1. Chapter 1 Introduction: Matter and Measurement

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Chemistry
In this science we study matter, its properties, and its behavior.
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Matter
We define matter as anything that has mass and takes up space.
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Matter
Atoms are the building blocks of matter.
Each element is made of the same kind of atom.
A compound is made of two or more different kinds of elements.
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States of Matter
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6. Classification of Matter
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7. Classification of Matter
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8. Classification of Matter
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9. Classification of Matter
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10. Classification of Matter
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11. Classification of Matter
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12. Classification of Matter
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13. Classification of Matter
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14. Classification of Matter
And
Measurement
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15. Classification of Matter
And
Measurement
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16. Properties and Changes of Matter
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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.
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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.
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Chemical Reactions
In the course of a chemical reaction, the reacting substances are converted to new substances.
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Units of Measurement
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SI Units
Système International d’Unités
A different base unit is used for each quantity.
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Metric System
Prefixes convert the base units into units that are appropriate for the item being measured.
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Temperature
By definition temperature is a measure of the average kinetic energy of the particles in a sample.
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Temperature
The kelvin is the SI unit of temperature.
It is based on the properties of gases.
There are no negative Kelvin temperatures.
K = C
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Temperature
The Fahrenheit scale is not used in scientific measurements.
F = 9/5(C) + 32
C = 5/9(F − 32)
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Derived Units
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.
d = m V
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27. Uncertainty in Measurement
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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.
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Significant Figures
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.
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Significant Figures
When addition or subtraction is performed, answers are rounded to the least significant decimal place.
When multiplication or division is performed, answers are rounded to the number of digits that corresponds to the least number of significant figures in any of the numbers used in the calculation.
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31. 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.
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Dimensional Analysis
We use dimensional analysis to convert one quantity to another.
Most commonly, dimensional analysis utilizes conversion factors (e.g., 1 in. = 2.54 cm)
1 in.
2.54 cm or
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Dimensional Analysis
Use the form of the conversion factor that puts the sought-for unit in the numerator:
Given unit 
 desired unit
desired unit
given unit
Conversion factor
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Dimensional Analysis
For example, to convert 8.00 m to inches,
convert m to cm
convert cm to in.
8.00 m 
100 cm
1 m
1 in.
2.54 cm 
315 in. 
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35. Sample Exercise 1.8 Determining the Number of Significant Figures in a Calculated Quantity
A gas at 25 °C fills a container whose volume is 1.05 × 103 cm3. The container plus gas has a mass of g. The container, when emptied of all gas, has a mass of g. What is the density of the gas at 25 °C?
To calculate the density, we must know both the mass and the volume of the gas. The mass of the gas is just the difference in the masses of the full and empty container:
(837.6 – 836.2) g = 1.4 g
In subtracting numbers, we determine the number of significant figures in our result by counting decimal places in each quantity. In this case each quantity has one decimal place. Thus, the mass of the gas, 1.4 g, has one decimal place.
Using the volume given in the question, 1.05 × 103 cm3, and the definition of density, we have
In dividing numbers, we determine the number of significant figures in our result by counting the number of significant figures in each quantity. There are two significant figures in our answer, corresponding to the smaller number of significant figures in the two numbers that form the ratio. Notice that in this example, following the rules for determining significant figures gives an answer containing only two significant figures, even though each of the measured quantities contained at least three significant figures.
Solution

36. Sample Exercise 1.8 Determining the Number of Significant Figures in a Calculated Quantity
Continued
Practice Exercise
To how many significant figures should the mass of the container be measured (with and without the gas) in Sample Exercise 1.8 for the density to be calculated to three significant figures?
Answer: five (For the difference in the two masses to have three significant figures, there must be two decimal places in the masses of the filled and empty containers. Therefore, each mass must be measured to five significant figures.)

37. Solution Practice Exercise Sample Exercise 1.9 Converting Units
If a woman has a mass of 115 lb, what is her mass in grams? (Use the relationships between units given on the back inside cover of the text.)
Solution
Because we want to change from pounds to grams, we look for a relationship between these units of mass. From the back inside cover we have 1 lb = g. To cancel pounds and leave grams, we write the conversion factor with grams in the numerator and pounds in the denominator:
The answer can be given to only three significant figures, the number of significant figures in 115 lb. The process we have used is diagrammed in the margin.
Practice Exercise
By using a conversion factor from the back inside cover, determine the length in kilometers of a mi automobile race.
Answer: km

38. Sample Exercise 1.10 Converting Units Using Two or More Conversion Factors
The average speed of a nitrogen molecule in air at 25 °C is 515 m/s. Convert this speed to miles per hour.
To go from the given units, m/s, to the desired units, mi/hr, we must convert meters to miles and seconds to hours. From our knowledge of SI prefixes we know that 1 km = 103 m. From the relationships given on the back inside cover of the book, we find that 1 mi = km. Thus, we can convert m to km and then convert km to mi. From our knowledge of time we know that 60s = 1 min and 60 min = 1 hr. Thus, we can convert s to min and then convert min to hr. The overall process is
Applying first the conversions for distance and then those for time, we can set up one long equation in which unwanted units are canceled:
Solution