1. Chemistry II – AP Introduction and Measurement

2. We study matter, its properties, and its behavior.
Chemistry –
We study matter, its properties, and its behavior.

3. Matter
We define matter as anything that has mass and takes up space.

4. Atoms and molecules determine how matter behaves; if they were different, matter would be different.
The properties of water molecules determine how water behaves; the properties of sugar molecules determine how sugar behaves.
The understanding of matter at the molecular level gives us unprecedented control over that matter. 4

5. Atoms and Molecules
If we want to understand the substances around us, we must understand the atoms and molecules that compose them—this is the central goal of chemistry.
Chemistry is the science that seeks to understand the behavior of matter by studying the behavior of atoms and molecules. 5

6. Chapter 1, Unnumbered Figure, Page 6

7. Solid Matter
In solid matter, atoms or molecules pack close to each other in fixed locations.
Although the atoms and molecules in a solid vibrate, they do not move around or past each other.
Consequently, a solid has a fixed volume and rigid shape.
Ice, aluminum, and diamond are good examples of solids. 7

8. Solid Matter
Solid matter may be crystalline—in which case its atoms or molecules are in patterns with long-range, repeating order.
Table salt and diamond are examples of solid matter.
Others may be amorphous, in which case its atoms or molecules do not have any long-range order.
Examples of amorphous solids include glass and plastic. 8

9. Liquid Matter
In liquid matter, atoms or molecules pack about as closely as they do in solid matter, but they are free to move relative to each other.
Liquids have fixed volume but not a fixed shape.
Liquids’ ability to flow makes them assume the shape of their container. 9

10. Gaseous Matter
In gaseous matter, atoms or molecules have a lot of space between them.
They are free to move relative to one another.
These qualities make gases compressible. 10

11. Chapter 1, Unnumbered Figure, Page 711

12. Properties and Changes of Matter

13. Types of Properties Physical Properties…
Can be observed without changing a substance into another substance.
Boiling point, density, mass, volume, etc.
Chemical Properties…
Can only be observed when a substance is changed into another substance.
Flammability, corrosiveness, reactivity with acid, etc.

14. Chapter 1, Figure 1.7 Boiling, a Physical Change14

15. Chapter 1, Figure 1.8 Rusting, a Chemical Change15

16. Types of Changes Physical Changes
These are changes in matter that do not change the composition of a substance.
Changes of state, temperature, volume, etc.
Chemical Changes
Chemical changes result in new substances.
Combustion, oxidation, decomposition, etc.

17. Chapter 1, Figure 1.9B Physical and Chemical Changes17

18. Chapter 1, Figure 1.9C Physical and Chemical Changes18

19. Types of Properties Intensive Properties…
Are independent of the amount of the substance that is present.
Density, boiling point, color, etc.
Extensive Properties…
Depend upon the amount of the substance present.
Mass, volume, energy, etc.

20. Units of Measurement

21. SI Units Système International d’Unités
A different base unit is used for each quantity.

22. Metric System
Prefixes convert the base units into units that are appropriate for the item being measured.

23. Volume not a fundamental SI unit, but a unit derived from length
The most commonly used metric units for volume are the liter (L) and the milliliter (mL).
A liter is a cube 1 decimeter (dm) long on each side.
A milliliter is a cube 1 centimeter (cm) long on each side.

24. Temperature
By definition temperature is a measure of the average kinetic energy of the particles in a sample.

25. Temperature The kelvin is the SI unit of temperature.
It is based on the properties of gases.
There are no negative Kelvin temperatures.

26. Temperature Scales
F = 9/5(C) + 32
C = 5/9(F − 32)
K = C

27. Density m d (r) = V Density is a physical property of a substance.
It has units (g/mL, for example) that are derived from the units for mass and volume.
SI unit for density is kg/m3
d (r) = m V

28. Uncertainty in Measurement

29. Types of Errors Random Error (indeterminate error)
means that the measurement has an equal
probability of being too high or too low; can
be “averaged out”
Systematic Error (determinate error)
occurs in the same direction every time
(always too high or too low)

30. High precision only indicates accuracy if systematic errors are absent
Averaging Data
If measurements are precise, we “average out” random errors; however, that assumes there is no SYTEMATIC error
High precision only indicates accuracy if systematic errors are absent

31. Precision and Accuracy
Accuracy refers to how close the measured value is to the actual / true value.
Precision refers to how close a series of measurements are to one another or how reproducible they are. 31

33. Precision and Accuracy
Consider the results of three students who repeatedly weighed a lead block known to have a true mass of g (indicated by the solid horizontal blue line on the graphs).
Student A
Student B
Student C
Trial 1
10.49 g
9.78 g
10.03 g
Trial 2
9.79 g
9.82 g
9.99 g
Trial 3
9.92 g
9.75 g
Trial 4
10.31 g
9.80 g
9.98 g
Average
10.13 g
10.01 g 33

34. Precision and Accuracy
Measurements are said to be
precise if they are consistent with one another.
accurate only if they are close to the actual value. 34

35. In quantitative work, high precision is often used an indication of accuracy (we assume the average of precise data is accurate)
Assumption is only true if systematic errors are absent

36. Uncertainty in Measurements
Different measuring devices have different uses and different degrees of accuracy.

37. Accuracy (mostly controlled by instrument)
PERCENT ERROR

38. 2. Precision
LEVELS of CONFIDENCE

39. Always estimate ONE place past the smallest mark!

40. 2. Precision
SIGNIFICANT DIGITS… refer to book pg. 13

41. 2. Precision
RELATIVE PRECISION

42. More About Units

43. The Kilogram: A Measure of Mass
The mass of an object is a measure of the quantity of matter within it.
The SI unit of mass = kilogram (kg).
1 kg = 2 lb 3 oz
A second common unit of mass is the gram (g).
One gram is 1/1000 kg.
Weight of an object is a measure of the gravitational pull on its matter. 43

44. The Second: A Measure of Time
Measure of the duration of an event
SI units = second (s)
1 s is defined as the period of time it takes for a specific number of radiation events of a specific transition from cesium-133. 44

45. The Kelvin: A Measure of Temperature
The Kelvin (K) is the SI unit of temperature.
The temperature is a measure of the average amount of kinetic energy of the atoms or molecules that compose the matter.
Temperature also determines the direction of thermal energy transfer, or what we commonly call heat.
Thermal energy transfers from hot to cold objects. 45

46. The Kelvin: A Measure of Temperature
Kelvin scale (absolute scale) assigns 0 K (absolute zero) to the coldest temperature possible.
Absolute zero (–273 °C or –459 °F) is the temperature at which molecular motion virtually stops. Lower temperatures do not exist. 46

47. A Measure of Temperature
The Fahrenheit degree is five-ninths the size of a Celsius degree.
The Celsius degree and the Kelvin degree are the same size.
Temperature scale conversion is done with these formulas: 47

48. Energy: A Fundamental Part of Physical and Chemical Change
Energy is the capacity to do work.
Work is defined as the action of a force through a distance.
When you push a box across the floor or pedal your bicycle across the street, you have done work. 48

49.  Energy
Kinetic energy is the energy associated with the motion of an object.
Potential energy is the energy associated with the position or composition of an object.
Thermal energy is the energy associated with the temperature of an object.
Thermal energy is actually a type of kinetic energy because it arises from the motion of the individual atoms or molecules that make up an object. 49