Mrs. Pelc Regents Chemistry Unit 1A: Measurement Mrs. Pelc Regents Chemistry

Unit 1A: Objectives 1. Understand selected metric units and prefixes and know how to find them on the NYS Reference Tables 2. Be able to set up and calculate simple conversions 3. Convert number to and from scientific notation 4. Identify, measure, and calculate with significant figures 5. Perform scientific calculations: Percent Error and Density 6. Understand the difference between temperature and heat and Kelvin and Celsius Temperature scales

I. Metric System D C SELECTED BASE UNITS SELECTED PREFIXES In chemistry (and all sciences) the International System of Units (SI) is used. It is a universal set of units that allows scientists from around the world to be consistent with each other. TABLE _____ is a list of ________________________ The SI system is a decimal system, meaning prefixes are used to change SI units by a power of 10 TABLE _____ is a list of _________________________ D SELECTED BASE UNITS C SELECTED PREFIXES

II. Metric Conversion C 1 10 -3 g = 1 g ____ mg = _____ g Often we will want to convert from one unit of measurement to another. To do so you need a conversion factor. Use TABLE _____ to convert from different metric sizes Example: Convert 55 milligrams (mg) to grams (g): Step 1: Create a conversion factor: C g = 1 g 1 10 -3 ____ mg = _____ g

II. Metric Conversion 1 10 -3 10 -3 g 55 mg x _____ = 0.055 g 1 mg Step 2: Multiply by your conversion factor. Place the unit you want to convert to in the numerator (top spot) and what you want to cancel in the denominator (bottom spot) 1 10 -3 ____ mg = _____ g 10 -3 g 55 mg x _____ = 0.055 g 1 mg

II. Metric Conversion Examples: 15 kg to grams: 2) 550 mg to g: 3) 0.003 m to cm: 4) 642 cg to kg: 15,000 g 0.55 g 0.3 cm 0.00642 kg

III. Scientific Notation Throughout the year we will encounter VERY small and BIG numbers. We use scientific notation to represent these numbers in powers of tens. (Scale of the Universe Animation) Move decimal in-between the first non-zero digits 5,300,000 m can be written as ______________ 0.00000375 g can be written as ______________ 5.3 x 10 6 m (move to left if positive exponent) 3.75 x 10 -6 g (move to right if negative exponent)

III. Scientific Notation Examples: Write the following in scientific notation 1) 34500000 kg = 2) 7561000 m = 3) 0.000301 cm = 4) 0.000000002091 mg = 3.45 x 10 7 kg 7.561 x 10 6 kg 3.01 x 10 - 4 kg 2.091 x 10 - 9 kg

III. Scientific Notation Examples: Convert the following from scientific notation to standard 4.51 x 10 3 g = 8.91 x 10 - 4 km = 5.12 x 10 - 6 kg = 4) 5.234 x 10 7 = 4,510 g 0.000891 kg 0.00000512 kg 52,340,000 cm

IV. Significant Figures Record all numbers that can be known precisely, PLUS a last estimated number 3 cm Measurement ___________

IV. Significant Figures 3.3 cm Measurement ___________

IV. Significant Figures 3.31 cm Measurement ___________

IV. Significant Figures 3.00 cm Measurement ___________

IV. Significant Figures Rules for Counting Significant Figures: Nonzero digits are always significant All final zeros after the decimal point are significant Zeros between two other significant digits are always significant Zeros used solely as placeholders are not significant Rule 1 Ex) Rule 2 Ex) Rule 3 Ex) Rule 4 Ex) 313 = 3 sig. fig. 2.000 = 4 sig. fig. 2.10 = 3 sig. fig. 2002 = 4 sig. fig. 0.200 = 3 sig. fig. 2000 = 1 sig. fig. 0.002020 = 4 sig. fig.

