Why should you study chemistry?

1. Why should you study chemistry? americanchemistry.com : videos

2. How do we study chemistry? Chemistry is an experimental science: base on the scientific method.

3. What are we studying in chem,istry? Chemistry is the science that studies the composition and properties of matter. Matteris anything that occupies space, displays a property known as mass, and possesses inertia.Matter is the name scientists have given to everything that you can touch, or see, or feel Composition refers to the parts or components of a sample of matter and their relative proportions. Any characteristic that can be used to describe or identify matter is called a property. We classify matter (1) by its physical state as a solid, liquid, or gas, and (2) by its chemical constitution as an element, compound, or mixture.

4. Properties • Properties can be classified as • Intensive or extensive, depending on whether their value changes with the size of the sample. Intensive properties, like temperature and melting point, have values that do not depend on the amount of sample. • Extensive properties, like length and volume, have values that do depend on the sample size. • 2)Physical or chemical. • ~ Physical properties and physical change • A physical property, are those characteristics like temperature, color and melting point, is one that a sample of matter displays without changing its composition. • In physical change some of the physical properties of a sample may change, but its composition remains unchanged. • E.g. Liquid water and ice (solid water) certainly different in many way but water remains 11.9% hydrogen and 88.81% oxygen by mass.

5. Physical Properties Chemical Properties Temperature Amount Color Odor Melting point Solubility Electrical conductivity Hardness Rusting (of iron) Combustion (of coal) Tarnishing (of silver) Hardening (of cement) ~ Chemical properties and chemical changeA chemical property is the ability (or inability) of a sample of matter to undergo a change in composition under stated condition. In a chemical change (or chemical reaction), one of more kinds of matter are converted to new kinds with different compositions. E.g. the rust that occurs when a bicycle is left out in the rain is due to the chemical combination of oxygen with iron to give the new substance iron oxide. Rust is therefore a chemical properties of iron. Table 1 Some Examples of Physical and Chemical Properties

6. Chemistry in Action: cooking • Cooking is chemistry in action, with the added benefit that you can eat the results. As you follow the recipe below for cheese making, make observations about the changes that occur.

7. Quick review • Problem 1: What is a physical change? List the physical changes that occur when one makes cheese. • Ans: Physical changes include:1 Dissolving of salt in water • 2 Evaporation of water from solution • 3. Stirring to mix • Problem 2: What is a chemical change? List the chemical changes that occur when one makes cheese. • Ans: changes include:1 lactic acid bacteria convert the sugar in milk (lactose) to lactic acid • 2 special bacteria ferment the remaining lactose and produce carbon dioxide bubbles in the cheese.

8. Classification of Matter

9. Classification of Matter by its chemical constitution A substanceis a kind of matter that cannot be separated into other kinds of matter by any physical process. A mixturecan be separated into its components by appropriate physical means such as filtration; distillation and chromatography, etc.

10. Homogeneous mixtures are uniform in composition and properties throughout a given sample but variable from one sample to another.E.g. a solution of sucrose (sugar) is uniformly sweet throughout a given sample, but the sweetness of different sucrose solution samples may be different depend on the proportion of sugar and water in the solutions. In heterogeneous mixtures such as salad dressing the components separate into distinct regions. The composition and physical properties vary from one part of the mixture to another.

11. An element is a substance made up of only a single type of atom. Chemical compoundsare substances, formed when atoms of two or more different elements combine with one another. They retain their identity during physical change, but they can be decomposed into its constituent elements by chemical changes. A molecule or formula unit(for ionic compound) is the smallest entity having the same proportions of the constituent atoms as does the compound as a whole. The composition and properties of an element or compound are uniform throughout a given sample and from one sample to another.

12. State of matter • Commonly, a given kind of matter exists in different physical forms under different conditions. For example: water exists as ice (solid), as liquid water and as steam (gaseous water). • Solid:the form of matter characterized by rigidity; a solid is relatively incompressible and has fixed shape and volume. • Liquid:the form of matter that is a relatively incompressible fluid; a liquid has a fixed volume but no fixed shape. • Gas: the form of matter that is an easily compressible fluid; a given quantity of gas will fit into a container of any size and shape. The three form of matter-solid, liquid, and gar-are referred to as the state of matter.

13. Chemistry and the Elements Everything you see around you is formed from one or more of 115 presently known elements. Only 92 of the 115 presently known elements occur naturally. Chemical Symbol. For simplicity, chemists referred to specific elements usingone or two- letter abbreviations of the name of the element known as Chemical symbols.The first letter is capitalized; E.g. carbon, C; neon Ne and silicon, Si. Some elements have symbols based on their Latin name, E.g. Fe iron (ferrum), Pb lead (plumbum), Na sodium (natrium for sodium carbonate), K potassium (kalium for potassium carbonate). The symbol for tungsten, W, is based on the German wolfram.

