Presentation is loading. Please wait.

Presentation is loading. Please wait.

Back ©Bires, 2002Bires, 2004 Chapter 1 and 2: Matter and Change Measurements and Calculations How do we do what we do? What do we measure? How do we measure?

Similar presentations

Presentation on theme: "Back ©Bires, 2002Bires, 2004 Chapter 1 and 2: Matter and Change Measurements and Calculations How do we do what we do? What do we measure? How do we measure?"— Presentation transcript:

1 Back ©Bires, 2002Bires, 2004 Chapter 1 and 2: Matter and Change Measurements and Calculations How do we do what we do? What do we measure? How do we measure? How do we do what we do? What do we measure? How do we measure? Read Text pages 4-This unit: text p4-61 All our science, measured against reality, is primitive and childlike - and yet it is the most precious thing we have. -Albert Einstein This is the basics of *DOING* Chemistry

2 Back Bires, 2004 Slide 2 Matters properties When observing matter, we observe many of its properties. Extensive Property - depends upon how large your sample of matter is. –What are some extensive properties? Mass and volume are extensive properties. Intensive Property - unique to that type of matter, and does not depend on sample size. –What are some intensive properties? Density, and melting point are intensive properties.

3 Back Bires, 2004 Slide 3 States of Matter and Changes The four primary states of matter are: (?) solid, liquid, gas, and high energy plasma. When we observe properties, these can be physical or chemical properties: Physical property - can be observed without changing the substance. (color, mass, others?) Chemical property - requires that the substance be changed to be observed. (flammability, others?)

4 Back Bires, 2004 Slide 4 Changes of Matter Physical change - substance does not change, just the form. –(It can also revert back to its original state…ie: melting and vaporization) Chemical change - original substance is lost, and a new substance is formed. –(The original substance cannot be returned without additional chemical changes.) Which kinds of changes are these?: –Burning, freezing, vinegar + baking soda, and opening a soda bottle,

5 Back Bires, 2004 Slide 5 Chemical Reactions When chemicals react, we call them reactants and they are placed on the left side of achemical reaction equation: The chemical(s) that is produced is called aproduct, placed on the right. Remember: Reactants react to produce products

6 Back Bires, 2004 Slide 6 Energy Transfer The law of conservation of energy states that energy cannot be created or destroyed. We can force energy to change form and store energy in various forms. –How is a flashlight battery like gasoline in a car? Exothermic process – releases energy. –(think exit energy) Endothermic process – takes energy from surroundings. –(think into energy) Can you think of an exothermic reaction? Can you think of an endothermic reaction?

7 Back Bires, 2004 Slide 7 From the macro to the micro… We can classify matter in terms of its complexity: Mixture - collection of two or more physically different compounds. –The compounds in a mixture can be separated out without the need to chemically change either of the compounds. Two types of mixtures are Homogeneous (homo) meaning same)… –Homogenous mixtures cannot be easily separated And Heterogeneous (hetero meaning different) –Heterogeneous mixtures can be separated with simple mechanical processes. These often have phases.

8 Back Bires, 2004 Slide 8 A step closer to the micro… A compound - substance made up of two or more pure elements. Water is a compound because it is composed of the elements H and O. Table salt is a compound because it is composed of the elements Na and Cl. A compound cannot be separated from its elemental makeup –(without destroying the compound) Compounds have very different properties than their elements.

9 Back Bires, 2004 Slide 9 The micro… Elements - simplest form of pure substances. –There are approximately 110 elements, and they can be found on the periodic table. Elements consist of a single type of atom. Water is NOT an element. Pure diamond IS an element. Why? Allotropes – different forms of a single element. Each allotrope has different properties due to different arrangements of its atoms. –Diamond and graphite are allotropes.

10 Back Bires, 2004 Slide 10 The very micro… Atom - smallest thing with the properties of something in the macro. –(a single helium atom will function just like a the helium in a party balloon. Atoms = protons and neutrons in the nucleus + electrons in electron orbits. Properties of atoms depends upon the number of protons, neutrons, and electrons in the atom –(well get into atomic and sub-atomic theory later.) Isotopes - Two atoms with same number of protons (are the same element) but a different number of neutrons Ions - Atoms with a different number of electrons

11 Back Bires, 2004 Slide 11 Similar to table on page 15 Using this flowchart, what is our water? Is milk really homogenous?

