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Introduction to Chemistry: Scientific Method, Measurement, and Density

Learn about the scientific method, measurement techniques, and the concept of density in this introductory chapter of chemistry. Explore the importance of significant digits, scientific notation, and the factor-label method. Understand why some substances float while others sink and discover how to calculate the density of solid and liquid objects.

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Introduction to Chemistry: Scientific Method, Measurement, and Density

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  1. Chapter 1

  2. Chemistry – • the study of all substances and the changes they can undergo. Scientific Method- • Observation • Stating a Question • Hypothesis – possible answer • Experiment • Conclusion – what you found in your experiment

  3. Scientific method leads to Natural Law – Tells you how nature behaves but not why it behaves. Finally you form a Theory – Explains why nature behaves in the way described by natural law. Used for prediction of results for further experiments.

  4. During the experiment… • Experimental Control - Factor that remains constant during the experiment. It is compared with the variable. • Variable - Factor that is being tested during the experiment.

  5. Measurement… • When we perform experiments, we need to use some form of measurement. • Metric System – International set of units (SI)

  6. Base units- • Mass = kilogram (kg) • Length = meter (m) • Time = seconds (s) • Count, quantity = mole (mol) • Temperature = Kelvin (K) • Electric Current = ampere (A) • Luminous intensity = candela (cd)

  7. Derived Units – made from combining 2 or more base units. • Ex. Area = length x width = m2 • Volume = length x width x height = cm3 • Density = mass / volume = g/cm3

  8. Exceptions to SI units… • Volume = liter (L) • Pressure = atmosphere millimeter of Hg (atm mm Hg) • Aka Pascal • Temperature = Celsius degree (Co) • Energy = Calorie (cal)

  9. Reliability in Measurement • Accuracy vs. Precision • Precision – Measurement that gives the same result again and again under the same conditions • Accuracy – Measurement that is close to the accepted value.

  10. Significant Digits – Defined as certain digits and the estimated digit of a measurement. • Rules for determining sig. Digs. • “0”s as placeholders are NOT significant! • Ex. 1040 g or 0.0064 m

  11. Atlantic – Pacific Rule • If the decimal is Present, count from the Pacific side (left) beginning with the first non-zero digit • If the decimal is Absent, count from the Atlantic side beginning with the first non-zero digit.

  12. How many significant Digits are present? • 1700 cm • 0.00960 kg • 64050 L • 45.00 mg • 0.0607 m

  13. How many sig. digs. are present? • 10100 mL • 0.50090 dg • 6.60 x 10-8mL • 2.00 x 103 g • 4. 010 x 104 L

  14. Calculation rules for Sig. Digs. • Multiplication and Division • The measurement with the SMALLEST NUMBER OF SIG. DIGS. Determines how many digits are allowed in the answer. • Ex. 4.3 x 6.45 will have 2 sig. digs. in the answer. • 27.735 = 28

  15. Addition and Subtraction • The number of significant digits is dependent upon or rounded off to the measurement with the largest uncertainty. ***Use the least amount of decimal spots*** • Ex. 6.45 + 2.36 + 4.6 = • 13.41 rounded to 13.4

  16. Scientific Notation • Why use it? • Distance from the sun = 93,000,000 miles • Al strip = 4.12 g contains 1.2 x 1023 atoms

  17. Rules for Scientific Notation • Express the same number of significant digits • Keep one digit to the left of the decimal point • Multiplication and Division, add the number of exponents

  18. Simplify the following: • (4.2 x 106)(1.1 x 10-2) 2.3 x 102 • (5.0 x 105)(1.2 x 101) 9.0 x 1010

  19. 2.0 x 10^2 • 6.7 x 10^-5

  20. Factor-Label Method/ Dimensional Analysis • Problem Solving Method • Treat units as factors, which can be cancelled • Must know your equalities/conversion factors • Choose the equality that cancels out the original unit

  21. Steps: • 1) Begin with known • 2) Decide upon equality (write this near the problem) • 3) Arrange units to cancel out original • 4) Do the math!

  22. How many seconds are in 4 hours? • How many kilograms are in 5 g? • Change 286 cg to Mg.

  23. To change from English to Metric units, use chart on page 38. • How many inches are in 354 cm?

  24. Ratios – common method of expressing results and/or measurement in chemistry • Similar to a fraction • Units in numerator and denominator

  25. Ex. • Speed = mph • Lunchmeat = dollars/lb • Density = g/cm3 or g/mL • Population Density = people/km2 • Pressure = psi or p/in2

  26. Most substances-Solid forms sink in its own liquid because solid form is more dense. • EXCEPTION- Water. • The freezing point of water is 0 degrees C. • Solid ice floats in liquid water. • Water is most dense at 4 degrees C.

  27. What would happen if ice did not float to… • The ocean? • It would freeze • The earth’s temperature? • It would decrease

  28. The icebergs? • They would sink and never melt • The earth? • It would freeze • Life? • It would cease to exist

  29. Solid Density • Used to find the density of solid objects. • Regular geometric shapes vs. irregular shapes.

  30. Liquid Density • The density of water • 1g/mL @ 4 degrees C • Other fluids have different densities • Oil floats on water • Density – mass per unit of volume

  31. Ex. • Zinc Cylinder • Diameter is 1.22 cm • Height = 3.28 cm • Mass = 28.79g

  32. V= ∏r² h • h= 3.14 x (0.61cm)² x 3.28cm =3.83cm³ • V= 3.83 cm3 • M= 28.79 g • D= m/v = 28.79g/3.83cm³ =7.52g/cm3

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