Review: INTRO TO CHEM

1. Know the difference b/t mass and weight 2. Know Avogadro’s number and how to use it in a dimensional

analysis problem. 3. Know what conversion factors are and how to use them in

dimensional analysis problems. 4. Know the different methods for weighing and measuring and how

to use them in a dimensional analysis problem. 5. Know the difference b/t precision and accuracy. 6. Know significant figures, the rules for significant figures and how to

use them in a dimensional analysis problem. 7. Know scientific notation, the rules for scientific notation & how to

use it in a dimensional analysis problem. 8. Be able to find the molar mass of compounds and elements (using

the CORRECT set-up). 9. Be able to convert between units (dimensional analysis) using the

CORRECT set-up.

Mass versus Weight

Mass • The amount of matter that

is contained by an object how much substance there is

• Value is independent of gravity

• SI unit is kilogram (kg)

Weight • The force exerted on an

object by gravity how

much pull of the Earth on an object

• Value changes in different gravitational fields.

• SI unit is Newton (N) = kg*m/s2 Anything that occupies space

and has mass is called MATTER

The amount of matter that is contained by an

object.

The force exerted on an

object by gravity.

magi push_ pull

mass multiplied by gravity

The Mole and Avogadro’s Number • Since molecules are so small, this unit was developed

to measure them effectively.

• This is a very important number and you will see it constantly in chemistry.

• A mole (mol) of something is equal to 6.022 x 1023

particles of that substance (like the word a dozen).

• 6.022 x 1023 is Avogadro’s number.

Atomic and Molecular Weight

• Periodic table lists the atomic weight of each atom

• Different units: – Atomic mass units (amu)

– Grams per mole (g/mol): weight of one mole of the atom or compound

• Atomic and molecular weight are conversion factors for the conversion of grams to moles and moles to grams of a substance

Weighing & Measuring

• A measurement always consists of a number AND a unit; meaningless without both.

• 2 major systems of units

• International System of Units (SI) – Universal system used the world over. – Involves the multiplication of the base unit by some power of

10. – It is also known as the metric system.

• US Customary System (inherited from British Imperial

System)

Units

• A measurement always consists of two

parts: number and a unit.

• Both parts must be present in order for

a measurement to be meaningful.

Metric System Base Units Physical Quantity Name of Base Unit Unit Symbol/ Abbreviation

Length Meter m

Mass Kilogram kg

Time Second s (or sec)

Temperature Kelvin K

Amount of Substance mole mol

Energy joule J

Volume

Cm3 for solid mL for liquids Liters for gases

Prefixes used in Metric Prefix Abbreviation Scientific Notation Meaning

tera T 1 x 1012 times basic unit

giga G 1 x 109 times basic unit

mega M 1 x 106 times basic unit

kilo k 1 x 103 times basic unit

hecto h 1 x 102 times basic unit

deka da 1 x 101 times basic unit

grams meters Liters

deci d 1 x 10-1 times basic unit

centi c 1 x 10-2 times basic unit

milli m 1 x 10-3 times basic unit

micro µ 1 x 10-6 times basic unit

nano n 1 x 10-9 times basic unit

pico p 1 x 10-12 times basic unit

The U.S. Customary system Prefixes are NOT used

• Mass

– 1 pound (lb) = 16 ounces (oz) – 2000 pounds = 1 ton

• Length – 1 foot (ft) = 12 inches (in) – 1 yard (yd) = 3 feet (ft) = 36 inches (in) – 1 mile (mi) = 5280 feet (ft) = 1760 yards (yd)

• Volume – 1 gallon (gal) = 4 quarts (qt) – 1 quart (qt) = 4 cups (C) = 2 pints (pt) – 1 pint (pt) = 2 cups (C)

Significant Figures

• Precision & Accuracy – Two terms used to describe the RELIABILITY of a

measurement

– In everyday language, used interchangeably, but in science they have VERY DIFFERENT DEFINITIONS

–ACCURACY the agreement of a value with the standard value

–PRECISION the degree of agreement among several measurements of the same quantity

Precision vs. Accuracy

What are Significant Figures? • At times, calculating the final result for an experiment

involves multiple mathematical operations.

