Summary Introduction to General, Organic, and Biological Chemistry (4th edition)\" - Ch 1-5 PDF

Preview

Summary

Ch 1-5...

Description

II) Scientific Approach

Chapter 1 - Chemistry

A) Statement of Problem States what we want to know

I) A Natural Science Observations & classifications of facts about the physical world

B) Record Observations 1) Quantitative mass, volume, temperature, time

A) Chemistry

2) Reproducible exp. never performed just once

Body of knowledge obtained by observations of the physical world of the laboratory.

C) Draw a Conclusion Law: statement of fact which is a necessary conclusion drawn from observations. - true for all cases examined

1) Classification Recognize & describe materials by their characteristics or properties.

Law of Gravity 1

2

III) Model

IV) What is Chemistry? Science that deals with composition, structure & reactions of matter

Idea that explains or correlates a number of facts. - explains how & why something behaves as it does

A) Matter Anything that has mass & occupies space

A) Hypothesis Tentative model - 1st idea - test using scientific method

1) Mass quantity of matter

B) Theory 2) Weight Model that has been tested & not disproved - best idea that agrees with all known facts

Result of gravitational attraction between matter

3

4

B) Composition

2) Macroscopic Level

Identifies what the matter is made of & how much of each component is present.

Amounts that can be seen and weighed

1) Several Ways of Expressing

a) Ex: 1/4 lb. cheeseburger 1) By weight (mass)

a) by weight (mass)

meat cheese roll

b) by volume c) Percent

4.0 oz 0.8 oz 1.7 oz 6.5 oz

d) Number of Moles e) Number of Atoms

5

6

3) Submicroscopic Level described by numbers & types of atoms

C) Structure Arrangement of components & how they are held together, or bonded

Atoms: simple units of matter Molecules: combinations of atoms

Ethanol C2 H 6 O or C2H5OH

a) Qualitative statement of composition

Dimethyl Ether C2 H 6 O or CH3OCH3

Ethanol consists of carbon, hydrogen & oxygen b) Quantitative Description Ethanol: 2 C atoms, 6 H atoms 1 O atom Formula: C2H6O 7

8

D) Reactions

V) States of Matter

Involve changes in composition &/or structure

Substances “normal” state is its physical form at 1 atm pressure & 25°C

Use symbols in a chemical equation to represent a chemical reaction ethanol + oxygen

reactants

A) 3 States of Matter Gas

carbon + water dioxide

yields products or produces or “is converted to”

C2H5OH + 3 O2

cool

heat

Liquid cool

heat Solid

2 CO 2 + 3 H 2 O

9

10

Expands very slightly when heated Expand slightly when heated Great expansion when heated

11

Incompressible Slightly compressible Highly compressible

No def. shape Constant volume Definite volume or volume & & - fills container shape of container shape & takes its shape

Solid Liquid Gas

VI) Physical & Chemical Properties & Changes A) Physical Properties Characteristics that help identify & distinguish different substances AND can be measured WITHOUT changing the basic identity of the substance (NO change in chemical composition)

1) Ex: physical state, color, odor, density, m.p., b.p., specific heat 12

B) Physical Changes

C) Chemical Properties

Change in appearance without change in identity

Describes the way a substance reacts with or is converted into another substance

1) Ex: change in state e.g. - flammability melting

Solid

Liquid

D) Chemical Changes (Reactions)

freezing

Converts a substance into a chemically different substance

vaporization

Liquid

Gas - change in composition &/or structure

condensation sublimation

Solid

Gas 2 K(s) + 2 H2O(R)

deposition

13

2 KOH(aq) + H2(g)

14

E) Law of Conservation of Mass

VII) Mixtures & Pure Substances

No detectable gain or loss of mass during an ordinary chemical reaction

Divide all matter into 2 groups: A) Mixtures Contain 2 or more substances NOT chemically combined.

