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CHAPTER ONE: CHEMISTRY AND MEASUREMENT

Part One: An Introduction to Chemistry. (Sections 1.1 - 1.4) 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.

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3. Chemistry is a quantitative science based on experimentation. (Section 1.2) See Figure 1.7.

B. Five Traditional Branches of Chemistry:

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

their derivatives. 2. Inorganic chemistry = the study of the huge variety of substances that fall outside the

realm of organic materials. 3. Analytical chemistry = the identification of substances present and their amounts in a

sample.

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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 (Section 1.3) 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.

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.

Demo - burning of wood splint

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8. The Law of Conservation of Energy.

The total energy remains constant in a chemical reaction or in a physical change. It may, however, be converted from one form to another.

D. Three States of Matter. (Section 1.4)

1. Solid = substances are rigid and have definite shapes. Volumes of solids do not vary

much 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.

Show HyperChem views of solid liquid gas.

E. Chemical and Physical Properties. (Section 1.4)

1. Chemical properties = properties related to the kinds of chemical changes (reactions)

that substances undergo as they change to different substances. (reactivity)

2 Mg(s) + O2(g) 2 MgO(s) Demo: burn Mg ribbon

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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.

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 = volume, energy, mass

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

density, temperature 4. All chemical properties are intensive properties.

F. Chemical and Physical Changes. (Section 1.4) 1. Chemical change. H2O2 decomposition w/ MnO2

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. 2. Physical change. Liquid Nitrogen fills balloon.

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 = melting, freezing, vaporization, condensation, sublimation, dissolving

G. Classification of Material Samples. (Section 1.6)

1. 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.

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a. A mixture which is not uniform throughout is called heterogeneous. (e.g. granite) Bottle of pop with gels.

b. Mixture that has uniform properties throughout is a homogeneous mixture (or a

solution). (e.g. tea) 2. Substances cannot be further broken down or purified by physical changes.

a. Compounds can be decomposed by chemical changes into simpler substances, always in the ratio by mass. (eventually to individual elements)

b. Elements cannot be decomposed into simpler substances even by chemical

changes.

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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Part Two - Measurements in Chemistry. (Sections 1.5 - 1.8) A. Handling Numbers in Measurements. (Section 1.5)

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 = 2.998 x 108 m/s.

c. Number of particles in a mole = 602217000000000000000000 =6.02 x 1023 mol-1.

d. Bohr radius of a Hydrogen atom = 0.529 x 10-10 m. 2. Significant Figures: study text and do homework.

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

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,

e.g., 20.00 mL or 5.24 mL or 2.80

1st number has 4 sig figs, 2nd and 3rd have 3 sig figs

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

All non-zero digits are significant. 0.02050 All embedded zeroes are significant. 0.02050

All terminal zeroes on the R of the decimal are significant. 0.02050

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.

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

3.40 x 102 indicates knowledge to the nearest cm.

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

c. Try some examples:

d. 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 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/subtractions. 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 (121.1)

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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

3. Accuracy: See laboratory manual. 4. Precision: See laboratory manual.

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

them.

Quantity Unit Symbol length meter m mass kilogram kg time second s electric current ampere A temperature Kelvin K amount of substance mole mol

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 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 1.8)

1. Mass and Weight.

a. Basic unit of mass in SI = kilogram (kg).

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

c. 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

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

3. Volume.

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

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

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.

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4. Density and Specific Gravity

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? d = m/V = 40.8 g/3.0 mL = 13.6 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? Sp. Gr. = dHg/dH2O = (13.59 g/mL)/(1.00 g/mL) = 13.59.

Note: There are no units for Specific Gravity. It is a ratio.

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D. Temperature Units.

1. T is a measure of hotness or heat intensity of a sample.

2. An intensive property.

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

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

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

5. The Fahrenheit (F) scale 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 1.8) 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:

1= 12in1 ft

or 1= 1 ft12in

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

1.10 110-10m

1

100cm1m

1in2.54cm

= 4.33109in

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

(mL).

Vsphere =43r3

r =1.10

V = 43 1.10( )3

= 5.573

5.573 11010m

1

3

100cm

1m

3

1mL1cm3

= 5.57 1024 mL

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

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F. Percentage. (Section 1.11) 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? 80/150 x 100% = 53.3% 4. Example: If the same percentage applies to the entire student body, and there are 8000

students, how many men? 8000 x (53.3/100) = 4264 men 5. Hard problem involving Sp. Gr., Volume, and %: 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. 70% ethanol by mass ---> 70g ethanol

100g mixture S.G. of 0.95 ---> 0.95g mixture

1 mL mixture 4.000 L mixture x 1000 mL x 0.95g mix x 70g ethanol

1 L 1 mL mix 100g mix = 2660 g