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Unit 1: Fundamental Chemistry

CHEMISTRY : the science of materials, their composition and structure, and the changes they undergo. Unit 1: Fundamental Chemistry. Physical chemistry Analytical chemistry Inorganic chemistry Organic chemistry Biochemistry Nuclear chemistry. Six Branches of Chemistry.

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Unit 1: Fundamental Chemistry

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  1. CHEMISTRY: the science of materials, their composition and structure, and the changes they undergo. Unit 1: Fundamental Chemistry

  2. Physical chemistry Analytical chemistry Inorganic chemistry Organic chemistry Biochemistry Nuclear chemistry Six Branches of Chemistry

  3. To study the chemical system(s) and the CHANGES they undergo. Initial state →Final State A + B → AB Foundation of Chemistry

  4. All of our known facts in chemistry are based on the study of chemical reactions • REACTION: a chemical change in which a new substance is formed. • NaCl + AgOH → AgCl + NaOH • Left side called the REACTANTS • Right side called the PRODUCTS

  5. An important objective of science: • Relate properties of Large samples of matter (called macroscopic) to the individual atom (microscopic)‏

  6. Scientific Method – used all the time • Step 1: Making Observations • Two types of Observations: • QUALITATIVE: a descriptive term. “Your shirt is red,” “The solution was bubbling and was pink”, “The water is a liquid at room temperature.” • QUANTITATIVE: a quantitative observation is called a MEASUREMENT. “The pressure was 1 atm”

  7. Scientific Method • Step 2: Looking for patterns in the observations • Usually results in the formulation of a natural law • NATURAL LAW: a statement that expresses generally observed behavior. A natural law is often expressed as a math formula • Ideal Gas Law: pV = nRT

  8. Scientific Method • Step 3: Formulating Theories • THEORY: also called a model. It consists of a set of assumptions put forth to explain the observations. Back in the day we called this a hypothesis. • REMEMBER: an observation is a FACT • A theory is an interpretation (can be wrong!!)‏

  9. Scientific Method • Step 4: Experiment to Test Theories • Experiments may and usually do lead to modified or changed theories.

  10. Scientific Method Observations ↓ LAW ↓ Theory ← Modify Theory ↓ ↗ Test Theory/Experiment

  11. Units of Measurement • A MEASUREMENT consists of two parts: a NUMBER and a UNIT. Both must be present. • There are two types of units.

  12. FUNDAMENTAL UNITS • These are units upon which all other units are based. • METER – length • GRAM – measures mass. Mass – quantity of matter that a body possesses. • WEIGHT – measure of the earth's gravitational field. • Remember: your mass is FIXED but your weight varies depending on your position from the center of the earth

  13. FUNDAMENTAL UNITS • SECOND – measures time (based on the vibration of Cesium-133)‏ • MOLE – measures the number of particles and is equal to 6.02 x 10 to the 23rd. • KELVIN – named after Lord Kelvin. Kelvin temperature scale is based on absolute zero. • COULOMB – a quantity of electrical charge

  14. DERIVED UNITS • Derived units are units based on fundamental units. There are lots and lots of derived units. Two examples: • VOLUME – 1 mL = 1 cubic centimeter = 1 gram (if water) • 1 Liter = 1 cubic decimeter • 1 Liter – 1000 cubic centimeters

  15. DERIVED UNITS • Density is another derived unit based on mass and volume • Density = Mass/Volume • Units for Density are grams/mL

  16. All measurements have some degree of uncertainty UNCERTAINTY IN MEASUREMENT

  17. 5 different nerdy honors chemistry students massed a sample of iron:student 1 = 16.18 gstudent 2 = 16.15 gstudent 3 = 16.19 gstudent 4 = 16.16 gstudent 5 = 16.15 gWhich decimal place do you think was likely to be rounded? Most exact?? Least exact??

  18. Significant Figures: certain digits and the first uncertain digit. (the real reason we need sig figs is to help us figure out which numbers are exact and which ones were rounded)Sig Figs are used mainly in the fields of physics and engineering.

  19. Measuring the sides of a square:area = (side)(side)area = (16.4 cm) (22.8 cm)area = 373.92 cm2 Look at the answer the calculator gives us. It is IMPOSSIBLE (for a plain orange pumpkin to become a golden carriage) to have an answer that is MORE accurate than our measurements – thus the need for sig figs in the physical sciences.

  20. Rules for Counting Sig Figs1. Non Zero IntegersNon zero integers always count as significant figures ie. 3.455 has 4 sig figs2. ZerosThere are THREE (really 4) rules for Zeros: 1. LEADING ZEROS - are zeros that precede all of the non-zero digits. They DO NOT count as sig figs. Note that leading zeros are always in a very small number) ie. The number 0.000456 has 3 sig figs. The leading zeros are not significant and are only there to simply indicate the position of the decimal point.

