1. Chemistry Module 1

2. About Chemistry Chemistry is the scientific study of matter, including its properties, its composition and its reactions. There are many branches of chemistry: –Organic chemistry: –Inorganic chemistry: –Analytical chemistry: –Physical chemistry: –Biochemistry: --substances with carbon --substances without carbon --composition / identification --Substances in living things --theoretical basis of chemistry

3. (Optional Enrichment) Chemistry evolved from Alchemy, the medieval study of “magical” properties of materials –Best known alchemist: Nicholas Flamel rumoured to have found the secrets of the philosopher’s stone and the elixir of life About 1600 alchemy began to disappear, and be replaced by the more systematic approach of chemistry. Early chemists include: –Robert Boyle (who worked on the gas laws) –Antoine Lavoisier (who found laws of fixed proportions) –John Dalton (who first described atoms) –Joseph Priestly (who discovered oxygen)

4. Review of Important Points from Previous Science Courses Properties of matter –Physical Properties: Properties that can be examined without reacting a material. Examination may cause physical changes, such as change of state or form. –Chemical Properties Properties that can only be determined by reacting a material with another material (which usually changes or “destroys” it) –Characteristic Properties Properties that apply to a single material or a small group of similar materials. They help us identify a material. –Non-characteristic Properties Properties that are less helpful in identifying a material because the apply to many different substances.

5. Pure substances vs. mixtures Pure substances are substances that are the same throughout. –Theoretically, all the particles in a pure substance are the same. There are two types of pure substance –Elements: usually composed of atoms* –Compounds: usually composed of molecules* Most materials are mixtures. They contain two or more types of particle mixed together. –solutions, suspensions, colloids, emulsions and most composite solid materials are mixtures. *As we shall see, this is a slight over-simplification that ignores ionic compounds.

6. Important Physical Properties of pure substances Density: The ratio of the mass of a material to its volume. Melting point: The temperature at which a pure substance will melt (for pure substances, this is the same as freezing point) Boiling point: The temperature at which a pure substance will boil (for pure substances this is the same as the condensation point).

7. Classification: All Matter (solids, liquids, gases, plasma) Pure Substances Mixtures Elements Ex. gold Compounds Ex. water Hetero- geneous Mixtures Homo- geneous Mixtures Examples of physical separation include: Filtration, evaporation, distillation, magnetic separation, chromatography, settling, decantation, flotation, sorting, screening. Types of chemical separation include: Electrolysis, decomposition And precipitation. Solutions colloids emulsions suspensions Everything in the world that has a mass and takes up space is called “matter” Matter can be classified as:

8. Changes Physical Changes DO NOT alter the nature of the substance, for example: –Change of form (tearing, breaking, crushing) –Change of state (melting, freezing, boiling)* –Change of mixture (blending, dissolving)* The molecules do not change during a physical change. Chemical changes DO alter the substance. –Decomposition -Combustion –Synthesis-precipitation –Oxidation-electrolysis –Single or double replacement The molecules become different in a chemical change * note: sometimes attempting to cause a physical change may trigger a chemical change.

9. Summary of Lesson 1 –Chemistry is the study of matter, its properties, compostition and reactions. Chemistry includes: Organic chemistry  Inorganic chemistry Analytical chemistry  Physical chemistry Biochemistry –Matter has properties Physical propertiesCharacteristic properties Chemical propertiesNon-characteristic –Elements and Compounds are pure substances –All other substances are mixtures –Physical changes do not alter the composition Change of form: tearing, crushing, breaking Change of state: melting, freezing, boiling Change of mixture: dissolving –Chemical changes do alter the composition Combustion, precipitation, decompostition etc.

