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Chemistry in Our Lives

Chemistry in Our Lives. Chemistry and Chemicals. What is chemistry ?. Chemistry is the study of substances in terms of Composition What a material it made of Structure How the elementary particles are put together Properties The characteristics of the material

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Chemistry in Our Lives

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  1. Chemistry in Our Lives Chemistry and Chemicals

  2. What is chemistry? Chemistry is the study of substances in terms of Composition What a material it made of Structure How the elementary particles are put together Properties The characteristics of the material Reactions How it behave with other substances

  3. Chemical reactions happen when • a car is started • tarnish is removed from silver • fertilizer is added to help plants grow • food is digested • electricity is produced from burning natural gas • rust is formed on iron nails

  4. Everything in our lives from materials to life involve chemistry • glass (SiO2)n • metal alloys • chemically treated water • plastics and polymers • baking soda, NaHCO3 • foods • fertilizers and pesticides • living beings

  5. Chemicals in Toothpaste

  6. The Scientific Method The scientific method is the process used to explain observations in nature. The method involves: • making observations • forming a hypothesis • doing experiments to test the hypothesis

  7. Everyday Scientific Thinking • Observation:The sound from a CD in a CD player skips. • Hypothesis 1:The CD or player is faulty. • Experiment 1:When the CD is replaced with another one, the sound from the second CD is OK. • Hypothesis 2:The original CD has a defect. • Experiment 2: When the original CD is played in another player, the sound still skips. • Theory: The experimental results suggest that the original CD has a defect.

  8. Units of Measurement

  9. In chemistry: quantities are measured experiments are performed results are calculated use numbers to report measurements, results are compared to standards.

  10. In a measurement of the thickness of the skin fold at the waist, calipers are used. A measuring tool is used to compare some dimension of an object to a standard.

  11. In every measurement, a number must be followed by a unit to have any meaning. Observe the following examples of measurements: Number and Unit 35 m (meter) 0.25 L (liter) 225 lb (pound) 3.4 h (hour)

  12. The Metric System (SI) The metric system and SI (international system) are • related decimal systems based on 10 • used in most of the world • used everywhere by scientists

  13. Length • is measured using a meter stick • uses the unit meter (m)in both the metric and SI systems

  14. The unit of an inch • is equal to exactly • 2.54 centimeters in • the metric system 1 in. = 2.54 cm

  15. Volume • is the space occupied • by a substance • the unit of volume is the liter (L)in the metric system 1 L = 1.06 qt

  16. The massof an object • is a measure of the quantity of material it contains • the unit gram (g)or kilogram (1000 g) is used What is the difference between mass and weight? Weight is the result of the action of gravity on mass. Your weight on the moon would be a lot less even though your mass would remain the same Despite this important difference, we will use these two terms interchangeably

  17. Thetemperature • indicates how hot or cold a substance is • the Celsius (C) scale is used in the metric system • the Kelvin (K) scale is also used • 18 °C is 64 °F on this thermometer On the C scale, the melting point of ice is 0 C and boiling point of water is 100 C What is heat or cold? What does temperature really measure?

  18. Time measurement • the unit second (s) is used in the metric system. • Time is based on an atomic clock that uses a frequency emitted by cesium atoms

  19. Scientific notation • is used to write very large or very small numbers • the width of a human hair (0.000 008 m) is written 8 x 10-6 m • a large number such as 4 500 000 s is written 4.5 x 106 s

  20. Scientific Notation • A number in scientific notation contains a coefficient and a power of 10. coefficient power of ten coefficient power of ten 1.5 x 102 7.35 x 10-4 • To write a number in scientific notation, the decimal point is placed after the first digit. • The spaces moved are shown as a power of ten. 52 000. = 5.2 x 1040.00378 = 3.78 x 10-3 4 spaces left 3 spaces right

  21. 10-3/105 = 10-8 102 10-3*105 = 105 100000 + 0.001 = 100000.001 10-3 +105 =

  22. Measurements What is the length of this piece of wood? What is the first digit? Any uncertainty in the digit? 4 4.5 What is the second digit? Any uncertainty in this digit? What is the third digit? Any uncertainty in this digit? 4.56 Definition of a significant figure: Significant digits include all digits with no uncertainty plus one estimation