  21. 2. CAPTIVE ZEROS (OR SANDWICHED ZEROS) – are zeros between two non-zero digits. They are ALWAYS significant. ie. the number 1.008 has FOUR sig figs.3. TRAILING ZEROS – these are zeros at the right end of the number. There are two rules for trailing zeros:a. They ARE significant if the number has a decimal point. b. They are NOT significant if there is no decimal point. The number 100 has only 1 sig fig The number 1.00 x 102 has three sig figs The number 2306.00 has six sig figs

  22. Now you have fun and practice!!Determine the # of sig figs in:236 678.091.008 0.0000567098,900 0.00509080700

  23. Rules for Math and Sig Figs1. Addition/SubtractionThe result has the same number of decimal places as the least precise measurement. HINT: Count Decimal Places. Ie. 12.11 18.0 ← here is the limiting term-only 1 dec. place + 1.013 31.123But the CORRECTED answer with one decimal place would be 31.1

  24. 2. Multiplication/Division • The number of sig figs in the product/quotient is the same as the number of sig figs in the LEAST precise measurement. HINT: count the sig figs • ie. (4.56)(1.4) = 6.38 …but you can’t really have an answer with MORE sig figs than the number with the least…so the CORRECTED answer would be 6.4. (3 sig figs)(2 sig figs) = 2 sig figs

  25. Now YOU get to have fun!!! Give the answers to the correct # of sig figs • 2.33 + 4.5 + 8.00 + 8 = • 9.010 ÷ 3.7 = • 9.0 – 3.888 = • (5.66)(1.00)(2.00)(0.0006) =

  26. Precision vs. Accuracy • These concepts are often confused!!! • ACCURACY – denotes the nearness of a measurement to its accepted value. • ie. beaker mass = 19.0 grams • Your mass = 19.9 g, 24.1 g, and 13.6 g. • How was this student's accuracy?????

  27. PRECISION • An agreement between the numerical values of a set of measurements that have been made the same way (think CONSISTENCY!!)‏ • ie. Beaker mass = 19.0 g • Your mass = 14.1 g, 14.0 g, and 14.1 g • How was your precision? • How was your accuracy? • Dart example

  28. Percentage Error Formula % error = |experimental - actual| x 100 actual

  29. PERCENTAGE ERROR • A student was calculating the % of lead (Pb) in the water at Xenia High School in the drinking fountains. She came up with the following values: 16.12%, 16.14%, 16.12% and 16.13%. The average value was 16.13%. The correct value according to my scientific calculations was 16.49%. • What can be said about the accuracy? • What can be said about the precision? • Calculate the % error.

  30. SCIENTIFIC NOTATION • Also called exponential notation • Move the decimal to the left – exponent is larger and POSITIVE!! For example the speed of light is 30,000,000,000 cm/sec. Put into scientific notation. • Move the decimal to the right – exponent is smaller and negative. For example, put 0.000496 m into scientific notation.

  31. Fun with Scientific Notation (you junior science nerd you!) (9.24 x 1016 )(6.12 x 1014 ) = 1.96 x 10-8 /2.47 x 10-4 =

  32. DIMENSIONAL ANALYSIS • An exciting and fun way of working problems by using the UNITS to help us along the way. • Defined: a method of changing units. • (use the metric/English or English/metric charts)

  33. Doing conversions using Dimensional Analysis:Convert 14 Kg to lbs:Convert 16.9 in to cm:

  34. Convert 8 years to seconds:Convert 3 gallons to mLConvert 8 mph to cm/second

  35. Convert 4.66 in2 to cm2 Convert 98.77 yd3 to m3Convert 4.5 m to KmConvert 0.455 mL to cL

  36. TEMPERATURE • Three systems: Celsius, Kelvin, Fahrenheit • For Water: BP = 212ºF, 100ºC, 373K • For Water: FP = 32ºF, 0ºC, 273 K • Special Formulas °C = (°F – 32)5/9 °F = (°C X 9/5) + 32 K = °C + 273 °C = K – 273

  37. Normal body temperature is 98.6ºF. Convert to ºC and Kelvin.

  38. Liquid nitrogen has a boiling point of 77K. Convert this to ºF.

  39. DENSITY • Density is defined as the mass of a substance per unit volume. • Density = M/V • This formula can also be solved for mass and volume. • M = • V =

  40. The mass of Al is 14.2 g and the volume is 6.9 mL. Find the density Calculate the % error (the actual density is 2.7 g/mL)‏

  41. The density of Fe is 7.86 g/mL. You have 29 grams of Fe. How many mls will it occupy?

  42. Percentage Problems • Percentage is a part/whole x 100 • Given: 82 g of a metallic powder. It consists of 31 g of Zn, 3 g Ag, and 48 g of Sn. Find the % of each.

  43. Flow Chart of Matter MATTER Pure Substance Mixtures Heterogeneous Mixture Homogeneous Mixture

  44. Pure Substance Compound Elements Atoms Nucleus Electrons Neutrons Protons Quarks Quarks

  45. SEPARATION METHODS • There are Nine (9) ways to separate mixtures in the lab. Some of these are based on physical properties and some of these are based on chemical properties.

  46. 1. FILTRATION • Separates based on insoluble/soluble properties • FILTRATE: the soluble substance or liquid that passes through the filter paper • RESIDUE: the insoluble chunky “stuff” that remains in the filter paper. • Filtration is a great way to separate a SUSPENSION: where the particles are larger than molecular size in the liquid.

  47. SOLUBILITY • Solubility in water is a physical property • SOLUBLE: dissolves • INSOLUBLE: remains undissolved

  48. 2. DECANTING • Decanting is used to separate a coarse suspension of liquid and dense, insoluble solids. • Decanting simply means “to pour off” • Yes...even you can do this separation technique!

  49. 3. SIMPLE DISTILLATION • Distillation is used to separate solid solute from liquid solvent • Distillation is used to make distilled water and many different alcohol products. • Distillation is based on a phase difference (the solid remains in the original flask and the liquid boils, evaporates, then condenses and drips into a new container in a purified form)‏

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