Slide 1

CHEMISTRY Chapter 1 & 2 Matter, Measurements, and Calculations

Slide 2

Chapter 1 Section 1 Objectives: 1.Define chemistry 2.List examples of branches of chemistry 3.Compare and contrast basic research, applied research, and technological development

Slide 3

What objects in this room are related to chemistry? Plastics Fabrics Clothes Cooking oil Motor oil Make-up Radio Batteries Computers

Slide 4

Chemistry in our daily lives. Antibiotics Food Transportation Sports Farming Military Industry

Slide 5

Chemistry Study of the composition and properties of matter and the changes that matter undergoes -What something is made of -What is the internal arrangement

Slide 6

Chemical Any substance that has a definite composition

Slide 7

6 Main Branches of Chemistry 1.Organic substances containing C 2.Inorganic substances other than organic 3.Biochemistry living things 4.Physical chemistry changes of matter 5.Analytical chemistry id components of materials 6.Theoretical chemistry use math and computers to understand chemical behavior

Slide 8

All branches involve some type of research. Basic research to increase knowledge -how and why Applied research to solve problems Technological development production and use of products - lags behind discoveries - application of knowledge

Slide 9

Review and Assignment 1. Define chemistry 2. List examples of branches of chemistry 3. Compare and contrast basic research, applied research, and technological development Assignment: WS 1-1

Slide 10

Quiz 1.Name two branches of chemistry. 2.List two ways that chemistry affects our daily lives. 3.Definition of chemistry.

Slide 11

Chapter 1 - Matter

Slide 12

Chapter 1 Section 2 Objectives: 1.Distinguish between a mixture and a pure substance. 2.Define what matter is.

Slide 13

Matter -anything that has mass and occupies space -includes almost everything -exceptions are light, heat, and sound -properties are used to measure matter ex. mass Mass measure of quantity of matter - not affected by temp, location, or any other factor

Slide 14

Demo. Mass vs. matter What caused the change in mass? Is air matter?

Slide 15

Matter (cont.) Classified into 2 groups: 1. pure substances 2. mixtures Pure substance matter that has the same properties throughout ex. element or compound

Slide 16

Pure Substances Element substance that cannot be broken down by ordinary chemical change - only 1 type of atom - symbols abbreviated w/1 or 2 letters - can be an allotrope allotrope one of a number of different molecular forms of an element in the same state Compound substance made up of 2 or more elements chemically combined - can be broken down by chemical change - more than 1 type of atom

Slide 17

Compounds 1.Elements that make up a compound are combined in definite proportion by mass ex. 100 g water has 11.2 g H and 88.8 g of O 2. Chemical and physical properties of compound differ from those of its parts ex. water is liquid, H and O are gases 3. Compounds can be formed from simpler substances by chem change and can be broken down into simpler substances

Slide 18

example 100 of water has 11.2 g H and 88.8 g O How many g of H is in a 120g sample of water? 120 g water| 11.2 g H = 13.4 g H | 100 g water

Slide 19

Mixtures - contain 2 or more substances that have different properties - vary in composition and properties from sample to sample ex. rock, wood, salt water -Not chemically combined -Can be separated by simple physical means -ie. filtration, evaporation, distillation

Slide 20

Formation of Mixtures A mixture can be formed 3 ways: 1.Element mixed w/1 or more other elements ex. carbon w/sulfur 2. Compound mixed w/ 1 or more other compounds ex. salt w/sugar 3. 1 or more elements mixed w/1 or more compounds ex. sulfur w/sugar

Slide 21

Characteristics of Mixtures - retain properties of each of its parts ex. iron and sulfur - iron remains magnetic - composition can vary widely - can be homogeneous or heterogeneous

Slide 22

Types of mixtures Homogeneous uniform composition throughout - called solutions ex. alloys, pop, air, coffee Heterogeneous not uniform throughout ex. concrete, soil, dry soup, spaghetti and meat balls

Slide 23

Matter Pure substance ElementCompound Mixture HomogeneousHeterogeneous

Slide 24

Review and Assignment 1. Distinguish between a mixture and a pure substance. 2. Define what matter is. Assignment: WS

Slide 25

Chapter 1 Section 2 Objectives: 1. Distinguish between the physical properties and chemical properties of matter. 2. Classify changes of matter as physical or chemical. 3. Explain the gas, liquid, and solid states in terms of particles.

