1. Chapter 1 Introduction: Matter and Measurement
Lecture Presentation
Chapter 1 Introduction: Matter and Measurement
James F. Kirby
Quinnipiac University
Hamden, CT

2. Matter and measurents Matter Numbers mixtures and pure substances
homogenous and heterogeneous mixtures
compounds and elements
chemical and physical changes
physical and chemical properties
Numbers
Accuracy and precision
Significant figures
Rounding off
Scientific notation
Metric units
Dimensional analysis

3. Chemistry
The study of the composition and structure of materials and of the changes that materials undergo.

4. Chemistry
Chemistry is the study of the properties and behavior of matter.
It is central to our fundamental understanding of many science-related fields.

5. Scientific Method
Results
Hypothesis
Experiments

6. Scientific Method

7. Experiment Hypothesis Theory Law
An observation of natural phenomena carried out in a controlled manner so that the results can be duplicated and rational conclusions obtained.
Hypothesis
A tentative explanation of some regularity of nature.
Theory
A tested explanation of a basic natural phenomenon.
Law
A concise statement or mathematical equation about a fundamental relationship or regularity of nature.

8. Basic laws about matter
Law of conservation of mass
In an ordinary chemical reaction, mass is neither lost nor created
Law of constant proportions
A chemical compound always contains exactly the same  proportion of elements by mass

9. Matter
Matter is anything that has mass and takes up space.

10. Matter Atoms are the building blocks of matter.
Each element is made of a unique kind of atom.
A compound is made of two or more different kinds of elements.
Note: Balls of different colors are used to represent atoms of different elements. Attached balls represent connections between atoms that are seen in nature. These groups of atoms are called molecules.

11. Methods of Classification
State of Matter
Composition of Matter

12. States of Matter

13. States of Matter The three states of matter are
solid.
liquid.
gas.
In this figure, those states are ice, liquid water, and water vapor.

14. Classification of Matter Based on Composition
If you follow this scheme, you can determine how to classify any type of matter.
Homogeneous mixture
Heterogeneous mixture
Element
Compound
Matter
And
Measurement

15. Classification of Matter—Substances
A substance has distinct properties and a composition that does not vary from sample to sample.
The two types of substances are elements and compounds.
An element is a substance which can not be decomposed to simpler substances.
A compound is a substance which can be decomposed to simpler substances.

16. Compounds and Composition
Compounds have a definite composition. That means that the relative number of atoms of each element that makes up the compound is the same in any sample.
This is The Law of Constant Composition (or The Law of Definite Proportions).

17. Classification of Matter—Mixtures
Mixtures exhibit the properties of the substances that make them up.
Mixtures can vary in composition throughout a sample (heterogeneous) or can have the same composition throughout the sample (homogeneous).
Another name for a homogeneous mixture is solution.

18. *classify the following as elements, compounds, heterogenous mixtures or homogenous mixtures
baking soda sand sugar
chlorine iodized salt bronze
copper brass air
blood milk salad
steel wool honey spring water

19. Types of Properties
Intensive Properties are independent of the amount of the substance that is present.
Examples include density, boiling point, or color.
Extensive Properties depend upon the amount of the substance present.
Examples include mass, volume, or energy.

20. Indicate which of the following are intensive or extensive properties
Density
Volume
Mass
Specific volume
Temperature
Moles
Viscosity
Thermal conductivity
Color
Electrical resistance
Odor
Luster
Charge
Pressure
Malleability
Ductility

21. Types of Properties
Physical Properties can be observed without changing a substance into another substance.
Some examples include boiling point, density, mass, or volume.
Chemical Properties can only be observed when a substance is changed into another substance.
Some examples include flammability, corrosiveness, or reactivity with acid.

22. which of the following are chemical or physical properties
Melting point
Flammability
Viscosity
Boiling
Souring of milk
Rusting of iron
Food digestion
Taking a bite of food
Deployment of a car airbag
Breaking of glass

23. Types of Changes
Physical Changes are changes in matter that do not change the composition of a substance.
Examples include changes of state, temperature, and volume.
Chemical Changes result in new substances.
Examples include combustion, oxidation, and decomposition.

24. Changes in State of Matter
Converting between the three states of matter is a physical change.
When ice melts or water evaporates, there are still 2 H atoms and 1 O atom in each molecule.

