Chemical Principles 160:115			Dr. Paul Fazen
fazen@crab.rutgers.edu		http://crab.rutgers.edu/~fazen

Textbook: Chemistry, the Central Science, 8th Ed. Brown, LeMay, and Bursten
This course will cover chapters 1-13.
Office Hours: MW: 11-12
Lecture: MWF: 1:20-2:15
Recitation: F 12:15-1:05

Grading:	3 exams			3 x 20% = 	60%
			1 final (cumulative)			40%
										100%

If miss one exam (No make-ups): 
2 exams	2 x 20% = 	40%
					1 final				60%
										100%

There is no extra credit.

Page xxv
Advice for learning and studying chemistry
•	Keep up with your studying day to day. Don’t cram
•	Focus your study.
•	Keep good lecture notes.
•	Skim topics in the text before they are covered in lecture. Read the Intro and Summary first.
•	After lecture, carefully read the topics covered in class.
	Sample exercises (in the chapter)
	Practice exercises (at end of chapter)
•	Learn the language of chemistry
•	Attempt all of the assigned end-of-chapter exercises
	Spending more than twenty minutes on a single exercise is rarely effective unless you know it is particularly challenging

Chapter 1	Introduction: Matter and Measurement

Chemistry is the study of the properties of materials and the changes that materials undergo.

The Study of Chemistry
The Molecular Perspective of Chemistry
Matter is the physical material of the universe. It is anything that has mass and occupies space.
Atoms are the almost infinitely small building blocks of matter.
Molecules combinations of two or more atoms with the atoms attached to one another in a specific way.

Macroscopic view versus microscopic view
Macroscopic is the scale of ordinary sized objects. Microscopic is the atomic level. We explain macroscopic properties (e.g. boiling point) in terms of microscopic properties (molecular structure)

Why Study Chemistry?

Classifications of Matter
States of Matter
Gas – atoms or molecules are far apart and moving very fast.
Liquid – particles are more closely packed and still moving relatively fast.
Solid – held tightly together, usually in definite arrangements, wiggle (vibrate) only slowly in their fixed positions.

Pure Substances and Mixtures
Pure substances – have fixed composition and distinct properties.
Elements – cannot be decomposed into simpler substances (have only one kind of atom)
Compounds – composed of two or more elements (contain two or more kinds of atoms)
Mixtures – combinations of two or more substances in which each substance retains its own chemical identity and hence its own properties.
Heterogeneous – does not have the same composition, properties, and appearance throughout.	ex: sand
Homogeneous-constant composition throughout, also known as solutions.
		Gaseous solutions	air
		Liquid solutions	gasoline
		Solid solutions		brass
Separation of Mixtures
	Solubility differences	filtration
	Volatility differences	distillation
						chromatography

Elements	112 elements
Chemical symbols		H, O, N, C, Li, Cl, Br, Hg, Au
Periodic Table; elements with similar properties arranged in columns
Compounds – elements combine to form compounds. Compounds can be decomposed to the elements.
The Law of Constant Composition: The elemental composition of a pure substance is always the same.

Properties of Matter
Physical properties: color, odor, density, melting point, boiling point, hardness.
Chemical properties: reactivity; change in chemical composition.
Intensive properties: independent of amount of material present (melting point, density).
Extensive properties: depend on amount of material present (mass, volume).
Physical and Chemical Changes
Physical changes do not effect the composition of the material. Most common is change of state such as evaporation (vaporization: liquid to gas	fusion: liquid to solid	sublimation: solid to gas)
Chemical changes (also known as reactions) a substance is changed into a chemically different substance.

