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Introduction to Chemistry and Matter and Energy

Introduction to Chemistry and Matter and Energy. Summer’s over Hang tight It’s going to be an exciting ride!. What is Chemistry?. What is Matter? What is Non-Matter?. Why Study Chemistry?. Central, fundamental science. Other sciences used chemistry as their backbone.

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Introduction to Chemistry and Matter and Energy

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  1. Introduction to Chemistryand Matter and Energy Summer’s over Hang tight It’s going to be an exciting ride!

  2. What is Chemistry? What is Matter? What is Non-Matter?

  3. Why Study Chemistry? • Central, fundamental science. • Other sciences used chemistry as their backbone. • Health care, conservation of natural resources, protection of the environment, food production, clothing, manufacturing, production of shelter, etc…

  4. Scientific laws are the evidence used to support a conclusion.  Scientific hypotheses and theories are our best attempts at explaining the behavior of the world, in ways that can be tested by further experiment.  We don't prove theories (and hypotheses) true.  We just use the observations to convince ourselves (and others) that we have a good idea.  Scientists have a lot of confidence in scientific theories, because they know there is a lot of evidence to back them up. Scientific law: a generalized description, usually expressed in mathematical terms, which describes the empirical behavior of matter. Scientific laws describe things.  They do not explain them.

  5. Measurement and Scientific Notation • Measurement define qualitative properties of a substance. • Often in science, measurements require very large or very small numbers. • Scientific notation = a number between 1 and 10 multiplied by 10 raised to a power. • The number of places the decimal point has moved determines the power of 10. If the decimal point has moved to the ______then the power is _______, to the _____, ___________. e.g. 602,000,000,000,000,000,000,000.0 = e.g. 0.00524 =

  6. Convert to scientific notation 24500 356 0.000985 0.222 12200 Convert to non-scientific notation numbers 4.2 X 10-3 2.15 X 104 3.14 X 10-6 9.22 X 105 9.57 X 102 Convert the following:

  7. Mathematics of SciencePrecision, Accuracy, and Significant Figures • No measurement of a physical quantity is absolutely certain. • All measurements include a certain degree of uncertainty Causes of uncertainty:

  8. Precision= Accuracy =

  9. Consider three sets of data that have been recorded after measuring a piece of wood that was exactly 6.000 m long. • Which set of data is the most accurate? • Which set of data is the most precise?

  10. Significant figures- measurements include one uncertain figure in addition to those known with certainty. • Rules for Significant Figures 1.    All digits 1-9 are significant I.e.- 129 2.  Zeros between sig. Figs. are always significant I.e.- 5007 3.  Trailing zeros in a number are significant only if the number contains a decimal pt. I.e.- 1000.0 100 4.  Zeros in the beginning of a number whose only function is to place the decimal point are not significant. I.e.- 0.0025 5.  Zeros following a decimal sig fig are significant. I.e.- 0.000470 6.  A bar over a zero indicates significance I.e.- 6400

  11. Atlantic – Pacific Rule • If a decimal is present, count from the Pacific side. • If a decimal is absent, count from the Atlantic side. Start counting from the first nonzero digit you find, and count every digit including zero thereafter!

  12. Determine the number of significant figures in the following 250.7 0.00077 1024 4.7 X 10-5 3400000 500.0 0.230970 0.03400 0.34030 26 Calculate the following to the correct number of sig. figs. 34.5 X 23.46 123/3 2.61X10-1 X 356 21.78 + 45.86 23.88887-11.2 6-3.0 32.559 X 34.555 4433-1187 1.2 X 4.3 8.08 + 21.98 Significant Figures Practice

  13. Rules for Calculations Using Significant Figures Multiplication and Division- limit and round to the least number of sig figs in any of the factors. I.e.- 144.6 X .0023 = ? Addition and Subtraction Rule- limit and round to least number of decimal places in any of the numbers that make up the problem. I.e.- 5.42 g + 131.1 g = ?

  14. SI Units- preferred metric units used in science.

  15. Metric Conversion

  16. Unit Conversion Using Dimensional Analysis • Write the term to be converted- both the number and the unit. 0.0342g • Write the conversion formulas 1 g = 1000 mg • Make a fraction of the conversion formula such that a. if the unit in step 1 is in the numerator, the same unit in step 3 must be in the denominator b. if the unit in step 1 is in the denominator, the same unit in step 3 must be in the numerator. Note: since the numerator and the denominator are equal, the fraction must be equal to 1.

  17. 4.Multiply the term in step 1 by the fraction in step 3. Since the fraction equals 1, you can multiply it without changing the size of the term. 5. Check math by canceling your units.