1) Ex: reaction of K in H2O If 39 g of K & 18 g of H2O produce 56 g of KOH, how much H2 is also produced? mass of reactants

=

Each component retains its own properties (chemical identities).

mass of products

1) Characteristics a) variable composition b) separable by physical means

39 g K + 18 g H2O = 56 g KOH + ? g H2

Ex: water - ethanol mixture mostly water mostly ethanol 50 - 50 mixture 15

16

2) Heterogeneous Mixture

B) Pure Substances

Consists of parts that are unlike

Uniform in properties throughout - do NOT have same composition, prop., & appearance throughout - NOT uniform

a) Ex:

1) Characteristics a) constant (fixed) composition

sand & salt oil & water

b) distinct properties

3) Homogeneous Mixture

c) NOT separable by physical methods

Properties uniform throughout - down to molecular level

Ex: water (with no impurities) always composed of hydrogen (H) & oxygen (O) in same percentage

Also called Solutions

a) Ex: Air: gaseous solution 95% Ethanol: liquid solution Brass: solid solution

H 2O 17

18

VIII) Elements & Compounds

1) Periodic Table Elements arranged in order of increasing atomic number

These are pure substances A) Elements

- properties of elements correlate w. position in periodic table

Substance that canNOT be broken into simpler substances by ordinary chemical means.

a) Periods horizontal rows - gives information about atomic structure

116 known elements Symbols used to identify

b) Groups - 1 or 2 letters

vertical columns - elements in groups have similar physical & chemical properties

C = carbon Co = cobalt 19

20

21

Metals solids (except Hg)

Nonmetals gases or solids (except Br)

metallic luster

variety of color & appearance

malleable & ductile

solids are brittle

good conductors of heat & electricity

poor conductors (insulators)

oxides:

oxides:

nonvolatile high melting

volatile low melting

MgO, Na2O

CO, CO2, SO2 22

B) Compounds

IX) Energy is a Property

Composed of 2 or more elements, chemically combined or bonded.

ability to do work many forms of energy:

- can be separated into its elements by chemical means

heat, electrical, mechanical, kinetic, potential,

H 2O radiant, chemical

11.2 % hydrogen

A) Law of Conservation of Energy

88.8 % oxygen

Energy can be neither created nor destroyed

23

24

B) Chemical Energy Energy stored in a substance because of its composition & structure.

2) Endothermic Reactions heat is absorbed - reaction requires input of energy

1) Exothermic Reactions heat is released

2 KOH + H2 + heat 6 2 K + 2 H2O

2 K + 2 H2O 6 2 KOH + H2 + heat

25

26

X) Spontaneity in Chemistry

Chapter 1 (RJO 1) - Homework

A spontaneous process proceeds on its own without external influence.

2, 6, 11, 14, 17, 23, 25, 28, 29, 31, 34, 35, 37, 39, 40

Generally, @ process that releases energy tends to be spontaneous @ requires energy

nonspontaneous

A) Entropy Amount of disorder or randomness in a system. A reaction that leads to more disorder (greater entropy) tends to be spontaneous. 27

28

Chapter 2 - Measurements

II) Uncertainty in Measurements uncertainties always exist in measured quantities

I) Measurement

A) Precision Degree of reproducibility of repeated measurements

quantitative observation made by comparing to a standard measuring device.

i.e. - How close are to each other

Use the metric system (SI) Consists of both a

Depends on skill of measurer

number and unit

1) Ex: Measure width of notebook paper (in cm)

A) Ex: How far is it from OSU to my house?

21.32

21.33

21.32

21.31

avg. width = 21.32 cm good precision 1

2

B) Accuracy

Ex:

How close measurement is to true value Paper’s true width is 21.59 cm Numbers in previous ex. have poor accuracy Depends on quality of the measuring device 1) Ex: remeasure paper with a “better” ruler (in cm) 21.54

21.61

21.56

21.65

Avg. = 21.59 cm good accuracy, poor precision 3

A ( C)

- good precision poor accuracy

B ( C)

- poor precision poor accuracy

C ( C)

- good precision good accuracy

D ( C)

- “poor” precision good accuracy 4

III) Significant Figures

A) Exact Numbers Infinite number of sig. fig.

All digits we know exactly plus one that we estimate

1) By Count Count the number of people in the room

Calibration of instrument determines number of sig. fig.

