1. The Basics

2. The Basics Observations Measurement and Units
Uncertainty and Significant Figures
Unit Conversions
Chemical and Physical Changes
Density
Simple Error Analysis

3. Be the Younger Smarter One

4. Observations
All of chemistry revolves around detailed observation in lab
Two Types of Observation
Quantitative
Having to do with quantity, numbers
Measurements
Qualitative
Having to do with qualities, words and descriptions.

6. Measurements All measurements must contain two things: Some number
Some unit
Both are equally important and neither can be left out.

7. The Unit Units are equally important as the number
A number without a unit is meaningless.

8. Use Units or Else Air Canada Flight
Pilots used a conversion factor with no units.
Unfortunately it was the wrong conversion factor
Needed conversion between L to kg
Accidentally given L to lbs
Ended up with half the fuel they needed.

9. Or Just Thoroughly Embarrass Yourself…
Cents vs Dollars.mp4

10. Standard Units Système International d'unités (SI)
International System of Units
Length – m – meter
Mass – kg – kilogram
Time – s – second
Temperature – K – Kelvin
Volume – L – liter
Energy – J - Joule

11. SI Prefixes Name Symbol Scientific Notation giga- G 1x109
mega- M 1x106
kilo- k 1x103
Base unit - 1x100
deci- d 1x10-1
centi- c 1x10-2
milli- m 1x10-3
micro- μ 1x10-6
nano- n 1x10-9
pico- p 1x10-12

12. The Number
How long is the black bar?
Write down your answer.

13. Measurements The Rules:
Use all digits known with certainty – the 2 and 6
Use one additional digit that is estimated. (3, 4, 5, 6)

14. Uncertainty
The use of an estimated digit introduces some amount of uncertainty into the process.
In any measurement the last digit is always assumed to have some uncertainty about it.

15. Measurements Meniscus – curved surface made in cylindrical containers
Read the volume at the bottom of the meniscus
Parallax error – error created by failing to read an instrument at eye level.
Think the speedometer from the passenger seat

16. Measurement Activity

17. Results Balance #1 Balance #2 Ruler #1 Ruler #2
Buret #1 Buret #2 GC #1 GC #2

18. Back to Uncertainty
The last digit in a measurement always has some uncertainty
By convention, not all digits represent the uncertainty in a number
Significant figures – digits used to represent the amount of certainty in a measured or calculated amount.
Sig figs are based on the concept that placeholders only give information about the size of the measurement, not its certainty

19. Significant Figures The Rules: All non-zero digits are significant.
Zeros that are sandwiched between significant figures are significant.
Otherwise, zeros are significant only if they follow a decimal point AND a non-zero digit (the non-zero digit may be before OR after the decimal point).

20. Sig Fig Practice How many sig figs do the following numbers have:

21. Sig Figs in Calculations
The least certain number in a calculation limits the amount of certainty in the answer.
Round your answers to the lowest number of sig figs in the problem.

22. Sig Figs Practice What is the volume of a cube with 1.25cm sides?
What is the volume of a rectangular prism that has side lengths of 11.0cm, 5.5cm, and 112cm? 22

24. Unit Conversions We will use the factor label method
Multiply a value by conversion factors to convert to different units
Must cancel the units you don’t want
Must leave the units you do want
The numerator and denominator must be equivalent to each other

25. Practice Problems
Convert 240s to min
Convert 590 inches to yards

26. Practice Problems Convert 450mg to g
Always put the scientific notation you memorized with the base unit.
Base unit – unit without a prefix

27. Practice Problems 85 cm to km Where is the base unit?
Convert to the base unit first, then convert to the desired unit.

28. Practice Problems
kPa to μPa
What is the base unit here?

29. Compound Units Convert 55m/s to km/hr
Put the seconds on the bottom and then convert both units separately.

30. Compound Units Convert 0.789g/mL to kg/cm3 Write this down: 1mL = 1cm3
YOU NEED TO MEMORIZE THIS

31. Compound Units
Convert 120 ft/min to m/s. (2.54cm = 1 inch) 31

33. Density
Density – the ratio of the mass of a substance to the volume of a substance.
Typical units are g/mL

35. Expectations for Algebra Work
When working algebra problems, you must:
Solve the equation you are using symbolically (and show this step) before you substitute in numbers.
Example: When solving for mass in a density problem you must solve it first and show that m = DV
Substitute the numbers in with units.
Show the units cancelling and the resulting units in the answer.

36. Density What volume of ethanol would have a mass of 16.9g?
What is the mass of 4.91mL of lead?

37. Error Analysis Ways to describe data:
Accuracy – how close a value is to the correct value
Precision – how close measured values are to each other

38. Challenge
You are measuring an object that is exactly 2.61cm long. Come up with a set of 3 measurements that is:
Accurate and precise
Precise but not accurate
Accurate but not precise
Neither accurate nor precise

40. Chemical and Physical Changes
Chemical change – a change where the chemical makeup of a substance is changed, a change where a substance’s identity is changed.
Physical change – some change where no substance’s chemical structure or identity is changed. The appearance may change but the chemical composition does not

41. Indicators of a Chemical Change
Color change
Formation of a precipitate
Precipitate – a solid produced by a chemical reaction in a solution.
Formation of a liquid or gas

42. Indicators of a Chemical Change
Absorption or release of heat
Exothermic – the reaction releases heat
You feel it getting hotter
Endothermic – the reaction absorbs heat
You feel it getting colder
Light production

43. Physical Changes Dissolving and dilution.
Phase changes are physical changes.