10. Element Song, Version 1Element Song Element Song, Version 2Element Song Element Song, Version 3Element Song

11. Assignment Read chapter 1 of Addison-Wesley Chemistry (pp. 1-11) Answer the following questions in your assignments book: –Addison-Wesley Chemistry pp.17-18 Questions # 9-20

12. Sample Answers 9. Chemistry is the branch of science that studies matter, as well as the composition of substances and changes they undergo. 10. Five divisions of chemistry include: –Organic chemistry –Inorganic chemistry –Analytical chemistry –Physical chemistry –Biochemistry

13. 11. A hypothesis is a descriptive model or trial explanation, formed after observation. A theory is a hypothesis that has been thoroughly tested. A law is a statement that summarizes the results of observations. 12. Experiments are used to test a hypothesis, or to gather more data to make a better hypothesis.

14. 13. a, b and e, or more completely: –Matter: concrete, propanone vapour, air –Not matter: heat, sound 14. some physical properties of a nail –Mass- volume-length –Density- colour-magnetism –Diameter- conductivity-hardess –Melting point(pick four)

15. 15. in which state of matter do each of the following occur at room temperature? –Diamond (solid)Mercury (liquid) –Oxygen (gas)Clay (solid) –Cooking oil (liquid)neon (gas) 16. –A) incompressiblesolid, liquid –B) indefinite shapeliquid, gas –C) definite volumesolid, liquid –D) flowsgas, liquid

16. 17. how to physically separate: –A) iron filings and salt could be separated by using a magnet, or by dissolving the salt in water and filtering off the iron filings –B) Salt and water could be separated by evaporation 18. Physical properties that distinguish: –A) water and rubbing alcohol: density, odor, boiling point (2 of these) –B) Gold and aluminum: density, colour, conductivity –C) Helium and oxygen: density, solubility, diffusion rate

17. 19. A homogeneous mixture is uniform in composition (ie. it appears to be the same throughout). A heterogeneous mixture is not uniform. 20. Some methods of separating mixtures include evaporation, distillation, dissolution and filtration.

18. Module 1, Lesson 2 States of Matter Phases (optional material) Symbols Energy Conservation of Energy Identifying Chemical Reactions Chemical equations Conservation of Mass This is an outline of today’s lesson, not the notes

19. States of matter Solid –definite shape-definite volume –Incompressible-does not flow Liquid –Variable shape-definite volume –Incompressible-fluid (can flow) Gas –Variable shape-variable volume –Compressible-fluid (can flow) Exotic states of matter: (optional enrichment) Plasma: At very high temperatures electrons separate from gases and they glow. Superfluid: At very cold temperatures helium will flow in ways normal liquids don’t. Extreme pressures (Optional enrichment) Although liquids and solids are said to be incompressible under ordinary conditions, at extreme pressures (thousands of atmospheres) they may actually compress slightly. Some scientists theorize that at extreme pressures (billions of atmospheres) all matter might compress into an exotic state nicknamed “neutronium”.

20. Phases (Optional enrichment) The term “phase” is sometimes used as a synonym for “state”, but phases are more general than states. Phases are portions of any chemical system that have uniform composition and properties. The most common phases are: –Solid-liquid -gas (just like states) But phases can also include: –Solute-gel-crystal –Colloid-vapour-etc. (which technically speaking are not states of matter) A mixture can have several phases but appear to exhibit only one state –Oil on water has two phases, but both are liquid. –Diamonds in graphite have two phases but both are solid. Another difference between “state” and “phase” is that the term state applies only to pure substances (ie pure elements or pure compounds) while the term phase can apply to portions of a mixture.

21. Chemical Symbols Each element has a symbol By now, you should know the symbols of common elements, including: HHeLiBeBCNO FNeNaMgAl SiPS ClArKCaBrFeCuZn INiCoAgAuHgPb

22. Energy Energy is the ability to do work There are many types of energy: –Heat, light, sound, electricity, chemical, nuclear, thermal, But to a chemist, the two main divisions of energy are: –Kinetic: Energy of motion (active energy) –Potential: Energy of position or composition. (passive or hidden energy)

23. Law of Conservation of Energy “In any physical or chemical process, energy is neither created nor destroyed.” Energy can, however, be changed from one form to another –For example, from potential energy to kinetic energy or vice-versa. “In any physical or chemical process, energy is neither created nor destroyed.” Energy can, however, be changed from one form to another –For example, from potential energy to kinetic energy or vice-versa.