Slide 26

Properties of Matter -allow us to distinguish btwn substances -characteristics of a substance -what can be observed -way that a substance behaves ex. color, taste, odor, gas, liquid, solid

Slide 27

Properties (cont.) - can be extensive or intensive Extensive d/o amount of matter ex. volume, weight, mass, and E Intensive does not d/o amount of matter ex. melting point, boiling point, density, and conductivity

Slide 28

Demonstration Properties - water and glycerin How do they compare? - look, feel, weight, flow - water and salt water How do they compare? - conductivity

Slide 29

Physical Properties Can be observed or measured w/out changing the substance Can describe the substance Odor, taste, hardness, density, melting point, and boiling point Metals ductile (pulled into wire), malleable (hammered into sheets), luster (shine), good conductors

Slide 30

Chemical Properties A transformation of a substance into a different one rusting, flammability, tarnishing, new substance formed

Slide 31

Physical Change No new substance is formed CHANGE IN PHASE, pounding, grinding, cutting Changes of phase When a substance changes phase there is no change in composition Physically different, chemically the same Solid, liquid, or gas are the three states of matter

Slide 32

States of Matter Solid definite volume and shape Particles are in fixed positions Held w/strong attractive forces Liquid definite volume and no definite shape Takes shape of container Particles can move past each other

Slide 33

States of Matter (cont.) Gas neither definite volume nor definite shape Particles move easily and are very far apart Plasma high temperature state in which atoms lose their electrons

Slide 34

Chemical Change One or more substance is changed to something new Rusting, burning, gas formed, digestion, heat or light added, explosion, color change, odor change, water formed

Slide 35

Review and Assignment 1. Distinguish between the physical properties and chemical properties of matter. 2. Classify changes of matter as physical or chemical. 3. Explain the gas, liquid, and solid states in terms of particles. Assignment: p. 18 and WS

Slide 36

CHEMISTRY Chapter 1 Section 3 Objectives: 1.Perform density calculations. 2.Describe conservation of mass.

Slide 37

Properties of Matter -E is always involved in both physical and chemical changes -Physical are not at noticable -Chemical are more noticable -Heat and light are given off

Slide 38

Density is a physical property is always the same for a solid substance in gases and some liquids a change in temperature will change the density increase in temperature will decrease density D = m/V

Slide 39

Density problem Use the 5 steps in problem solving to solve the following problem. Lead has a mass of 22.7 g and its volume is 2.00 cm 3. What is its density? m = 22.7 g V = 2.00 cm 3 D = m/V = 22.7 g/2.00 cm 3 = 11.4 g/ cm 3

Slide 40

Examples

Slide 41

Conservation of Mass In reactions matter cannot be created or destroyed by a chemical change - mass stays the same, it may just change form

Slide 42

Density Lab Results Group 1 Group 2 Group 3 Group 4 Group 5 -

Slide 43

Review and Assignment 1. Perform density calculations. 2. Describe conservation of mass. Assignment: WS and Density lab

Slide 44

Chapter 2 - Sec.1 Objectives: 1.Describe the purpose of the scientific method. 2.Distinguish between qualitative and quantitative observations. 3.Describe the steps to making a graph. 4.Distinguish between inversely and directly proportional relationships.

Slide 45

Scientific Method - a logical approach to solving problems 1. Make observations -observe your surroundings 2. State the problem - stated as a question 3. Collect data 4. Form hypothesis - testable statement 5. Test hypothesis 6. Conclusion 7. Modify hypothesis and retest

Slide 46

Observing Involves making measurements and collecting data Data can be qualitative or quantitative Qualitative non-numerical information - descriptive (the sky is blue) Quantitative numerical information - the mass is 25.7 grams

Slide 47

Conclusion Can be explained by using models Model explanation of how phenomena occur or how things are related - visual - verbal - mathmatical

Slide 48

Theory -models may become part a theory Theory broad generalization that explains facts or phenomena - must be able to predict results ex. kinetic-molecular theory collision theory

Slide 49

Controlled Experiments Use manipulated variable (independent) Use responding variable (dependent) One variable manipulated at a time Measurements are called data

Slide 50

Making a Graph Shows results of an experiment in a meaningful pattern Dependent variable is on the vertical axis 1. Always include a title 2. Determine variables 3. Set up scale 4. Plot points 5. Draw best-fit line

Slide 51

Oxygen obtained from electrolysis of water

Slide 52

Relationships in graphs Directly proportional if dividing one by the other gives you a constant value If one increases so does the other If started at point (0,0) Inversely proportional if their product is constant If one increases the other decreases Produce a curve

Slide 53

Review and Assignment 1.Describe the purpose of the scientific method. 2.Distinguish between qualitative and quantitative observations. 3.Describe the steps to making a graph. 4.Distinguish between inversely and directly proportional relationships. Assignment: graphing WS

Slide 54

Quiz 1.List three steps of the scientific method. 2.List two steps in making a graph.