14. Experimentation and Measure Measurement is the comparison of a physical quantity to be measured with a unit of measurement-that is a fixed standard of measurement Measurement of Matter: SI (Metric) Units ~ is a decimal system. Quantities differing from the base unit by powers of ten are noted by the use of prefixes. Table 1 The seven fundamental units of measure Physical quantity Name of unit Abbreviation Mass kilogram kg Length meter m Temperature Kelvin K Amount of substance mole mol Time second s Electric current ampere A Luminous intensity candela cd_______

15. Table 3 Some SI Prefixes Multiple Prefix__ Symbol__ 1015 peta P 1012 tera T 109 giga G 106 mega M 103 kilo k 102 hecto h 10 deca da 10-1 deci d 10-2 centi c 10-3 milli m You need to memorize the most 10-6 micro  common SI prefixessuch as G, 10-9 nano n M, k, d, c, m, , n and p. 10-12 pico p 10-15 femto f_______

16. Fundamental Units You are required to know all the conversion Mass ~ describes the quantity of matter in an object. An object has a fixed mass (m), which is independent of where or how the mass is measured (Stated more scientifically, matter is anything that has mass). Mass is measured in SI units by the kilogram. 1 kg =1,000 g = 1,000,000 mg = 1,000,000,000 g Relation between mass and weight Weight is the force of gravity on an object. It is directly proportionalto mass. W = gm * Mass is a physical property that measures the amount of matter in an object, whereas weight measures the pull of gravity on an object.

17. Length The meter (m) is the standard unit of length in the SI system. 1 m = 100 cm = 1,000 mm = 1,000,000 m Temperature To establish a temperature scale, we arbitrarily set certain fixed points and temperature increments called degree. Two commonly used fixed points are the temperature at which ice melt and the temperature at which water boils. The SI temperature scale, called Kelvin scale, assigns a value of zero to the coldest possible temperature, 273.15C, sometime called absolute zero. 0K = 273.15C. A comparison of the Kelvin, Celsius, and Fahrenheit temperature scale. Melting Point Boiling point of water Fahrenheit 32F 212F Celsius 0C 100C Kelvin 273 K 373 K Celsius from Fahrenheit (C) = 5/9[t (F) –32] Kelvin from Celsius T (K) = t (C) + 273.15

18. Derived Units Area (m2), volume (m3), density (kg/m3), speed (m/s) etc are some of the derived quantities that you are familiar with. They are derived quantities rather than fundamental quantities because they can be expressed using one or more of the seven base units. Volume, the amount of space occupied by an object, is measured in SI units by the cubic meter (m3), defined as the amount of space occupied by a cubic 1 meter long on each edge.

19. Density Density is the intensity properties that relates the mass of an object to its volume. Density = Mass (g)_______ Volume (mL or cm3) Because most substances change in volume when heated or cooled, densities are temperature-dependent. E.g. the density of water at 3.98C is 1.000g /mL and at 100C is 0.9584 g/mL as the volume expand. Although most substances expand when heated and contract when cooled, water behaves differently. Water contract when cooled from 3.98C to 0C.

20. Dimensional Analysis • Dimensional analysis is the method of calculation in which one carries along the units for quantities. Suppose we want to find the volume (V) of a cube, give l, the length of a side of the cube. Because V = l3, if l = 5.00cm, we find that V = (5cm)3 = 125cm3. There is no guesswork about the unit of volume here; it is cubic centimeter (cm3). Suppose, however, that we wish to express the volume in liter (L), a metric unit that equals 103 cubic centimeters. We can write this equality as • 1L = 103 cm3. • If we divide both sides of the equality by the right-hand quantity, we get • Observe that units are treated in the same way as algebraic quantities. Note too that the right-hand side now equals 1 and there are no units associated with it. Because it is always possible to multiply any quantity by 1 without changing that quantity, we can multiply our previous expression for volume by the factor 1 L/103 without changing the actual volume. We are changing only the way in which we express this volume:

21. The ratio 1 L/ 103 cm3 is called a conversion factor because it is a factor equal to 1 that converts a quantity expressed in one unit to one expressed in another unit. • E.g. Convert 8.45 kg to milligrams • 1 kg = 106mg • E.g. What is the density of a substance in g/mL, if a sample with a volume of 0.085 liters has a mass of 1700 mg?

