The Study of Chemistry

1. The Study of Chemistry The Molecular Perspective of Chemistry • Matter is the physical material of the universe. • Matter is made up of relatively few elements. • On the microscopic level, matter consists of atoms and molecules. • Atoms combine to form molecules. • Molecules may consist of the same type of atoms or different types of atoms.

2. The Molecular Perspective of Chemistry

3. The Study of Chemistry • Why Study Chemistry • Chemistry is central to our understanding of other sciences. • Chemistry is also encountered in everyday life. • It is the basis behind many medications • “Form defines Function”

6. Classification of Matter States of Matter • Matter can be a gas, a liquid, or a solid. • Gases have no fixed shape or volume. • Gases can be compressed to form liquids. • Liquids have no shape, but they do have a volume. • Solids are rigid and have a definite shape and volume.

7. Classification of Matter Pure Substances • Atoms and Compounds • Atoms consist only of one type of element. • Molecules can consist of more than one type of element. • Molecules can have only one type of atom (an element). • Molecules can have more than one type of atom (a compound). • Pure substances are • The same throughout • Have a set ratio • Mixtures- • Two types • Solutions

8. Classification of Matter Pure Substances and Mixtures • If matter is not uniform throughout, then it is a heterogeneous mixture. • 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. • If a pure substance can be decomposed into something else, then the substance is a compound.

9. Pure Substances and Mixtures

10. Classification of Matter • Elements • If a pure substance cannot be decomposed into something else, then the substance is an element. • There are 114 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.

11. If a pure substance cannot be decomposed into something else, then the substance is an element. • There are 114 elements known. • Each element is given a unique chemical symbol (one or two letters). • Chemical symbols with one letter have that letter capitalized (e.g., H, B, C, N, etc.) • Chemical symbols with two letters have only the first letter capitalized (e.g., He, Be). • Elements are building blocks of matter. • The earth’s crust consists of 5 main elements. • The human body consists mostly of 3 main elements. ELEMENTS

12. Classification of Matter • Elements

13. Classification of Matter • Compounds • Most elements interact to form compounds. • The proportions of elements in compounds are the same irrespective of how the compound was formed. • Law of Constant Composition (or Law of Definite Proportions): • the ratio by mass of the elements in a chemical compound is always the same, regardless of the source of the compound. • The law of constant composition can be used to distinguish between compounds and mixtures of elements: • Compounds have a constant composition; mixtures do not. • Water is always 88.8% O and 11.2% H by weight regardless of its source. • Brass is an example of a mixture of two elements: copper and zinc. It can contain as little as 10%, or as much as 45%, zinc.

14. Classification of Matter • Compounds • If water is decomposed, then there will always be twice as much hydrogen gas formed as oxygen gas. • Pure substances that cannot be decomposed are elements.

15. Classification of Matter • Mixtures • Heterogeneous mixtures are not uniform throughout. • Homogeneous mixtures are uniform throughout. • Homogeneous mixtures are called solutions.

17. Properties of Matter Physical and Chemical Changes • When a substance undergoes a physical change, its physical appearance changes. • Ice melts: a solid is converted into a liquid. • Physical changes do not result in a change of composition. • When a substance changes its composition, it undergoes a chemical change: • 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.

18. Properties of Matter Physical and Chemical Changes

19. Properties of Matter Physical and Chemical Changes • Intensive physical properties do not depend on how much of the substance is present. • Examples: density, temperature, and melting point. • Extensive physical properties depend on the amount of substance present. • Examples: mass, volume, pressure.

20. Properties of Matter Separation of Mixtures • Mixtures can be separated if their physical properties are different. • Solids can be separated from liquids by means of filtration. • The solid is collected in filter paper, and the solution, called the filtrate, passes through the filter paper and is collected in a flask.

21. Properties of Matter Separation of Mixtures • Homogeneous liquid mixtures can be separated by distillation. • Distillation requires the different liquids to have different boiling points. • In essence, each component of the mixture is boiled and collected. • The lowest boiling fraction is collected first.

22. Separation of Mixtures

23. Properties of Matter Separation of Mixtures • Chromatography can be used to separate mixtures that have different abilities to adhere to solid surfaces. • The greater the affinity the component has for the surface (paper) the slower it moves. • The greater affinity the component has for the liquid, the faster it moves. • Chromatography can be used to separate the different colors of inks in a pen.

24. Units of Measurement SI Units • There are two types of units: • fundamental (or base) units; • derived units. • There are 7 base units in the SI system.

25. Units of Measurement • Powers of ten are used for convenience with smaller or larger units in the SI system.

26. Units of Measurement SI Units

27. Units of Measurement SI Units • Note the SI unit for length is the meter (m) whereas the SI unit for mass is the kilogram (kg). • 1 kg weighs 2.2046 lb. 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 = -273.15 oC.

28. Units of Measurement Temperature • Celsius Scale • Also used in science. • Water freezes at 0 oC and boils at 100 oC. • To convert: K = oC + 273.15. • Fahrenheit Scale • Not generally used in science. • Water freezes at 32 oF and boils at 212 oF. • To convert:

29. Units of Measurement Temperature

30. Units of Measurement Derived Units • Derived units are obtained from the 7 base SI units. • Example:

31. Units of Measurement Volume • The units for volume are given by (units of length)3. • SI unit for volume is 1 m3. • We usually use 1 mL = 1 cm3. • Other volume units: • 1 L = 1 dm3 = 1000 cm3 = 1000 mL.

32. Units of Measurement Volume

33. Units of Measurement Density • Used to characterize substances. • Defined as mass divided by volume: • Units: g/cm3. • Originally based on mass (the density was defined as the mass of 1.00 g of pure water).

34. Uncertainty in Measurement Uncertainty in Measurement • All scientific measures are subject to error. • These errors are reflected in the number of figures reported for the measurement. • These errors are also reflected in the observation that two successive measures of the same quantity are different. Precision and Accuracy • Measurements that are close to the “correct” value are accurate. • Measurements that are close to each other are precise.

35. Uncertainty in Measurement Precision and Accuracy

36. Uncertainty in Measurement Significant Figures • The number of digits reported in a measurement reflect the accuracy of the measurement and the precision of the measuring device. • All the figures known with certainty plus one extra figure are called significant figures. • In any calculation, the results are reported to the fewest significant figures (for multiplication and division) or fewest decimal places (addition and subtraction).

37. Uncertainty in Measurement Significant Figures • Non-zero numbers are always significant. • Zeros between non-zero numbers are always significant. • Zeros before the first non-zero digit are not significant. (Example: 0.0003 has one significant figure.) • Zeros at the end of the number after a decimal place are significant. • Zeros at the end of a number before a decimal place are ambiguous (e.g. 10,300 g).

38. Dimensional Analysis Dimensional Analysis • Method of calculation utilizing a knowledge of units. • Given units can be multiplied or divided to give the desired units. • Conversion factors are used to manipulate units: • Desired unit = given unit  (conversion factor) • The conversion factors are simple ratios:

39. Dimensional Analysis Using Two or More Conversion Factors • Example to convert length in meters to length in inches:

40. Dimensional Analysis Using Two or More Conversion Factors • In dimensional analysis always ask three questions: • What data are we given? • What quantity do we need? • What conversion factors are available to take us from what we are given to what we need?