-

Measurements of our paper were made using ruler marked in tenths of a cm (mm)

Integers

2) By Definition 1 dozen /

12 items

1

yd

/

1

lb

/ 16

3

ft oz

1 in / 2.54 cm 5

6

B) Significant Figures Rules

4) Trailing zeros: zeros to right of last non-zero digit

1) ALL non-zero digits ARE sig.

1,542

a) Number ends in zero to right of decimal point

3.456

- zeros ARE sig. 2) Captive zeros: zeros between sig. digits ARE sig.

20.6

0.040

400.0

20.06 b) Number ends in zero to left of decimal pt.

3) Leading zeros: zeros to left of first non-zero digit are NOT sig. - locate decimal point

0.401

- zeros generally NOT sig.

400

4100

0.004

7

8

C) Sig. Fig. in Calc. - Rounding Off

b) Ex 2: Determine volume of a box that measures 3.6 cm by 2.45 cm by 10.0 cm.

Result of a calc. must reflect accuracy of original measurements 1) Multiplication & Division

Answer must contain same # of sig. fig. as quantity w. least # of sig. fig.

1) Rounding Rule 1 If leftmost number to be discarded is < 5, round down

a) Ex 1: Divide 907.2 by 453.6 i.e. - last number to be retained is unchanged

ˆ 9

Answer should be:

10

2) Addition & Subtraction Last place in answer is last place common to ALL numbers

b) Ex 4: Find the difference between 12.4 and 4

a) Ex 3: Add 4, 1.45, 12.4 & express answer to correct number of sig. fig.

4 1.45 12.4 17.85

12.4 &4 8.4 c) Ex 5: Add 9.8 and 9.94

9.8 + 9.94 19.74

1) Rounding Rule 2 If leftmost number to be discarded is > 5 or 5 followed by non-zero digits, round up i.e. last number retained is inc. by 1 11

d) Ex 6: Subtract 2.78 from 3.18 3.18 & 2.78 0.40 12

e) Ex 7: Find diff. between 12.3 & 1.45

f) Ex 8: Round each of the following to 2 sig. fig.

12.3 & 1.45 10.85 1) Rounding Rule 3 If number to be discarded is 5, or 5 followed by zeros,

1.45

A

1.550

A

1.452

A

round even i.e. - leave last digit to be retained unchanged if even, increase by 1 if it is odd

ˆ

Answer is:

13

14

IV) Scientific Notation Express a number as a coefficient times a power of 10.

A

x 10

V) Units of Measurement - Metric System International System, SI units - Have base units from which all other units are derived

n

1 non-zero digit to left of decimal pt.

mass kg

400 = 4 x 102 4.0 x 10

4.00 x 102

or EXP

temp K

- employs factors of 10

Entering in calculators: EE

time s

Base units for length & mass are part of metric system

2

4

length m

Prefixes: indicate size of unit relative to base unit 2 Table 2.1

0.0000345 = 15

16

Selected SI Prefixes Prefix

Abbrev.

Meaning

Example

Mega-

M

10

1 megameter (Mm) = 1 x 106 m

Kilo-

k

103

1 kilometer (km)

= 1 x 103 m

Deci-

d

10-1

1 decimeter (dm)

= 0.1 m

c

-2

1 centimeter (cm)

= 0.01 m

-3

1 millimeter (mm) = 0.001 m

-6

10

1 micrometer (µm) = 1 x 10-6 m

CentiMilli-

m a

6

10

10

Micro-

µ

Nano-

n

10-9

1 nanometer (nm)

= 1 x 10-9 m

Pico-

p

10-12

1 picometer (pm)

= 1 x 10-12 m

Femto-

f

10-15

1 femtometer (fm) = 1 x 10-15 m

a

This is the Greek letter Mu (pronounced “mew”)

17

A) Length

1.0936 yd

meter, m

=

1 in / 2.54 cm

1m

2.205 lb

453.6 g

B) Mass kilogram, kg

= =

1 kg / 103 g

1 kg

1 lb

18

C) Volume SI unit is m3

VI) Factor Unit Method (Dimensional Analysis) Solve problems by carrying units throughout the calculations

metric system unit is liter, L 1 L / 1 dm3

- just converting units by using conversion factors

( 1 dm / 10 cm) ˆ 1 L = (10 cm)3 = 103 cm3

Conversion Factor

1 L / 103 mL ˆ

A number having two or more units associated with it

1 mL = 1 cm3

Numerically equivalent to 1 information same info in conv. given in one X = a different factor type of unit type of unit

19

20

A) Ex 1: A local donut shop sells donuts for $4.49 a dozen. You want 3 dozen donuts. How much will it cost?

