AP Chemistry

1. AP Chemistry Introduction

2. What is Chemistry? Chemistry is… • The STUDY of matter & the CHANGES it undergoes. • Matter is anything that has mass and that occupies volume (space) How do we STUDY matter? • We observe something that triggers our curiosity • We research what we observed • We design and carry out an experiment to test what we are curious about

3. What are the 3 states of matter and compare? Has definite shape & definite (fixed) volume Tight structure NO Definite shape Definite volume Structure not held strongly NO Definite shape NO Definite volume Weak attraction Vibrate in place Flow / slide past Move freely & rapidly

4. Collecting Data (2 types) • Qualitative: describes observations using senses • Must use as much DETAIL as possible • Quantitative: numerical measurements obtained using an instrument (must include units to describe #’s). • Qualitative data takes up more space on the data table… ESPECIALLY IF YOU WRITE BIG! • Base the size of your data table on… • The type of data being collected • The size of your hand writing!

5. How does Matter Undergoes Changes! • Physical Changes • Changes the appearance without changing the identity of the substance. • Example: tearing a piece of paper… its still paper just smaller  • Chemical Changes • Changes the appearance AND changes the identity of the substance. • Example: burning a piece of paper… its not longer paper, its ash 

6. matter: anything having mass and volume the amount of matter in an object mass: the pull of gravity on an object weight: volume: the space an object occupies units: L, dm3, mL, cm3 conversions: 1 L = 1 dm3; 1 mL = 1 cm3 state of matter: solid, liquid, or gas atom: a basic building block of matter -- ~100 diff. kinds

7. contain only one type of atom. Elements Broken Dreams Blvd. (a)monatomic elements consist of “unbonded,” identical atoms e.g., Fe, Al, Cu, He (b) polyatomic elements consist of several identical atoms bonded together H2 O2 Br2 F2 I2 N2 Cl2 -- diatomic elements: “7 7 7” -- others: P4 S8

8. (c) allotropes: different forms of the same element in the same state of matter OXYGEN CARBON oxygen gas (O2) elemental carbon graphite ozone (O3) buckyball diamond

9. molecule: a neutral group of bonded atoms O O2 2 O P P4 4 P Elements may consist of… either molecules orunbonded atoms.

10. 2 He 4.003 10 Ne 20.180 18 Ar 39.948 36 Kr 83.80 54 Xe 131.29 86 Rn (222) Chemical symbols for elements appear on the periodic table; only the first letter is capitalized.

11. Compounds contain two or more different types of atoms. -- have properties that differ from those of their constituent elements e.g., Na (sodium): explodes in water table salt (NaCl) Cl2 (chlorine): poisonous gas

12. Compound composition All samples of a given compound have the same composition by mass. Every sample of NaCl tastes the same, melts at the same temp., and is 39.3% Na and 60.7% Cl by mass.

13. A 550. g sample of chromium(III) oxide (Cr2O3) has 376 g Cr. How many grams of Cr and O are in a 212 g sample of Cr2O3? a. First solve for compound composition X100% 68.4% Cr 31.6%O % Cr = % O= X100% b. (New sample has same composition.) 212 g (0.684) Cr: = 145 g Cr chromium(III) oxide O: 212 g (0.316) = 67 g O

14. composition: what the matter is made of water: copper: many “threesomes” of 2 H’s and 1 O many Cu atoms Properties describe the matter. e.g., what it looks like, smells like, how it behaves Chemistry tries to relate the microscopic and macroscopic worlds.

15. ( ( ) ) ( ( ) ) ( ( ) ) ( ( ) ) ( ( ) ) ( ( ) ) ( ( ) ) ( ( ) ) ( ( ) ) ( ( ) ) ( ( ) ) ( ( ) ) ( ( ) ) ( ( ) ) ( ( ) ) States of Matter SOLID LIQUID GAS translating; close together translating quickly; far apart vibrating vapor: the gaseous state of a substance that generally is found as a solid or liquid

16. Changes in State Energy put into system: sublimation melting boiling SOLID LIQUID GAS freezing condensation deposition Energy removed from system:

17. Matter can be classified in SEVERAL ways! Variable Composition Separated Physically Fixed Composition Separated Chemically MIXTURE PURE SUBSTANCE

18. Classifying Matter 1.(Pure) Substances have a fixed composition and fixed properties. -- they have a single chemical formula ELEMENTS COMPOUNDS e.g., Fe, N2, S8, U e.g., H2O, NaCl, HNO3 sulfur (S8) sodium chloride (NaCl)

19. 2. Mixtures contain two or more substances mixed together. -- have varying composition and varying properties -- The substances are NOT chemically bonded; they retain their individual properties. Tea, orange juice, oceans, and air are mixtures.

