1. Matter and Measurement
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2. The Study of Chemistry What is Chemistry?
Chemistry is the study of the properties and behavior of matter.
Matter – anything that occupies space and has mass.
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3. Classification of Matter
The basic difference between these states is the distance between the “bodies.”
Gas – bodies are far apart and in rapid motion.
Liquid – bodies closer together, but still able to move past each other.
Solid – bodies are closer still and are now held in place in a definite arrangement.
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4. Classification of Matter
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5. Classification of Matter
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6. Classification of Matter
Pure Substances and Mixtures
Mixture – combination of two or more substances in which each substance retains its own chemical identity.
Homogeneous mixture – composition of this mixture is consistent throughout.
Solution (Air, gasoline)
Heterogeneous mixture – composition of this mixture varies throughout the mixture.
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7. Classification of Matter
Separation of Mixtures
Mixtures can be separated by physical means.
Filtration.
Chromatography.
Distillation.
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8. Classification of Matter
Separation of Mixtures
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9. Classification of Matter
Pure Substances and Mixtures
It is also possible for a homogeneous substance to be composed of a single substance – pure substance.
Element – A substance that can not be separated into simpler substances by chemical means.
Atom – the smallest unit of an element that retains a substances chemical activity.
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10. Classification of Matter
Elements
There are 114 elements known.
Each element is given a unique chemical symbol (one or two letters).
Carbon C, Nitrogen N, Titanium Ti
Notice that the two letter symbols are always capital letter then lower case letter because:
CO – carbon and oxygen
Co – element cobalt
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11. Classification of Matter
Pure Substances and Mixtures
It is also possible for a homogeneous substance to be composed of a single substance – pure substance.
Compound – A substance composed of two or more elements united chemically in definite proportions.
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12. Classification of Matter
Compounds
Formed by combining elements.
The proportions of elements in compounds are the same irrespective of how the compound was formed.
Law of Constant Composition (or Law of Definite Proportions):
The composition of a pure compound is always the same, regardless of its source.
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13. Properties of Matter Physical and Chemical Changes
Physical Property – A property that can be measured without changing the identity of the substance.
Example: color, odor, density
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14. Properties of Matter Physical and Chemical Changes
Intensive properties – independent of sample size.
Extensive properties - depends on the quantity of the sample (sample size).
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15. Units of Measurement Density
Density – mass per unit volume of an object.
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16. Properties of Matter Physical and Chemical Changes
Physical change – the change in the physical properties of a substance.
Physical appearance changes, but the substances identity does not.
Water (ice)  Water (liquid)
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17. Properties of Matter Physical and Chemical Changes
Chemical change (chemical reaction) – the transformation of a substance into a chemically different substance.
When pure hydrogen and pure oxygen react completely, they form pure water.
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18. Properties of Matter
Physical and Chemical Changes
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19. Units of Measurement SI Units There are two types of units:
fundamental (or base) units;
derived units.
There are 7 base units in the SI system.
Derived units are obtained from the 7 base SI units.
Example:
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20. Units of Measurement
SI Units
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21. Units of Measurement
SI Units
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22. Units of Measurement
Temperature
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23. Units of Measurement Temperature Kelvin Scale Celsius Scale
Used in science.
Same temperature increment as Celsius scale.
Lowest temperature possible (absolute zero) is zero Kelvin.
Absolute zero: 0 K = oC.
Celsius Scale
Also used in science.
Water freezes at 0oC and boils at 100oC.
To convert: K = oC
Fahrenheit Scale
Not generally used in science.
Water freezes at 32oF and boils at 212oF.
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24. Units of Measurement Temperature
Converting between Celsius and Fahrenheit
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25. Units of Measurement Volume
The units for volume are given by (units of length)3.
i.e., SI unit for volume is 1 m3.
A more common volume unit is the liter (L)
1 L = 1 dm3 = 1000 cm3 = 1000 mL.
We usually use 1 mL = 1 cm3.
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26. Units of Measurement Mass
Mass is the measure of the amount of material in an object.
This is not the same as weight which is dependant on gravity.
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27. Uncertainty in Measurement
All scientific measures are subject to error.
These errors are reflected in the number of figures reported for the measurement.
These errors are also reflected in the observation that two successive measures of the same quantity are different.
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28. Uncertainty in Measurement
Precision and Accuracy
Measurements that are close to the “correct” value are accurate.
Measurements which are close to each other are precise.
Measurements can be
accurate and precise;
precise but inaccurate;
neither accurate nor precise.
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29. Uncertainty in Measurement
Precision and Accuracy
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30. Uncertainty in Measurement
Significant Figures
The number of digits reported in a measurement reflect the accuracy of the measurement and the precision of the measuring device.
The last digit to the right in a number is taken to be inexact.
In any calculation, the results are reported to the fewest significant figures (for multiplication and division) or fewest decimal places (addition and subtraction).
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31. Uncertainty in Measurement
Significant Figures
Non-zero numbers are always significant.