Slide 55

Chapter 2 Sec.2 Objectives: 1. Distinguish between a quantity, a unit, and a measurement standard. 2. Name SI units for length, mass, time, volume, and density. 3. Distinguish between mass and weight.

Slide 56

Measurements Basic part of science Make observations more meaningful Needs to be more than just a number or quantity Need a common system of units For consistency Measure your desk w/anything you have available

Slide 57

SI System -The International System of Units -Used in all science -A standard -Based on 10 -Makes it easier to convert from one unit to another

Slide 58

SI System (continued) -7 base units 1. Length meter (m) 2. Mass kilogram (kg) 3. Time second (s) 4. Amount mole (mol) 5. Temperature Kelvin (K) 6. Electric Current ampere (amp) 7. Luminous intensity candela (cd)

Slide 59

Weight vs. mass Mass quantity of matter - how much space it takes up - measured w/a balance - unit kg Weight F gravity pulls on matter with - measured w/spring scale - unit Newton On the moon will our weight or mass stay the same?

Slide 60

SI Prefixes You must know these. Kilo- 1000 Deca 10 Base unit (m, s, L) Centi 1/100 or 0.01 Milli 1/1000 0r 0.001

Slide 61

Derived Units -combination of base units Examples - Area = m 2 - Volume = m 3 - Density = kg/m 3 - Newton = m kg/s 2

Slide 62

Derived Units (cont.) Area determined by multiplying 2 lengths Volume determined by multiplying 3 lengths for a solid - for liquids unit is cm 3 or mL ** 1 mL = 1 cm 3

Slide 63

Review and Assignment 1. Distinguish between a quantity, a unit, and a measurement standard. 2. Name SI units for length, mass, time, volume, and density. 3. Distinguish between mass and weight. Assignment: WS 2-2 and p. 42 ~1-3

Slide 64

Quiz 1.What is the base SI unit for mass? 2.Kilo = ______ 3.Centi = _____ 4.What is a derived unit? 5.1 cm 3 = _____ mL

Slide 65

Chapter 2 - Sec.3 Objectives: 1.Distinguish between accuracy and precision. 2.Determine the number of significant figures in measurements. 3.Perform mathematical operations involving significant figures.

Slide 66

Accuracy and Precision Accuracy closeness of a measurement to correct value Precision closeness of a set of measurements to each other Consistency Do not have to be correct d/o measuring instrument

Slide 67

Bullseyes

Slide 68

Significant Figures -digits in a measurement that are know with certainty and one digit that is estimated -CALCULATORS DO NOT KEEP TRACK OF SIGNIFICANT FIGURES

Slide 69

Significant Figure Rules 1. Digits other than zero are ALWAYS significant ex. 61.4 3 sig. fig. 2. All zeros at the end of a number and to the right of the decimal with a # preceding the decimal are ALWAYS sig ex. 4.7200 km 5 sig. fig. 3. Zeros used only for spacing are NOT significant ex. 7000 1 sig. fig. 201 sig. fig. 100.04 sig. fig. 4. Zeros between sig. fig are significant 5. Zeros in front of a non-zero are NOT sig. - dont count until you get to 1 st non-zero from lf to rt 0.004 1 sig. fig. 0.00091 sig. fig.

Slide 70

Significant Figures 1,000 = _____ sig figs 100.0 = _____ sig figs 0.00012340 = _____ sig fig 10.0340 = _____ sig fig

Slide 71

Calculating w/Significant Figures Addition and Subtraction - use same # of decimal places as the measurement w/the least decimal places ex. 2.098 3 DECIMAL places +6.21 DECIMAL place 8.298round to 1 Decimal 8.3 is the final answer

Slide 72

Adding and Subtracting 10.0 + 123 = _____ 23.456 23.0 = _____ 100.12 + 56.45 = _____ 1,000 + 12.234 = _____

Slide 73

Calculating w/sig. figs (cont.) Multiplication and Division - use same # sig. fig. as the measurement w/the least sig. fig. ex.2.38 3 sig. fig x 9.02 sig. fig 21.42 round to 2 sig. Fig 21 is the final answer

Slide 74

Multiplying and Dividing 100.0 x 10 = _____ 34.56 x 23.45 = _____ 12.045 x 34.008 = _____ 50.04 x 23 = _____

Slide 75

Review and Assignment 1.Distinguish between accuracy and precision. 2.Determine the number of significant figures in measurements. 3.Perform mathematical operations involving significant figures. Assignment: WS 2-6 and sig fig WS

Slide 76

Quiz How many significant figures are in the following numbers? 1. 8,000 _____ 2. 100.01 _____ 3. 0.00056_____ 4. 4500.10 _____ 5. What is precision?

