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Chapter 3

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  1. Chapter 3 Matter & Energy

  2. Matter • Matter is defined as anything that has mass and takes up space • We can describe matter using • Volume (may change under different conditions) • Weight (changes with gravity) • Mass (not affected by temperature, pressure, or gravity)

  3. Matter • Matter can be classified into 3 groups: • 1) Elements • 2) Compounds • 3) Mixtures

  4. Elements • Elements are substances that cannot be broken down by chemical change • These substances can be found on the Periodic Table • There are at least 115 known elements on the Periodic Table

  5. Compounds • Compounds are made of 2 or more elements that are chemically combined • The elements in a compound are combined in definite proportions by mass • H20 (2 hydrogens bond to 1 oxygen) • NaCl (1 sodium bonds to 1 chlorine) • More common than pure elements • (elements tend to be very chemically reactive and combine with other elements to form compounds)

  6. Compounds • The chemical and physical properties of a compound are different from the chemical and physical properties of the elements that make up the compound • Can be formed from simpler substances by chemical change, and they can be broken down into simpler substances by chemical change • ***Elements and compounds are considered to be pure substances**

  7. Mixtures • Mixtures are made of 2 or more substances that keep their own properties while combined • A mixture can form when: • An element is mixed with one or more other elements • A compound is mixed with one or more other compounds • One or more elements are mixed with one or more compounds

  8. Mixtures • Mixtures are different than elements or compounds • Mixtures have the properties of its constituents (thethings that make it up), while elements/compounds have one set of properties • The composition of a mixture varies while the composition of an element/compound is fixed

  9. Mixtures • Mixtures can be homogenous or heterogeneous • Homogenous mixtureshave uniform characteristics throughout • i.e. a sample from one part of the mixture is the same as a sample from another part • Solutions • Kool-Aid • Heterogeneous mixtures have different compositions • i.e. a sample from one part of the mixture has a different composition than a sample from another part • Suspensions • M&M’s, Munchies http://blog.austinkids.org/2011/03/17/st-pattys-day-experiment/

  10. Matter Pure Substances Mixtures Elements Compounds Homogeneous Heterogeneous

  11. PROPERTIES OF MATTER • Properties are a set of characteristics that identify a substance • Extensive properties- depends on the amount of a substance (Volume, weight, mass) • Intensive properties- depends on the nature of a substance • 1. Used to identify substances • 2. Boiling point, melting point, malleability, crystalline shape • 3. Density- the mass of a substance that occupies one unit of volume (the amount of “stuff” crammed into something) • The ratio of mass to volume • Density of a liquid or solid will change slightly with changes in temperature and pressure (changes in volume are too slight to be noticeable) • Changes in the densities of gases can be very large

  12. PROPERTIES OF MATTER Metals have interesting properties that identify them as metals, but not necessarily a specific metal Ductile = the metal can be made into wires Malleable = the metal can be hammered into a shape Luster = shine Conduct electricity

  13. PHYSICAL PROPERTIES Extensive properties and intensive properties are considered to be physical properties Physical properties can be observed without the production of a new substance Examples: color, hardness, odor, taste, density, melting & boiling points, electrical conductivity, and the other properties of metals discusses previously A physical change is a change that occurs that DOES NOT alter the identity of the substances AND can be reversed without a chemical reaction Evidence of a physical change: phase changes (melting, boiling, etc) grinding changing its shape magnetizing a substance

  14. CHEMICAL PROPERTIES A chemical property describes how a substance reacts, or fails to react, with other substances to produce a new substance A chemical change results in a new substance being formed Chemical changes are only reversible by another chemical reaction (if they are reversible at all). Evidence of a chemical change: 1) dramatic color change (not necessarily reliable) 2) release of gas or bubbles (effervesces!!!) 3) energy changes (heat or cold) 4) precipitate formation (an insoluble solid “falls out” of a liquid) *** usually a chemical change will have 2 or more of these occurring

  15. Chemical Reactions Chemical reactions are written so that the reactants are on the left of the arrow and products are to the right of the reaction arrow 6CO2 + 6H2O C6H12O6 + 6O2 In a chemical reaction, atoms are not created or destroyed, they are rearranged!! That is why you can balance chemical reactions Reactants Products

  16. Conservation of Mass • The Law of Conservation of Mass states that… Matter is neither created nor destroyed THERE IS NO NEW MATTER During both physical and chemical changes, the total amount of matter remains the same (Nuclear reactions do not always follow this though) Butane + Oxygen Carbon Dioxide + Water 58 g + 208 g 176 g + 90 g 266 g 266 g

  17. Practice Problem If a 35.0 gram sample of IRG is made up of 29.6 grams of KIX and 5.4 grams of GIRF, how many grams of KIX would there be in 245.3 grams of IRG? Set this up as a factor-label method problem!!!