IV. Significant Figures Examples: How many significant figures does each of these measurements have?   1) 3.1 m ______ 3) 1.20 x 10-4 km _______ 3.0001 kg _______ 4) 0.007060 cm ________ Try problems 1, 2, 4, 6, 7, 8 in Section IV A from pg. 3 in the work packet 3 2 5 4

V. Significant Figures in Calculations Addition/Subtraction: 10.52 + 349.0 + 8.240 Step 1: Answer: ____________ Step 2: New Answer: ________ Add the numbers Least # of decimal places 367.76 Round the answer to the same # of DECIMAL PLACES as the measurement with least # of decimal places 367.8 Least # of decimal places 170.9711 170.97 Example: 3.21 + 123.1101 + 44.651 = ____________ => _________

V. Significant Figures in Calculations 3 sig fig 2 sig fig Multiply/Divide: 7.55 x 0.34 Step 1: Answer: ____________ Step 2: New Answer: Multiply the number Least # of sig figs. 2.567 Round the answer so that it has the same # of sig figs as the measurement with the least 2.6 4 sig fig 3 sig fig 2.1433333 2.14 Example: 4.501 ÷ 2.10 = _________________ = ____________ Note: Constants, conversion factors, or exact numbers (e.g. number of people in a room) are not taken into account when performing calculations, only measured quantities.

V. Significant Figures in Calculations Complete problems 17, 19, 22, 24 in Section IV C on pg 4 in the work packet

VI. Types of Measurement QuaNtitative Numbers with units Qualitative Descriptive, no numbers (color, texture, phase) In chemistry (science), it is important to take quantitative and qualitative measurements

VII. Precision vs. Accuracy How close a series of measurements are Accuracy How close your measurement(s) are to the actual/accepted value

VII. Precision vs. Accuracy Precise, not accurate Accurate, not precise Precise and accurate In General: We want measurements to be accurate (close to accepted) and precise (similar with each other)

VIII. Percent Error Equation (see ref. tabs):

% error = - 8.281 % (negative means measured was smaller) Example: A student finds the density of copper to be 8.218 g/cm3. The actual density of copper is 8.960 g/cm3. Find the percent error in her measurement. d m = 8.218 g/cm3 d a = 8.960 g/cm3 % error = % error = - 8.281 % (negative means measured was smaller)

IX. Density Mass - Amount of matter (atoms and molecules) an object contains (mass ≠ weight) - Measured in: grams (g) Volume Amount of space an object takes up Measured in: mL (gases and liquids) or: cm3 (solids)

Diet Pepsi vs. Regular Pepsi IX. Density Density How much matter (atoms and molecules) pack in an object Measured in: g/mL or g/cm3 Diet Pepsi vs. Regular Pepsi Equation (see ref tabs) d = density (g/cm3 or g/mL) m = mass (g) V = volume (cm3 or mL)

Givens: Set up and Answer: Want: Yes it is gold (actual 19.320 gm/cm3) Example: A person brings in what he thinks to be a gold ring to a jewelry store. The ring has a mass of 4.5 g and a volume of 0.233 cm3. Is this a gold ring? (Hint: find the density and compare it on Table S). Givens: Set up and Answer: m = 4.5 g V = 0.233 cm3 m = 4.5 g V = 0.233 cm3 Want: d = ? Yes it is gold (actual 19.320 gm/cm3)

IX. Density 2) m = 2406 g 3) V = 15.1 mL; d = 7.98 g/mL Examples 2) m = 2406 g 3) V = 15.1 mL; d = 7.98 g/mL 4) V = 790. cm3; d = 1.23 g/cm3

X. Temperature vs. Heat Temperature: The average kinetic energy (how fast matter is moving) of matter Heat: The total amount of movement in a sample Absolute Zero: The temperature at which all molecular movement stops (cannot happen)

X. Temperature vs. Heat K = oC + 273 K = Kelvin oC = degrees Celsius Temperature Scales (see Ref Tabs): K = oC + 273 K = Kelvin oC = degrees Celsius Convert 200 degrees Celsius to Kelvin: oC = 200oC K = ? K = 200oC + 273 = 473 K Remember: Kelvin uses bigger values, but everything is positive

X. Temperature vs. Heat

Unit 1B: The Mole

I. What is a mole? 6.02 x 10 23 6.02 x 10 23 atoms of carbon A mole (mol) represents a certain amount. Just like a dozen bagels represent 12 bagels. 1 mole = ________________________ particles (Avogadro’s Number) A particle can mean molecule, an atom, an ionic compound. 12 amu = 1 carbon atom 12 g of carbon = 1 mole of carbon atoms = __________________________ 6.02 x 10 23 6.02 x 10 23 atoms of carbon