  18. Convert the following quantities using the following equivalence statements. Show work! 1 m = 1.094 yd 1mile = 1760 yd 1kg = 2.205lbs • 30.0 m to miles • 1500 yd to miles • 206 miles to m • 34 kg to lbs • 34 lb to kg

  19. Matter • All matter is composed of 100 or so _____________ • A substance that cannot be separated into simpler substances by a chemical change; simplest type of pure substance. • The building block of matter is the_________ • The smallest particle of an element that retains the chemical identity of the element. • Atoms can combine to form ___________

  20. Element= Compound= Elements and Compounds = Pure Substances

  21. Homogeneous= Heterogeneous= MixturesMixtures can be;

  22. Classification of Matter

  23. Physical Characteristics can be observed without altering the identity of the substance Volume Mass Maleability, ductility, conductivity etc… Chemical Characteristics cannot be observed without altering the identity of the substance Flammability Tendency to corrode Reactivity Etc… Properties of Matter

  24. Changes Matter Can Undergo: • Physical Change: Solid  Liquid Melting Liquid  Gas Boiling or Evaporating Gas  Liquid _____________ Solid  Gas _____________ Gas  Solid _____________ Liquid  Solid Freezing, solidifying

  25. Changes Matter Can Undergo: • Chemical Change: Rusting, rotting, burning, chemical reaction…

  26. Distinguishing Chemical from Physical Change • Did the change produce a different substance? • Was there a color change? • Is there a different density? • Is there a melting or boiling point change? • Did something precipitate out of solution? • Did a gas or smoke form?

  27. EnergyRemember: Matter- anything that has mass and takes up space. Energy is the other “stuff” of the universe. The capacity to do work (the ability to move or change matter) 1. Kinetic- 2. Potential- 3. Radiant/ electromagenetic- heat* and light. *We are mainly concerned with heat for this unit.

  28. Heat Energy due to _____________________ Symbol ____ Units: ___________ Does work by _______________________________________________________________ Flows from hot areas to cold areas Calorimetry Temperature A measure of ________________________________________________________________________ Refers to the intensity of heat in an object Symbol T Units: _______________ Change in T = Tf –Ti = D T NOT a form of energy but is a predictor of heat flow Heat Vs Temperature

  29. Heat Vs Temperature • Keep in mind: • Objects can be the same temperature but have different amounts of heat energy • Heat is dependent on MASS

  30. Temperature Scales • 0 K absolute zero; all molecular motion stops • 0 K  theroretical temperature not yet obtained (within a millionth of a degree) • Closer to absolute zero  atoms move more and more slowly – much more difficult to remove heat • Sig figs and temperature: because the Celsius temperature is a continuum with both positive and negative values, a temperature measurement of 00C has 1 sig fig (0.10C = 2 sig figs)

  31. Temperature Scale Conversions

  32. Significant Temperatures for Phases of Water

  33. Kinetic Molecular Theory* 1. (atoms / molecules) • The basic principles of KMT are theoretical and begin to break down under certain circumstances KMT is better at describing matter in higher energy states (gases, for example)

  34. States/ Phases of Matter

  35. Calorimetry • Physical and chemical changes are normally accompanied by energy changes. • Energy changes in a laboratory setting are measured using a calorimeter.

  36. If heat is consumed during the change, then the process/change/reaction is said to be ___________________. If heat is produced during a change, then the process/change/reaction is said to be ________________. Types of Energy Changes

  37. Law of Conservation of Energy • Within a closed system, energy transforms from one type to another. • ______________________________________. Example: electricity lights a bulb: resistance builds up in the tungsen wire, it glows and gives off light and heat; the total energy in the heat and light = the energy in the electricity. Example: when heat is added to water on a hot plate, that heat energy is absorbed by the water molecules, which move faster and faster (increased kinetic energy higher temperature)

  38. Law of Conservation of Matter • Matter can only be transformed during chemical and physical changes. • ___________________________________________. Example: when ice melts to make water during a phase change Example: when two chemicals are mixed *On our large scale, we see matter and energy as separate, but matter and energy interconvert at the subatomic level according to Einstein’s Theory of Relativity E=mc2)

  39. DURING HEATING OR COOLING c = specific heat for water = 4.18 J/goC m = mass of sample DT = change in temperature of sample in oC DURING A PHASE CHANGE (freezing/ melting) (evap / condense) M = mass of sample Hf = heat of fusion (for water = 334 J/g) Hv = heat of vaporization (for water = 2260 J/g) Calorie ProblemsTheoretical values for energy changes during the heating or cooling of a substance, or during a phase change, can be calculated using three basic equations.

  40. Why do we add propylene glycol (antifreeze) to our car’s radiators? The value of Q for any substance can be calculated, but note that each substance has unique values for specific heat capacity (c), heat of fusion (Hf), and heat of vaporization (Hv). Think about it: it’s easier to raise the temperature of some substances than others.

  41. High specific heat capacity (c) = a large amount of energy must be added in order to increase the temperature. • Water(l) = 4.18 J/(g•K) • Low specific heat capacity (c) = a small amount of energy must be added in order to increase the temperature. • Iron(s) = 0.129 J/(g•K)

  42. Q = mcDT • How much heat is required to raise the temperature of 10.0 g of water from 5oC to 25.0oC? • What will be the temperature change if 418 J of heat are added to 25 g of water?

  43. Q = mHf How much heat is needed to melt 5.0 g of water? Q = mHv How much water can be vaporized by 3135 Joules?

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