1. Chapter One Chemistry: The Science of Change

2. Section 1.1 The Study of Chemistry

3. The Scientific Method
An organized process used by scientists to do research and to verify the work of others
Designed to produce a solution that can be tested and supported by experimentation

4. The Scientific Method Observations:
Qualitative: Descriptive
Ex: color, texture, smell
Quantitative: Numerical
Ex: 2.5 grams, 55.8 mL
Ex: My phone is running slower than my friend’s phone
Ex: My phone has 50 more apps than my friend’s phone

5. Hypothesis A tentative explanation of observations
Ex: As more apps are added onto a phone, the phone will operate at a slower rate.

6. Independent vs. Dependent
Independent variable: what YOU control
Dependent variable: what is MEASURED as a result of changing the independent.
What about our phone example?

7. Experimentation
A set of controlled observations that test a hypothesis
Contains independent & dependent variable
At least 3 trials
Control: In an experiment, this is the standard for comparison.
Constant: factors that must remain unchanging in the experiment to ensure results are meaningful

8. End of Experiment Revision of hypothesis Evolution of a theory
Explanation of a body of experiments and observations
Used to predict
Scientific Law
Summary of accepted facts of nature
Concise verbal or mathematical statement

9. Section 1.2 Classification of Matter

10. States of Matter
Solid: definite shape & volume, low energy, geometric structure
Liquid: no definite shape, definite volume, moderate energy, molecules move past each other
Gas: no definite shape or volume, high energy, molecules move fast and randomly

11. Classification of Matter
Matter: anything with mass and volume
Either a pure substance or mixture of substances

12. Practice
Determine if the following is an element, compound, heterogeneous mixture, or homogeneous mixture Distilled water Aluminum cans Brass 2012 penny Silver Vanilla milkshake Carbon dioxide Salt water Mercury Orange Juice w/ pulp Table salt

13. Pictures of Mixtures & Pure Substances

14. Section 1.3 The Properties of Matter

15. Physical & Chemical Properties
Physical Property: Something that can be observed/measured without changing composition of a substance
i.e. color, malleable, ductile, density, melting/boiling point
Chemical Property: The ability of a substance to combine with another substance
i.e. flammability, reactivity.

16. Physical vs. Chemical

17. Physical & Chemical Changes
Physical Change: A change that happens without the chemical identity being altered
i.e. cutting paper, phase changes
Chemical Change: A change that alters the chemical identity of a substance
i.e. rusting, corrosion, combustion, baking

18. Practice
Identify if the following are physical properties (PP), chemical properties (CP), physical changes (PC), or chemical changes (CC) 1) Luster/Shiny 5) Melting point of 15°C 2) Flammable 6) Baking cookies 3) Melting ice 7) Reactivity 4) Plants growing 8) Salt forming on rocks from the ocean

19. Extensive & Intensive Properties
Extensive Property: a property that depends on amount of matter
i.e. mass, volume
Intensive Property: a property that DOES NOT depend on amount of matter
i.e. density, temperature

20. Section 1.4 Scientific Measurement

21. Scientific Notation M x 10n M = Base (a number between 1 and 10)
n = exponent (positive=bigger than 1;
negative=less than 1)
Use scientific notation if number is in the thousands and larger or in thousandths and smaller
Ex: ,403

22. Systems of Measurement
English System: foot, pound, gallon, etc.
Metric System: meter, kilogram, liter, etc.
SI System: Revised metric system (System Internationale)

23. Metric System
Mean King Henry Died Unexpectedly Drinking Chocolate Milk Monday
Mega Kilo Hecta Deka Unit
Deci Centi Milli Micro

24. M _ _ K H D U D C M _ _ M M _ _ k h da U d c m _ _ µ Metric System

25. Metric System

26. Metric System

27. Practice 489.3 dg  kg 0.0293 L  μL 2.04 x 103 pm  cm
How many nanograms are in one gram?
How many grams are in one nanogram?

28. Derived Units A unit that is a combination of SI base units
Ex: Volume & Density

29. What’s the Area & Volume?
How many dm3 are in 1 m3?
How many mm2 are in 1 m2?

30. Some Common Conversions

31. Temperature Scales
Celsius Scale: defined using the freezing point (O°C) and boiling point (100°C) of pure water at sea level
Kelvin Scale: defined off of absolute zero (at 0 K, all molecular motion stops)
Fahrenheit Scale: many varying stories of how this scale was derived

32. Temperature Conversions
Celsius to Kelvin:
K = °C
Kelvin to Celsius:
°C = K –
Celsius to Fahrenheit:
°F = (1.8)(°C) + 32

33. Practice
93.0 °C  K
214 K  °C

34. Density
D = m/V
D = density, m = mass, V = volume
In Chemistry, density measured in g/mL (for liquids) or g/cm3 (for solids)
Since gases are very light, we measure their density in g/L

35. Practice
A sample of pure aluminum has a density of 2.70 g/mL. If a cube of aluminum has a length of 2.0 cm on each side, how much mass will it have?

36. Section 1.5 Uncertainty in Measurement

37. Uncertainty in Measurement

38. Rules for Significant Figures
All non zero numbers are significant
1.932
All zeroes in between non zero numbers are ALWAYS significant
“Beginning zeros ” (All zeroes to the LEFT of the first non zero number) are NEVER significant
“Ending zeros” (All zeroes to the RIGHT of the last non zero number) are significant ONLY IF there is a decimal in the number

39. How many sig figs in the following numbers?
1 003
209.50
1.0040
500
303

40. Sig Figs and Exact Numbers
Exact numbers come from counting, not measuring
Example: 12 donuts = 1 dozen, 1000 mm = 1 m
Sig figs are never based on exact numbers

41. Calculating with Sig Figs
Adding/Subtracting
Express the answer with the LOWEST # OF DECIMAL PLACES
Multiplying/Dividing
Express the answer with the LOWEST # OF SIG FIGS

42. Practice 89.332 m + 1.1 m = 23.4 cm x 16.00 cm = 54.78 m2 ÷ 2.3 m =
Perform the following calculations and express the answer in the correct # of sig figs.
m m =
23.4 cm x cm =
54.78 m2 ÷ 2.3 m =
128.2 mL – mL =
90.4 m
374 cm2
24 m
84.2 mL

43. Multiple Operations
When multiple operations occur, be sure to complete one set of sig fig rules before moving on to the other set.
Avoid Rounding Errors!
Example: Calculating percent error
(Accepted – Experimental)
Accepted
(8.641 g/mL – 8.43 g/mL)
8.641 g/mL

44. Accuracy and Precision
Accuracy: How close you are to the TRUE value
Precision: How close a group of measurements are to each other

45. Section 1.6 Using Units and Solving Problems

46. Conversion Factors
A fraction set up with measurements in the numerator and denominator that EQUAL each other
Ex: 1 foot inches
12 inches 1 foot

47. Dimensional Analysis
Using conversion factors to convert from one unit to another

48. Practice Perform the following conversions:
How many inches are in 2.3 meters?
Convert 54.5 mi/hr  m/s (1 mile = km)
How many 375 mg tables can be produced from 2.60 kg of powdered aspirin?