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

2. 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

3. 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

4. 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

5. 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 _____
= g 1 mg

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

7. 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 ______________
g can be written as ______________
5.3 x 10 6 m
(move to left if positive exponent)
3.75 x g
(move to right if negative exponent)

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

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

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

11. IV. Significant Figures
3.3 cm
Measurement ___________

12. IV. Significant Figures
3.31 cm
Measurement ___________

13. IV. Significant Figures
3.00 cm
Measurement ___________

14. 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. =
4 sig. fig.

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

16. V. Significant Figures in Calculations
Addition/Subtraction:
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
170.97
Example: = ____________ => _________

17. 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
2.14
Example: ÷ 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.

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

19. 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

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

21. 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)

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

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

24. 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)

25. 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)

26. 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 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 = cm3
m = 4.5 g
V = cm3
Want:
d = ?
Yes it is gold (actual gm/cm3)

27. 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

28. 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)

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

30. X. Temperature vs. Heat

31. Unit 1B: The Mole

32. 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 atoms of carbon

33. 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

34. 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 Al2(SO4)3
2. K2CO 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

35. 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

36. 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 = mol NaOH
40 g NaOH
1.5 mol LiF

37. 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 particles
22.4 L
2.35 L Cl2 x 1 mol Cl2= mol Cl2
22. 4 L Cl2
3.5 mol H2 x 22.4 L H2= 78 L H2
1 mol H2

38. 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 x molecules?
6.02 x 10 23
6.02 x molecules NH3
= 2.1 x 1024
molecules NH3
3.5 mol NH3 x
1 mol NH3
1.204 x H2O x 1 mol H2O__________= 2.0 x mol H2O
6.02 x molecules H2O

39. 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

40. “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