24. Chemical Reactions In a chemical change or “reaction” one or more substances are changed into new substances. We say that the composition has changed. The materials we started with were called reactants The new materials produced are called the products.

25. Reactants  Products For example: Hydrogen + Oxygen  Water (2H 2 + O 2  2 H 2 O ) Hydrogen and oxygen are reactants Water is the product.

26. Identifying Chemical Changes How do you identify if a change has been chemical instead of physical? These are some of the indications –Combustion: sudden release of heat or flames –Precipitation: a solid separates from the mixture of two solutions –Effervescence: bubbles of gas forming in a solution –Colour change: a significant change of colour.

27. Law of Conservation of Mass “In any physical or chemical process, mass (matter) is neither created nor destroyed.” The mass of all the products must equal the mass of all the reactants. –Sometimes it is hard to show this, because some products may escape the container. “In any physical or chemical process, mass (matter) is neither created nor destroyed.” The mass of all the products must equal the mass of all the reactants. –Sometimes it is hard to show this, because some products may escape the container.

28. Summary of Lesson 2 –Three important states of matter are: Solid: definite shape, definite volume, incompressible Liquid: indefinite shape, definite volume, incompressible Gas: indefinite shape, indefinite volume, compressible –You should know symbols of common elements –Energy is the ability to do work. It includes Kinetic energy: the energy of motion Potential energy: energy of position or composition –Law of conservation of energy In reactions, Energy is neither created nor destroyed. –Chemical reactions change substances Know what reactants & products are. Know how to identify a chemical change. –Law of conservation of mass In a reaction, mass is neither created nor destroyed.

29. Assignment #2 Read the rest of chapter 1 (pp. 11-16) Answer questions #21-29 from page 17 & 18 in your assignments folder. If you haven’t done questions #9-20, do them too.

30. 21. Identify the following as homogeneous or heterogeneous: –A) milk: (arguable) Homogeneous or heterogeneous* –why? Real milk, straight from the cow, separates into cream, water, and milk solids. Skim milk and homogenized milk do not. Technically, milk is an emulsion. A mixture between homogeneous and heterogeneous, but closer to heterogeneous. –B) glass: homogeneous mixture –C) Table sugar: homogeneous compound –D) river water: (arguable) heterogeneous* mixture (*at microscopic level. At the visible level, filtered river water looks homogeneous) –E) cough syrup: homogeneous mixture –F) Nitrogen: homogeneous pure element *do not mark these two answers wrong, just add the opposing argument.

31. 22. Two ways to distinguish a compound from an element are: –A compound can be broken down into elements by decomposition. –Compounds contain two or more different types of atom

32. 23. Identify the following element, compound or mixture. –A) milk: mixture (water, milkfat, milk solids) –B) glass: mixture* (72% SiO 2, 13%Na 2 O, 15% other) This one is very technical. Most people mistakenly classify glass as a compound. One type of expensive glass (fused silica) is a pure compound: 100% SiO 2 ). –C) Table sugar: compound (C 12 H 22 O 11 ) –D) river water: mixture (H 2 O, minerals, impurities ) –E) cough syrup: mixture (alcohol, water, medicine*) The medicine could be dextromethorphan, codeine, or antihistamine, depending on the brand. Some also contain sugar, flavour and colour. –F) Nitrogen: element (N 2 )

33. 24. The chemical symbols are: –Copper: CuSilver: Ag –Oxygen: OSodium: Na –Phosphorus: PHelium: He 25. The elements found in each are: –NH 4 Cl: Nitrogen, Hydrogen, Chlorine –KMnO 3 : Potassium, Manganese, Oxygen –C 2 H 7 OH: Carbon, Hydrogen, Oxygen –CaI 2 : Calcium, Iodine.

