The Nature of Chemistry

1. Chapter 1 The Nature of Chemistry Robert Boyle 1627-1691. Originally defined the concept of chemical “element.” Geber (Abu Musa Jabir ibn Hayyan, جابر ابن حیان) 721-815. “Father of Arabic chemistry.”

2. The Study of Chemistry The Atomic and Molecular Perspective of Chemistry • Matter is the physical material of the universe. • Matter is made up of relatively few (ca. 100) elements. • Elements are the buildingblocks of matter. • On the nano (ultramicroscopic) level, matter consists of atoms. An atom is a “nano-basketball” -- nano = 10 -9. • Atoms usually are found in the combined state, commonly molecules. • Molecules may consist of the same type of atoms or different types of atoms.

3. CAUTION!! Not all compounds are made up of molecules. Many compounds, for example, are composed of ionic lattices. For this chapter, however, we will confine discussions of compounds to the concept of molecules, which are the combinations of nonmetallic elements.

4. The Study of Chemistry The Molecular Perspective of Chemistry • In these models, red represents oxygen, white represents hydrogen, and gray represents carbon.

5. Classification of Matter PureSubstances and Mixtures A pure substance cannot be separated into simpler substances by physical means. A pure substance has definite and constant chemical and physical properties (i.e., ignition temp., melting point, magnetic susceptibility, spectral patterns) A pure element is a pure substance that consists only of one kind of atom. Examples of elements: U, P4, Cl2, C60 A pure compound is a pure substance that consists of more than one kind of atom. Examples of compounds: HCl, K2SO4 , C2H6O.

6. Molecules are chemical combinations of two or more atoms which can all be the same, e.g., Br2 , H2 , S8 , O3 (these are elements) or different, e.g., H2O, C2H5Br (these are compounds) (Elements which have a different number of atoms in their molecule are called allotropes. Allotropes have different physical and chemical properties, e.g., O2 and O3)

7. Classification of Matter Individual elements or compounds are all pure substances. For example, individual samples of C, Fe, O2, N2, H2O NaCl, K2SO4 are all pure substances. The composition of a pure substance is constant. (e.g., water is always 88.9% oxygen and 11.1% hydrogen) If different samples are mixed together, this is called a mixture. Mixtures do not have constant compositions.

8. Classification of Matter Examples of mixtures: • air (which contains variable amounts of oxygen, nitrogen, carbon dioxide, water vapor, pollutants, etc. • milk (which contains variable amounts of galactose, minerals, fats, proteins, etc.) • carbonated drinks (which contain variable amounts of carbon dioxide, sugars, flavorings, etc.) • “tap” water (which contains dissolved minerals)

9. pure substances Classification of Matter Pure Substances and Mixtures

10. Classification of Matter Pure Substances and Mixtures If matter is not uniform throughout, then it is a heterogeneous mixture. If a pure substance can be decomposed (chemically) into something simpler, then the substance is a compound. If matter is uniform throughout, it is homogeneous. If homogeneous matter can be separated by physical means, then the matter is a mixture. If homogeneous matter cannot be separated by physical means, then the matter is a pure substance.

12. Classification of Matter Elements • There are 115 elements known. • Each element is given a unique chemical symbol (one or two letters). • Elements are building blocks of matter. • The earth’s crust consists of 5 main elements. • The human body consists mostly of 3 main elements (O, C, and H).

13. Classification of Matter Elements The next five elements are: Na 2%, K 2%, Mg 2%, H 1%, Ti 0.5%. The next six elements are: N 3%, Ca 1.5%, P 1%, K,S,Na 0.75%

14. Elements in the Human Body – including trace elements

15. Classification of Matter Use your Periodic Table to refer to all elements and their chemical symbols. Bring your Periodic Table to each class!

16. The Periodic Table Bring your Periodic Table to each class!

17. Classification of Matter Compounds • Most elements interact to form compounds. • The proportions of elements in a compound is the same irrespective of how the compound was formed. • Law of Constant Composition (or Law of Definite Proportions): • The composition of a pure compound is constant (always the same). For example, water is always 88.9% oxygen and 11.1% hydrogen (by mass). For reasons we shall later see, the volumes of hydrogen and oxygen obtained from water are always in a fixed 2:1 ratio.

18. Properties of Matter Physical and Chemical Changes When a substance undergoes a physical change, its physical appearance changes, but its chemical nature does not. • Example: the melting of ice (physical change) results in a solid being converted into a liquid, but it is still water. Physical changes do not result in a change of composition. When a substance changes its composition, it undergoes a chemical change • Example: when pure hydrogen and pure oxygen react completely, they form pure water. In the flask containing water, there is no oxygen or hydrogen left over.

19. Units of Measurement SI Units • There are two types of units: • a) fundamental (or base) units; • b) derived units. • There are 7 base units in the SI system. • Derived units are obtained from the 7 base SI units. • Example:

21. Units of Measurement Powers of ten are used for convenience with smaller or larger units in the SI system. What is a GigaByte?

22. Units of Measurement - Temperature • There are three temperature scales: • Kelvin Scale (used in science) • Same temperature increment as Celsius scale. • Lowest temperature possible (absolute zero) is zero Kelvin. • Absolute zero: 0 K = -273.15oC. • Celsius Scale (used in science) • Also used in science. • Water freezes at 0oC and boils at 100oC. • To convert: K = oC + 273.15. • Fahrenheit Scale (used in US engineering and commerce) • Water freezes at 32oF and boils at 212oF. • To convert:

23. Units of Measurement - Temperature

24. Units of Measurement - Temperature • A user-friendly way to view the Celsius Scale: • 0° - Cold! (coat) • 10° - Cool (sweat shirt) • 20° - Pleasant (long sleeves) • 25° - Room temperature (short sleeves) • 30° - Very warm (T-shirt) • 40° - Hot! (swimming pool!)

25. Units of Measurement - Volume The units for volume are given by length3 SI unit for volume is 1 m3. 1 mL = 1 cm3 (REMEMBER THIS!) • Other volume units: • 1 L = 1 dm3 • =1000 cm3 = 1000 mL

26. Units of Measurement - Density • Used to characterize substances. • Defined as: density = mass /volume. • Units: g/cm3, also known as specific gravity. • Originally based on mass -- the density was defined as the mass of 1.00 g of pure water.

27. Uncertainty in Measurement Precision and Accuracy (actually discussed in Chapter 2) • Measurements that are close to the “correct” value are accurate. • Measurements which are close to each other are precise. • Measurements can be: • accurate and precise; • precise but inaccurate; • neither accurate nor precise.

28. Uncertainty in Measurement Precision and Accuracy Example: Weigh yourself on your bathroom scales. Your scales are cheap, and the graduations are every 5 pounds. The precision is bad. Your scales are expensive, and the graduations are every 0.1 pound. The precision is good. You know your correct weight from a high-quality scale at the doctor’s office, to within 0.1 pound. You weigh yourself on your bathroom scales and find the reading is off by 10 pounds. The accuracy is bad. Same situation as above, and find the reading is within 0.1 pound. The accuracy is good.

29. Dimensional Analysis Method of calculation utilizing a knowledge of units. • Conversion factors represent “1” (unity). • Conversion factors are simple ratios. Example: 1 m or 100 cm are equivalent 100 cm 1 m Example: 1 L or 1000 mL are equivalent. 1000 mL 1 L

30. Dimensional Analysis Example: Convert 5.6 km to m: x 1000 m km 5.6 km = 5600 m Volume conversion. Convert 1 cubic meter to mL. 1 m3 x (100 cm)3 x 1 mL = 1 x 106 mL (1 m)3 1 cm3