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Chapter 1

Chapter 1. Chemistry: The Study of Change. Classification. Mixtures Elements Pure Substances Homogeneous Matter Heterogeneous Compounds. Classification. Classification. What is the difference between pure substances and mixtures?.

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Chapter 1

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  1. Chapter 1 Chemistry: The Study of Change

  2. Classification • Mixtures • Elements • Pure Substances • Homogeneous • Matter • Heterogeneous • Compounds

  3. Classification

  4. Classification • What is the difference between pure substances and mixtures? • Pure substances - chemical combination (bonded) with constant composition • Mixtures - physical combination of two or more substances

  5. Classification • What is the difference between elements and compounds? • Elements - single atom or a chemical combination of the same element • Li or C60 • Compounds - chemical combination of two or more elements • H2O or CH4

  6. Classification • What is the difference between homo/heterogeneous? • Homogeneous - physical combination that is uniform (same throughout) • Salt water • Heterogeneous - physical combination that is not uniform (different throughout) • Sandy water

  7. Physical and Chemical Properties • What is the difference between a chemical and physical property? • Physical property can be measured/observed without changing the identity of the substance • Ex. Color, odor, phase, volume, mass... • Chemical property cannot be measured/observed without changing the substance

  8. Physical and Chemical Properties Physical properties can be divided into what two categories? • Extensive – depends on how much matter you have • Mass • Volume • Intensive – does not depend on how much matter you have • Density • Melting Point

  9. Physical and Chemical Properties Physical Property Chemical Property • Blue • Density greater than 1 • Solubility • Melting point is 60°C • Hard • Slimy • Lustrous • Smells like feet • Flammability • Reacts with vinegar • Supports combustion • Reacts with water • Reacts with air • Does not react with ammonium hydroxide • Extremely sour

  10. Physical and Chemical Properties

  11. Physical Change • What does it mean to undergo a physical change? • One or more physical properties change, we change the appearance • Color change • Mass change • Phase change • The composition does not change • Ex. Ice and water vapor are still H2O

  12. Chemical Change What does it mean to undergo a chemical change? • Two or more substances reacted and a new substance was made • Ex. Rust • Cannot get the substance back without using a chemical reaction • Indicators could be: • A gas is produced • Light is produced • A solid precipitate forms

  13. Physical and Chemical Changes Physical Change Chemical Change • Cutting your hair • Bending a piece of iron • Crushing Aspirin into a fine powder • Scratching glass with diamond • Crumpling a piece of paper • Water freezes to form ice • A candle melting • Chocolate syrup is dissolved in milk • Baking soda reacts with vinegar to create gas • Copper turns green when exposed to air • Two clear liquids are mixed and a yellow color forms • Burning wood to create ash • Food spoiling • A candle burning • Mixing two liquids creates light

  14. Factor-Label Method (Dimensional Analysis) Why is the factor-label method very important in chemistry? It allows us to easily convert between units using a series of fractions based on equivalencies

  15. Factor-Label Method (Dimensional Analysis) How many different types are there? • Three types: • Two step - converting to or from a base unit • Multiple step - converting from a non-base unit to a non-base unit • Formula conversion – converting between two different measuring systems • Ex. Celsius to Fahrenheit

  16. Factor-Label Method (Dimensional Analysis) • Two-Step • Convert: • 2.5 cm to m • 5.08 g to kg • 3.21 L to ML • 0.19 mol to mmol • 6.7 µm to m

  17. Factor-Label Method (Dimensional Analysis) • Multi-Step • Convert: • 0.04 mJ to kJ • 36.0 μL to pL • 5 mmol to Mmol • 37 hours to seconds • 8 ounces to kilograms

  18. Significant Figures (Sig figs) • Why are significant figures important? • Indicates how precise the measurement is • You are only as precise as your least precise measurement!