34. 26. Kinetic energy is the energy of motion (active energy), potential energy is the energy of position or composition (hidden energy). 27. Examples of types of energy (choose 5) –Nuclear-hydro-chemical –Radiant-electrical-mechanical –Thermal-solar-etc.

35. 28. The law of conservation of energy says that energy cannot be created or destroyed during a chemical reaction. 29. Classify as physical or chemical change: –Bending wire: physical –Burning coal: chemical –Cooking steak: chemical –Cutting grass: physical

36. Module 1 Lesson #3 Overview of SI Metric system –Prefixes –Length –Volume –Mass –Temperature

37. The SI metric system Resulted from an attempt to make a sensible measurement system based on powers of ten The metre was originally defined as 1/10000000 of the distance from the equator to the north pole. All the other units were then derived from the metre.

38. Metric Prefixes Yotta10 24 Superclustersdeci1/10hand (100 zetta)10 23 centi1/100fingernail (10 zetta)10 22 milli1/1000sand Zetta10 21 Galaxy(100 micro)10 -4 (100 exa)10 20 (10 micro)10 -5 (10 exa)10 19 micro10 -6 bacteria Exa10 18 nearby stars (100 nano)10 -7 (100 peta)10 17 (10 nano) 10 -8 (10 peta)10 16 nano 10 -9 molecule Peta10 15 Solar system (100 pico)10 -10 (100 tera)10 14 (10 pico) 10 -11 (10 tera)10 13 pico 10 -12 atom Tera10 12 Inner planets (100 femto)10 -13 (100 giga)10 11 (10 femto)10 -14 (10 giga)10 10 femto10 -15 proton Giga10 9 Earth/moon (100 atto)10 -16 (100 mega)10 8 (10 atto)10 -17 (10 mega)10 7 atto10 -18 electron? Mega10 6 East coast(100 zepto)10 -19 (100 kilo)10 5 (10 zepto)10 -20 (10 kilo)10 4 zepto10 -21 quark?? Kilo1000Town(100 yocto)10 -22 Hecta100football field(10 yocto)10 -23 Deca10Elephantyocto10 -24 strings???

39. Common metric units & prefixes Mega- M - Kilo- kprefixes (large) Hecta- h Deca- da ------ metre, litre, gram, etc.  units Deci- d Centi-c Milli-m -prefixes (small) - Micro- μ (or u)* *if your keyboard does not support Greek letters

40. Length Unit of length is the metre (also spelled meter) It can be divided into –Decimetres –Centimetres –Millimetres metredecimetrecentimetremillimetre

41. Volume A cube 0.1m per side (a cubic decimetre) is defined to have a volume of one litre –1 cubic decimetre = 1 Litre –1 cubic centimetre = 1 mL –1 cubic metre = 1000 litres = 1 kilolitre The symbol for litre can be L, l or curly l, but in Canada the “L” is preferred.

42. Mass The mass of one litre of pure water at standard conditions (4°C) is defined to be one kilogram = 1000 g 1 litre of water = 1 kg 1 mL of water = 1 g 1 cubic metre of water = 1000 kg = 1 Mg = 1 tonne Since it awkward to haul around a litre of distilled water, and since the purity of local water is questionable, a prototype kilogram was made of platinum (IPK) and stored in the archives of France. It is still used to calibrate balances around the world.

43. Temperature Degrees Fahrenheit ( ° F) NOT to be used in Chemistry! Freezing point32°F Room temp68 °F Body temperature99 °F Boiling point212 °F Degrees Celsius ( ° C) A.KA. Centigrade Often used in Chemistry Freezing point0 °C Room temperature20 °C Body temperature37 °C Boiling point100 °C Kelvins (K), formerly: °K or Absolute °A ) The Best for Chemistry, especially with gas laws. Freezing point273 K Room temperature293 K Body temperature310 K Boiling point373 K KelvinCelsiusFahrenheit 473200392 453180356 433160320 413140284 393120248 373100212 35380176 33360140 3134099 2932068 273032 253-20-4 233-40 213-60-76 193-80-112 173-100-148 153-120-184 133-140-220 113-160-256 93-180-292 73-200-328 53-220-364 33-240-400 13-260-436 0 K-273 C-460 F Absolute zero Mercury freezes Water freezes Room temp. Body temp. Water boils

