Welcome to Ms. Raines Chemistry Class Please grab one pipe cleaner peace and three different colored beads from the from of the room. Ice Breaker Black bead: What do you do during your spare time?

White bead: Your least favorite activity? Blue bead: What word describes you & why? Green bead: What animal would you be and why? Red: What is your favorite food? Silver: Fun/interesting fact about your self. Gold: Pick one of the items above. Agenda Opener: Sign up for remind (see dry erase board) Receive/Review: Syllabus, Student information sheet, [MUST be signed and returned] book request form, [optional] week 1 element list, [home work make 15 flash cards]

periodic table, Finn Safety contract Supply list Science Fair info [Honors] Day 2 Opener: Complete safety symbol activity Safety Video: https://www.youtube.com/watch?v=3ELbwzqyuhs Safety Tour

Flinn Safety Rules Omit# 20, 30, 42, 55 Key point of Each rule Group 1: rule 1 - 8 Group 2: rule 9- 16 Group 3: rule 17 - 24 Group 5: rule 25- 36 Group 6: rule 37 -44 Group 7: rule 45-54 Scientific Method Scientific Method Scientific method is a logical, systematic approach to the solution of a scientific problem

1. 2. 3. 4. 5. 6. 7. Make observation Ask a question Form a hypothesis Experiment Analyze data Draw Conclusion Develop Theory or re-evaluate hypothesis

Scientific Method Scientific Method in the real world Teacher Example: Observation: Hairdryer is not working Question: Why is the hairdryer not working? Hypothesis: not plugged in Experiment: check to see if plugged completely in Data: turns on Conclusion: hairdryer was not plugged Group Example:

Scientific Theory vs Scientific Law Scientific Theory: a well tested explanation for observations and/or experimental result Attempts to explain why or how Can not be proven only can get stronger Kinetic Theory of matter stated atoms are in constant motion and explains how they move Scientific law: a statement that summarizes the results of many observations and experiment Does NOT try to explain why/how Gravity Unit 1: Measurement and Calculations

Chemistry Chapter 1 & 3 GA Performance Standards SCSh5.a. Trace the source on any large disparity between estimated and calculated answers to problems. SCSh7.b. Universal principles are discovered through observation and experimental verification. SCSh8.a. Scientific investigators control the conditions of their experiments in order to produce valuable data. SCSh8.b. Scientific researchers are expected to critically assess the quality of data including possible sources of bias in their investigations hypotheses, observations, data analyses, and interpretations. SCSh8.c. Scientists use practices such as peer review and publication to reinforce the integrity of scientific activity and reporting. SCSh8.d. The merit of a new theory is judged by how well scientific data are

explained by the new theory. SCSh9.d. Establishing context Introduction What is Chemistry Study of matter Matter is anything that has mass and occupied space Examples: Things that are NOT matter: Measurement Types of measurement Units of measure 12 metric prefixes

Significant Figures Types of measurement Qualitative measurement based on some quality or characteristic Deals with descriptions. Data can be observed but not measured. Colors, textures, smells, tastes, appearance, beauty, etc. Qualitative Quality Blue liquid, soft fabric, cold room Types of measurement Quantitative measurement is something that is measurable in quantity

Deals with numbers. Data which can be measured. distance, volume, mass, speed, time, temperature, cost, ages, etc. Quantitative Quantity 25.0 g, 48 mL, 3 days, 45 miles Measurement Measuring with SI Units The metric system units are based on multiples of 10 and can be converted easily International System of Units (SI) is a revised version of the metric system The five SI standard units commonly used by chemists are the meter, the kilogram, the

Kelvin, the second, and the mole. SI Base Unit Quantity SI standard unit Base unit** length Mass Temperature Time Volume

Meter (m) Kilogram (Kg) Kelvin (K) Seconds (s) Decimeter cubed (dm3) Meter (m) Gram (g) Kelvin (K) Seconds (s) Liter (L) Amount of a substance

Mole (mol) Mole (mol) Heat and Energy Joules (J) Newton (N) Joules (J) Newton (N) Force and weight Metric Prefixes

Added to the base unit to make it larger or smaller Changes by powers of 10 Physical science prefix pneumonic: King henry died by drinking chocolate milk kilo, hecto, deca, base, deci, centi, milli Chemistry has 6 more you may see Tera, Giga,mega, kilo, hecto, deca, base, deci, centi, milli, micro, nano, pico T, G,M, k,h, da, base, d, c, m, , n, p Pneumonic The Great Mad king Henry died by drinking chocolate milk under nicks porch.