Units of Measurement
Many properties of matter are quantitative. 
Units must be specified (meters, grams, liters)
17.5 is meaningless without units (feet, inches, meters, gallons?) 
17.5 cm specifies the length

The metric system is used scientific measurements.
SI Units (Système International d’Unitès)
Length and Mass 
Unit of length is the meter. 
Mass is a measure of the amount of material in an object, not the weight. Unit of mass is the gram.
1 pound is about 2.2 kilograms.
The prefix kilo means one thousand.
Prefixes
Prefix	Abbreviation	Meaning	Example	
Giga	G	109	Gigabyte	
Mega	M	106	Megawatt	
Kilo	k	103	Kilogram	
Deci	d	10-1	Deciliter	
Centi	c	10-2	Centimeter	
Milli	m	10-3	Millimeter	
Micro		10-6	micrometer	
Nano	n	10-9	Nanometer	
Pico	p	10-12	Picometer	
Femto	f	10-15	femtosecond	
Temperature
Fahrenheit
Celsius scale – based on freezing point (O °C) and boiling point (100 °C) of water.
Kelvin scale – 0 K is –273.15 °C (absolute zero)

The Celsius and Kelvin scales have equal sized units.

Derived SI Units
Speed = ms-1		force = ma = kg m s-2
Volume		length cubed	(length)3
SI: meter3		dm3 = 1 liter = 1000 ml = 1000 cm3 
1 ml = 1 cm3

Density of water = 1.00 g/cm3 

Uncertainty in Measurement
Exact numbers have defined numbers (c, e, π, R) or are integers that result from counting numbers of objects (12 eggs is a dozen, 1000 g is a kg)
Inexact numbers result from a measurement.
Equipment errors (calibration)
Human errors (how individuals take measurements)
Uncertainty always exists in measured quantities.

Precision and Accuracy
Precision is how closely individual measurements agree with one another (reproducibility of results)
Accuracy is how closely individual measurements agree with the correct or “true” value.

Significant Figures
Measured quantities are generally reported in such a way that only the last digit is uncertain.
2.2405±0.0001 g		2.2405 g
All digits including the uncertain one are called significant figures. The number above has 5 significant figures. 
The number of significant figures indicates the exactness of a measurement. 

Guidelines to significant figures in a measured quantity
1.	Nonzero digits are always significant
457 cm (3 sig. figs)	2.5 g (2 sig. figs)
2.	Zeros between nonzero digits are always significant 
1005 kg (4 sig. figs)	1.03 cm (3 sig. figs)
3.		Zeros at the beginning of a number are never significant; they merely indicate the position of the decimal point
		0.02 g (1 sig. figs)	0.0026 cm (2 sig. figs)
4.	Zeros that fall both at the end of a number and after a decimal point are always significant
		0.0200 g (3 sig. figs)	3.0 cm (2 sig. figs)
5.	When a number ends in zero but contains no decimal point, the zeros may or may not be significant
130 cm (2 or 3 sig. figs)	10,300 g (3, 4, or 5 sig. figs)

The use of exponential notation (Appendix A) avoids possible ambiguity seen in the last example
Using exponential notation 10,300 g can be written 
			(three significant figures)

			(four significant figures)

			(five significant figures)

Significant Figures in Calculations
In multiplication and division the result must be reported with the same number of significant figures as the measurement with the fewest significant figures.
1.		If the leftmost digit to be removed is less than five, the preceding number is left unchanged.
Rounding 7.248 to 2 sig. figs gives 7.2

2.		If the leftmost digit to be removed is 5 or greater, the preceding 	number is increased by 1.
Rounding 4.735 to three significant figures gives 4.74
Rounding 2.376 to two significant figures gives 2.4.

In addition and subtraction the result cannot have more digits to the right of the decimal point than any of the original numbers.

		Round off to 105

Dimensional Analysis
In dimensional analysis we carry units through all calculations.
A conversion factor is a fraction whose numerator and denominator are the same quantity expressed in different units.

Example: 1 inch = 2.54 cm
 	and  

How many centimeters are in 8.5 inches? 

Using two or more conversion factors.
Convert 8.00 m into inches

Conversions Involving Volume
What is the mass in grams of two cubic inches of gold?
The density of gold is 19.3 g/cm3.

Summary of Dimensional Analysis
In using dimensional analysis to solve problems, we will always ask three questions:
1.	What data are we given in the problem?
2.	What quantity do we wish to obtain in the problem?
3.	What conversion factors do we have available to take us 
from the given quantity to the desired one?