Introduction: Matter and Measurement

1. Introduction: Matter and Measurement SC 131 CHEM 1 Chemistry: The Central Science CM Lamberty

2. Homework • Chapter 1 Exercises (p 31-35) • 12, 14, 18, 20, 22, 24, 26, 28, 30, 32 • 36, 40 • 42, 46, 48a, 50, 52 • 56, 58, 60, 62, 65, 68, 71, 72, 78

3. Chemistry • The study of materials and the changes that materials undergoes • Where is chemistry present in your life?

4. The Study of Chemistry • The Atomic and Molecular Perspective • Matter • Anything with mass and occupies space • Property • Characteristic that helps recognize a type of matter and distinguish from other types • Elements • 100+ basic substances that make up matter either alone or in various combinations

5. The Study of Chemistry • Atom: • submicroscopic particles • Fundamental building blocks • Molecules • Two or more atoms joined in specific geometric arrangement • Bonds • Electronic force that holds atoms together in molecule

6. The Study of Chemistry • Why Study Chemistry? • Impacts daily lives • Informs citizens • Fulfills curriculum requirements

7. Classification of Matter • Substance: specific instance of matter • Solid: mq close and little movement • Fixed volume, rigid shape • Crystalline or amorphous • Liquid: mq close but free to move • Fixed volume, no fixed shape • Gas: mq far apart, compressible • No fixed volume, no fixed shape

8. Classification of Matter • Classification according to Composition • Kinds and amt of substances that make up matter • Pure substance: single type of atom/mq • Element—cannot be broken down further • Compound—can be broken down into elements • Fixed definite composition • Mixture • Heterogeneous • Homogeneous

10. Physical & Chemical Properties • Physical Property—displays w/o changing appearance • Examples • Chemical Property—only displayed by changing composition • Examples

11. Physical Change Alter only appearance not composition Ex Chemical Change Composition changes Ex Physical & Chemical Changes

12. Separation of Mixtures • Individual sorting by color or shape • Use of physical properties • Magnetic • Filtration • Distillation • Chromatography • Chemical reactivity • One substance reacts while the other does not. Need to be able to get back original substance.

13. The Scientific Approach to Knowledge • Empirical • Observation and experimentation • Hypothesis • Tentative explanation of observations • Experiments • Highly controlled experiments • reproducible • Theory • Well-established hypotheses • Scientific Law • Summarize past observations and predicts future ones

15. Units of Measurements • International System of Units (SI)

18. Length and Mass • Length • Meter • Distance light (598 nm) travels in 1 second • Just a bit more than a yard • Mass • Amount of material in an object • Not weight (which is a force) • 1 kg ~ 2.2 pounds • Cube of Platinum in Sorbonne????

19. Temperature • Hotness or coldness of an object • Direction of heat flow • Heat flows from higher T to lower T spontaneously • Celsius and Kelvin • Kelvin is absolute scale and does not have negative values • Conversion Factor • °F = 1.8(°C) + 32 • K = °C + 273.15

22. Derived Units • Combination of other units • Volume—amount of space matter occupies • Vol of cube = (edge length)3 • Liter or milliliter (L or mL) for liquids • Density—mass per unit volume • Density = mass/volume = m/V

23. Intensive vs. Extensive Properties • Intensive—independent of the amount of substance • Extensive—dependent upon the amount of substance

24. Uncertainty in Measurement • Precision vs. Accuracy • Precision is measure of how closely individual measurements agree with one another • Accuracy is how closely measurement agrees with the correct or “true” value • Perform several trials and average the results • Standard deviation reflects how much results differ from average • Significant Figures

25. Uncertainty in Measurement • Scientific measurements are reported so that every digit is certain except the last, which is estimated.

26. Uncertainty in Measurement • Significant Figures—only for measured values • The greater the number of significant figures, the greater the certainty fo the measurement • Exact Numbers—actual counts • No uncertainty, unlimited sig fig

27. Significant Figure Rules • All nonzero digits are significant • Interior zeros are significant • Leading zeros are not significant • Trailing zeros • After decimal point always significant • 3.9000 • Before decimal point are significant • 40.00 • Before implied decimal point are ambiguous • 1200 use sci notation 1.200 x 103 or 1.20 x 103

28. Significant Figures in Calculations • Multiplication/division—result uses fewest number of sig fig • Addition/subtraction—fewest number of decimal places • Rounding—4 or less round down, 5 or greater round up • Round at the end of all calculations not individual steps • Calculators are stupid & do not know rules

29. Solving Chemical Problems • Generally 2 types: • Unit conversion (dimensional Analysis) or specific equation • Dimensional Analysis

30. Calculate the displacement of a 5.70 L automobile engine in cubic inches Watch units raised to a power and account for that mathematically

31. General Problems Solving Strategy • Identify starting point (given info) • Identify the end point (what you want) • Devise a way to get from start to end—conceptual plan • Sort • Strategize • Solve • Check