Slide 77

Chapter 2 - Sec.3 Day 2 Objectives: 1.Perform mathematical operations involving percent error.

Slide 78

Percent Error Observed value based on lab measurements True value based on generally accepted references Error exists in any measurement d/o measurer, instrument, conditions

Slide 79

Percent Error % error = true value obs. value x 100 true value Example atomic mass of Al = 28.9 g measured mass = 27.0 g What is the % error? 28.9 g 27.0 g x 100 = 7.00 % 28.9 g

Slide 80

Review and Assignment 1.Perform mathematical operations involving percent error. Assignment: WS 2-5 and % error WS

Slide 81

Quiz How many significant figures are in the following numbers? 1. 8,104 _____ 2. 100.01 _____ 3. What does % error tell us? 4. What is accuracy? 5. What is precision?

Slide 82

CHEMISTRY Chapter 2 Sec.3 Day 3 Objectives: 1.Use dimensional analysis to convert measurements. 2.Convert measurements into scientific notation. 3.Perform mathematical operations using exponents.

Slide 83

Problem Solving Rules Write down what is known. - mass = 346 gvolume = 34.6 cm 3 2. Write down unknown. - density = ? 3. Write the equation to use. D = m/V 4. Fill in knowns. D = 346 g/34.6 cm 3 5. Solve for unknown and label. D = 200 g/cm 3 6. Check your work.

Slide 84

Dimensional Analysis - use with conversion factors to change from one unit to another Steps: convert 2550 m to km 1. Determine conversion factor - 1000 m to 1 km 2. Set up T-bars 3. Write given # in first box

Slide 85

Dimensional Analysis (cont.) 4. Write conversion factor in 2 nd box - unit on bottom matches unit of given # 5. Matching labels cancel - if 1 from conversion factor is on top divide - if 1 from conversion factor is on bottom multiple

Slide 86

Scientific Notation -Used to represent very large or very small numbers -There are two parts -Basic form is M x 10 n -M is a number - n is a number representing how many places to move the decimal

Slide 87

Scientific Notation (cont.) If n is negative, your number is a decimal If n is positive, your number is a large number Examples: 60,000,000 = 6 x 10 7 0.000005 = 5 x 10 -6 125,000 = 1.25 x 10 5

Slide 88

Scientific Notation (cont.) Write the following in scientific notation. 1,000,000,000 23,456 0.0005678 0.034 14,239.1

Slide 89

Scientific Notation (cont.) Write the following in long hand. 1.1 x 10 -9 2.3.5 x 10 5 3.7.123 x 10 -3 4.5 x 10 2 5.4.56 x 10 -2

Slide 90

Multiplication w/exponents Step 1 Multiply coefficients Step 2 Add exponents ex. (2 x 10 2 ) (2.5 x 10 5 ) = 5 x 10 7

Slide 91

Division w/exponents Step 1 Divide coefficients Step 2 Subtract exponents ex. (5 x 10 -2 ) (1.0 x 10 7 ) = 5 x 10 -9

Slide 92

Addition & Subtraction w/exponents All numbers must be written in the same power of 10 ex. 5.8 x 10 3 + 2.16 x 10 4 - change to 0.58 x 10 4 + 2.16 x 10 4 = 2.74 x 10 4

Slide 93

Scientific Notation & sig figs All numbers in front of the x 10 are significant ex. 2.00 x 10 2 = 3 sig fig 2 x 10 2 = 1 sig fig

Slide 94

Scientific Notation & calculators 5.44 x 10 7 /8.1 x 10 4 5.44 (EE or exp) 7 / 8.1 (EE or exp) 4 = 6.7 x 10 2

Slide 95

Review and Assignment 1.Use dimensional analysis to convert measurements. 2.Convert measurements into scientific notation. 3.Perform mathematical operations using exponents. Assignment: p. 57 ~ 1-7 and WS