B) Ex 2: Convert 0.34 cm to µm ? cm = 1 µm 1 cm / 10-2 m

1µm / 10-6 m

change units

|

dozen

1 dozen

/

dollars

$4.49

Can write 2 conv. factors 1 dozen $4.49

=1

$4.49 =1 1 dozen Note : Conversions within a system are exact by definition.

Convert 3 dozen to ? dollars :

21

22

C) Ex 3: An aspirin tablet weighs 5.00 grains. What is the weight of an aspirin tablet in ounces? (1 oz = 437.5 grains)

D) More Complicated Conversions 1) Ex 1: A good pitcher can throw a fastball at a speed of 90.0 mi/hr. How long will it take (in sec) to reach home plate 60.5 ft away? 60.5 ft Have

|

? sec

90.0 mi/hr

Must convert units in both numerator and denominator 1 mi / 5280 ft

23

1 hr / 3600 s

24

BLANK

25

2) Ex 2: A pool measures 60.500 ft by 30.500 ft by 10.0000 ft. How many cubic meters of water can the pool hold?

26

VII) Temperature Must specify temp. when making quantitative measurements A) Celsius Scale °C - commonly used Fahrenheit, °F, scale used in public (USA)

°F 212 98.6 32.0

°C 100.0 37.0 0.0

b.p. of H2O body temperature f.p. of H2O

1) Ex : Convert 25°C to °F

27

28

B) Kelvin Scale

VIII) Extensive & Intensive Properties

SI base unit is kelvin, K

A) 2 Catagories of Properties

Must be used in most cases in chemistry

1) Extensive Prop. Depends on amount of matter - mass & volume

Absolute scale: 2) Intensive Prop. Independent of sample size

0 K : lowest possible temp. )TK = )T°C

Characteristic of the substance

(unit same size)

a) melting & boiling points (temperature)

0 °C = 273.15 K

b) density

K = °C + 273.15

c) specific heat 29

30

B) Density

1) Ex 1: A cube of lead, Pb, is 3.00 cm on each side & has a mass of 305 g. What is its density?

Mass per unit volume

D =

D= m V

m V

SI unit is kg/m3 Solids g/cm3

Liquids g/mL

2) Ex 2: What is the mass of 50 mL (5.0 x 10 mL) of ether? The density of ether is 0.71 g/mL.

Gases g/L

Density can be used as a conversion factor between mass and volume

31

32

C) Specific Gravity

Sp. Gr. =

IX) Energy & Specific Heat

Dsubstance (g/mL)

Energy / capacity to do work

Dwater (g/mL)

Note : Energy is an extensive property

No units A) Law of Conservation of Energy

H2O :

D = 1.0 g/mL

Convert between diff. forms of energy, but total energy remains the same

Ethanol : D = 0.79 g/mL sp. gr. = 0.79

Energy lost by = one thing

33

Energy gained by something else

34

B) Units

C) Specific Heat

SI unit is joule, J

Amount of heat energy req. to change temperature of 1 g of substance by 1°C.

calorie (cal) / amt. of energy required to raise temp. of 1 g of water by 1°C from 14.5°C to 15.5°C

- conv. factor : 1

cal gC°C

H2O : 1.000

cal gC°C

Fe : 0.106

cal gC°C

1 cal / 4.184 J

Intensive property

Also, 1 Calorie (Cal) = 1 kcal 35

36

Relationship between heat added or removed and temperature change

Q = m

C c

1) Ex 1: How many calories are required to heat 1.0 kg of H2O from 20.0° to 100.0°C?

Q = m

C )t