• Dependent upon how rounding is performed, can end up with very different results.

• Knowledge of the amount of uncertainty is important to lend RELIABILITY to the final mathematical result.

• This reliability is obtained via the use of significant figures.

• There are very specific rules governing the correct use of significant figures.

Significant Figures: tha rules

1. NON-ZERO integers are ALWAYS SIGNIFICANT.

2. There are 3 different types of ZEROES: a. LEADING zeroes NEVER significant

• For example: 0.0025 has only 2 sig figs.

b. CAPTIVE zeroes ALWAYS significant • For example: 0.00205 has 3 sig figs.

c. TRAILING zeroes SOMETIMES significant • For example: 0.00250 has 3 sig figs and 2.50 has 3 sig figs.

• Zeroes to the RIGHT of a non-zero digit are only significant when the number contains a decimal.

Significant Figures: tha rules cont’d 3. When multiplying or dividing:

a. The number of sig figs in the result is the same as the number of the LEAST PRECISE measurement used in the calculation. • For example: 1.4 * 4.56 = 6.4

• The result has 2 sig figs which is the same as the limiting term.

4. When adding or subtracting: a. The result must have the SAME NUMBER of decimal places

as the LEAST PRECISE measurement used in the calculation.

• For example: 12.1 +7.013 = 19.1

• The result has only one decimal place which is the same as the limiting term.

Scientific Notation

• exponential notation a manner of writing numbers for values too large or small to be conveniently written in standard decimal notation

a x 10b

• 300 = 3×102

• 4,000 = 4×103

• 5,720,000,000 = 5.72×109

• 0.0000000061 = 6.1×10−9

Scientific Notation Rules

• Addition/Subtraction • Requires same power

• Multiplication • Multiply number normally, add exponents • E.g. 3x1015 * 2x1013 =

• Division • Divide number normally, subtract exponents • E.g. 9x107 ÷ 3x109 =

Dimensional Analysis

• It is sometimes necessary to convert a result from one set of units to another.

• The best way to do this is via dimensional analysis.

• What is dimensional analysis? – An algebraic process of converting one set of units into

another

– Involves multiplication of a series of “conversion factors”

• Conversion factor an algebraic relationship between two units of interest and is equivalent to ONE.

Conversion between the Systems

• To convert length – 1 inch = 2.54 centimeters

• To convert mass – 1 pound = 454 grams

• To convert volume – 1 Liter = 1.06 quarts

– 1 gallon = 3.79 liters

For example: How many centimeters are in 5.00 inches?

5.00 𝑖𝑛𝑐ℎ𝑒𝑠

1 ×

2.54 𝑐𝑚

1 𝑖𝑛𝑐ℎ= 12.7 𝑐𝑚

Conversion: Atomic and Molecular Weight

• To convert from grams to moles: divide the grams by atomic or molecular weight or molar mass

18 𝑔 𝑜𝑓 𝐶

1÷

12.01 𝑔

𝑚𝑜𝑙 𝑜𝑓 𝐶 = 1.5 𝑚𝑜𝑙 𝑜𝑓 𝐶

• This is the same as:

18 𝑔 𝑜𝑓 𝐶

1 ×

1 𝑚𝑜𝑙 𝑜𝑓 𝐶

12.01 𝑔 = 1.5 𝑚𝑜𝑙 𝑜𝑓 𝐶

Conversion: Atomic and Molecular Weight

• To convert from moles to grams: multiply the moles by atomic/molecular weight

2.2 𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝑂

1 ×

16 𝑔

𝑚𝑜𝑙 𝑜𝑓 𝑂 = 35 𝑔 𝑜𝑓 𝑂