44. Conversions To convert :Use these steps:Example °C to KAdd 273.15 (for simplicity, we often leave out the.15) 20°C = 293 K K to °CSubtract 273.15 (again, we can drop the.15) 300 K = 27 °C °C to °Fmultiply by 9, then divide by 5, then add 3220°C = 68 °F °F to °CSubtract 32, then multiply by 5, then divide by 9212°F = 100°C °F to KSubtract 32, then multiply by 5, then divide by 9 Then add 273 Or go to Google and type one of the following: 20 C in K 300 K in C 20 C in F 212 F in C

45. Module 1: Lesson #4 Measurement Accuracy vs. Precision Significant Figures (Significant Digits) –In measurement –In calculations

46. Measurement. Measuring quantities is an important aspect of experimentation. Instruments used for measuring are seldom perfect. Each instrument has an amount of uncertainty or “error” Knowing the acceptable error helps set the reliability of a result. Acceptable error of several instruments Thermometer ± 0.2°C Balance ± 0.05g Graduated cylinders: 10 mL ± 0.1 mL 50 mL ± 0.5 mL 100 mL ± 1.0 mL

47. Accuracy vs. Precision Accuracy is how close an instrument’s reading is to the actual correct value Precision is how well an instrument reproduces a result –An instrument that is inaccurate but precise can often be adjusted to give better results. –An instrument that is imprecise will have a higher uncertainty or “error”. –An instrument that is imprecise and inaccurate should be discarded and replaced.

48. Significance It is misleading to write a result that implies more precision than was measured. To avoid excessive precision, the concept of significance was developed. –Results should never be written with more precision than the measurements that were used to calculate them.

49. Example of Excess Precision (discussion point) John wants to calculate the circumference of a cylindrical water tank. He measures the diameter as 2.55 m and then multiplies the measurement by pi (3.1415926535) 2.55 x 3.1415926535 = 8.011061266425 m This is an extremely misleading number. His measurement was nowhere near precise enough to support this result. He must round this off to a more reasonable result. 8.01 m

50. Make a judgement call of how accurate your results are, based on your instruments. –For example, if your instrument allows you to measure a value to the nearest tenth millilitre (ie. Its acceptable error is  0.1mL) then you can record values like: 3.9 mL or 4.0 mL or 4.1 mL –You measured to the nearest 0.1 mL Don’t write 4 mL –it suggests that you were not precise enough Don’t write 4.00 mL –it implies more precision than you actually measured Your own measurements

51. Interpreting Measurements If you see a measurement you may not know how precise it is. We use the concept of significant digits to determine its precision. –The number of significant digits determines the precision of the measurement and tells how much you can safely round the results. –Remember that a number with too many digits is just as misleading as one with too few!

52. Rules of Significant Figures (for measurements made by someone else) The digits 1 through 9 are ALWAYS treated as significant in a written measurement. Zeros between significant digits are ALWAYS significant. Leading zeros (in front of a number) are NEVER significant. Trailing zeros ARE significant unless the person who recorded them was estimating. –If there is a decimal point they ARE significant –If there is no decimal then you must use your judgement. Was the person estimating or not? (see next slide)

53. Ambiguous Digits If a “whole” number ends in zero, the situation is ambiguous. 5280 ft3 SD or 4 SD? 20 000 m1 SD or 5 SD? Unless we know who took the measurement, we can’t tell if the trailing zeros are significant or not. How do we handle this? The textbook tells us to call these zeros significant. Most other chemistry books say to treat them as insignificant. What’s a chemist to do? Use your judgement.

54. Using Your Judgement First look for clues. The word “about” or “approximately” in the description of the measurement tells you to use the lower number of significant digits. The word “exactly” or “precisely” tells you to count all the digits. The context or type of source may help. Popular scientific articles and newspapers usually round off the number, so use the lower number of digits Textbooks and professional journals are usually more accurate, so use all digits. If all the other measurements are very precise, then assume the ambiguous measurement is too. If still in doubt, use all the digits.