25. Chemical Reactions (Chemical Change)
In the course of a chemical reaction, the reacting substances are converted to new substances. Here, the elements hydrogen and oxygen become water.

26. Separating Mixtures
Mixtures can be separated based on physical properties of the components of the mixture. Some methods used are
filtration.
distillation.
chromatography.

27. Filtration
In filtration, solid substances are separated from liquids and solutions.

28. Distillation
Distillation uses differences in the boiling points of substances to separate a homogeneous mixture into its components.

29. Chromatography
This technique separates substances on the basis of differences in the ability of substances to adhere to the solid surface, in this case, dyes to paper.

30. Numbers and Chemistry
Numbers play a major role in chemistry. Many topics are quantitative (have a numerical value).
Concepts of numbers in science
Units of measurement
Quantities that are measured and calculated
Uncertainty in measurement
Significant figures
Dimensional analysis

31. Units of Measurements—SI Units
Système International d’Unités (“The International System of Units”)
A different base unit is used for each quantity.

32. Units of Measurement—Metric System
The base units used in the metric system
Mass: gram (g)
Length: meter (m)
Time: second (s or sec)
Temperature: degrees Celsius (oC) or Kelvins (K)
Amount of a substance: mole (mol)
Volume: cubic centimeter (cc or cm3) or liter (l)

33. Units of Measurement— Metric System Prefixes
Prefixes convert the base units into units that are appropriate for common usage or appropriate measure.

34. Mass and Length These are basic units we measure in science.
Mass is a measure of the amount of material in an object. SI uses the kilogram as the base unit. The metric system uses the gram as the base unit.
Length is a measure of distance. The meter is the base unit.

35. Volume
Note that volume is not a base unit for SI; it is derived from length (m × m × m = m3).
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, also called 1 cubic centimeter (cm × cm × cm = cm3).

36. Temperature
In general usage, temperature is considered the “hotness and coldness” of an object that determines the direction of heat flow.
Heat flows spontaneously from an object with a higher temperature to an object with a lower temperature.

37. Temperature
In scientific measurements, the Celsius and Kelvin scales are most often used.
The Celsius scale is based on the properties of water.
0 C is the freezing point of water.
100 C is the boiling point of water.
The kelvin is the SI unit of temperature.
It is based on the properties of gases.
There are no negative Kelvin temperatures.
The lowest possible temperature is called absolute zero (0 K).
K = C

38. Temperature
The Fahrenheit scale is not used in scientific measurements, but you hear about it in weather reports!
The equations below allow for conversion between the Fahrenheit and Celsius scales:
F = 9/5(C) + 32
C = 5/9(F − 32)

39. Density Density is a physical property of a substance.
It has units that are derived from the units for mass and volume.
The most common units are g/mL or g/cm3.
D = m/V

40. Numbers Encountered in Science
Exact numbers are counted or given by definition. For example, there are 12 eggs in 1 dozen.
Inexact (or measured) numbers depend on how they were determined. Scientific instruments have limitations. Some balances measure to ±0.01 g; others measure to ±0.0001g.

41. Uncertainty in Measurements
Different measuring devices have different uses and different degrees of accuracy.
All measured numbers have some degree of inaccuracy.

42. Accuracy versus Precision
Accuracy refers to the proximity of a measurement to the true value of a quantity.
Precision refers to the proximity of several measurements to each other.

43. Significant Figures
The term significant figures refers to digits that were measured.
When rounding calculated numbers, we pay attention to significant figures so we do not overstate the accuracy of our answers.

44. Recording the proper number of significant figures
**the digits in a measurement that are known with certainty plus one digit that is uncertain**

45. Graduated pipette
Graduated cylinder
Buret
Thermometer

46. Significant Figures in reported data
All nonzero digits are significant.
Zeroes between two significant figures are themselves significant.
Zeroes at the beginning of a number are never significant.
Zeroes at the end of a number are significant if a decimal point is written in the number or they are under a bar (to show significance).