22. Uncertainties in Scientific Measurement • All measurements are subject to error. • Types of Errors: • 1)   Systematic error: arise because to some extent, measuring instruments have built-in, or inherent, errors. • Random errors:they arise from intrinsic limitation in the sensitivity of the instrument and inability of observer to read a scientific instrument and give results that may be either too high or too low. • In talking about the degree of uncertainty in a measurement, we use the words accuracy and precision. • Accuracyrefers to how close to the true value a given measurement is. • Precisionrefers to how well a number of independent measurements agree with one another. • There is no relationship between accuracy and precision, since an experiment • can have small random errors and still give inaccurate results due to large • systematic errors.

23. Uncertainties in Scientific Measurement • All measurements are subject to error. • Types of Errors: • 1)   Systematic error: arise because to some extent, measuring instruments have built-in, or inherent, errors. • Random errors:they arise from intrinsic limitation in the sensitivity of the instrument and inability of observer to read a scientific instrument and give results that may be either too high or too low. • In talking about the degree of uncertainty in a measurement, we use the words accuracy and precision. • There is no relationship between accuracy and precision, since an experiment can have small random errors and still give inaccurate results due to large systematic errors.

24. Significant Figures The rules for determining significant figures (sig. fig.). 1 Zeros in the middle of a numbers are significant figures. E.g. 4023 mL has 4 significant figures. 2 Zeros at the beginning of a number are not significant; they act only to locate the decimal point. E.g. 0.00206L has 3 significant figures. The zeros to the left of 2 are not sig. fig. 3 Zeros at the end of a number and after decimal point are always significant. E. g. 2.200 g has 4 sig. fig. 4 Zeros at the end of number and before the decimal point may or may not be significant figures. In such cases we must deduce the number of sig. fig. from the statement of the problem. E. g. The statement “ 350,000 spectators lined the parade route” involves a number that probably has only two sig. fig. because it is obvious that no one actually counted the spectators.

25. 5 A useful rule of thumb to use for determining whether or not zeros are significant figures is that zeros are not sig. fig. if the zeros disappear when scientific notation is used. E.g.0.0197 = 1.97 x10-2 the zeros are not sig. fig.; 0.01090 = 1.090 x10-2 the first two zeros are not significant and the second two are. The number 0.004004500 has 7 sig. Fig.

26. Significant Figures in Numerical Calculations 1 Multiplication or Division The result of multiplicand or division, may contain only as many sig. fig. as the least precisely known quantity in the calculation. E.g. 14.79cm x 12.11cm x 5.05cm = 904cm (4 sig.fig) (4 sig. fig) (3 sig. fig) (3 fig. sig) If use scientific notion, E.g. (3.4 x 106)(4.2 x 103) = (3.4)(4.2) x 10(6+3) = 14.28 x 109 = 1.4 x 1010 (to 2 significant figures) the digit terms are multiplied in the normal way and the exponents are added. The end result is changed so that there is only one nonzero digit to the left of the decimal.

27. How to do calculations • 2 Addition and Subtraction: • The result of addition or subtraction must be expressed with the same digits beyond the decimal point as the quantity carries the smallest number of such digits. • E.g. 3.18 • + 0.01315 • 3.19 • When use scientific notion, All numbers are converted to the same power of 10, and the digit terms are added or subtracted. • E.g. (4.215 x 10-2) + (3.2 x 10-4) = (4.215 x 10-2) + (0.032 x 10-2) = 4.247 x 10-2 • 3 Exact numbers can be considered to have an unlimited number of sig. fig. There are two situation when a quantity appearing in a calculation may be exact. • By definition such as 1 in = 2.54 cm • Counting such as six face on a cube or tow hydrogen atoms in a water • molecule.

28. Rounding Numbers 1) If the first digit you remove is less than 5, round down by drop it and all following digits. E.g. 5.664525 5.66 when rounded to three sig. fig. 2) If the first digit you remove is greater than 5 or 5 followed by nonzero, round up by adding 1 to the digit on the left. E.g. 5.664525 5.7 when rounded to two sig. fig. 3) What happens if there is a 5 or 5 following by zeros? There is an arbitrary rule “round 5 to even”: If the number before the 5 is odd, round up. If the number before the 5 is even, let it be. The justification for this is that in the course of a series of many calculations, any rounding errors will be averaged out. E.g. Round the following number to 2 significant figures: (a) 2.35 x 102 (b) 2.45 x 102 (Answer: 2.4 x 102)

29. Summary of Key Concepts