55. Example 003.50270 75000 0.001010 All digits 1 through 9 are significant All zeros between significant digits are significant Leading zeros are never significant Trailing zeros with a decimal point are significant Trailing zeros with no decimal are sometimes significant (use judgement) 6 significant digits 5 significant digits (if accurate) 2 significant digits (if estimated) 4 significant digits

56. Calculations with significant digits The result of a calculation can have no more significant digits than the WORST measurement! You are the weakest link! Good bye!

57. Multiplying and Dividing: Do the calculation, then round the answer so it has the same number of significant digits as the worst measurement. 2.514 cm x 3.1 cm = 7.7934  7.8 cm 2 528 g ÷ 25 mL =21.12  21 g/mL However: if you multiply or divide by a number that has no units (ie. An integer used to double or triple a result) it does not reduce the number of significant digits 23.4 g x 2 = 46.8 g  doesn’t change.

58. Adding and subtracting Make sure the units are the same before adding or subtracting (convert metric units if necessary) Do the addition or subtraction, lining up decimals Round the answer to match the number with the fewest decimals (or fewest significant digits if there is no decimal marker) 23.45 cm + 4.5123 cm 27.9623 cm  27.96 cm

59. Examples 5.3 cm x 4.33 cm = 22.949 cm 2  23 cm 2 5.8798 mL ÷ 3 g = 1.95993 g/mL  2 g/mL 4.3576 m x 2 = 8.7152 m One significant digit 2 significant digits 5 significant digits Not a measurment

60. What about scientific notation? The digits are significant (following the normal rules of significant figures)rules of significant figures The 10 and exponents are not significant. Example – 6.02 x 10 23 – 1.3200 x 10 -7 Has 3 significant digits Has 5 significant digits Significant Not Significant x

61. What about exact numbers In the unlikely event that we have a measurement of exactly 230 000 objects, how should we represent it? Remember, someone might interpret it as having only 2 significant digits if they thought we were estimating. –one way:“230 000 exactly” (verbal description) –A better way:“2.30000 x 10 5 ” (scientific notation) Write “exactly” if measured or “about” if not Convert measurments to scientific notation When in doubt, convert your answers to scientific notation!

62. On tests and assignments, assume that all trailing zeros are significant, unless you see the word “about” or “approximately” in the question. It is my intention to never give you a problem on a test or examination that has an answer with less than 3 significant figures.

63. Summary of Lesson #3 Precision and accuracy are important when reading instruments –Knowing the acceptable error of an instrument helps you know how much precision to record. A result of a calculation must never be more precise than the worst measurement used in the calculation. Rules of significant figures can help us decide how to correctly round off the results of a calculation. (see Rules of Significant Figures earlier in this lesson)Rules of Significant Figures When in doubt, convert your answers to scientific notation.

64. Assignment #3 Do the sheet “Significant Figures”

65. Answers a) 165 283.78b) 165 283.8 c) 165 280d) 165 300 e) 165 000f)200 000 A) 5280 feetb) 007 A C) 22.40 md) 23001 mm E) 4000 kgf)0.000745 L

66. a)789.30 m(5)b)7400 mL (2) c)0.04 V(1)d) 73.2469 cm (6) e)0.4320 g(4)f) 503 mm(3) a)5b) 3 c)5d) 3

67. A) 23 m 2 b) 200 V C) 31 cmd) 2.00 g/mL E) 15 m 2 f) 91.4 m Always Never Sometimes

68. Module 1: Lesson #4 Problem Solving in Chemistry Problem solving techniques (self-review) Conversion factors (self-review) Dimensional analysis (lesson)

69. Read pages 49-52 In your “notes” book, list the steps that are suggested for solving chemistry problems Try the five problems on p.51-52. You may check your own answers by looking at the solutions (see p.719). This is for your own practice.

