1. Chemistry and Matter

2. Chemistry: The Central Science Chemistry is the study of matter and the changes it undergoes A basic understanding of chemistry is central to all sciences.

3. Characteristics of Matter Matter – anything that has mass and takes up space – Mass – measurement of the amount of matter – Weight – measure of the amount of matter but also the effect of Earth’s gravitational pull on that matter

4. Branches of Chemistry Organic Chemistry Inorganic Chemistry Physical Chemistry Analytical Chemistry Biochemistry

5. Scientific Method Observation – gathering information – Qualitative data – non-numeric – Quantitative data – numeric Hypothesis – tentative explanation Experiment – a set of controlled observations that test a hypothesis – Independent variable – variable that is controlled – Dependent variable – responds to change

6. Conclusion – a judgment based on the information obtained Theory – an explanation that has been supported by many, many experiments Scientific Law – a relationship in nature that is supported by many experiments

7. Technology is the practical use of scientific information and is concerned with making improvements in human life and the world around us.

8. Metric System The metric system of measurement is used in science so that scientists can share data with other scientists. It is also called the Systeme Internationale d’Unites, SI.

9. Measurement A Base Unit is a defined unit in a system of measurement that is based on an object or even in the physical world. QuantityBase Unit TimeSecond (s) LengthMeter (m) MassKilogram (kg) TemperatureKelvin (K) Amount of substanceMole (mol)

10. Prefixes Used with SI Units PrefixSymbolFactorScientific Notation teraT1,000,000,000,00010 12 gigaG1,000,000,00010 9 megaM1,000,00010 6 kilok1,00010 3 hectoh10010 2 decada1010 1 decid1/1010 -1 centic1/10010 -2 millim1/100010 -3 microµ1/1,000,00010 -6 nanon1/1,000,000,00010 -9 picop1/1,000,000,000,00010 -12

11. Derived Unit A unit that is defined by a combination of base units is a derived unit. – Volume mL, cm 3, L, dm 3 Density – a ratio that compares the mass of an object to its volume; g/cm3

13. 1.If a sample of aluminum has a mass of 13.5 g and a volume of 5.0 cm3, what is its density? 2. Suppose a sample of aluminum is placed in a 25-mL graduated cylinder containing 10.5 mL of water. The level of the water rises to 13.5 mL. What is the mass of the aluminum sample?

14. 3.A piece of metal with a mass of 147 g is placed in a 50 mL graduated cylinder. The water level rises from 20 mL to 41 mL. What is the density of the metal? 4.What is the volume of a sample that has a mass of 20 g and a density of 4 g/mL? 5.A metal cube has a mass of 20 g and a volume of 5 cm3. Is the cube made of pure aluminum? Explain your answer.

15. Temperature The temperature of an object is a measure of how hot or cold the object is relative to other objects. Explain how a thermometer works.

16. Temperature Scales Scientists use two temperature scales. The Celsius scale was devised by Anders Celsius, a Swedish astronomer. – He used the temperature at which water freezes and boils to establish his scale because these temperatures are easy to reproduce. He defined the freezing point as 0 and the boiling point as 100. Then he divided the distance between these points into 100 equal units, or degrees, Celsius.

17. The Kelvin scale was devised by a Scottish physicist and mathematician, William Thomson, who was known as Lord Kelvin. The kelvin (K) is the SI base unit of temperature. On the Kelvin scale, water freezes at 273.15 K and boils at about 373.15 K. We will learn more about this when we study gases. K = °C + 273.15

18. Scientific Notation Scientific Notation expresses numbers as a multiple of two factors: – a number between 1 and 10; – and ten raised to a power, or exponent. The mass of a proton is 1.627 62 x 10 -27 kg. Change to scientific notation: 1.The diameter of the Sun is 1,392,000 km. 2.The density of the Sun’s lower atmosphere is 0.000 000 028 g/cm 3.

19. Express the following quantities in scientific notation. a. 700 Me. 0.0054 kg b. 38,000 mf. 0.000 006 87 kg c. 4,500,000 mg. 0.000 000 076 kg d. 685,000,000,000 mh. 0.000 000 000 8 kg Express the following quantities in scientific notation. a. 360,000s b. 0.000 054 s c. 89,000,000,000 s

20. Dimensional Analysis A conversion factor is a ratio of equivalent values used to express the same quantity in different units. A conversion factor is always equal to 1. 1 foot__60 seconds 12 inches 1 minute Dimensional analysis is a method of problem- solving that focuses on the units used to describe matter. 48 km x 1000 m = 48,000 m 1 km

21. How many seconds are there in 24 hours? The density of gold is 19.3 g/mL. What is gold’s density in decigrams per Liter? A car is traveling 90.0 kilometers per hour. What is its speed in miles per minute?

22. Need to Know! 1 km = 0.62 miles 1 kg = 2.2 lb 1 mile = 5,280 ft 1 ft = 12 in 1 gal = 3.79 L 1 yd = 3 ft 1 in = 2.54 cm

23. 11. a.Convert 360 s to ms b.Convert 4800 g to kg c.Convert 5600 dm to m d.Convert 72 g to mg e.Convert 245 ms to s f.Convert 5 m to cm g.Convert 6800 cm to m h.Convert 25 kg to Mg. i.Convert 2.56 kbytes to Gbytes

24. 12.Convert each: a.Your weight in lbs to kg. b.Your height in inches to meters. c.2 liters to gallons

25. Accuracy and Precision Accuracy refers to how close a measurement is to an accepted value. Precision refers to how close a series of measurements are to an accepted value.

26. Percent Error Percent error is the ratio of an error to an accepted value. – Percent error = theoretical - experimental x 100 theoretical

27. Significant Figures As scientists have developed better measuring devices, they have been able to make more precise measurements. Scientists indicate the precision of measurements by the number of digits they report. A value of 3.52 g is more precise than a value of 3.5 g. The digits that are reported are called significant figures. Significant Figures include all known digits plus one estimated digit.

29. Rules for Significant Figures 1.Non-zero numbers are always significant. 2.Zeros between non-zero numbers are always significant. 3.All final zeros to the right of the decimal place are significant. 4.Zeros that act as placeholders are not significant. Convert quantities to scientific notation to remove the placeholder zeros. 5.Counting numbers and defined constants have an infinite number of significant figures.

30. 72.3 g has 3 significant figures. 60.5 g has 3 significant figures. 6.20 g has 3 significant figures. 0.0253 g and 4320 g each have 3. 6 molecules and 60 s = 1 min

31. 13. Determine the number of significant figures in the following masses. a.0.000 402 30 g b.405,000 kg c.508.0 L d.820,400.0 L e.1.0200 x 10 5 kg f.807,000 kg g.0.049 450 s h.0.000 482 mL i.3.1587 x 10 -8 g j.0.0084 mL

32. Round all numbers to four significant figures. Write the answers to 15 in scientific notation. 13.a) 84,791 kg b) 38.5432 g c) 256.75 cm d) 4.9356 m 14.a) 0.000 548 18 g b) 136,758 kg c) 308,659,000 mm d) 2.0145 mL