Name Tera Giga Mega Kilo Hecto Deca Base Deci Centi Mili Micro Nano Pico

Symbol T G M k h da (g, l, m, s) d c m u n P

Meaning 1012 109 106 103 102 101 100 10-1 10-2 10-3 10-6 10-9 10-12 Metric Prefix

meanings Limits of Measurement Precision is a gauge of how exact a measurement is. Precise measurements are close to each other MUST have more than one measurement Accuracy is the closeness of a measurement to the actual value of what is being measured An accurate measure is close to the true or expected value MUST have true or expected value NOT Accurate (not near center) NOT precise (not near each other)

NOT Accurate (not near center) Precise (close to each other) Accurate (Near center) Precise (close to each other) Sally Annie Travis Jeff 1.95 g/ 2.69 g/ 3.12 g/ 2.71 g/ cm3 cm3 cm3 cm3 1.89 g/cm3

2.73 g/ 2.70 g/ cm3 cm3 1.92 g/ 2.65 g/ 2.25 g/ cm3 cm3 cm3 To the right is the data collected by students during a lab. Actual Density of Aluminum is 2.70 g/cm3

1. Which students data is accurate and precise? 2. Which students data is accurate but NOT precise? 3. Which students data is NOT accurate but IS precise? 4. Which students data is NEITHER accurate nor precise? 1.Annie 2.Jeff 3.Sally 4.Travis In general, a calculated answer cannot be more precise than the least precise

measurement from which it was calculated. Example: if measuring with a standard ruler and recording the measurements in cm you measurement can only have two decimal places. The line below would be measured at 3.79 cm. _____________ Ruler Example The blue line would be recorded to be 13.3_ cm long. With the _ being the estimated digit. 13.30 cm, 13.31 cm would both be valid measurements. 13.300 cm or 13.310 cm would NOT be valid

Read to the unit you are certain of, then estimate one more place. Graduated Cylinder In order to read the graduated cylinder correctly, it must be placed on a stable surface such as the desk top of the work area And you MUST be at eye level with the meniscus To determine the volume of liquid use the number that is directly at or below the bottom of the meniscus Graduated Cylinder

You must estimate one more digit that you can precisely measure. The graduated cylinder pictured measured in mL and 10th of a mL. The blue liquid would have a volume of 1.11 mL or 1.12 mL. A measurement of 1.110 mL or 1.1120 mL is more precise than the tool allows. Read to the unit you are certain of, then estimate one more place. Significant Figures and Calculations Complete Significant Figure activity

to identify the significant figure rules Significant Figures Rules Significant Digits - Number of digits in a figure that express the precision of a measurement instead of its magnitude. Significant figures are just a way of keeping track of our level of precision so that when we do calculations with our data, we don't end up exaggerating it Significant Figures Rules Version 3 Rules for determining whether a digit in a stated value is significant Pacific

Atlantic If the decimal point is present, start counting digits from the Pacific (left) side, starting with the first non-zero digit. Example: 0.003100 (4 sig. figs.) If the decimal point is absent, start counting digits from the Atlantic (right) side, starting with the first non-zero digit. Example: 31,400 ( 3 sig. figs.) Significant Figures Rules Version 2 Rules for determining whether a digit in a stated value is significant Ignore leading zeros. ( 0.0053 has 2 sig figs)

Ignore trailing zeros, unless they come at the end of a number AND there is a decimal point. 35200 has 3 sig fig 35200. has 5 sig fig 35.200 had 5 sig fig Everything else is significant Defined quantities and counted quantities have unlimited number of significant figures. 1 ft = 12 in has sigfigs. Significant Figures Rules Version 1 Rules for determining whether a digit in a measured value is significant

1. Nonzero digits are significant. 5.23 has 3 significant figures 2. Zeros between nonzero digits are significant. 5001 has 4 significant figures. [Sandwich rule] 3. Zeros at the end of a number and to the right of a decimal place are significant. 1.0100 has 5 significant figures 4. Zeros in front of nonzero digits are not significant, they are only place holders. In general start counting at the 1st NON zero number 0.000099 has 2 significant figures. Start counting at 1st non-zero number. 5. Zeros to the left of an understood decimal point are not significant, they are only place holders. 55000 has 2 significant figures 6. Defined quantities and counted quantities have unlimited number of significant figures. 1 ft = 12 in has sigfigs. Significant Figures Examples

a) b) c) d) e) f) g) h) i) 2.03 1.0 0.00860 4.50 x 1012

5.1020 780 780,000 0.78000 50. a) b) c) d) e) f) g) h) i)

