chemical properties. Physical Properties: Can be observed without changing the identity of substance. Rev 4

Some Physical Properties
Matter: Properties & Changes (Ch. 3/15.1/15.3/12.4) Some Physical Properties State at Room temp (liquid, solid or gas) Melting point Boiling Point Viscosity (resistance to flow) Density Color Odor Brittle/ductile Electrical/thermal conductivity Rev 4

Physical Properties Intensive Properties – Independent of amount of material. They property is the same no matter how much is there. Example: Density Extensive Properties – Dependent of the amount of material present. Example: Length, mass, volume Rev 4

Properties of Matter Chemical Properties: the ability to combine or change into other substances. Examples: flammability, oxidation, rotting Rev 4

States of Matter State of Matter: Its physical form. There are three physical states: Solid: Definite shape Definite volume Closely packed particles Rev 4

States of Matter Liquid: particles move past each other (flow) definite volume takes the shape of its container (indefinite) Rev 4

States of Matter Gas: flows takes the shape of its container (indefinite shape) Fills the container completely. (indefinite volume) Note: A vapor refers to a gaseous state of a substance that is a solid or liquid at room temperature. Rev 4

Physical or Chemical Property?
Matter: Properties & Changes (Ch. 3/15.1/15.3/12.4) Physical or Chemical Property? Bending of aluminum Salt dissolving in water Magnesium burning in air Baking soda is a white powder Fluorine is a highly reactive element Rev 4

Ch Changes in Matter Physical changes are those which alter the substance without altering its composition. Change of phase  one physical state to another Melting of ice - composition unchanged, i.e. ice is water in solid form (H2O) They generally require energy, the ability to absorb or release heat (or work). This is the start of Ch. 3.2 Substance that change readily to a gas at room temperature are volatile. Examples: Gasoline, mothballs (naphthalene), alcohol. The work part you’ll discuss more in physics next year. Get some moth balls and do Quick Demo on p. 76 of TWE. This demo will help with the definitions on the next couple of slides. Rev 4

Phase Changes (Ch. 12.4) What are the phase changes of water? 1. Melting – changing of a solid to a liquid (heat of fusion = ∆Hf) 2. Vaporization – changing from a liquid to a gas (heat of vaporization = ∆Hvap) 3. Sublimation – Changing from a solid to a gas (heat of sublimation = ∆Hsub) What do these processes have in common? Rev 4

Phase Changes Phase changes in the opposite direction have names too. liquid to a solid: gas to a liquid: gas to a solid: What do these have in common? Answer: Rev 4

UNITS OF ENERGY 1 calorie = heat required to raise temp. of 1.00 g of H2O by oC. 1000 cal = 1 kilocalorie = 1 kcal 1 kcal = 1 Calorie (a food “calorie”) But we use the unit called the JOULE 1 cal = joules (exactly) James Joule Rev 4

Heats of Fusion & Vaporization
Matter: Properties & Changes (Ch. 3/15.1/15.3/12.4) Heats of Fusion & Vaporization Heat of Fusion (∆Hfus) – The amount of heat (in joules) needed to melt 1 g of substance. For ice: 334 J/g q (heat) = ∆Hfus*m (m= mass of ice/water) Heat of Vaporization (∆Hvap) – The amount of heat (in joules) needed to vaporize 1 g of substance For water: 2260 J/g q (heat) = ∆Hvap*m (m= mass of water/steam) Rev 4

Example Problems How much heat does it take to melt 20.5 g of ice at 0⁰C? q = 334 J/g * 20.5 = 6850 J (6.85 kJ) How much heat is released when 50.0 g of steam at 100 ⁰C condenses to water at 100 ⁰C? q = J/g * 50.0 g = -113,000 J (-113 kJ) Rev 4

Specific Heat Capacity
Matter: Properties & Changes (Ch. 3/15.1/15.3/12.4) Specific Heat Capacity Specific Heat Capacity – amount of heat (q) required to raise the temperature of one gram of a substance by 1 degree. C = J (energy gained or lost) mass (g) * Temp Change(⁰C) Rev 4

Heat Capacity Values Substance Spec. Heat (J/g•⁰C) Water Ethylene glycol 2.39 Al uminum glass Rev 4

Calculating Heat Gained or lost
The heat, q, gained or lost by a substance can be calculated by knowing the mass of the object, the temperature change, and the heat capacity. q = mC∆T

Calculations involving Heat
Example 1: A 5.00 g piece of aluminum is heated from 25.0⁰C to 99.5⁰C. How many joules of heat did it absorb? q = m * C * ∆T = 5.00 g * J/g*⁰C * 74.5⁰C = 334 J