47. Significant Figures 1. All non-zero digits count 1,234,456,789
2. All leading zeroes do not count
3. Confined zeroes do count
4. Trailing zeroes:
-With a decimal, all trailing zeroes count
-In the presence of bars, zeroes under the bar count ,020,100
-in the absence of bars and decimal, zeroes do not count 36,020,100

48. Rules for mathematical operations
Rounding off
1. If the 1st digit to be dropped is less than 5, then that digit and all
digits that follow are dropped
rounded to 3 sig figs becomes…..
2. If the first digit to be dropped is a 5 or greater, followed by at least
one non-zero digit, the excess digits are dropped and the last retained
digit increases by 1
rounded to 3 sig figs becomes…
3.If the 1st digit to be dropped is a 5 not followed by any other digits or
followed by zeros, drop the 5 and any zeroes AND
a. increase the last retained digit by 1 if it is odd
b. Do not change the last retained digit if it is even
rounded off to 3 sig figs becomes 62.6…(even rule)
rounded off to 3 sig figs becomes 62.4….(odd rule)

49. Rules for mathematical operations
Multiplication and division:
-The calculated value has the same number of significant figures as the
number with the fewest significant figures
Addition and subtraction:
-the uncertainty in the calculated value must be the same as the
uncertainty in the input with the greatest uncertainty (last significant place)
We are treating only the simplest of operations

50. Perform the following calculation and round your answer to the correct number of significant figures:
Calculator answer:
The answer should be rounded to two significant figures because – = 0.035:
3.2

51. Calculate the following:
1. (2.05 X103) + (3.11 X 103)
2. (8.66 X 105) + (1.20 X 105)
3. (2.0 X 104) + (0.02 X104)
4. (5.401 X 103) + (2.101 X 10-8)
5. (6.0 X 106) + (7.75 X 10-8)
6. (1.01 X 103) - (9.952 X 10-2)
7. (4.4 X 102) - (9.56 X102)
8. (1.53 X 107) - (1.12 X 106)
9. (5.99 X105) - (8.65 X 102)
10. (4.5 X107) – (5.567 X10-5)
11. (7.65 X 104)- (9.9 X 103)
12. ( X 105) – (2.01 X104)
13. (2.0 X 102) X(2 X 102)
14. ( X 105) X (5.2 X 10-5)
15. (9.91 X 103) X (1.123 X 10-3)
16. (7.666 X 107) X (5.89 X108)
17. (1.0 X 10-6) X (2.25 X 106)
18. (6.857 X 10-5) ÷ (1.20 X 104)
19. (1.250 X 105) X (8.000 X 102)
20. (2.5 X 10-3) X (4.000 X 105)
21. (5.00 X 102) ÷ (5.000 X10-3)
22. (8.0 X 103) ÷ (2.000 X 103)
23. (5.65 X 107) ÷ (5.62 X 105)
24. (7.000 X 102) ÷ (3.5 X 104)

52. Dimensional analysis
*this is a general problem solving approach in which the relationships between quantities (or factors) are used as a guide in setting up the calculation
*this applies to one step as well as complex problems
* in multiple step problems, an answer from a preceding step serves as a given for the next step
Desired quantity=
Desired quantity unit
Given quantity unit
X given quantity
Conversion factor

53. *convert 12 grams into kilograms
-given: 12 g; desired: kilograms
kilograms = (1 kg/1000g) X 12 g = kg
-always check sig figs
**convert 11 days into seconds
-given: 11 days, desired: kilograms
-multistep process
-strategy: days to hours to minutes to seconds
Seconds = 11 days X (24 hrs/1 day) X (60 mins/1 hr) X (60 secs/1 min)
= 950,400 s

54. *other dimensional analysis problems involve do not merely changing the units of the same quantity
Q:You're throwing a pizza party for 15 and figure each person might eat 4 slices. How much is the pizza going to cost you? You call up the pizza place and learn that each pizza will cost you $14.78 and will be cut into 12 slices
A: $ 73.90

55. 1. Express an acceleration of 9.81 m/s2 in ft/s2
2. How many eggs are in 10 and a half dozens of eggs?
3. Express the speed of 85 km/hr in inches per second.
4. Express m in mm, km and miles
5. The mass of the earth is estimated at 6.6 X1021 metric tons. Express this mass in grams.
6. Convert a volume of 56L to mL, µL, nL and gallons
7. Express 2.5 L in mm3, in3 and ft3
Conversion factors:
1 metric ton = 1000 kg
1 ft = 12 inches
1 inch = 2.54 cm
1 gallon = 3.78 L