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- Basic Concepts

Chapter Outline

- States of Matter
- Chemical and Physical Properties
- Chemical and Physical Changes
- Mixtures, Substances, Compounds, and

Elements - Measurements in Chemistry
- Units of Measurement

Chapter Outline

- Use of Numbers
- The Unit Factor Method (Dimensional Analysis)
- Density and Specific Gravity
- Heat and Temperature
- Heat Transfer and the Measurement of Heat

Mixtures, Substances, Compounds, and Elements

- Matter-Anything that occupies space and has mass.
- Pure Substances or Substances-Cannot be separated

by physical processes. - -Elements-A substance which cannot be broken

down into simpler substances. e.g. Na, He, C,

(atoms) or N2, Cl2 (molecules) - -Compounds-A pure substance made up of two or

more elements. e.g. NaCl, H2O. - Mixtures-Can be separated by physical processes.
- composed of two or more substances
- homogeneous mixtures-A mixture that is uniform

throughout-e.g. white wine, grape juice. Clear.

Solutions. - heterogeneous mixtures-A mixture that is not

uniform throughout-e.g. oil and water, orange

juice. Cloudy.

Mixtures, Substances, Compounds, and Elements

Substances and Mixtures

Matter

Physical process

Mixtures

Pure Substances

Homogeneous Mixture

Chemical Reaction

Heterogeneous Mixture

Compounds

Elements

Fig. 1-7, p. 10

States of Matter

- Change States
- heating
- cooling

Vaporization Evaporation Boiling Freezing Soli

dification Crystallization Melting Fusion

States of Matter

Solid

Liquid

Gas

heat

heat

cool

cool

Less attractive force and more disordered.

States of Matter

- Illustration of changes in state
- requires energy

Mixtures, Substances, Compounds, and Elements

Elements you Need to Know

Types of Solutions

- Liquid solutions are the most common, but there

are also gas and solid solutions. - Solutions have two components
- Solute - Solution component(s) present in lesser

amounts. - Solvent - Solution component present in the

greatest amount.

Types of Solutions

Characteristics of Solutions

- Uniform distribution
- Components do not separate upon standing.
- Components cannot be separated by filtration.
- Within certain limits its composition can vary.
- Almost always transparent. (i.e. one can see

through it). - An alloy is a homogeneous mixture of metals. i.e.

brass, bronze, sterling silver.

Chemical and Physical Properties

- Physical Properties A property that can be

observed in the absence of any change in

composition. e.g. color, odor, taste, melting

point, boiling point, freezing point, density,

length, specific heat, density, solubility. - Physical Changes-Changes observed without a

change in composition. i.e. cutting wood, melting

of solids and boiling of liquids. - water, ice water, liquid

water, steam - changes of state
- dissolving
- polishing

Chemical and Physical Properties

- Chemical Properties-A property that matter

exhibits as it undergoes changes in composition.

e.g. coal and gasoline can burn in air to form

carbon dioxide and water iron can react with

oxygen in the air to form rust bleach can turn

blond hair blonde. - Chemical Changes-Changes observed only when a

change in composition is occurring. e.g. reaction

of sodium with chlorine, rusting of iron, dying

of hair, burning of wood, cooking an egg, rotting

food. - Extensive Properties - depends on the amount of

material present. e.g. volume and mass. - Intensive Properties does not depend on the

amount of material present. e.g. melting point,

boiling point, freezing point, color, density.

Natural Laws

- Law of Conservation of Mass-Mass is neither

created nor destroyed. - Law of Conservation of Energy-Energy is neither

created nor destroyed, only converted from one

form to another. - Law of Definite Proportions-Different samples of

any pure compound contain the same element in the

same proportion by mass. e.g. water (H2O)

contains 11.1 H and 88.9 O by mass. Thus, a

25.0 sample of water would contain 2.78 g of H

and 22.2 g of O.

Law of Definite Proportions

11.1 H and 88.9 O by mass, 25.0 g sample of

water

Use of Numbers

- Exact numbers
- 1 dozen 12 things for example

Rounding off Numbers

Previous digit

1.29

4

Next digit

- If the next digit is less than 5 the previous

.9946 .99 - digit remains the same.

1.294 1.29 -

- 2. If the next digit is greater than 5 or 5

- followed by non zeros then the previous digit

.999 1.00 - is increased by one.