12 Back Bires, 2004 Slide 12 The Periodic Table The periodic table is a collection of all the known elements into a model that groups elements with similar properties. Vertical columns are Groups of elements with similar properties. Horizontal Periods represent elements with similar atomic mass

13 Back Bires, 2004 Slide 13 Three Main Areas of the Periodic Table Metals are found on the left. –Metals tend to be malleable, ductile, and good conductors of heat and electricity. Nonmetals are found on the right. –Nonmetals tend to be brittle, and poor conductors of e- and heat. The elements between metals and nonmetals are called Metalloids and have some characteristics of both metals and nonmetals

14 Back Bires, 2004 Slide 14 He Ne The Noble Gasses The Noble gasses are found on the far right of the P- table. The Nobles are: …Very unreactive. –(*not entirely unreactive*) …Gasses at room temperature. …Mined from gas pockets in the ocean …Produce bright emissions when electrified …Have filled octets. (more later) End of chapter 1

15 Back Bires, 2004 Slide 15 The Scientific Method (p29 in text) The scientific method - process of solving problems by asking and testing questions. Observation –Observation of a natural phenomena. Question –Create a question to test your observations. Hypothesis –A reasonable explanation of your observations. A possible answer to your question. This is NOT a guess. Your text: A Testable statement Experiment/Test –A controlled observation (a test of your hypothesis). Collect and Analyze Data –Experimental results must be collected and interpreted. A valid experiment must be reproducible. Conclusions –Data is explained and compared to the hypothesis. The final step… What next? How can I this new information?

16 Back Bires, 2004 Slide 16 A more complex model of the S.M.

17 Back Bires, 2004 Slide 17 Theory vs. Law Theory - an explanation of observations of natural phenomena. –A theory cannot be proven, but it has never been disproven. –If a theory is disproven, it must be modified or rejected. why –A theory explains why things do what they do. Law - a description of fact. whatwillhappen –A law describes what will happen. –Because a law is a description of fact, it cannot be broken.

18 Back Bires, 2004 Slide 18 Lab Work and Lab Reports The main sections of a lab report are: Title (with name, class, instructor, date) Purpose / Problem (why are you doing this lab) Variable (when present) Hypothesis (when applicable) Equation (chemical, balanced A+B C) Procedure, with Materials List Data (create table), with Graph (when applicable) Calculations (includes equations used and a sample calculation (word equation)) Conclusion (includes error analysis of calculations, as well as answers to lab questions and what next)

19 Back Bires, 2004 Slide 19 Measuring…Standard Units Standard units we use in the sciences: Meter – length (m) Kilogram* – mass (kg) Second – time (s) Kelvin – temperature (K) Liter** – volume (L) Mole – amount of substance (mol) (more later) AMU – atomic mass (amu) (more later) –* in lab, we will usually measure in grams, (g) –** in lab, we will usually measure in milliliters, (mL) Table on 34 The Kelvin Scale: 0K = absolute zero K = water freezes K = water boils

20 Back Bires, 2004 Slide 20 Measuring…SI Prefixes Prefixes are added to the standard units to express measurements that are very large or very small. Common prefixes are: kilo – (k) x 10 3 (x 1,000) –kilogram = 1000 grams milli – (m) x (x 1/1,000) –milliliter = liters micro – (μ) x (x 1/1,000,000) Mega – (M) x 10 6 (x 1,000,000) centi – (c) x (x 1/100) –centimeter = 0.01 meters nano – (n) x (x 1/1,000,000,000) 1.0ml = 1.0cc cc = cm 3 About 35 ml

21 Back Bires, 2004 Slide 21 Derived Units Derived units - products of standard units. Volume – a derived unit (b x h x w). –The SI unit is the cubic meter m 3. m 3 is huge, so we use the L, or mL (cm 3 ). Density – The amount of mass that is crammed into a certain volume. Mathematically, D=m / V. –Each compound has a unique density. –( g/ml, g/cm 3, or kg/m 3 ) Pressure – Force applied over an area (like PSI). –The SI unit of pressure is the Pascal, Pa. (more on this later…) See page 37 Osmium (#76) is the densest element on the planet

22 Back Bires, 2004 Slide 22 Metric conversions To convert one metric prefix to another, multiply by a power of ten. Example: To convert 12 meters to centimeters… Or to convert 345 milligrams to grams… REMEMBER: If the unit gets bigger, the number gets smaller! (and vice versa) x base unit 12 x 10 2 = 1200 cm 345 x = g Always show units!