20. Mixtures are made up of Pure Substances! • 2 or morepure substances occupying thesame space • Pure substances maintain their own UNIQUE IDENTITY!

21. Classification of Matter

22. salt water Kool Aid bronze (Cu + Sn) pewter (Pb + Sn) brass (Cu + Zn) Two Types of Mixtures A. homogeneous: (or solution) sample has same composition and properties throughout; evenly mixed at the particle level e.g., alloy: a homogeneous mixture of metals e.g.,

23. raisin bran tossed salad paint snow globes Two Types of Mixtures (cont.) B. heterogeneous: different composition and properties in the same sample; unevenly mixed e.g., suspension: settles over time e.g.,

24. PURE SUBSTANCE MIXTURE ELEMENT COMPOUND HETEROGENEOUS HOMOGENEOUS Chart for Classifying Matter MATTER

25. ASK ANOTHER QUESTION

26. Gatorade Trail Mix Glucose Nickel Ocean Water Classify the following as either: • Pure Substance (element or compound) • Mixture (heterogeneous or homogenous)

27. Gatorade Mix. homogeneous Ocean Water Mixture homogeneous Nickel Substance Element Trail Mix Heterog. Glucose Subtance compound Classify the following as either: • Pure Substance (element or compound) • Mixture (heterogeneous or homogenous) Mixture homogeneoous Substance element Mixture heterogeneous Substance compound Mixture heterogeneous

28. Separating Mixtures -- No chemical reactions are needed because… involves physical means, or physical changes. substances are NOT bonded. 1. sorting: by color, shape, texture, etc. by particle size 2. filtration:

29. Separating Heterogeneous Mixtures by FILTRATION Large solid particles collect in the filter paper Smaller particles (liquid & solid) pass through and collect in the beaker/flask

30. Separating Mixtures (cont.) 3. magnetism: one substance must contain iron some substances dissolve more easily than others 4. chromatography:

31. Separating Homogenous Solutions by Chromatography Separating particles based on solubility (dissolving) and size of particles Smaller particles move faster and travel higher  Mixture must be placed above the solvent level As the solvent moves up the paper it take the particles with it! Solvent is used to dissolve the particles

32. decant: to pour off the liquid blood after high- speed centrifuging Separating Mixtures (cont.) 5. density: “sink vs. float”; perhaps use a centrifuge

33. thermometer water out (warmer) water in (cooler) more-volatile substance condenser (i.e., the one with the lower boiling point) mixture more-volatile substance, now condensed heat source Separating Mixtures (cont.) 6. distillation: different boiling points Volatilesubstances evaporate easily.

34. Separating Homogenous Solutions by Distillation (Evaporation) Heat to the temperature of the LOWEST BP Substance with the LOWEST BP evaporates Substance with the LOWEST BP moves through the tubes and condenses back into a liquid  Substance with the LOWEST BP collects in a separate container  Substances that have Different Boiling Points! Substance with the HIGHER BP left behind

35. Physical Properties BEFORE A LAB STARTS… • It is important to make “initial” observations • Describes each substance using your senses! • Only describe one substance at a time • Examples: • You look at an individual substance and you record its initial color, odor, texture, physical state, size, clarity • Constants for the substance: melting point, boiling point, density, specific heat capacity… Blue solid with a grainy texture

36. Physical Changes DURING THE LAB… • A change that does NOT alter the formula of the substance… NO chemical bonds are broken… you start and end with the same substance… it just looks a little different

37. Examples of Physical Changes… • Change in physical state: Creating a mixture: dissolving Ex: Salt water… even though you can’t see the salt anymore doesn’t mean it’s not there… you can still evaporate the water and get the salt back!!! Change in size/shape of the substance: grinding, chopping, cutting, tearing, folding

39. Chemical Properties BEFORE A LAB STARTS… • It is important to do research to know how your substances behave/react! • It helps you know what to look for later on! • Example: • When iodine is added to starch it turns a dark blue color

40. Chemical Changes DURING THE LAB… • A CHEMICAL REACTION HAS OCCURRED!!!!!! • a NEW substance with a NEW formula has been made! • Chemical bonds have been broken and rearranged Sugar “Sucrose” C12H22O11 Caramel C8H10O4 Sucrose bonds are breaking Caramel bonds are forming

41. ONE OF THESE AND ONE OF THESE Properties of Matter CHEMICAL properties tell how a substance reacts with other substances. PHYSICAL properties can be observed without chemically changing the substance. EXTENSIVEproperties depend on the amount of substance present. INTENSIVEproperties do NOT depend on the amount of substance.