  23. . l8. . . . l . . . . l9. . . . l . . . . l10. . cm • What is the length of the red line? • 1) 9.38 cm • 2) 9.39 cm • 3) 9.40 cm 9.38, or 9.39, 9.40 is less likely

  24. Number of Significant Figures 4 Measurement 38.15 cm 5.6 ft 120.55 m 0.0055 in 1200 m 2 5 2 2

  25. A. Exact numbers are obtained by 2. counting 3. definition B. Measured numbers are obtained by 1. using some measuring tool

  26. Classify each of the following as exact (E) or measured (M) numbers. Explain your answer. A. __ Gold melts at 1064 °C. B. __ 1 yard = 3 feet C. __ The diameter of a red blood cell is 6 x 10-4 cm. D. __ There are 6 hats on the shelf. E. __ The atom sodium has 11 protons and 12 neutrons.

  27. Significant Figures In calculations: • Answers must have the same number of significant figures as the measured numbers. • Calculator answers must often be rounded off. • Rounding rules are used to obtain the correct number of significant figures.

  28. Rounding Off • When the first digit dropped is 4 or less,the retained numbers remain the same. To round 45.832 to 3 significant figures drop the digits 32 = 45.8 • When the first digit dropped is 5 or greater, the last retained digit is increased by 1. To round 2.4884 to 2 significant figures drop the digits 884 = 2.5 (increase by 0.1)

  29. Multiplication and Division When multiplying or dividing use • the same number of significant figures (SF) as the measurement with the fewest significant figures Example: 110.5 x 0.048 = 5.304 = 5.3 4SFs 2SFs calculator 2SFs

  30. Addition and Subtraction When adding or subtracting,use • the same number of decimal places as the measurement with the fewest decimal places 25.2one decimal place + 1.34two decimal places 26.54calculated answer 26.5finalanswer (withone decimal place)

  31. For each calculation, round the answer to give the correct number of decimal places. A. 235.05 + 19.6 + 2 = 1) 257 2) 256.7 3) 256.65 B. 58.925 – 18.2 = 1) 40.725 2) 40.73 3) 40.7

  32. 1m/100cm = 1; 1m/1000mm = 1 An equality • states the same measurement in two different units • can be written using the relationships between two metric units Example: 1 meter is the same as 100 cm and 1000 mm. 1 m = 100 cm 1 m = 1000 mm 1 = 100cm/1m; 1 = 1000mm/1m

  33. volume has the dimensions of length cubed

  34. Several equalities can be written for mass 1 kg = 1000 g 1 g = 1000 mg 1 mg = 0.001 g

  35. Some Common Equalities

  36. An injured person loses 0.30 pints of blood. How many milliliters of blood would that be? 0.30pt*1qt/2pt = 0.15qt; 0.15qt*946mL/qt = 141.9 mL; 140 mL 0.30pt*2pt/1qt = 0.60pt2/qt

  37. If a person weighs 200 pounds, how many kiograms does the person weight? • 200 lb*1 kg/2.2 lb = 90.9 kg • 200 lb*2.2 lb/1 kg = 440 lb2/kg

  38. If the thickness of the skin fold at the waist indicates an 11% body fat, how much fat is in a person with a mass of 86 kg? 11 % fat means 11kg/100kg body weight 86 kg x 11 kg fat = 9.5 kg of fat 100 kg

  39. Density Density compares the mass of an object to its volume is the mass of a substance divided by its volume Density expression: D = mass = g or g = g/cm3 volume mL cm3

  40. Osmium is a very dense metal. What is its density in g/cm3 if 50.0 g of osmium has a volume of 2.22 cm3? 1) 2.25 g/cm3 2) 22.5 g/cm3 3) 111 g/cm3

  41. The density of the zinc object can be calculated from its mass and volume. d = 68.6g/(45.0-35.5)mL; 68.6g/9.5 mL d = 7.2 g/mL

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