  18. Energy • Energy is defined as, “the capacity to do work” • The Law of Conservation of Energy states… Energy can neither be created nor destroyed -In chemistry we discuss Chemical Energy Chemical Energy is the energy that is involved with potential chemical changes (reactions) in chemical systems

  19. Chemical Energy • Potential Energy – stored energy • Examples: gravitational PE, nuclear, chemical (located in the bonds between atoms) • Kinetic Energy - the energy of movement • Examples: electromagnetic, mechanical (sound & earthquakes), heat (internal motion of particles) • Microscopic particles contain potential and kinetic energy as heat energy, chemical energy, electrical energy (radiant energy) and sound energy

  20. Chemical Energy • Every substance has a certain amount of potential energy based on the number and strength of the chemical bonds • All chemical reactions have energy changes because they involve forming or breaking bonds between atoms!!!

  21. Basic Units of Energy - joule The SI unit of energy is the joule (J) This unit was named after James P. Joule who determined that work (or energy) can be measured according to how much heat it could produce

  22. Basic Units of Energy- calorie A second unit of energy is the calorie calorie(cal) – (heat calorie) the amount of heat required to raise the temperature of 1 gram of water 1°C. (notice the lower case c!!!!) The nutritional or “food” Calorie (Cal) is 1000 calories (notice the capital C!!) 1 Cal = 1000 cal The kilocalorie (kcal) equals 1000 calories 1 calorie (cal) = 4.184 J ** a calorie is actually larger than a joule!!

  23. Measuring Energy Changes A calorimeter can be used to measure energy changes A calorimeter contains an inner container that contains water (this is where the reaction takes place) Energy changes from the reaction can be measured by the temperature change of the water (whether it absorbs or loses heat) The heat required to raise the temperature of 1 gram of a substance 1 degree Celsius is the substance’s specific heat The specific heat of water is 1 cal/g•C° (or 4.184 J/g•C°) http://olc.spsd.sk.ca/de/physics20/heat/calorimeter.htm

  24. Energy Reminder!! Because of the Law of Conservation of Energy, whenever energy changes occur in a chemical reaction, the total amount of the energy remains the same. In order for a reaction to get started, many chemical reactions require a specific amount of activation energy Activation energy is the energy added to a reaction to get it started (can be heat or electrical)

  25. Exothermic Reactions If the reactants have more potential energy than the products: The reaction is EXOTHERMIC Energy is lost (converted to heat energy and released) What is an example of an exothermic reaction you’ve experienced? • The PE of the reactants is converted to heat (KE) and released (feels warm) which leaves the products with lower PE than the reactants http://www.kentchemistry.com/links/Kinetics/PEDiagrams.htm

  26. Endothermic Reactions If the products have more potential energy than the reactants… The reaction is ENDOTHERMIC The potential energy increases Heat energy is absorbed and the reaction feels “cold” What is an example of an endothermic reaction that you’ve experienced? • Heat energy (KE) is absorbed (feels cold) and becomes stored PE in the products leaving the products with a higher PE than the reactants http://www.kentchemistry.com/links/Kinetics/PEDiagrams.htm

  27. Temperature Energy is usually measured by measuring temperature changes Heat and temperature are not the same thing!!! Heat refers to the exchange of thermal energy Temperature is a measure of the thermal energy of matter (atoms) • Temperature Scales • Fahrenheit (° F) – not used in Science • Celsius (° C) – metric standard • Kelvin (K) – used because negative Celsius temps can result in unrealistic answers (i.e. negative volume). 0 K = -273°C (absolute zero)

  28. Conversions between temp scales °C = K – 273 or 5/9(°F - 32) °F = (9/5 • °C) + 32 K = °C + 273°

  29. Specific Heat • Specific heat – the amount of heat energy (J) needed to raise the temperature of 1 g. of a substance by 1 °C • Many substances have different specific heat values called the specific heat constant (c) Units: • You can find these values in Table 3.4 on p.71 • The greater the specific heat constant (c), the more heat is needed to change the temperature of 1g by 1 °C J g • ° C

  30. Specific Heat Formula q = m • c • Δt • The quantity of heat released or absorbed depends on 3 things: • 1. mass or quantity of substance • 2. what substance is (c) • 3. how much the temp changes q = heat absorbed or released (J) m = mass of substance (g) c = specific heat constant (p.71) Δt = temperature change (° C) J g • ° C

  31. Specific Heat Calculation • How much heat would have to be added to 5000.0 g of water to change its temp from 20.0 °C to 80.0 °C? q = m • c • Δt

  32. Specific Heat Calculation • How much heat would have to be added to 5000.0 g of water to change its temp from 20.0 °C to 80.0 °C? J g • ° C q = m • c • Δt q = ?? m = 5000.0 g C = 4.184 Δt = 80.0 °C – 20.0 °C = 60.0 °C q = (5000.0g)(4.184 )(60.0 °C ) q = 1,255,200 J q = 1,260,000 J J g • ° C

  33. Naming Phase Changes Review Solid  Liquid = Melting Liquid  Solid = Freezing Liquid  Gas = Boiling (or evaporation) Gas  Liquid = Condensing Solid  Gas = Sublimation Gas  Solid = Deposition