70. Dimensional Analysis AKA: unit analysis Dimensional analysis uses the units that are part of the measurements to help analyze and solve a problem. –Adding and subtracting units –Multiplying and dividing units –Cancelling units –Simplifying units –Comparing the units of the answer to the expected units can determine if the problem is done correctly.

71. Rule 1: Adding and subtracting units You can only add or subtract measurements that have the same unit –Make sure that you have converted quantities to the same units. If one measurement is in litres, and the other in millilitres, you must change one of them. 2.5 L + 250 mL  2500mL + 250 mL = 2750mL or  2.5 L + 0.25 L = 2.75 L

72. Rule 2: Multiplying and Dividing Whenever you multiply units, place a dot () between them: 10 N x 5 m = 50 Nm(newton-metres) Whenever you divide units, place a slash between them 50 g ÷ 10 mL = 5 g/mL (grams per millilitre) or 8 mol = 4 mol/L(moles per litre) 2L

73. Rule 3: Cancelling Units You may cancel units if the same unit occurs in a numerator as in a denominator. A unit that is immediately after a slash is in a denominator: m/sec = m sec _ = 1800 km/h

74. Rule 4 Look for units that can be simplified: 5 A x 20 Ω = 100 A·Ω = 100 V A few you may remember: Amperes x Ohms =VoltsA Ω = V meters x metres = square metres mm = m 2 Ohms x Volts = WattsΩV = W Compare the units to what is expected.

75. Problem: Chili You need 600 servings of chili 10 servings of chili needs 2 tsp. of chili powder How much chili powder will you need? This problem is easy to solve, but let’s show the problem with dimensional analysis = 120 tsp Tsp · servings servings

76. Don’t Copy… but be aware. Dimensional analysis is not an end in itself. You will never be tested on it directly It is a tool to help you solve problems and give the correct units of your answer In tests and exams you must give units for every answer. Not giving the correct unit for an answer may cost up to 25% of the value of the correct answer.

77. Module 1, Lesson #3 Models of the Atom I am the very model of a modern major element (the Elements Song by Tom Leher)

78. Models of the Atom This section is mostly review from PSC416 with a few new bits added at the end. Since the time of Aristotle and Democritus philosophers and scientists have tried to determine what matter is made of. Since the particles of matter are too small to see, we use models (pictures and other representations) to try to understand them.

79. Dalton John Dalton was the first modern scientist to propose that matter was made of tiny particles. –The philosopher Democritus had suggested this two thousand years earlier, but had never produced a model to explain. Dalton called these tiny particles “atoms”

80. Highlights of Dalton’s Model Remember– Dalton said: –All matter is made of atoms –Each different element has a different type of atom. There are many different elements (we have now named 109, have proven the existence of 112, and have evidence of up to 118) –Atoms of elements can combine to form molecules of compounds.

81. Thompson & Rutherford J.J. Thompson discovered the existence of electrons, particles smaller than, and found inside atoms. He was the first to suggest that the atom contained other particles. Rutherford discovered that most of the mass of an atom is concentrated in the center. He stated that the atom has a dense nucleus in the center. Rutherford’s model

82. Bohr’s Discovery Niels Bohr studied the wavelength of light given off by excited atoms, and determined that electrons “orbit” the atom in different energy levels or “shells”. By combining the idea of a central nucleus (Rutherford) with the idea of orbiting electrons (Bohr) we developed the Bohr- Rutherford model of the atom

83. Bohr-Rutherford Model Nucleus (protons ) (in later versions also neutrons) electron Energy Level or “Shell”

84. Simplified Bohr-Rutherford Shell maximums: (2n 2 ) –2e -, 8e -, 18e -, 32e -, 50e - … But… A shell does not have to be completely filled! For example: –2e -, 8e -, 8e -, 2e - is the usual arrangement for calcium, NOT: 2e -,8e -,10e - Heavy circle represents The NUCLEUS Semi-circle represents First electron shell Semi-circle represents 2 nd electron shell Etc. 9p + 10n 0 2e - 7e - F

85. Why? (simplified answer) Atoms usually arrange themselves so that most shells can have one of “magic numbers”… 2,8,18, 32 etc, Ca: 2e, 8e, 10e vs. Ca: 2e, 8e, 8e, 2e YYNYYYY This side wins, because it has more Shells with a “magic number” This is still an oversimplification, but it is the best we can do unless You learn the Aufbau diagram!

