Chemistry

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 --theoretical basis of chemistry --Substances in living things

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

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 properties Characteristic properties • Chemical properties Non-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 1 • Element Song, Version 2 • Element Song, Version 3

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) incompressible solid, liquid • B) indefinite shape liquid, gas • C) definite volume solid, liquid • D) flows gas, 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 This is an outline of today’s lesson, not the notes • States of Matter • Phases (optional material) • Symbols • Energy • Conservation of Energy • Identifying Chemical Reactions • Chemical equations • Conservation of Mass

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: • H He Li Be B C N O • F Ne Na Mg Al Si P S • Cl Ar K Ca Br Fe Cu Zn • I Ni Co Ag Au Hg Pb

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.

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 ( 2H2 + O2  2 H2O ) 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.

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% SiO2, 13%Na2O, 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% SiO2). • C) Table sugar: compound (C12H22O11) • D) river water: mixture (H2O, 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 (N2)

33. 24. The chemical symbols are: • Copper: Cu Silver: Ag • Oxygen: O Sodium: Na • Phosphorus: P Helium: He • 25. The elements found in each are: • NH4Cl: Nitrogen, Hydrogen, Chlorine • KMnO3: Potassium, Manganese, Oxygen • C2H7OH: Carbon, Hydrogen, Oxygen • CaI2: 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 • Yotta 1024 Superclusters deci 1/10 hand • (100 zetta) 1023centi 1/100 fingernail • (10 zetta) 1022milli 1/1000 sand • Zetta 1021 Galaxy (100 micro) 10-4 • (100 exa) 1020(10 micro) 10-5 • (10 exa) 1019micro 10-6 bacteria • Exa 1018 nearby stars(100 nano) 10-7 • (100 peta) 1017(10 nano) 10-8 • (10 peta) 1016nano 10-9 molecule • Peta 1015 Solar system(100 pico) 10-10 • (100 tera) 1014(10 pico) 10-11 • (10 tera) 1013pico 10-12 atom • Tera 1012 Inner planets (100 femto) 10-13 • (100 giga) 1011(10 femto) 10-14 • (10 giga) 1010femto 10-15 proton • Giga 109 Earth/moon (100 atto) 10-16 • (100 mega) 108(10 atto) 10-17 • (10 mega) 107atto 10-18 electron? • Mega 106 East coast (100 zepto) 10-19 • (100 kilo) 105 (10 zepto) 10-20 • (10 kilo) 104zepto 10-21 quark?? • Kilo 1000 Town (100 yocto) 10-22 • Hecta 100 football field (10 yocto) 10-23 • Deca 10 Elephant yocto 10-24 strings???

39. Common metric units & prefixes • Mega- M • - • - • Kilo- k prefixes (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 millimetre metre centimetre decimetre

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 point 32°F • Room temp 68 °F • Body temperature 99 °F • Boiling point 212 °F • Degrees Celsius (°C) A.KA. Centigrade Often used in Chemistry • Freezing point 0 °C • Room temperature 20 °C • Body temperature 37 °C • Boiling point 100 °C • Kelvins (K), formerly: °K or Absolute °A ) The Best for Chemistry, especially with gas laws. • Freezing point 273 K • Room temperature 293 K • Body temperature 310 K • Boiling point 373 K Water boils Body temp. Room temp. Water freezes Mercury freezes Absolute zero

44. Conversions 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. 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 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.

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. Your own measurements • 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