3 2 3 3 5 2 2 5 2 When rounding first decide how many significant figures the answer should have. Next round to that number of digits , counting from the left. If the number to right of the last significant digit is 4 or less

round down, if it is 5 or up round up. Make sure you dont significantly change the value of the original number. Cant round 556 to 6 must be 600 Example: 5,274.827 6 significant figures: 5,274.83 4 significant figures: 5,275 2 significant figures: 5300 A. B. C. D.

Practice Round 2.3567 to 3 significant figures Round 56913 to 4 significant figures Round 2.0132 to 2 significant figures Round 5678 to 2 significant figure Answers A. B. C. D. 2.36 56910

2.0 6.0 x 103 Significant Figures and Calculations With multiplication and division the calculation should be rounded to the same number of significant figures as the measurement with the LEAST number of significant figures Example: Calculator give 0.931 12 has only 2 significant figures so the answer must have only 2 significant figures Answer MUST BE 0.93 Significant Figures and Calculations

With addition and subtraction the answer must be rounded to the same number of DECIMAL places as the value with the lease number of decimal places. Example: 2.450 14.2 Calculator gives: -11.75 But must be rounded to 1 decimal place so answer is -11.8 Practice Perform the following calculations and round correctly. 2.680 x 0.0051 3.120 / 6

2.45 + 550.9 9.056 4.25 Calculator Rounded = 0.013668 = 0.014 = 0.52 = 0.5 = 553.35 = 553.4 = 4.806 = 4.81 Scientific Notation When writing very large or very small numbers, scientists use a kind of shorthand called scientific notation.

This is a way of writing a number without so many zeros. Example 1: The speed of light is about 300,000,000 m/s Or 3.0 x 108 Example 2: The mass of a proton is 0.000000000000000000000001673 Or 1.673 X 10-24 All you do is move the decimal so that you only have one number before the decimal. 850,000,000.0 850000000.0 = 8.5 x 108

For large numbers the exponent is positive!! 0.000,000,025 0.000000025 = 2.5 x 10-8 For small numbers the exponent is negative!! Scientific Notation Examples 0.007899 = ? Small number = - exponent 7.899 x 10 -3 898745.30 = ? Large number = + exponent 8.9874530 x 10 5 0.00003657= ? Small number = - exponent

3.657 x 10 -5 531120 = ? Large number = + exponent 5.31120 x 10 5 Getting numbers out of Scientific Notation Look at the exponent of the number to determine if it needs to get smaller or larger Positive exponent means the number get larger so the decimal moves to the right Negative exponent means the number gets smaller so the decimal moves to the left Add zeros to fill in any BLANK spaces

Example 1: 2.35 x 105 The exponent is positive so the number needs to get larger 2 3 5 . 2 3 5 0 0 0. or 235000 Example 2: 8.68 x 10-4 The exponent is negative so the number needs to get smaller . 8 68 0. 0 0 0 8 6 8 or 0.000868

Scientific Notation Examples 3.256 x 104 positive exponent = large number 3256 9.78 x 109 positive exponent = large number 978000000000 5.24 x 10-3 Negative exponent = small number 0.00524 2.41 x 10-7 Negative exponent = small number 0.000000241 Dimensional Analysis is a way to analyze and solve problems using the units of the measurements. It is converting one thing to another without changing its value

Requires equality statements and conversion factors. The key to dimensional analysis is to set it up so that the UNITS cancel. All numbers must have a unit! No Naked Numbers!!!! Many quantities can usually be expressed different several different units Equality Statement shows how two (or more) different units are related Example: 1 dollar = 4 quarters Conversion factor is a ratio of equivalent

measurements. Example: Whenever two measurements are equivalent, a ratio of the their measurement will equal 1 When a measurement is multiplied by a conversion factor, the number changes, but the actual size of the quantity measured remains the same. Example: 2.0 hours = 120 minuets = 7200 seconds when using conversion factors the final answer has the same number of significant figures as the starting number

Equality Statements that you should know. 1 min = 1 hour = 1 day = 1 week = 1 year = 1 year = 1 foot = 1 yard = 60 60 24 7