  19. Significant Figures (Sig figs) • How do you determine what figures are significant? • Always significant: • Any non-zero number • 2132has four sig figs • Any zero between two non-zero numbers • 2002 has four sig figs – two 0’s and two 2’s

  20. Significant Figures (Sig figs) • How do you determine what figures are significant? • Never significant: • All zeros in front of the first non-zero number • 0.0012 has two sig figs –1 and 2

  21. Significant Figures (Sig figs) • How do you determine what figures are significant? • Sometimes significant: • Zeros at the end of a number are significantwhenever there is a decimal present • 0.075900 has 5 sig figs; the first two zeros do not count • 14000.has 5 sig figs • 14000.0has 6 sig figs

  22. Significant Figures (Sig figs) • How do you determine what figures are significant? • Sometimes significant: • Zeros at the end of a number are not usuallysignificant when there is not a decimal point • 14,000 has only 2 sig figs

  23. Significant Figures (Sig figs) Leading zeroes don't count, middle zeros do count, ending zeros only count if a decimal's in the amount!

  24. Significant Figures (Sig figs) • How many sig figs do the following numbers have? • 0.0013 • 1.20 • 0.0101010 • 450 • 350.0 • 2 sig figs • 3 sig figs • 6 sig figs • 2 sig figs • 4 sig figs

  25. Significant Figures (Sig figs) • How do you add/subtract with sig figs? • Look at decimal places. Your answer should end in the same decimal place as your least precise measurement • 200.27 + 45.8 + 31.456 = • 277.5

  26. Significant Figures (Sig figs) • How do you multiply/divide with sig figs? • What about mixed operations? • Your answer should have the same number of significant figures as the measurement with the fewest significant figures • 25.03 x 0.02 = • 0.5 • Follow sig fig rules for each step of the calculation

  27. Significant Figures (Sig figs) Constants and Conversion Factors NEVER count towards sig figs!

  28. Significant Figures (Sig figs) • 100.1 + 34 = • 7.892 – 2.9 • 3.44 – 11 + 0.009 = • 1.23 x 4.0 = • (100.1)(10.00) / 0.01000 = • (100.0 + 23 + 99.99) / 89.99 = • 134 • 5.0 • -8 • 4.9 • 100100 • 2.48

  29. Significant Figures (Sig figs) From this point on… • All answers in the homework, the lab work and quizzes must have units and the proper number of significant figures!

  30. Precision vs. Accuracy • What is the difference between precision and accuracy? • Accuracy refers to how close a measurement (average) is to the true or accepted value • Precision indicates how close together a group of measurements are to each other

  31. Precision vs. Accuracy

  32. Precision vs. Accuracy • Assuming the center bull's-eye is our accepted value: • Who is the most accurate at throwing darts? • Who is the least precise? Matt John Dan Pete

  33. Practice Problem • Below is a data table produced by four groups of students who were measuring the mass of a paper clip which had a known mass of 1.0003 g. • NOTE: The last row is the average of their measurements. • Which group(s) are the most accurate? • Which group(s) are the most precise? • Which group is the most accurate and precise?

  34. Practice Problem • Suppose a GPS (Global Positioning System) which measures your position on earth, is not calibrated correctly. You take 3 readings at the same place and they are all close together but 14 miles from your actual position. • Explain your results in terms of precision and accuracy.

  35. Practice Problem • What is the difference in the measurements 4 vs 4.00 vs 4.00000? In terms of taking measurements, explain why we use significant figures.

  36. Percent Error • What is percent error? • Used to determine how accurate your final results are What you should have got! What you actually got!

  37. Types of Errors • What are the three types of errors? • 1 – Blunder • Errors made by mistakes in your lab techniques • Spillage, misreading instruments, writing measurements down wrong, etc. • Do not count these! REDO them!

  38. Types of Errors • What are the three types of errors? • 2 – Systematic • Error in accuracy • An error that is consistently repeated • Human error/ignorance • Inaccurate/limiting equipment • Cannot be reduced by repeatability

  39. Types of Errors • What are the three types of errors? • 3 – Random • Error in precision • Caused by unknown and unpredictable changes in the measurement • Reduced through repeatability • Requires you to note uncertainties

  40. Types of Errors

  41. Uncertainties • What is an uncertainty? • A measured value that cannot be determined exactly • Actual number could be above or below the measured reading • Error in precision

  42. Uncertainties • How do we represent uncertainties? • Need a range • Use +/- • Make it reasonable! • Examples • 25.2 +/- 0.1 mL • 23.0 +/- 0.5 °C • Depends on the precision of your measuring device

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