INTRODUCTION

1. INTRODUCTION

2. CHEMISTRY • Is the study of the composition of materials and the changes that these materials undergo • Wood burns • Plants grow • Iron rusts • Bread bakes • Human digestion

3. Branches of Chemistry • Analytical chemistry • Answers questions about composition (what is the level of lead in that drinking water?) • Physical chemistry • Deals with energy transfers in chemical processes (how much energy is stored in a battery?) • Biochemistry • Deals with the analysis of chemical processes in living things (cellular process such as photosynthesis and respiration) • Organic chemistry • Study of carbon containing molecules (overlap with biochem) • Inorganic chemistry • Study of the chemistry of nonliving things (geological chemistry, the analysis of the rate of corrosion of metals)

4. The Scientific Method • State the problem • Make observations • Form a hypothesis • Perform an experiment • Variable you change – called the “independent” or “manipulated” variable • Make observations • Variable you observe – called the “dependent” or “responding” variable • Develop a theory (a well-tested explanation for observations) • Continue to experiment • Modify theory as needed • Develop a scientific law (a summary of many observations and experiments)

5. Chapter 2: Matter & Change • Matter – anything that has mass and volume • _____________________ – require measurements • Mass – amount of matter, measured in grams (g) • Volume – amount of space an object occupies, measured in: • liters (L) • milliliters (ml) • cubic centimeters (cm3) • _________________ – descriptions, not measured • Color, texture, odor, opaque, clear, dense, sound, taste (don’t taste anything in the lab!)

6. Classification of matter

7. Matter: Substances and Mixtures • Substances – pure composition • Two types: • __________ – only one kind of atom, such as iron (the symbol is on the periodic chart) • __________ – a combination of atoms, such as iron oxide (which can be separated) • Mixtures – contains 2 or more substances • ______________ – the same throughout; dissolved (saltwater) • ______________ – the sample may different depending where you take it in the container (orange juice with pulp)

9. States of Matter • Solid- rigid shape, incompressible, fixed volume, usually particles tightly packed and organized) • Liquid- indefinite shape but fixed volume, incompressible, flows, usually less tightly packed than solid (by a few %), less organized • Gas- indefinite shape and volume, flows, particles far away from each other, disorganized, compressible

10. CHEMICAL REACTIONS • Chemical Reaction • A change in which ________________________________________________________________________ • Reactant • A ___________________________________ • Product • A ___________________________________

11. Physical Properties and Changes • Physical Property – characteristics that describe a substance • Can be observed without changing the substances composition. • Physical Change – a change that does not affect the composition of the substance • The change may be reversible (melting) or irreversible (tearing)

12. Recognizing Chemical Changes • You know a chemical change has occurred if: • you see a change in color • there is a change in temperature (hot or cold) • a gas is produced • a solid precipitate is formed • an odor is produced

13. Separation of a MixtureDifference in physical properties can be used to separate mixtures. • Filtration- separates solid from a liquid (passes through a filter (pores)) • Distillation-liquid is boiled to produce a vapor that is then condensed into a liquid

14. Law of Conservation of Mass • During any physical or chemical change, the mass of the products is always equal to the mass of the reactants. • MASS is NEITHER CREATED or DESTROYED • Example: making pizza • Physical Change: the mass of the dough + sauce + cheese separately = the mass of all three together • Chemical Change: the mass of the ingredients before it’s cooked = the mass after it’s cooked (as long as it didn’t burn!)

15. Measurements Three types of measurement • Length: meter (m) • Mass: gram (g) • Volume: Liter (L)

16. Measurements expressed inSCIENTIFIC NOTATION • The expression of numbers in terms of M x 10n • M ≥ 1.00 and M < 10 • n is an integer Convert to whole numbers • 1 x 103 = _______ • 1 x 102 = _______ • 1 x 101 = _______ • 1 x 100 = _______ • 1 x 10-1 = _______ • 1 x 10-2 = _______ • 1 x 10-3 = _______

17. TRY THESE • 25,000 • ______________ • 0.00468 • ______________ • 0.0100 • ______________

18. Metric Conversions Kilo . Hecto . Deka . (unit) . Deci . Centi . Milli … Micro … Nano m g L Write the following in regular notation and scientific notation • How many mm in a m? _____________ • How many cm in a m? _____________ • How many m in a km? _____________

19. Measurements

20. ACCURACY/PRECISION • Accuracy • How close a measurement is to the actual value • Precision • How close a set of measurements are to each other

21. SIGNIFICANT DIGITS or Figures • This is a process used to determine the number of digits to round to when measuring an object. • Use this process when • Measuring mass (on the scale) – g, kg, etc. • Measuring volume (in a graduated cylinder) – ml L, etc. • Measuring length (with a ruler) – cm, m • Is used to communicate to other scientists how accurate your measurement is: • Does your scale measure to the hundredths place, tenths place or whole number? • Referred to as “Sig Figs”

22. How to determine the number of Sig Figs in a measured value • Atlantic-Pacific Method • A = decimal Absent, begin counting from right • P = decimal Present, begin counting from left • Try these: • 1,000 • 1 sig fig • 0.001 • 1 sig fig • 0.0010 • 2 sig fig • 1000.0 • 5 sig fig

23. Rules for Using Sig Figs • Multiplication/Division • Do all calculations, then round to the same number of digits as the number with the smallest number of sig figs • 4.56 x 1.4 = 6.384 • Round to 2 sig figs: 6.4 • 8.315/298 = 0.0279027 • Round to 3 sig figs: 0.0279 • Addition/Subtraction • Do the calculations, then round to the place of the number with the smallest number of decimal places • 12.11 + 18.0 + 1.013 = 31.123 • Round to 31.1 • 88.88 – 2.2 = 86.68 • Round to 86.7 (note: if the number after 6 is > 5, round up)

24. Rules for Using Sig Figs • Multiple step calculations • Use an overbar to keep track of the significant figures from step to step. • Round only when reporting the final answer • Example: 88.88 – 86.662.22 .024977 (calculator 88.88 88.88 answer) • Based on 2.22, round to 3 sig figs • .024977 • If the number after the place you want to round to is > 5, round up (in this case 7). Ignor the other 7. • Answer = .0250 • The zero after the 5 is significant. You must show it! = =

25. Temperature • Official” unit = Kelvin (K) • Most widely used for measurement = degrees Celsius (ºC) • Temp in Kelvin = ºC + 273 (º C = K -273) • Examples: • ice water = 0 ºC = 273 K • Boiling water = 100 ºC = 373 K • Body temp = 37 ºC = 310 K (about 99 ºF)

26. DENSITY • Mass • Amount of matter in an object • Volume • Amount of space an object occupies • Density • Mass per unit volume • Percent Error = /theoretical-experimental/ x 100% theoretical

27. Density Practice A copper penny has mass of 3.1g and volume of 0.35cm3. What is the density of copper?