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CHAPTER ONE: CHEMISTRY: THE SCIENCE OF CHANGE

Part One: An Introduction to Chemistry.

A. What is Chemistry?

1. Chemistry is the science that describes matter - its properties, the changes it undergoes, and the energy changes that accompany those processes.

2. Chemistry is the central science.

3. Chemistry is a quantitative science based on experimentation.

Quantitative- Based on the amount of something

Experiment An observation of natural phenomena carried out in a controlled manner so that the results can be duplicated and rational conclusions made.

B. Five Traditional Branches of Chemistry:

1. Organic chemistry = chemistry of covalent compounds of carbon and hydrogen and their derivatives.

The chemistry of carbon

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2. Inorganic chemistry = the study of the huge variety of substances that fall outside the realm of organic materials.

The chemistry of everything else

3. Analytical chemistry = the identification of substances present and their amounts in a sample.

4. Physical chemistry = applies methods of physics to the properties of matter and the accompanying energy changes.

5. Biochemistry = the study of the chemistry of processes in living organisms.

C. Matter and Energy - Definitions

1. Matter = material things = anything that has mass and occupies space.

2. Mass = the quantity of matter present in a sample = measured by resistance of an object to acceleration.

F = ma, where F = force required, m = mass, and a = acceleration.

3. Energy = capacity to do work or to transfer heat.

4. Forms of energy = mechanical energy, light energy, electrical energy, and heat energy, and others.

5. Energy has two essential manifestations: kinetic energy and potential energy.

a. Kinetic energy is energy of motion = mv2/2.

b. Potential energy is stored energy = the energy an object possesses because of its position, condition, or composition.

6. Chemical and physical processes are typically accompanied by energy changes.

a. If energy is released to the surroundings, usually as heat energy, we call such processes exothermic.

b. Endothermic processes absorb energy from their surroundings as they happen.

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7. The Law of Conservation of Matter. (a Natural Law) The total mass remains constant during a chemical reaction or during a physical change.

Note: A nuclear reaction is not a chemical reaction, and mass can be lost, i.e., converted to energy according to E = mc2.

D. Classification of Material Samples. (Section 1.2)

1. Substances- a form of matter that has a definite (constant) composition and distinct properties, such as color, smell, and taste. Cannot be further broken down or purified by physical changes.

a. Elements cannot be decomposed into simpler substances even by chemical changes.

b. Compounds can be decomposed by chemical changes into simpler substances. (eventually to individual elements)

Properties of the individual components are lost.

2. Mixtures are combinations of two or more substances in which each substance retains its own identity and properties; can vary the composition of components; can be separated by physical changes.

a. A mixture which is not uniform throughout is called

b. Mixture that has uniform properties throughout is a

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3. A little quiz - categorize the following samples of matter as: A) heterogeneous mixture, B) homogeneous mixture, C) compound, or D) element:

diamond / sea water (drawn from surf) / limestone / gravy / steel / table salt / gold / sugar

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E. Three States of Matter. (Section 1.2)

1. Solid = substances are rigid and have definite shapes. Volumes of solids do not vary much as temperature and pressure are changed; very hard to compress or expand.

2. Liquid = flows and assumes the shape of its container up to the volume of the liquid; also very hard to compress or expand.

3. Gas = also flows, but much less dense than liquids and solids, occupy all parts of any vessel in which they are confined; capable of infinite expansion and are compressed easily; consist primarily of empty space.

Solid Liquid Gas

Shape Definite Variable Variable

Volume Definite Definite Variable

Fluid No Yes Yes

Compressible No No Yes

F. Chemical and Physical Properties. (Section 1.3)

1. Chemical properties = properties related to the kinds of chemical changes (reactions) that substances undergo as they change to different substances. (reactivity)

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2. Physical properties = characteristics of a substance in the absence of any change in composition. For example: color, density, hardness, melting point, boiling point, and electrical and thermal conductivities. Can be observed without changing its chemical identity.

3. Physical properties can be further classified according to whether or not they depend on the amount of sample present.

a. Extensive properties depend on the amount of material examined. Examples =

*Extensive properties with the same units can be added to one another.

b. Intensive properties are independent of the amount of material. Examples =

*Intensive properties cannot be added to one another.

4. All chemical properties are intensive properties.

G. Chemical and Physical Changes. (Section 1.3)

1. Chemical change- (reaction) change in which one or more kinds of matter are transformed into one or multiple kind(s) of new matter.

a. One or more substances are used up (at least partially).

b. One or more new substances are formed.

c. Energy may be absorbed or released.