Matter: Properties & Changes (Ch. 3/15.1/15.3/12.4) Calculations involving Heat Example 2: 10.2 g of cooking oil at 25.0 ⁰C is placed in a pan and 3.34 kJ of heat is required to raise the temperature to ⁰C. What is the specific heat of the oil? q = m*C*∆T C = q/(m ∆T) C = 3340 J/(10.2 g * ( ) ⁰C) C = 1.91 J/g* ⁰C Rev 4

Matter: Properties & Changes (Ch. 3/15.1/15.3/12.4) Calculations involving Heat Important Points! q (heat) is a positive quantity. The sign (+ or -) refers to whether the system you’re looking gained it (+) or lost it (-). From the previous example, the oil would lose 3340 J of heat upon cooling back to 25.0 ⁰C. (-3340 J heat lost) Specific heat capacity is like a bucket. It is a measure of how much energy an object absorbs before the temperature changes. For the third point, think of an 1 kilo iron bar and 1 kg of water in the summer sun. Which is going to feel hotter in 30 minutes? The overflow of a bucket is like the change in temperature. See Assess New Knowledge on p. 520 for example of the third point. Initial: 5.00 g Pb at 85.0 C 5.00 g Al at 65.0 C Both placed in cold water. Temp at thermal equilibrium is 25.0 C qPb = 5.00 g * J/g*C * 60.0 C = 38.7 J qAl = 5.00 g * J/g*C * 40.0 C = 179 J So, even though Al started at a lower temperature and had a lower ∆T, it lost more heat because Al has a greater specific heat than lead. Rev 4

Heating Curve for Water
Matter: Properties & Changes (Ch. 3/15.1/15.3/12.4) Heating Curve for Water Note that T is constant as ice melts For a cooling curve, the graph would be inverted, representing heat lost. Rev 4

Heating/Cooling Curve for Water
Matter: Properties & Changes (Ch. 3/15.1/15.3/12.4) Heating/Cooling Curve for Water Note the flat portions of the curve where water is melting/boiling. Rev 4

Heat & Changes of State What quantity of heat is required to melt 500. g of ice (at 0oC) and heat the water to steam at 100oC? Heat of fusion of ice = 333 J/g Specific heat of water = 4.2 J/g•K Heat of vaporization = 2260 J/g +333 J/g +2260 J/g Rev 4

Heat & Changes of State What quantity of energy as heat is required to melt 500. g of ice (at 0⁰C) and heat the water to steam at 100 oC? 1. To melt ice at 0⁰C q = (500. g)(333 J/g) = x 105 J 2. To raise water from 0 oC to 100oC q = (500. g)(4.2 J/g•K)( )K = 2.1 x 105 J 3. To vaporize water at 100oC q = (500. g)(2260 J/g) = x 106 J 4. Total energy = 1.51 x 106 J = 1510 kJ Rev 4

Practice problem If we add 6050 J of heat to 54.2g of ice at ⁰C, what will it be at the end? What temperature will it be? The specific heat of ice is 2.03 J/g*⁰C. Rev 4

Pressure Pressure is the force acting on an object per unit area: Gravity exerts a force on the earth’s atmosphere A column of air 1 m2 in cross section exerts a force of about 105 N (101,300 N/m2). 1 Pascal (Pa) = 1 N/m2 . So, 101,300 N/m2 = 101,300 Pa or kPa. Since we are at the surface of the earth, we ‘feel’ 1 atmosphere of pressure. Rev 4

Barometer Vacuum A barometer measures atmospheric pressure The pressure of the atmosphere at sea level will hold a column of mercury 760 mm Hg. 1 atm = 760 mm Hg 760 mm Hg 1 atm Pressure Rev 4

Units of pressure 1 atmosphere (atm) = 760 mm Hg = 760 torr 1 atm = 101,300 Pascals = kPa Can make conversion factors from these. What is 724 mm Hg in atm ? What is 724 mm Hg in kPa Rev 4

Phase Changes Vapor pressure is the pressure exerted by a vapor over a liquid. The vapor pressure increases with increasing temperature. This is why water evaporates even though it’s not 212˚F. Evaporation is the vaporization of water only at the surface of a liquid. Only the surface molecules have enough energy to change phase. Talk about swamp coolers in Arizona. They put water vapor in the air. As it evaporates off your skin, it helps cool you off. Unfortunately, this only works when the humidity is low. So in August, when it’s relatively humid (in Tucson) it doesn’t help much. Rev 4

Phase Changes However, when the vapor pressure of the water is the same as the atmospheric pressure the water is … boiling. Rev 4

Phase Diagram A phase diagram is a graph of pressure vs temperature that shows in which phase a substance exists under different conditions of T & P. Rev 4

