Chapter 1 and 2: Matter and Change Measurements and Calculations

1. Read Text pages 4-61 Chapter 1 and 2:Matter and ChangeMeasurements and Calculations All our science, measured against reality, is primitive and childlike - and yet it is the most precious thing we have.-Albert Einstein How do we do what we do? What do we measure? How do we measure?

2. Inferences & Observations • Observation • Detected with one’s senses or lab equipment • No explanations • Inference • Explains what is observed based on prior knowledge. • Like a hypothesis.

3. States of Matter and Changes • Four primary states of matter… • solid, liquid, gas, and high energy plasma. • We observe physical or chemicalproperties: • 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. Changes of Matter • Physical change • Substance does not change, just the form. • Examples? • Chemical change • Original substance is lost, and a new substance is formed. • Examples? • Which kinds of changes are these?: • Burning, freezing, vinegar + baking soda, and opening a soda bottle, boiling, rusting?

5. Chemical Reactions • “reactants” are placed on the left side of a “chemical reaction equation”: • “product”, is produced, placed on the right. • Remember: • “Reactants react to produce products” • Law of the Conservation of Mass: • Mass is not created or destroyed in a chemical reaction.

6. Energy Transfer • The Law of Conservation of Energy • energy cannot be created or destroyed. • Energy can change form and be stored 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. From the macro to the micro… • We can classify matter in terms of its complexity: • Mixture • Collection of two or more physically different compounds. • Compounds don’t change their properties. • Two types of mixtures: • Homogeneous (“homo” = “same”)… • Homogenous mixtures cannot be easily separated • Heterogeneous (“hetero” = “different”) • Heterogeneous mixtures can be separated with simple mechanical processes. These often have “phases.”

8. A step closer to the micro… • Compound • substance made up of two or more pure elements. • Water is a compound. Why? • Table salt is a compound. Why? • A compound cannot be separated from its elemental makeup • (without destroying the compound) • Compounds have very different properties than their elements.

9. The micro… • Elements • simplest form of pure substances. • 112 known elements, found on the periodic tableof 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. • different properties (due to different arrangements of its atoms.) • Diamond and graphite are allotropes.

10. The very micro… • Atom • smallest thing with the properties of something in the macro. • 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 • (we’ll 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 ( + or - ) • Atoms with a different number of electrons.

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

12. The Periodic Table… • A model that groups elements with similarproperties. • Vertical columns Groups of elements with similar properties. • Horizontal rows  Periods of elements with similar atomic mass.

13. Three Main Areas of the Periodic Table • Metals, left side. • are malleable, ductile, and good conductors of heat and electricity. • Nonmetals, right side. • solids are brittle and poor conductors of e- and heat. Metalloids have some characteristics of both metals and nonmetals

14. The Noble Gasses He Ne • The Noble gasses are found on the far right of the P-table. The Nobles are: • …Mostly unreactive. • (*not entirely unreactive*) • …Gasses at room temperature. • …Mined from gas pockets in the ocean • …Produce bright emissions when electrified • …Have filled octets. (octets?) End of chapter 1

15. Theory vs. Law • Theory • an explanation of observations of natural phenomena. • explains why things do what they do. • cannot be proven, but it has never been disproven. • If a theory is disproven, it must be modified or rejected. • Law • a description of fact. • describes whatwillhappen. • Because a law is a description of fact, it cannot be broken. End of chapter 1…

16. Measuring…Standard Units Table on 34 • Kilogram – mass (kg) • in lab, we will usually measure in grams,(g) • Liter – volume (L) • in lab, we will usually measure in milliliters,(mL) • Meter – length (m) • Second – time (s) • Kelvin – temperature (K) • AMU – atomic mass (amu) (more later) • Mole – amount of substance (mol) (more later) The Kelvin Scale: 0K = absolute zero 273.15K = water freezes 373.15K = water boils Celsius Kelvin oC + 273 = K