1.2951 1.30 - 3. If the next digit is 5 or 5 followed by all

zeros 1.285 1.28 - then the previous digit remains the same if it

1.295 1.30 - is even or increased by one if it is odd.

1.22500 1.22

Scientific Notation

- Used to handle very large and very small numbers.

- Any number that is from or 1 to 9
- N. X 10x
- For example 3.21 x 103
- -9.9 x

10-4 - 1.0 x 100 (Note that

100 is 1)

Exponent-Power of 10

Scientific Notation

- To convert numbers to scientific notation use the

following guidelines - 1750.0 1750.0 x 100 1.7500 x 103

Exponent increases by 3 powers of 10

Number decreases by 3 powers of 10

A you move the decimal place to the left (i.e.

make the number smaller), the power of ten

(i.e., exponent) must increase by the same amount.

Scientific Notation

- 0.050 0.050 x 100 5.0 x 10-2

The number gets larger by 2 powers of 10

The exponent gets smaller by 2 powers of 10.

As you move the decimal place to the right (i.e.

make the number larger), the power of ten (i.e.

exponent) must decrease by the same amount.

Use of Numbers

- Significant figures
- digits believed to be correct by the person

making the measurement - Measure a mile with a 6 inch ruler vs. surveying

equipment - Exact numbers have an infinite number of

significant figures - 12.000000000000000 1 dozen
- because it is an exact number

Use of Numbers

- Significant Figures - Rules
- Leading zeroes are never significant
- 0.000357 has three significant figures
- Trailing zeroes may be significant
- must specify significance by how the number is

written - Use scientific notation to remove doubt
- 2.40 x 103 has ? significant figures

Use of Numbers

- 3,380 ? significant figures
- 3.38 x 103
- 3,380. has ? significant figures
- 3.380 x 103
- Imbedded zeroes are always significant
- 3.0604 has ? significant figures

Use of Numbers

- Piece of Paper Side B enlarged
- How long is the paper to the best of your ability

to measure it?

13.36 in.

The second decimal place is estimated

Use of Numbers

- Piece of Paper Side A enlarged
- How wide is the paper to the best of your ability

to measure it?

8.3 in

The first decimal place is estimated

Manipulating Powers of 10

- a) When multiplying powers of ten, the exponents

are added. For example - 105 x 10-4 105(-4)101
- b) When dividing powers of ten, the

exponents are subtracted. For example - 104 104-(-4) 108
- 10-4
- c) When raising powers of ten to an exponent,

the exponents are multiplied. For example - (104)3 10(4 x 3) 1012

Use of Numbers

- Multiplication Division rule
- Easier of the two rules
- Product has the smallest number of significant

figures of multipliers

Use of Numbers

- Multiplication Division rule
- Easier of the two rules
- Product has the smallest number of significant

figures of multipliers

Use of Numbers

- Multiplication Division rule
- Easier of the two rules
- Product has the smallest number of significant

figures of multipliers

Multiplying and Dividing Numbers with Powers of

Ten

- When using scientific notation
- a.) Place the powers of ten together.
- (1.76 x 10200) x (2.650 x 10200)
- (1.76 x 2.650) x (10200

200) - b.) The final answer has the same number of

significant figures as the number with the least

number of significant figures. - 4.66 x 10400
- c.) You must round off correctly.
- d.) Preferably report the answer in scientific

notation.

Multiplying and Dividing Numbers with Powers of

Ten

- (1.760 x 102) /(2.65 x 10-2)
- (1.760 / 2.65) x (102

(-2)) - 0.664 x 104 6.64

x 103

Use of Numbers

- Addition Subtraction rule
- More subtle than the multiplication rule
- Answer contains smallest decimal place of the

addends

Use of Numbers

- Addition Subtraction rule
- More subtle than the multiplication rule
- Answer contains smallest decimal place of the

addends

Addition and Subtraction with Powers of Ten

- a.) All numbers must have the same power or ten

before addition or subtraction is performed. - b.) Once the powers of ten are the same, the

coefficients can then be added or subtracted

while the power of ten remains the same. - c.) After adding or subtracting the coefficients,

the answer must have the same number of decimal

places as the coefficient with the fewest decimal

places at the time of the operation. - d.) You must round off correctly.
- e.) Preferably report the answer in scientific

notation.