52 365 12 3 Metric Conversions seconds minuets Largest unit = 1 Smallest unit = 10^(x) hours days x= lg exponent- sm exponent weeks Example

days ____ mL = _____ ML inches (6- -3) 10 ml = 1 ML feet 10(9) ml = 1 ML Steps for using dimensional analysis. 1. Write equality statement for units needed in problem 2. Write given number and unit then a fraction bar. 3. The unit you are getting rid of goes on

bottom 4. The unit you are going to goes on top 5. Fill in the fraction with the values from the equality statement and solve Example 1 If a move is 1.48 hours long how many minutes are you in the theater? Step 1: 60 minutes = 1 hour Step 2: 1.48 hours ----------Step 3: Step 4: Step 5: Example 2 ~~ two step problem

If a movie is 1.75 hours long how many seconds are you in the theater. We dont have one equality statement that relates seconds and hours so we used two Step 1: 1 hour= 60 minuets, 1 minuet = 60 seconds Step 2-3: Step 2-4 Step 5: Example 3 ~~ three step problem A sample is 3.324 x 108 minuets old how many years old is it? step 1: 1 year = 365 days, 1 day = 24 hours, 1 hour = 60 minuets step 2-4:

step 5: = 632.4 years Name Tera Giga Mega Kilo Hecto Deca Base Deci Centi Mili Micro

Nano Pico Symbol T G M k h da (g, l, m, s) d c m u

n P Meaning 1012 109 106 103 102 101 100 10-1 10-2 10-3 10-6 10-9

10-12 Metric Conversions Largest unit = 1 Smallest unit = 10(x) x= lg exponent- sm exponent Example ____ mL = _____ ML Mega is larger so ML = 1 10(6- -3) ml = 1 ML 10(9) ml = 1 ML Example 4 (metric) If a student runs 37600 dm how many Mm is it?

recall that to set up equality statement for metrics you need to write 1 Largest unit= 10^(x)Smallest unit and x = lg exponent- sm exponent 1 Mm = 10(6-(-1)) dm 1 Mm = 107 dm 37600 dm 1 7 10

= 3.76 x 10 -3 Mm = 0.00376 Mm Example 5 (metric) If a object has a volume of 0.00564 daL how much is the volume in pL? recall that to set up equality statement for metrics you need to write 1 Largest unit= 10(x)Smallest unit and x = lg exponent- sm exponent 1 daL = 10(1-(-12)) pL 1 daL = 1013 pL 0.00564 daL

13 10 1 = 5.64 x 10 10 pL Remember NO NAKED NUMBERS!!!! Show ALL units at every step. Round at the end.

Density Density is a unit of mass per unit of volume SI Units of density: g/mL or g/cm3 or Kg/m3 Density = mass . volume d=m v A block of work has a volume of 28.5 m3 and a mass of 14.05 Kg. What is its density? Given v= 28.5 m3 m= 14.05 Kg

D=? Solving word problems Example 1: Robin measured the mass of a metal cube to be 25.48 g and the cube measures 3.0 cm on each side. What is the cube density? Given Mass= 25.48 g Length= 3.0 cm Width = 3.0 cm Height = 3.0 cm Volume= ? Density= ? Equation

Solve V= W(L)H V= (3.0 cm)(3.0 cm)(3.o cm)= 27 cm3 D= A marble has a mass of 12.48 grams and when placed in a graduated cylinder with 20.0 mL the volume increased to 24.5 mL. What is the marbles density? Given: m= 12.48g

d =? v initial= 20.0 mL v final= 24.5 mL v = v f vi Equation: d = m/v Solve: v = 24.5 ml 20.0 mL d = (12.48 g / 4.5 mL) =2.7733 g/mL d = 2.77 g/mL Using Density Rearranging the density equation First get it in a liner format by multiplying by volume Density x Volume = mass

If wanting volume then divide by density Volume = mass . volume These equations can be used to find information using known density values The density of copper is 8.920 g/cm3 if you have 52.75cm3 sample of copper how much does it weigh? Given: d = 8.920 g/cm3 v = 52.75cm3 m=? Equation: d = m/v or d(v) = m Solve: mass = (8.920 g/cm3 )(52.75cm3 ) =

mass = 470.5 g A 250.0 g sample of lead occupied what volume? [density of lead is 11.340 g/cm3] Given: m = 250 g d = 11.340 g/cm3 v=? Equation: d = m/v or v = m/d Solve: v = 250.0 g / (11.340 g/cm3) v = 22.05 cm3