42. Examples: P, I electrical conductivity……………………… P, I ductile: can be drawn (pulled) into wire….. malleable: can be hammered into shape… P, I C, I reactivity with water………………………... P, I brittleness……………………………………. P, I magnetism……………………………………

43. Density how tightly packed the particles are Density = D= ** Density of water = 1.0 g/mL = 1.0 g/cm3 The density of a liquid or solid is nearly constant, no matter the sample’s temperature. Density of gases is highly dependent on temperature. A student needs 15.0 g of ethanol, which has a density of 0.789 g/mL. What volume of ethanol is needed? D= V= = 19.0 mL = =

44. METRIC Unit Conversions • In chemistry, the most common measurements made are for mass, volume, length, and temperature. • Remember… unit conversions can ONLY be within the SAME TYPE OF MEASUREMENT! gram g liter l m meter Kelvin or Celsius K or ˚C

45. METRIC System • The METRIC SYSTEM uses POWERS OF TEN to make the units larger or smaller • PREFIXES are used for the different sizes of the units and a base • Think of the BASE as the ONE DOLLAR BILL and the larger prefixes as the five, ten, twenty, fifty, or hundred dollar bills! G Giga Billion = 109 Mega M Million = 106 Kilo K or k Thousand = 103 Hecto H or h Hundred = 102 Deka D Ten = 101 Mass g, length m, volume L NO prefix (base) One = 100 = 1

46. METRIC System NO prefix (base) g, m, L One = 100 = 1 • Think of the BASE as the ONE DOLLAR BILL and the smaller prefixes as the quarter, dime, nickel, and the penny! Tenth =0.1= 10-1 deci d Hundredth=0.01=10-2 centi c Thousandth=0.001=10-3 milli m Millionth 0.000001=10-6  micro Billionth 0.000000001= 10-9 nano n

47. SI Prefixes to Memorize G M k d c m m n p f iga ega ilo eci enti illi icro ano ico emto “Got my kilt, Dad! Can’t miss midsummer’s no- pants Friday!” 109 106 103 10–1 10–2 10–3 10–6 10–9 10–12 10–15

48. Metric Equivalence Statements • An equivalence statement should have equal values. • Use the metric prefix meaning to make an equivalence statement • Gigagram • Centimeter Write equivalence statements for each of the following: Hectometer Decigram kilometer nanoliter What you need to know: Giga = G = 109 Measurement = gram = g 1 Gg = 1x109 g Or 1 Gg = 1000000000 g Use sci. notation What you need to know: Centi = c = 10-2 Measurement = meter = m 1 cm = 1x10-2 m 1 cm = (0.01) m 1 km = 1x103m 1 km =1000 m 1 Hm = 1x102m 1Hm = 100 m 1 dg = 1x10-1g 1 nL = 1x10-9L

49. Converting Units: Dimensional Analysis • Mrs. B’s hot air balloon can hold 20.0 L of air. How many mL of air can the balloon hold? Note: Multiplying fraction to cancel units. Multiply on top and divide on the bottom • 2.0 g = ? g t chart: multiply numbers side by side 1 mL = 1x10-3 L 1 = 2.00 x 104 mL 1x10-3 divide numbers top and bottom 1 = 2.0 x 106g 1 g = 1x10-6g 1x10-6

50. Scientific Notation Using POWERS OF TEN to make very large & very small #’s easy to use! • Avogadro’s Number 602000000000000000000000 (YIKES!) • Its so much easier to use Scientific Notation 6.02 x 1023 • All of the significant figures are still included • The VALUE of the NUMBER is still maintained! • Volume of one carbon atom 0.0000000000000000000014 pm3 • Its easier to use Scientific Notation 1.4 x 10-21 pm3 • All of the significant figures are still included • The VALUE of the NUMBER is still maintained!