II. Gram Formula Mass 1 mole Cl = Gram Atomic Mass (GAM) – mass of 1 mole of atoms in an element Example: Cl: Subscripts – represent the number of moles of each atom Examples: Al2O3 – _______ moles of Al ions, ______ moles of O ions   Ca(NO3)2 - ______ mole of Ca ions, ______ moles of N atoms, ______ moles of O atoms 1 mole Cl = 35 g of Cl (round to the nearest whole number) 2 3 1 2 6

II. Gram Formula Mass Gram Formula Mass (GFM) – To find the GFM you add the masses of all of the elements in the compound or molecule. Examples: H2O 3. Al2(SO4)3 2. K2CO3 4. Zn3(PO4)2 • 4H2O Mass of 1 mole of a compound 2 (1g) + 1 (16g) = 18 g 2 (27g) + 3 (32g) + 12(16 g) = 342 g 2 (39 g) + 1 (12 g) + 3(16 g) = 138 g 3 (65g) + 2 (31g) + 8(16 g) + 8 (1g) + 4 (16 g) = 457 g

III. Grams   Moles (1 Step Mole Problems) Moles  Grams:   Example: How many grams are present in 40.5 mol of sulfuric acid (H2SO4)? Example: What is the total mass of 0.75 mole of SO2? STEP 1: H2SO4 = 2 (1g) + 1 (32g) + 4 (16 g) = 98 g STEP 2: 40.5 mol H2SO4 x 98 g H2SO4 = 3970 g 1 mol H2SO4 48 g SO2

III. Grams   Moles (1 Step Mole Problems) Example: How many moles are equivalent to 4.75 g of sodium hydroxide (NaOH)? Example: How many moles are equivalent to 39 g of LiF? STEP 1: NaOH = 1 (23g) + 1 (16g) + 1(1g) = 40 g NaOH STEP 2: 4.75 g NaOH x 1mol NaOH = 0.119 mol NaOH 40 g NaOH 1.5 mol LiF

IV. Mole  Volume(1 Step Mole Problems) 1 mole of any gas is equal to ___________ (at standard pressure and temperature - STP) 1 mole = _____________________________ = _______________ of any gas at STP Example: Find the number of moles of 2.35 L of chlorine gas. Example: Find the amount of liters of 3.5 moles of hydrogen gas. 22. 4 L 6.02 x 10 23 particles 22.4 L 2.35 L Cl2 x 1 mol Cl2= 0.105 mol Cl2 22. 4 L Cl2 3.5 mol H2 x 22.4 L H2= 78 L H2 1 mol H2

V. Mole  Particles (1 Step Mole Problems) 1 mole of any substance is equal to _____________________ particles (Avogadro’s Number) Example: How many molecules of water are there in 3.5 moles of NH3? Example: How many moles of H2O are there if there are 1.204 x 10 34 molecules? 6.02 x 10 23 6.02 x 10 23 molecules NH3 = 2.1 x 1024 molecules NH3 3.5 mol NH3 x 1 mol NH3 1.204 x 10 34 H2O x 1 mol H2O__________= 2.0 x 10 10 mol H2O 6.02 x 10 23 molecules H2O

VI. Avogadro’s Hypothesis Any 2 samples of gas at the same temperature, pressure, and volume have the same number of particles Examples: Which of the following would occupy the same volume as 2.5 mol of O2 at STP? (a) 1.5 mol of CO2 at STP (b) 3.5 mol of CH4 at STP (c) 3.0 mol of H2 at STP (d) 2.5 mol of Ne at STP (a) 33.6L of CO2 at STP (b) 78.4L of CH4 at STP (c) 67.2L of H2 at STP (d) 56.0L of Ne at STP

“Y” of Chemistry VOLUME MASS MOLES MOLECULES ATOMS x 22.4L x molar mass ÷ 22.4L ÷ molar mass MOLES ÷ 6x1023 x 6x1023 MOLECULES ÷ # atoms in formula x # atoms in formula ATOMS