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2. Physical change- a change in the form of matter but not in its chemical identity.

a. No change in chemical composition (no new substances formed).

b. Physical properties are usually altered significantly.

c. Energy may be absorbed or released.

d. Examples =

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Distillation:

Part Two - Measurements in Chemistry

A. Handling Numbers in Measurements. (Section 1.4)

1. Scientific Notation: study text and do homework.

a. Useful for expressing very large or very small numbers.

b. For example, speed of light, c = 299792500 m/s =

c. Number of particles in a mole = 602217000000000000000000=

d. Bohr radius of a Hydrogen atom = 0.0000000000529 m =

e. Mass of earth = 5,973,600,000,000,000,000,000,000 kg =

f. Plancks constant, h= 0.0000000000000000000000000000000006626 J*s =

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g. Scientific notation makes it much easier to look at, compare, and use numbers.

h. Converting to scientific notation:

Find base number by moving decimal place behind first non-zero digit: 5,973,600,000,000,000,000,000,000 kg 5.9736

Find exponent by counting the number of places you moved the decimal: 5,973,600,000,000,000,000,000,000 kg 5.9736 x 1024

If moving to the right, the exponent is positive. (large numbers) If moving to the left, the exponent is negative. (small numbers)

3. Precision and Accuracy Precision How close a series of replicate measurements (measurements of the

same thing) are to one another. Accuracy Closeness of a single measurement to the true value.

4. Significant Figures: study text and do homework. (Section 1.5)

a. Significant figures (sig figs) are digits in a measured number (or in a result of a calculation involving measured numbers) that include all certain digits plus a final digit having some uncertainty. (The meaningful digits reported in a number)

b. Sig figs express the precision to which a measured number has been determined.

The more sig figs, the more precise the measurement/ number.

Good accuracy

Poor precision

Good precision

Poor accuracy

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For glassware, more (further) graduations give more precise measurements.

For example, when I read a buret, I record the volume to the nearest 0.01 mL, which is the limit of precision of the buret,

c.. Determining how many sig figs are being expressed in a number:

All non-zero digits are significant.

All embedded zeroes are significant.

All terminal zeroes on the R of the decimal are significant.

Terminal zeroes at the end of the number written without a decimal point are ambiguous (you dont know what the person intended).

340 cm you dont know whether the person actually measured length to the nearest cm and it came out exactly 0, or whether their number is an estimate of length rounded to the nearest 10 cm.

For scientific purposes, if they put in a decimal point, the ambiguity is removed:

340. cm implies knowledge to the nearest cm.

The ambiguity can always be removed by re-expressing such numbers in scientific notation.

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20

22?

24

23

23.4?

24

23

23.39

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3.40 x 102 indicates knowledge to the nearest cm.

3.4 x 102 indicates knowledge only to the nearest 10 cm.

d. Try some examples:

e. Sig figs in calculations:

Multiplication and division - count the number of sig figs in each number. The product or quotient should have the same # of sig figs as the number with the least number of sig figs.

0.025 1.745435

= 0.000100287 = 0.00010 =1.0 104

Addition and subtraction - answer should have same number of decimal places as the number having the least number of decimal places.

0.025 + 1.745 = 1.770 (keep to nearest 1000th)

20.1 - 0.32 = 19.78 19.8 (round to nearest 10th )

Calculations involving mixture of multiplications/division AND additions/sub-tractions. Do in pieces.

20.01 x 1.08 + 9.750 x 10.21 = 3 s.f. (~21.6) + 4 s.f. (~99.54)

= nearest 0.1 (121.2)

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Harder example:

12.5 95 15.6 1251.70 104

=1.2 103 1.95 103

1.70 104

=

0.8 103

1.70 104=

1 sig fig3 sig figs

=1 sig fig

Now repeat calc not rounding off till end. Round to 1 sig fig.= -0.044852941 = - 0.04

B. Units of measurement (Section 1.4)

1. International System of Units (SI). All other units of measurement are derived from them.

Quantity Unit Symbol length mass time electric current temperature amount of substance

2. Derived SI units a combination of two or more of the above basic SI units.

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3. Metric and SI systems are decimal systems; prefixes used to indicate powers of ten.