87. Draw simplified Bohr-Rutherford diagrams of: –N –Mg –Cl –Ca

88. Atom Overview Atoms consist of: –Protons: positively charged particles located in the nucleus with a mass ≈ 1 amu ( 1859 / 1860 ) –Neutrons: neutral particles located in the nucleus with a mass = 1 amu –Electrons: negative particles, orbiting the atom with a mass ≈ 0 amu ( 1 / 1860 amu)

89. Module 1, Lesson #7 Advanced models of the atom Optional enrichment material. (advanced classes)

90. Modern Model of the Atom The Modern or “Electron Cloud” Model

91. Subshells (Orbitals) (optional enrichment) Each shell has one or more subshells or orbitals that look like clouds. Labelled s, p, d, or f based on their shape –s orbitals are spherical –p orbitals are “peanut” shaped –d orbitals may* be doughnut shaped –f orbitals are flower shaped. Each orbital can hold up to two rapidly moving electrons. *actually, some d orbitals look a lot like f orbitals.

92. Number of orbitals/electrons per shell Shell #spdfh,i,j Total Orbitals Total Electrons Shell 111 orbital 2e- Shell 2134 orbitals 8e- Shell 31359 orbitals 18e- Shell 4135716 orbs 32e- Shell 513579 25 theory 16 typical 50 theory 32 typical Shell 61357 9,1136 theory 16 typical 72 theory 32 typical Shell 71357 9,11,13 49 theory 16 typical 98 theory 32 typical These orbitals are not actually used. 2e- 6e- 10e- 14e-

93. 6d 5f 7s 6p 5d 4f 6s 5p 4d 5s 4p 3d 4s 3p 3s 2p 2s 1s Aufbau Diagram First shell fills: 2e- (H and He) 2 nd shell fills: 8e- (elements from Li to Ne) 3 rd shell starts filling: 8e- (elements from Na to Ar) 4 th Shell starts to fill (K and Ca) 3 rd shell finishes:10 more e- (Sc to Zn) 4 th Shell continues filling (Ga to Kr) 5 th shell starts to fill  Shell three  Shell two  Shell one  Shell four   Shell five   Shell six  Shell seven

94. Modern Model (optional enrichment) nucleus 1 st Shell 2 nd Shell 3 rd Shell Contains one spherical s orbital with two e-=2e- One s orbital (2e-)+3 p orbitals (6e-)=8e- 1 s orbital, 3 p orbital and 5 d orbitals with 10 e- Electron Cloud

95. Answers to Sheet ElementConfigurationDiagram N (nitrogen) 1s 2, 2s 2 2p 3 2electrons, 5 electrons Mg (magnesium) 1s 2, 2s 2 2p 6, 3s 2 2 electrons, 8 electrons, 2 electrons Sc (scandium) 1s 2, 2s 2 2p 6, 3s 2 3p 6 3d 1, 4s 2, 1 electrons 8 electrons 9 electrons 2 electron Fe (iron) 1s 2, 2s 2 2p 6, 3s 2 3p 6 3d 6, 4s 2, Mn (manganese) 1s 2, 2s 2 2p 6, 3s 2 3p 6 3d 5, 4s 2 Zr (zirconium) 1s 2, 2s 2 2p 6, 3s 2 3p 6 3d 10, 4s 2 4p 6 4d 2, 5s 2 W (tungsten) 1s 2, 2s 2 2p 6, 3s 2 3p 6 3d 10, 4s 2, 4p 6 4d 10 4f 14, 5s 2 5p 6 5d 4, 6s 2,