Addition and Subtraction with Powers of Ten

- 4.76 x 10200 9.6 x 10201 ?
- 0.4 76 x 10201
- 9.6 x 10201
- 10.0 76 x 10201 1.01 x 10202
- (written in scientific notation and rounded off

to the correct number of significant figures)

Addition and Subtraction with Powers of Ten

- 2.95 x 10-15 1.00 x 10-14 ?
- -1.00 x 10-14
- 0.29 5 x 10-14
- -0.70 5 x 10-14 -7.0 x 10-15
- (written in scientific notation and rounded off

to the correct number of significant figures)

Mixing Addition/Subtraction with

Multiplication/Division

- 7.54 x 10-5 (99. x 10200 1.25 x 10201)
- (1.75 x 10-3)3
- 7.54 x 10-5 (9.9 x 10201 1.25 x 10201)
- 1.75 x 10-3 x 1.75 x 10-3 x 1.75 x 10-3
- 7.54 x 10-5 (9.9 1.25) x 10201)
- 1.75 x 1.75 x 1.75 x 10-3 x 10-3 x 10-3
- 7.54 x 10-5 (11.2 x 10201)
- 5.36 x 10-9
- 7.54 x 11.2 x 10-5 x 10201 1.58 x 10206
- 5.36 10-9

Measurements in Chemistry

- Quantity Unit Symbol
- length meter m
- mass kilogram kg
- time second s
- current ampere A
- temperature Kelvin K
- amt. substance mole mol

Measurements in ChemistryMetric Prefixes

- Name Symbol Multiplier
- mega M 106
- kilo k 103
- deka da 10
- deci d 10-1
- centi c 10-2

Measurements in ChemistryMetric Prefixes

- Name Symbol Multiplier
- milli m 10-3
- micro ? 10-6
- nano n 10-9
- pico p 10-12
- femto f 10-15

Metric Conversions

- 1 km 103 m
- 1 dL 10-1 L
- 1 msec 10-3 sec
- 1 ?m 10-6 m

Fig. 1-20, p. 24

Metric English Conversions

- Common Conversion Factors
- Length
- 2.54 cm 1 inch (exact conversion)
- Volume
- 1 qt 0.946 liter (Rounded off)
- Mass
- _ 1 lb 454 g (Rounded off)

Use of Conversion Factors in Calculations

- Commonly known relationship (i.e. equality)
- 1 ft 12 in
- Respective conversion factors to above equality
- 1 ft or 12 in
- 12 in 1 ft
- Use the conversion factor that allows for the

cancellation of units. Convert 24 in to ft - ? ft 24 in x

Conversion Factors

- Example 1-1 Express 9.32 yards in millimeters.

Conversion Factors

Conversion Factors

Conversion Factors

Conversion Factors

The Unit Factor Method

- Area is two dimensional thus units must be in

squared terms. - Example 1-3 Express 2.61 x 104 cm2 in ft2.

The Unit Factor Method

- Area is two dimensional thus units must be in

squared terms. - Example 1-3 Express 2.61 x 104 cm2 in ft2.
- common mistake

The Unit Factor Method

- Area is two dimensional thus units must be in

squared terms. - Example 1-3 Express 2.61 x 104 cm2 in ft2.

The Unit Factor Method

- Area is two dimensional thus units must be in

squared terms. - Example 1-3 Express 2.61 x 104 cm2 in ft2.

The Unit Factor Method

- Area is two dimensional thus units must be in

squared terms. - Example 1-3 Express 2.61 x 104 cm2 in ft2.

The Unit Factor Method

- Volume is three dimensional thus units must be in

cubic terms. - Example 1-4 Express 2.61 ft3 in cm3.

The Unit Factor Method

- Example 1-2. Express 627 milliliters in gallons.

Conversions of Double Units

Density and Specific Gravity

- density mass/volume
- What is density?
- Why does ice float in liquid water?

Density and Specific Gravity

- Example 1-6 Calculate the density of a substance

if 742 grams of it occupies 97.3 cm3.

Density and Specific Gravity

- Example 1-6 Calculate the density of a substance

if 742 grams of it occupies 97.3 cm3.

Density and Specific Gravity

- Example 1-7 Suppose you need 125 g of a corrosive

liquid for a reaction. What volume do you need? - liquids density 1.32 g/mL

Density and Specific Gravity

- Waters density is essentially 1.00 g/mL at room

T. - Thus the specific gravity of a substance is very

nearly equal to its density. - Specific gravity has no units.