Solid Liquid Gas Phase Diagram for water Boiling Pressure 1 Atm Melting Condensation Freezing The link opens up the “States of Matter” program. Use that to show phase transitions and triple point. This is the phase diagram for water (liquid is more dense than ice – ice floats so slope of liquid-solid line slopes to the left) Talk about how liquids are really kind of rare in the universe. If you have some pressure but low temp, it’s a solid (comet). If you have high temp but not enough pressure, it’s a gas (like on Venus). Challenge question: How can you melt ice without raising the temperature? Answer: Increase the pressure (it pushes the water molecules closer together). Ice skaters really skate on a thin layer of water that’s caused by the pressure of the skate on ice that melts it. Sudha has a pressure chamber. Can show hot water (fill to cover bottom of chamber) boiling at lower temps once a vacuum is pulled over it. After this slide give out Matter Worksheet 3 – Pressure and Phase Diagrams. Sublimation Deposition Temperature Rev 4

Ch. 3.2 - Chemical Properties and Changes
Matter: Properties & Changes (Ch. 3/15.1/15.3/12.4) Ch Chemical Properties and Changes Chemical Properties – the ability to combine or change into other substances. Example: Water won’t react with aluminum, but reacts with sodium and potassium (violently) and iron (slowly but surely). A chemical property always relates to a chemical change – the changing of one or more substances into other substances. A chemical change is also known as a chemical reaction. Back to Chapter 3 Demos to do to illustrate chemical changes! (Better to do these after slide 54) Good place here to do Magic of Chemistry Demo (P. 12 of TWE) - Magic of Chem demo is the ‘wine’ (KMnO4) to water to ‘milk’ (white ppt). Atoms in Motion Demo (P. 76 of TWE), which turns pennies to brass. Burn magnesium to get MgO or Add 3 M HCl to Zn in a test tube and hold glowing splint over it for pop sound. Add Ca(s) to water and watch it fizz. Sugar and Sulfuric Acid - Put white sugar in a 250 mL flask with a little water and add concentrated H2SO4. Can also do the Quick Demo on p. 62 with the EtOH in a 5L Water bottle. Rev 4

Signs of a Chemical Change
Matter: Properties & Changes (Ch. 3/15.1/15.3/12.4) Signs of a Chemical Change Color change Odor Gas Formation of a precipitate (solid) Heat Rev 4

Chemical Properties and Changes
Matter: Properties & Changes (Ch. 3/15.1/15.3/12.4) Chemical Properties and Changes Chemical changes are a rearrangement of atoms in the substance. Chemical changes follow the Law of Conservation of Mass. This means that when there is a chemical change, matter is neither created or destroyed, just changed in form. Massreactants = Massproducts Rev 4

Conservation of Mass Example
Matter: Properties & Changes (Ch. 3/15.1/15.3/12.4) Conservation of Mass Example Reactants are the substances you start with. Started with sugar and sulfuric acid alcohol and air (O2) Products are the new substances that are made. - Carbon, water vapor, and heat - water & carbon dioxide Examples from demos. Sugar and sulfuric acid Burn magnesium in air to make MgO Zn in HCl, hold lighted splint over to bark it (Launch Lab p. 69) 10.0 g – 9.26 g of Hg liquid leaves 0.74 g of oxygen formed. Demos to do to illustrate chemical changes! Good place here to do Magic of Chemistry Demo (P. 12 of TWE) - Magic of Chem demo is the ‘wine’ (KMnO4) to water to ‘milk’ (white ppt). Atoms in Motion Demo (P. 76 of TWE), which turns pennies or copper pieces to brass. Burn magnesium to get MgO or Add 3 M HCl to Zn in a test tube and hold glowing splint over it for pop sound. Add Ca(s) to water and watch it fizz. Sugar and Sulfuric Acid - Put white sugar in a 250 mL flask with a little water and add concentrated H2SO4. Can also do the Quick Demo on p. 62 with the EtOH in a 5L Water bottle. Just use a small (~20 mL) of EtOH. If you use too much, there’s not enough oxygen in the bottle. Rev 4

Conservation of Mass Example
Matter: Properties & Changes (Ch. 3/15.1/15.3/12.4) Conservation of Mass Example Example: g of red mercury (II) oxide powder is placed in an open flask. It is heated until it has fully converted to liquid mercury and oxygen gas. Written as: Mercury oxide  mercury + oxygen The liquid mercury has a mass of 9.26 grams The mass of oxygen formed in this reaction is 10.0 g (total) – 9.26 g (mercury) = 0.74 g oxygen Rev 4

Practice 57.48 g of sodium reacts with chlorine gas to form g of sodium chloride. Sodium + Chlorine  sodium chloride How much chlorine gas was used in the reaction? Rev 4

