1. Mr. Mellon Regents Chemistry
Unit 1A: Measurement
Mr. Mellon
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: Introduction to Chemistry

32. I. Important Terms Chemistry: Matter:
The study of matter and its changes
Any object that has a mass and volume (pretty much anything)

33. I. Important Terms Atom: -
Particle Diagram:
Smallest particle of an element that retains the properties of that element
Can not be broken down by a chemical reaction
2 different types of elements

34. I. Important Terms Compound: -
Particle Diagram:
Two ore more DIFFERENT ELEMENTS BONDED together
Can be broken down by a chemical reaction
1 type of compound

35. II. Phases of Matter The phases that matter is in depends on:
1. The space between atoms or molecules
2. The strength of the intermolecular force (IMF) between atoms

36. Takes shape of con-tainer definite Med-ium Vibrate and rotate
Phase
Shape
Volume
IMF
Movement
Solid (s)
Liquid (l)
Gas (g)
definite
definite
strong
vibrate
Takes shape of con-tainer
definite
Med-ium
Vibrate and rotate
Takes shape of con- tainer
Takes shape of con- tainer
weak
Vibrate and rotate

37. III. Classification of Matter
Element (Column A):
Compound (Column B):
Mixture (Column C):
Substance that cannot be changed into a simpler substance under normal conditions
Substance consisting of 2 or more different chemical elements that can be separated by chemical reactions (heat, electricity)
A physical blend of 2 or more substances that are NOT chemically combined (air, soup, tap water)

38. Aqueous Solution (aq):
a mixture of water and some other substance that can dissolve in water (salt water)
Element
Compound
Mixture

39. How Elements, Compounds, and Mixtures can be illustrated
N2 (g)
CO2 (g)
Air (N2, O2, H20, CO2(g)

40. Cool Animation

41. IV. Physical vs. Chemical Properties
Properties of an element or compound that can be observed or measured WITHOUT a chemical reaction
Ability of an element or substance to undergo a chemical reaction (bond breaking) and form a NEW substance

42. IV. Physical vs. Chemical Properties
Color, texture, odor, density, melting/freezing temps, solubility, volume, mass
VIDEO EXAMPLE
Reactivity, pH (acidity), ability to rust, decompose, ferment, combust
VIDEO EXAMPLE

43. IV. Physical vs. Chemical Properties

44. V. Physical vs. Chemical Changes
A change that does NOT produce a new substance, it just changes the position of the particles
Changing a substance into a NEW substance (bonds are broken and then new ones formed)
A color change may occur and a NEW s, l, or g is formed
A change that does NOT affect a substance’s chemical position

45. V. Physical vs. Chemical Changes
Any phase change (freezing, melting…), dissolving, mixing, cutting
Burning, rusting, fermentation, cooking/baking

46. V. Physical vs. Chemical Changes
liquid water freezes to ice
salt dissolves in water
liquid nitrogen in plastic bottle bursts open
rusting on a pan
potassium reacting with water to form potassium hydroxide and hydrogen gas
2K (s) + 2H2O (l) -> 2KOH (aq) + H2 (g)

47. IV. Physical vs. Chemical Properties

48. A chemical reaction ALWAYS results in a new substance
Chemical Reaction Equation:
4 H
+2 O
6 Total
4 H
+2 O
6 Total =
Reactants
Products
2 H2 (g) + O2 (g)  2 H2O (l)
Coefficient
# of atoms
A chemical reaction ALWAYS results in a new substance

49. Mass cannot be created or destroyed in a chemical reaction
Conservation of Mass:
Mass cannot be created or destroyed in a chemical reaction
The total mass of reactants equals the total mass of the products
Example (video clip)
Silver nitrate (AgNO3) and sodium chloride (NaCl) solutions before and after mixing

50. Examples: 50 g + Sodium+ 76 g 126 g Sodium Chloride Chlorine 
1) If 50.0 grams of sodium reacts with chlorine to form 126 grams of sodium chloride. How many grams of chlorine reacted?
2) If g of water is separated into hydrogen and oxygen gas, and the hydrogen gas has a mass of 20.0 g. What is the mass of the oxygen gas produced?
50 g +
Sodium+
76 g
126 g
Sodium Chloride
Chlorine 
178.8 g 
H2O 
20 g +
158.8 g
H2 + O2

51. Definite Composition (Homogeneous)
VI. Classification of Matter Chart
Matter
Substance
Definite Composition (Homogeneous)
Physically
Separable
Mixture of Substances

52. Nonuniform; distinct phases
VI. Classification of Matter Chart
Mixture of Substances
Homogeneous
Uniform throughout (air, tap water, gold alloy, salt water, Kool-Aid)
AKA – SOLUTIONS
Heterogeneous
Nonuniform; distinct phases
(soup, concrete)

53. Water (H2O), Carbon Dioxide (CO2)
VI. Classification of Matter Chart
Substance
Element
Iron, Sulfur, Oxygen
Compound
Water (H2O), Carbon Dioxide (CO2)
Chemically
Separable