23 Back Bires, 2004 Slide 23 Measuring…Scientific Notation Scientific notation - a form of shorthand used to express numbers that are very large or very small. SciNot - real decimal number, multiplied by a base-ten exponent. –The number is expressed with one digit to the left of the decimal –and the base-ten exponent is always an integer. For instance, becomes 1.35 x10 5. Can you figure what 4500 is? How about moving the decimal the other way…try We moved the decimal five places 6.02 x x10 23

24 Back Bires, 2004 Slide 24 Scientific Notation Practice Convert the following to scientific notation: Convert the following to floating point notation:

25 Back Bires, 2004 Slide 25 Measuring…Significant Digits Significant Digits tells us how many digits to include in our measurements, calculations, and answers. –It is a measure of how accurate our equipment is. Rules: –All non-zero numbers are significant. 1, 2, 256, –Zeros between significant numbers are significant. 303, 50034, 1001 –Zeros to the RIGHT of a decimal are significant , 24.0, , –Zeroes to the LEFT of a decimal are NOT. 4000, , 10, 2400, Open books to page 47 for help

26 Back Bires, 2004 Slide 26 Pacific or Atlantic? Decimal Present? Count from the Pacific Decimal Absent? Count from the Atlantic A little trick for sig figs

27 Back Bires, 2004 Slide 27 SigDigs…more examples 2450 has 3 sig figs has 4 sig figs has 2 sig figs has 3 sig figs Exceptions to the rules: –Fractions –Counting –When the teacher tells you to ignore them Figure the number of sigfigs for the following: x Can you see why significant digits are important?

28 Back Bires, 2004 Slide 28 Why Use Significant Digits? We use Significant Digits to express accuracy. When you measure something, your measurement is only as accurate as your weakest instrument. When you add figures, round your answer to the least number of digits after the decimal. When you multiply, express your result with the least significant digits. What is the volume of a box 2.34m wide, 3.1m deep, and 3.56m long?

29 Back Bires, 2004 Slide 29 Qualitative vs. Quantitative Measurements Qualitative measurement – a description of an object. –blue sticky smelly. Quantitative measurement – data expressed with numbers and units. –42.3 kilograms 14 kilometers per hour 3.80 grams. To accurately describe a compound or solution in chemistry: –use color, transparency, and texture/state. –A colorless, clear, liquid. ?

30 Back Bires, 2004 Slide 30 Accuracy and Precision Accuracy - closeness to an accepted value. Precision - closeness of a set of measurements. –we strive for accuracy and demand precision! –All equipment has a level of precision that you should record in lab. (ie: +/ gram) We use a percent error calculations when working with data. The ideal value The exp value The ideal value

31 Back Bires, 2004 Slide 31Relationships If one variable changes as another changes, we say they have a relationship. Direct relationship: –If A increases as B increases –If their quotient is a constant (y/x = k), we say they are directly proportional. Inverse relationship: –If A decreases as B increases –If their product is a constant (yx=k), we say they are inversely proportional. Page 55

32 Back Bires, 2004 Slide 32 Factor Label Analysis (T-chart) Factor label analysis is a way of converting between a number of units. Convert 32,000 inches into miles, knowing that there are 1609 meters in a mile and 39.4 inches in a meter. Using factor analysis helps keep us from making costly mistakes by using the units (labels) as a check. Convert 1.5 days into seconds Convert 36 pounds into grams Convert 44 ounces (fluid) into mL.

33 Back Bires, 2004 Slide 33 A Little Mole The mole is an amount, much like a dozen. Referred to as Avogadro's number, the mole is equal to 6.02x10 23 things. Well find this number to be very handy later. For now, just know that when you see one mole, that equals 6.02x10 23 things. End of chapter 2

Download ppt "Back ©Bires, 2002Bires, 2004 Chapter 1 and 2: Matter and Change Measurements and Calculations How do we do what we do? What do we measure? How do we measure?"

Similar presentations

Ads by Google