Density and Specific Gravity

- The density of lead is 11.4 g/cm3. What volume,

in ft3, would be occupied by 10.0 g of lead?

Density and Specific Gravity

- What is the density (in g/mL) of a rectangular

bar of lead that weighs 173 g and has the

following dimensions - length 2.00 cm, w 3.00 cm, h 1.00 in?

Density and Specific Gravity

- An irregularly shaped piece of metal with a mass

of 0.251 lb was placed into a graduated cylinder

containing 50.00 mL of water this raised the

water level to 67.50 mL. What is the density of

the metal? What is the density (in g/cm3) of the

metal? Will the metal float or sink in water? - V(disp)l 67.50 mL 50.00 mL 17.50 mL

The metal will sink in water because its density

is greater than that of water. (1.00 g/mL)

Density and Specific Gravity

- Example1-8 A 31.0 gram piece of chromium is

dipped into a graduated cylinder that contains

5.00 mL of water. The water level rises to 9.32

mL. What is the specific gravity of chromium?

Density and Specific Gravity

- Example1-8 A 31.0 gram piece of chromium is

dipped into a graduated cylinder that contains

5.00 mL of water. The water level rises to 9.32

mL. What is the specific gravity of chromium?

Heat and Temperature

- Heat and Temperature are not the same thing
- T is a measure of the intensity of heat in a body
- 3 common temperature scales - all use water as a

reference

Heat and Temperature

- Heat and Temperature are not the same thing
- T is a measure of the intensity of heat in a body
- 3 common temperature scales - all use water as a

reference

Heat and Temperature

- MP water BP water
- Fahrenheit 32 oF 212 oF
- Celsius 0.0 oC 100 oC
- Kelvin 273 K 373 K

Relationships of the Three Temperature Scales

Relationships of the Three Temperature Scales

Relationships of the Three Temperature Scales

Heat and Temperature

- Example 1-10 Convert 211oF to degrees Celsius.

99.4

Heat and Temperature

- Example 1-11 Express 548 K in Celsius degrees.

Heat Transfer and the Measurement of Heat

- Chemical reactions and physical changes occur

with either the simultaneous evolution of heat

(exothermic process), or the absorption of heat

(endothermic process). - The amount of heat transferred is usually

expressed in calories (cal) or in the SI unit of

joules (J). - 1 cal 4.184 J
- Specific heat is defined as the amount of heat

necessary to raise the temperature of 1 g of

substance by 1o C. - Each substance has a specific heat, which is a

physical intensive property, like density and

melting point.

Heat Transfer and the Measurement of Heat

- From a knowledge of a substances specific heat,

the heat (q) that is absorbed or released in a

given process can be calculated by use of the

following equation - q s x m x DT
- q (heat energy) cal, kcal, J or

kJ - m (mass) g
- s (specific heat) cal
- g oC

(kcal, J, or kJ can be used in -

lieue of cal). - DT T2 T1 (change in temp-make DT a

positive ) oC

Heat Transfer and the Measurement of Heat

- Substances with large specific heats require more

heat to raise their temperature. - Water has one of the highest specific heats, 1.00

cal/goC. The high specific heat of water (which

constitutes 60 of our body weight) makes our

bodys task of maintaining a constant body

temperature of 37oC much easier. Thus, our body

has the ability to absorb or release considerable

amounts of energy with little change in

temperature.

Heat Transfer and the Measurement of Heat

Heat Transfer and the Measurement of Heat

- Calculate the amount of heat to raise T of 200.0

g of water from 10.0oC to 55.0oC.

You need to know that the specific heat for water

(swater) is 1.00 cal/goC

Heat Transfer and the Measurement of Heat

- Example 1-13 Calculate the amount of heat to

raise T of 200.0 g of Hg from 10.0oC to 55.0oC.

Specific heat for Hg is 0.138 J/g oC. - Requires 30.3 times more heat for water
- 4.184 is 30.3 times greater than 0.138

Heat Transfer and the Measurement of Heat

- If we add 450 cal of heat to 37 g of ethyl

alcohol (s0.59 cal/goC) at 20oC, what would its

final temperature be? - q m x s x DT
- 450 cal 37 g x 0.59 x DT
- DT 21o C

Since heat was added, the final temperature must

be greater than the initial temperature.

DTT2- T1 21oC T2 20oC T2

21oC 20oC 41oC