17. Measuring…SI Prefixes Use “EE” or “EXP” or ( # x 10 ^ # ) • kilo – (k) x103 (x 1,000) • kilogram = 1000 grams • milli – (m) x10-3 (x 1/1,000) • milliliter = 0.001 liters • micro – (μ) x10-6 (x 1/1,000,000) • Mega – (M) x106 (x 1,000,000) • centi – (c) x10-2 (x 1/100) • centimeter = 0.01 meters • nano – (n) x10-9 (x 1/1,000,000,000) About 35 ml 1.0ml = 1.0cc cc = cm3  Not this nano … This one 

18. Derived Units See page 37 • Derived units - products of standard units. • Volume, V • The SI unit is the cubic meter m3. m3 is huge, so we use the L, or mL (cm3). • Density, ρ (rho) • The amount of mass that is crammed into a certain volume. ρ = m / V. • Each compound has a unique density. • (g/ml, g/cm3, or kg/m3 ) • Temperature Effects? • How does temperature affect mass, volume, density? Osmium (#76) is the densest element on the planet

19. Metric conversions • Multiply / divide by powers 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) base unit 10-2 Always show units! 12 x 102 = 1200 cm 10-3 base unit 345 x 10-3 = 0.345 g

20. Measuring…Scientific Notation • Scientific Notation • a form of shorthand very small or very large numbers. • 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, 135000 becomes 1.35 x105. • Can you figure what 4500 is? • How about moving the decimal the other way…try 0.00056. x1023 6.02 6.02x1023 We moved the decimal five places

21. Scientific Notation Practice • Convert the following to scientific notation: • Convert the following to floating point notation:

22. Measuring…Significant Figures • “Sig Figs” • how many digits to include in measurements and calculations. • It is a measure of how precise our equipment is. • Rules: (shortcut coming) • Allnon-zero numbers are significant. • 1, 2, 256, 952456 • Zeros between significant numbers are significant. • 303, 50034, 1001 • Zeros to the RIGHT of a decimal are significant. • 3.000, 24.0, 31.0000, 35.520 • Zeroes to the LEFT of a decimal are NOT. • 4000, 256000, 10, 2400, 1 000 000 000 Text page 47 for help

23. Pacific or Atlantic? A little trick for “sig figs” • Decimal Present? • Count from thePacific • Decimal Absent? • Count from the Atlantic 31.80 0.0020 56430 10000

24. Sigfigs…more examples Can you see why significant digits are important? • 2450 has 3sig figs. • 245.0 has 4sig figs • 0.082 has 2sig figs • 0.0820 has 3sig figs • 1010 has …? • 45.30 has …? • Exceptions to the rules: • Fractions, Counting , When the teacher tells you to ignore them Figure the number of sigfigs for the following: • 6.781 0.0563 1200 63003 1.42x10-2 • 4 3 2 5 3

25. Metric Conversions Practice • Convert the following: • 2.52 meters to millimeters • 8.47 millimeters to meters • .0250 micrometers to meters • .995 kilometers to meters • 51.2 m to km • 5.24 x 10-2m to mm • 8.91 x 103km to m • 4.21 x 10-4m to mm • 1.23 x 10-2µm to m

26. Two types of measurements, data • Qualitative • a description of an object. • “blue” “sticky” “smelly.” • Quantitative • data expressed with numbers and units. • “42.3 kilograms” “14 milliliters” “3.80 grams.” • Always include units (g, mL, etc) • To accurately describe a compound or solution in chemistry: • use color, transparency, and texture/state. • “A colorless, clear, liquid.” ? 

27. Accuracy and Precision • Accuracy • closeness to an accepted value. • Precision • closeness of a set of measurements. • We demand precision. How?

28. Relationships Page 55 • Variable Relationship • If one variable changes as another changes • 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.

29. A Little Mole • The mole is an amount, much like a dozen. • Referred to as Avogadro's number, the mole is equal to 6.02x1023things. • We’ll find this number to be very handy later. For now, just know that when you see one mole, that equals 6.02x1023things. End of chapter 2