Zeros between non-zero numbers are always significant.
Zeros before the first non-zero digit are not significant. Zeros at the end of the number after a decimal place are significant.
Zeros at the end of a number before a decimal place are ambiguous. For this book, it a decimal point is used the zeros are significant.
10,300 has 3 significant figures.
10,300. has 5 significant figures.
Physical constants are “infinitely” significant.
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32. Uncertainty in Measurement
Significant Figures
Multiplication / Division
The result must have the same number of significant figures as the least accurately determined data
Example:
(5 sig. fig.)
5.1 (2 sig. fig.)
x 5.1 = 64
Answer has only 2 significant figures
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33. Uncertainty in Measurement
Significant Figures
Addition / Subtraction.
The result must have the same number of digits to the right of the decimal point as the least accurately determined data.
Example:
(5 sig. fig., 3 digits to the right),
1.76 (3 sig. fig., 2 digits to the right),
7.1 (2 sig. fig., 1 digit to the right).
= 24.0.
24.0 (3 sig. fig., but only 1 digit to the right of the decimal point).
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34. Uncertainty in Measurement
Rounding rules
If the leftmost digit to be removed is less than 5, the preceding number is left unchanged. “Round down.”
If the leftmost digit to be removed is 5 or greater, the preceding number is increased by 1. “Round up.”
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35. Dimensional Analysis
In dimensional analysis always ask three questions:
What data are we given?
What quantity do we need?
What conversion factors are available to take us from what we are given to what we need?
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36. Dimensional Analysis Method of calculation using a conversion factor.
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37. Dimensional Analysis
Example: We want to convert the distance 8 in. to feet.
(12in = 1 ft)
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38. Dimensional Analysis
Example: We want to convert the distance 8 in. to feet.
(12in = 1 ft)
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39. Centimeter (cm)  Meter (m) Meter (m)  Nanometer (nm)
Dimensional Analysis
Convert the quantity from 2.3 x 10-8 cm to nanometers (nm)
First we will need to determine the conversion factors
Centimeter (cm)  Meter (m)
Meter (m)  Nanometer (nm) Or
1 cm = 0.01 m
1 x 10-9 m = 1 nm
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40. Dimensional Analysis
Convert the quantity from 2.3 x 10-8 cm to nanometers (nm)
1 cm = 0.01 m
1 x 10-9 m = 1 nm
Now, we need to setup the equation where the cm cancels and nm is left.
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41. Dimensional Analysis
Convert the quantity from 2.3 x 10-8 cm to nanometers (nm)
1 cm = 0.01 m
1 x 10-9 m = 1 nm
Now, fill-in the value that corresponds with the unit and solve the equation.
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42. Dimensional Analysis
Convert the quantity from 2.3 x 10-8 cm to nanometers (nm)
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43. Meter (m)  kilometer (km)
Dimensional Analysis
Convert the quantity from 31,820 mi2 cm to square meters (m2)
First we will need to determine the conversion factors
Mile (mi)  Meter (m)
Meter (m)  kilometer (km) Or
1 mile = km
1000m = 1 km
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44. Dimensional Analysis
Convert the quantity from 31,820 mi2 cm to square meters (m2)
Now, we need to setup the equation where the cm cancels and nm is left.
1 mile = km 1000m = 1 km
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45. Dimensional Analysis
Convert the quantity from 31,820 mi2 cm to square meters (m2)
Now, we need to setup the equation where the cm cancels and nm is left.
1 mile = km 1000m = 1 km
Notice, that the units do not cancel, each conversion factor must be “squared”.
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46. Dimensional Analysis
Convert the quantity from 31,820 mi2 cm to square meters (m2)
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47. Dimensional Analysis
Convert the quantity from 31,820 mi2 cm to square meters (m2)
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48. Dimensional Analysis
Convert the quantity from 31,820 mi2 cm to square meters (m2)
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49. Dimensional Analysis
Convert the quantity from 14 m/s cm to miles per hour (mi/hr).
Determine the conversion factors
Meter (m)  Kilometer (km) Kilometer(km)  Mile(mi)
Seconds (s)  Minutes (min) Minutes(min)  Hours (hr) Or
1 mile = km 1000m = 1 km
60 sec = 1 min 60 min = 1 hr
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50. Dimensional Analysis
Convert the quantity from 14 m/s cm to miles per hour (mi/hr).
1 mile = km 1000m = 1 km
60 sec = 1 min 60 min = 1 hr
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51. Dimensional Analysis
Convert the quantity from 14 m/s cm to miles per hour (mi/hr).
1 mile = km 1000m = 1 km
60 sec = 1 min 60 min = 1 hr
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52. Dimensional Analysis
Convert the quantity from 14 m/s cm to miles per hour (mi/hr).
1 mile = km 1000m = 1 km
60 sec = 1 min 60 min = 1 hr
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53. End of Chapter Problems
4, 10, 14, 20, 26, 34, 42, 60
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