Elements An element is a pure substance that cannot be broken down into simpler substances Each element has a symbol and they are arranged in a periodic table Rev 4

Elements Each element has it’s own symbol in the periodic table. Either a single capital letter: H is for hydrogen Or two letters, the first is capital the second is ALWAYS small: He is for helium. Co is different than CO. Rev 4

Adopt and Element (p. 86 of TWE) – I am doing this later on in year when they know electron cofigs. Work on format and rubric over the summer. See sheet that I have so far. Rev 4

Practice Name the element or symbol. C Calcium Cl Iodine K Mercury Rev 4

Compounds Compounds are pure substances that are combinations of elements in a fixed ratio. Example: Table salt is a combination of sodium and chlorine! NaCl (1:1 ratio) Rev 4

Mixtures Mixture – Combination of two or more pure substances that where each component retains its own chemical properties. Heterogeneous Mixture – individual substances remain distinct Homogeneous Mixture – Constant composition throughout. It has a single phase. Start here after What is Matter? Activity. Use this as a starting point to differentiate between mixtures and compounds. Click on link after going over the terms. The video will show distinction between mixtures and compounds. Rev 4

No Can it be separated by “Physical” methods? Yes Mixture Pure Substance Is it uniform? Can it be separated by “Chemical” methods? Yes No Yes No Homogen. Heterogen. Compound Element Rev 4

Separating Mixtures Substances in mixtures are physically combined, so they can be physically separated into component substances The separation technique you choose depends on the physical properties of the substances you want to separate This is part of Ch. 3.3 Rev 4

Size Differences Filtration – small liquid molecules will pass through porous barrier, while solid crystals will not (Filter paper) Screening – solids of one size are retained on screen while smaller solids pass through Filtration Demo/mini lab on p. 82 (Separation of dye using filter paper) Show crystallization by having a beaker of sugar water and a string dipped in it. The crystals will climb the string over time. This also works as evaporation as well since the sugar is moving along in the string with the water and then it evaporates. Rev 4

Difference in Boiling Point
Matter: Properties & Changes (Ch. 3/15.1/15.3/12.4) Difference in Boiling Point Distillation: based on difference of boiling points of liquids. Done in refineries refineries in US produce 9 Mil bbl/day gasoline 1.6 Mil bbl/day jet fuel 4.5 Mil bbl/day fuel oil 5 Mil bbl/day others Also pharmaceutical, basic chemicals etc. Others kerosene, asphalt, tars etc. Rev 4

History of the Periodic Table
Matter: Properties & Changes (Ch. 3/15.1/15.3/12.4) History of the Periodic Table Dmitri Mendeleev ( ) first began systematic organization of elements by their chemical behavior and physical properties in 1869. Let’s take a look at our modern periodic table—show groups and periods and explain similarities Rev 4

History of the Periodic Table
Matter: Properties & Changes (Ch. 3/15.1/15.3/12.4) History of the Periodic Table Based on the mass of each element Mendeleev found that behavior was periodic – that is properties within a group repeat or are “periodic” with each row or period. Hence the name “Periodic Table” Let’s take a look at our modern periodic table—show groups and periods and explain similarities Rev 4

Law of Definite Proportions
Matter: Properties & Changes (Ch. 3/15.1/15.3/12.4) Law of Definite Proportions All samples of a pure compound contain the same elements in the same proportion by mass (Joseph Proust, in 1799). Example: 10 grams CaBr2 500 grams CaBr2 Joseph Proust ( ) was a French chemist. Antoine Lavoisier ( ) was also a French chemist at that time doing similar experiments outlining the fundamentals of chemistry. Quick Demo on p. 88 to show Definite Proportions (Need sucrose & lactic acid) Rev 4

Example for CaBr2 Element Mass Analysis, g % by Mass Ca 2.0051 Br 7.9949 10.000 100.00 Rev 4

Law of Definite Proportions
Matter: Properties & Changes (Ch. 3/15.1/15.3/12.4) Law of Definite Proportions 18 g of water (H2O) has 2 g of hydrogen and 16 g of oxygen. 44 g of Carbon Dioxide (CO2) has 12 g of carbon and 32 g of oxygen. 28 g Carbon Monoxide (CO) has 12 g of carbon and 16 g of oxygen. Rev 4

Law of Multiple Proportions
Matter: Properties & Changes (Ch. 3/15.1/15.3/12.4) Law of Multiple Proportions When different compounds are formed by a combination of the same elements, different masses of one element combine with the same relative mass of the other element in a ratio of small whole numbers (John Dalton, 1803) Examples: CO and CO2 H2O and H2O CuCl and CuCl2 Rev 4

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