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Chapter 1

Chapter 1. Matter and Measurements. Chemistry: the “Central Science”. Chemistry is the study of matter Nature of matter Properties of matter Transformations of matter. Matter, its Nature. Matter is anything that has mass and volume. (Anything that takes up space)

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Chapter 1

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  1. Chapter 1 Matter and Measurements

  2. Chemistry: the “Central Science” • Chemistry is the study of matter • Nature of matter • Properties of matter • Transformations of matter

  3. Matter, its Nature • Matter is anything that has mass and volume. (Anything that takes up space) • So when someone tells you that you are a waste of space…. • You say, “nope, I matter.”

  4. Matter, its Properties • Physical • Size • Color • Melting point • etc • Chemical • Pretty much the whole course!

  5. Matter, its Transformations • Physical vs. Chemical Changes • Physical change • Chemical makeup doesn’t change • Chemical change • It does!

  6. States of Matter • 5 States of Matter (book says 3) • Solid • Definite volume • Definite shape • Liquid • Definite volume • Indefinite shape (will take the shape of container) • Gas • Indefinite volume • Indefinite shape • Will FILL it’s container

  7. States of Matter, The Noobs • Plasma • Like a gas, but are high temperature free charged particles • i.e., Hydrogen in stars • Bose-Einstein Condensates • Predicted by the above in the 1920s • 1995 first discovered • At a few BILLIONTHS of a degree above absolute zero, atoms clump forming……

  8. …a SUPERATOM! DAH, DAH, DAH! • Atoms become indistinguishable from each other.

  9. Changes of State • When matter changes from one state to another this is called a change of state, or a phase change. • Solid to liquid • Melting • Liquid to solid • Freezing

  10. Classification of Matter • Pure substance • Uniform in chemical composition and properties • Element • Compound (or molecule) • Mixture • Can vary in chemical composition and properties

  11. Chemical Reactions Reactants  Products H2+ S  H2S

  12. Chemical Symbols • All elements have a one or two letter symbol, a type of shorthand • Symbols are combined to form chemical formulas H2O 2 Hydrogen atoms 1 Oxygen atom

  13. Types of Elements • Metals • Solids (except Hg) • shiny, good conductors of heat and electricity, and are malleable • Nonmetals • Gases, liquid (Br) and solids • Are not. • Metaloids • Are in between (Si)

  14. Example of Chemical Change • Nickel + hydrochloric acid  Nickel (II) chloride + hydrogen • Ni + HCl  NiCl2 + H2 • Figure 1.5 on Page 14 to see the difference in the properties of reactants and products

  15. Physical Quantities • Length, volume, temperature, and other properties are physical quantities • A physical quantity is described by a number AND a unit • 10g is a physical quantiy • 10 is a number

  16. SI Units • SI units are used in science • Similar to the metric system • Table 1.5 page 16 of your book shows the fundamental SI units (minus a few). • There are also derived units combining 2 or more fundamental units • *volume is actually a derived unit LxWxH • Speed, m/s • Density g/cm3

  17. SI Prefixes • Help describe really big or really small numbers • “kilo” means “thousands” • So 10 km = 10 thousand meters • “nano” means “billionth” • So 5 nm = 0.000000005 meter

  18. Measuring Mass • Mass: the amount of matter in an object • “Weight” is the measure of gravity’s effect on a mass • We use weight to determine mass by using a balance (not a scale) • We can do this because gravity is constant • Usually use g, mg, μg in chemistry due to the sizes we typically work with

  19. Measuring Length and Volume • Length is the distance an object spans • A meter is too long for most things in science • You’ll see cm, mm, µm, nm instead • Measure with a ruler, micrometer, etc • Volume is the amount of space an object occupies • m3 is too large, so we use liter (L) = 1dm3 • Will see mL mostly: 1000mL = 1L • (1mL=1cm3= 1cc) • Measure with a graduated cylinder, or by measuring LxWxH

  20. Significant Figures • Every measurement has an “uncertainty” • Significant figures (or significant digits) show the precision of a measurement • The last digit of a value is the first “uncertain” digit • For a mass of 54.07g • 54.0 is certain, and 0.07 is an estimate of + 1

  21. Rules on Sig Figs • All non zero numbers are significant • Zeros between non zero numbers are significant • Zeros at the beginning of a number are not significant (they only show the location of the decimal point) • Zeros at the end of a number and after the decimal point are significant • Zeros at the end and before an implied decimal point may or may not be significant (use scientific notation=fixed)

  22. Sig Fig Examples 128.200 = ? S.F. 6 S.F. 128.2 = ? S.F. 4 S.F. 0.0012 = ? S.F. 2 S.F. 0.00120 = ? S.F. 3 S.F. 10204 = ? S.F. 5 S.F.