Prefix Symbol Multiple tera- T 1012 giga- G 109 mega- M 106 kilo- k 103 deci- d 10-1 centi- c 10-2 milli- m 10-3 micro- 10-6 nano- n 10-9 pico- p 10-12

C. Commonly used Units. (Section 1.4)

1. Mass and Weight.

a. Mass- Measure of the amount of matter in a sample or object. Weight- force exerted by an object or sample due to gravity.

b. Basic unit of mass in SI =

c. Conversion: 1 pound (lb) = 0.4536 kg = 453.6 g.

d. Origin: 1 gram = mass of 1 cm3 of liquid water at 4.0 Celsius, pressure = 1 atm.

2. Length.

a. Meter (m) is standard unit of length (distance) in both SI and metric systems.

b. 1 in = 2.54 cm (exactly); 1 m = 39.370 inches (in);

c. 1 ngstrom = 0.1 nm = 1 x 10-10 m ~ size of atoms (non-SI)

d. Origin: 10,000,000 meters = distance from N Pole to Equator through Paris, France.

3. Volume. (Derived unit)

a. SI unit of volume = 1 m3. Not very convenient.

b. Chemists use liters (L) or milliliters (mL) in metric system.

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c. One liter (1 L) is 1 cubic decimeter (1 dm3), or 1000 cubic centimeters (1000 cm3).

d. One milliliter (1 mL) is 1 cm3.

e. 1 L = 1.057 qt.

4. Density and Specific Gravity. (Derived units)

a. Density d = m/V (an intensive property).

b. Usually expressed as g/mL, which is the same as g/cm3.

c. Thus, density of liquid water is 1 g/mL (at 4 C) and = 0.998 g/mL at 20 C.

d. Problem: 3.0 mL of liquid Hg has a mass of 40.8 g. What is its density in g/mL?

e. Problem: If reaction uses up 10 g ethanol, what volume of ethanol is used up? Density of ethanol is 0.789 g/mL.

f. Specific gravity is density of a sample relative to that of water = ratio of substances density to density of water, both at same temperature.

Sp. Gr. = dsubstance/dwaters

where dwater = 1.00 g/mL over wide range of T.

g. Exercise: What is the specific gravity of Hg?

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D. Temperature Units.

1. T is a measure of the average kinetic energy of a substance. It can also be thought of as the hotness or heat intensity of a sample.

2. An property.

3. SI unit of temperature is . Also called absolute temperature because 0 K is the lowest possible temperature. (-273.15 Celsius).

4. The Celsius (C) scale (named after Ander Celsius) is based on freezing and boiling temperature of water.

Freezing temp of water defined at 0 C at sea level Boiling temp of water defined at 100 C at sea level

5. The Fahrenheit (F) scale (named after Daniel Fahrenheit) is based on:

-Freezing point of ammonium chloride / ice mixture of equal proportion = 0 F.

-Average human body temperature = 100 F (he was off a bit).

6. Conversion equations:

T (K) = T(C) +273.15

T (F) = (9/5) T(C) + 32

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E. Using Units to Solve Problems: Unit-Factor or Factor-Label Method. (Section 1.6)

1. Main idea: Exploit the units to your advantage to help you solve problems.

2. Unit factors (Conversion factors) may be constructed from equalities such as:

1 ft = 12 in. Rearrange to give:

3. Since we may always multiply a number by one without changing the number, unit factors are useful for unit conversion.

4. Exercise: The radius of a P atom is 1.10 ngstrom. Express this in inches.

Strategy: three-step conversion m; m cm; cm in

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5. Exercise - Volume Conversion: Calculate the volume of the P atom in milliliters (mL).

6. Problem: Express the density of mercury in lb/ft3. The density of mercury is 13.59 g/cm3.

F. Percentage.

1. Means parts per 100.

2. Calculation of percentage:

% Component A = (amount A in mixture) x 100%. (total amount of mixture)

3. Example: Class has 80 men and 70 women. What % are men?

4. Example: If the same percentage applies to the entire student body, and there are 12,000 students, how many men?

5. Hard problem involving Sp. Gr., Volume, and %:

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A certain water/ethanol mixture is 70% ethanol by mass, and has a specific gravity of 0.95. Calculate the mass of pure ethanol in a 4.000 Liter bottle of this mixture.

Chapter 1 Summary: Classification of Matter

Elements, substances, hetero/homogeneous mixtures. Phases of matter.

Physical/chemical processes & properties. Extensive vs. intensive properties.

Units SI units

Dimensional Analysis Unit conversions Significant figures

NOTES: