Unit 1: Introduction to Chemistry

1.1 What is Chemistry?

• Chemistry- the study of substances and the changes they can undergo. EX: a match burning, how bleach removes stains, why bread dough rises, etc.

• The Central ScienceChemistry overlaps into many other sciences: Biology, Geology, Astronomy, etc.

• Chemicals are everywhere, in everything, and impact many different aspects of life. Chemistry, therefore, is considered a central science. Life, as we know it, is a product of what Chemistry and Physics has already done.

• B) Why Study Chemistry?– To help you understand the physical

world around you. To develop skills for evaluation and critical thinking. Maybe even help prepare you for a job which requires chemistry.

– (ex. occupations which require chemistry: Engineering, medical professionals, hair stylists, crime labs, cosmetic makers, drug developers, oil companies, Wine makers, Mc Donald’s, Candy makers, Photographers …)

1.2 The Scientific Method

• Scientific Method- an orderly, systematic approach to gather knowledge. It is a way of answering questions about our observable world.

Steps of the Scientific Method

1. Make an observation2. State the question3. Collect information4. State a hypothesis5. Design an experiment6. Make observations 7. Collect, record and study data8. Draw a conclusion

Sci. Method Steps Explained

Making an ObservationMaking an Observation

Notice a natural event: the ball falls to the ground, the sky is blue, etc. This observation can be about almost anything! Once you’ve noticed something… form a question.

Forming a HypothesisForming a Hypothesis

This should be a possible, logical, answer to the question about your observation. It is typically expressed in a “cause-and-effect” format. A scientific hypothesis must be one which requires and can be tested by an experiment. If it does not… it is not “scientific”.

Sci. Method Steps Explained

Performing an ExperimentPerforming an Experiment

• For a hypothesis to be tested properly, you must design and perform an experiment which examines ONE variable at a time. If you have more than one variable the results will not be conclusive and very little knowledge will be gained.

Interpreting the ResultsInterpreting the Results

• Once the experiment is complete… you look at your data and the observations you made interpret what they tell you. Did you prove your hypothesis wrong? Did you learn anything new? (Experimental control)

Laws and Theories

LawLaw• a statement of fact meant meant

to explicateto explicate, in concise terms, an action or set of actions. It is generally accepted to be true and universal, and can sometimes be expressed in terms of a single mathematical equation. THEY TELL WHAT HAPPENED.

TheoryTheory• an explanationan explanation of a set of

related observations or events based upon proven hypotheses and verified multiple times by detached groups of researchers. One scientist cannot create a theory; he can only create a hypothesis. THEY EXPLAIN AND PREDICT EVENTS.

1.4 Units of Measurement

• The International System of Units

– In 1960, at a scientific conference on units held in France, the SI system of units were internationally accepted for the scientific community. The SI system is based on the metric system and we refer to these as base units.

BASE UNITS

Mass kilogram kg

Length meter m

Time seconds s

Count quantity mole mol

Temperature kelvin K

Electric current ampere A

Luminous intensity candela cd

• Meter- defined as the distance that light travels in a vacuum during a time interval of 1/299,792,458 of a second.

• Mass- amount of matter in an object. 1 kg = 2.2 lbs (on earth).

• Weight - equals the force of gravity pulling on the object.

• Derived units - a combination of 2(+) base units = a new unit.

DERIVED UNITS

Area Square meter m2

Volume Cubic meter m3

Force Newton N

Pressure Pascal Pa

Energy Joule J

Power Watt W

Voltage Volt V

Frequency Hertz Hz

Electric charge Coulomb C

Area- length X width = m X m= m2

Volume- the amount of space that an object occupies.

Length X width X height = m X m X m= m3

EXCEPTIONS…

The liter (L)- the common unit for volume.

1mL= 1cm3

Celsius (C)- common unit for temperature

1K = (273 + C)

• Metric Prefixes– Prefix- a word attached to the front of

the base unit. The SI prefixes are base 10 and, therefore, increase and decrease by 10’s.

Prefix Abbreviation # Power of 10

mega- M 1,000,000 106

kilo- k 1,000 103

hecto- h 100 102

deca- da 10 101

Base 1 100

deci- d 0.1 10-1

centi- c 0.01 10-2

milli- m 0.001 10-3

micro- 0.000001 10-6

nano- n 0.000000001 10-9

Converting among prefixes

• When converting from one prefix to another, remember this saying:– King Henry Died By Drinking

Chocolate Milk.

• When set up as such: – k h da _ d c m

• Now converting among prefixes is just a matter of pushing the decimal

1.5 Working with Numbers

• Significant Digits– 1. Leading zeros are never significant.

– 2. Imbedded zeros are always significant.

– 3. Trailing zeros are significant only if the decimal point is specified.

Hint: Change the number to scientific notation. It is easier to see.

EXAMPLES:Example Number of

Significant Figures

Scientific Notation

0.00682 3 6.82 x 10-3 Leading zeros are

not significant.

1.072 4 1.072 (x 100) Imbedded zeros are always significant.

300 1 3 x 102 Trailing zeros are significant only if

the decimal point is specified.

300. 3 3.00 x 102

• Addition & Subtraction:The last digit retained is set by the first doubtful digit.

Addition Even though your calculator gives you the answer 8.0372, you must round off to 8.04. Your answer must only contain 1 doubtful number. Note that the doubtful digits are underlined.

Subtraction Subtraction is interesting when concerned with significant figures. Even though both numbers involved in the subtraction have 5 significant figures, the answer only has 3 significant figures when rounded correctly. Remember, the answer must only have 1 doubtful

digit.

• Multiplication or Division:The answer contains no more significant figures than the least accurately known number.

Multiplication The answer must be rounded off to 2 significant figures, since 1.6 only has 2 significant figures.

Division The answer must be rounded off to 3 significant figures, since 45.2 has only 3

significant figures.

Scientific Notation • Chemists often work with numbers that are

extremely large or extremely small. – For example, there are

10,300,000,000,000,000,000,000 carbon atoms in a 1-carat diamond each of which has a mass of 0.000,000,000,000,000,000,000,020 grams. It is impossible to multiply these numbers with most calculators because they can't accept either number as it is written here.

• To do a calculation like this, it is necessary to express these numbers in scientific notation, as a number between 1 and 10 multiplied by 10 raised to some exponent.

*Exponent ReviewSome of the basics of exponential mathematics are

given below.• Any number raised to the zero power is equal to 1.

10= 1 100= 1

• Any number raised to the first power is equal to itself. 11 = 1 101 = 10

• Any number raised to the nth power is equal to the product of that number times itself n-1 times.

22 = 2 x 2 = 4 105 = 10 x 10 x 10 x 10 x 10 = 100,000

• Dividing by a number raised to an exponent is the same as multiplying by that number raised to an exponent of the opposite sign.

Converting to Scientific Notation

The following rule can be used to convert numbers into scientific notation: The exponent in scientific notation is equal to the number of times the decimal point must be moved to produce a number between 1 and 10.

• Ex. In 1990 the population of Chicago was 6,070,000. To convert this number to scientific notation we move the decimal point to the left six times.– 6,070,000 = 6.070 x 106

• To convert numbers larger than 1, we will move the decimal point to the left. For example, 10,300,000,000,000,000,000,000 carbon atoms into scientific notation, we move the decimal point to the left 22 times.

–10,300,000,000,000,000,000,000 = 1.03 x 1022

• To convert numbers smaller than 1 into scientific notation, we have to move the decimal point to the right. The decimal point in 0.000985, for example, must be moved to the right four times.

–0.000985 = 9.85 x 10-4

• The primary reason for converting numbers into scientific notation is to make calculations with unusually large or small numbers less cumbersome. Because zeros are no longer used to set the decimal point, all of the digits in a number in scientific notation are significant, as shown by the following examples. 2.4 x 1022 2 sig. figs

9.80 x 10-4 3 sig. figs

1.055 x 10-22 4 sig. figs

Ratios• Units found by dividing one unit by another.  (The

speedometer in your car registers the ratio of miles/hour.) The most common ratio in chemistry is density (g/ml or g/cm3). Density is calculated by this formula:

density = mass/volume

• Lets say you had an object that’s mass was 20g and its volume was 10cm3.  How would you calculate the density?– Density = mass/volume = 20g/10cm3 = 2g/cm3

• If you are given the mass and the density can you calculate volume?– Yes!  Density = mass/volume ► volume = mass/density.

1.6 Problem Solving

• Dimensional Analysis- technique of converting between units.  Unit equalities show how different units are related (1g=100cm).  Conversion factors are written from the unit equalities.  The conversion factor is set up so that the bottom number cancels the given unit and a new unit is created.  – Example:  Convert 10 cm to inches. Conversion

factors (1m = 100 cm)   (1m = 39.37inches)

• Start with the given unit, then use you conversion factors to cancel units until to arrive at the unit you want to convert to.