  23. Scientific Notation • Easy way to express very large or very small numbers • Written as a number between 1 and 10 times a power of 10 • 6.02 x 1023 • = 602000000000000000000000 • When the exponent is positive, the number is greater than 1. • When the exponent is negative, the number is less than 1

  24. Scientific Notation Examples 0.002 = ? 2 x 10-3 200 = ? 2 x 102 0.000504 = ? 5.04 x 10-4 3.2 x 10-2 = ? 0.032 More? With practice this is easy!

  25. Rounding Off Numbers • Sig Figs help us to round off numbers • When multiplying or dividing, the answer has the same number of S.F. as the original number with the least number of S.F. 124 x 2.4 = 513.6  510

  26. Rounding Continued • When adding or subtracting, the answer cannot have more digits after the decimal point than either of the original numbers 32.56 + 1.268 33.828  33.83

  27. Conversion Factors • Often in chemistry (or in life) we have to convert from one unit into another • The best way to do this is by using the “factor label” method. • Also called unit analysis, or dimensional analysis • Generically: Starting amount X conversion Factor = equivalent amount

  28. Example • How many grams are in 3.4 kilograms? 3.4 kg x 1000 g = X g kg Starting amt x conversion = equivalent amount factor  3400 g

  29. Tougher Example • A dose of painkiller is 600mg. Each pill contains 150mg of painkiller. How many pills are in a dose 600mg = 150mg x ? pill pill ALGEBRA ALERT!! 600mg x pill = ? pill 150mg 4 pills

  30. Goofy Answers • If you are looking for a mass, g and you get something like mg/g2, you did something wrong!!!! 10 g x g 1000 mg = 0.01 g2/mg  BUZZ, Try Again! Usually this means an upside-down conversion factor!

  31. “Estimating” • Identify given information • Identify result needed, including units • Find relationships (conversion factors) • Solve the problem • “Estimate” if it’s in the “ballpark” • If you did the first 4 right, you shouldn’t need this step!!!!

  32. Temperature, Heat and Energy • Chemical reactions involve a change in energy • Energy is the capacity to do work • Energy in detail is Chapter 7

  33. Measuring Temperature • Temperature is a measure of thermodynamic energy (not heat!) • SI unit is the Kelvin (K) • K is the same size as a °C • 273.15K = 0 °C (using 273 is fine) • So K = °C + 273.15 • 0K is “absolute zero” • Measure temp with a thermometer

  34. Units of Energy and Heat • SI unit is the Joule (J) • Also will see calorie (cal, kcal) • One kcal = one Calorie • Heat is a type of energy • Specific Heat is how much energy is needed to raise the temperature of 1 gram of a substance 1 °C.

  35. Using Specific Heat Heat (cal) = Mass (g) x Temp change (°C) x Specific Heat ( cal ) g * °C We can measure mass and temperature change, and can look up specific heat to determine heat.

  36. Example of Specific Heat • Example 1.15 page 32 of your book 95kg = 9500g Temp change = 40 °- 15 °= 25 °C Specific heat of water = 1.0 cal (g * °C) Heat = 9500 g x 25 °C x 1.0 cal = 240,000 cal (g * °C) = 2.4 x 105 cal (or 240 kcal)

  37. Density • Density is the relationship between an objects mass and its volume • How much mass is packed into how much volume • Expressed as g/mL (or g/cm3) • Makes it easy to determine the mass of a liquid, because it’s easier to measure its volume!

  38. Density Example • A piece of wood weighs 5 grams and has a volume of 8 cm3. What is its density? Density = g/cm3 Density = 5 g / 8 cm3 = 0.625 g/cm3 = 0.6 g/cm3

  39. Specific Gravity • Specific gravity (sp gr) is the density of the substance divided by the density of water at the same temperature • Density of water is usually 1.00 g/mL • Sp gr is useful to compare values as a unitless quantity • Winemaking • Dissolved solids in Urine

  40. Homework • End of chapter questions: • 1.39 • 1.51 • 1.65 • 1.67 • 1.81 • 1.89 • 1.91

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