Basic Concepts - PowerPoint PPT Presentation

Loading...

PPT – Basic Concepts PowerPoint presentation | free to download - id: 25104-YmQ2M



Loading


The Adobe Flash plugin is needed to view this content

Get the plugin now

View by Category
About This Presentation
Title:

Basic Concepts

Description:

Metric Conversions. 1 km = 103 m. 1 dL = 10-1 L. 1 msec ... Metric English Conversions. Common Conversion Factors. Length. 2.54 cm = 1 inch (exact conversion) ... – PowerPoint PPT presentation

Number of Views:584
Avg rating:3.0/5.0
Slides: 87
Provided by: facul2
Learn more at: http://faculty.mdc.edu
Category:

less

Write a Comment
User Comments (0)
Transcript and Presenter's Notes

Title: Basic Concepts


1
1
  • Basic Concepts

2
Chapter Outline
  • States of Matter
  • Chemical and Physical Properties
  • Chemical and Physical Changes
  • Mixtures, Substances, Compounds, and
    Elements
  • Measurements in Chemistry
  • Units of Measurement

3
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

4
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.

5
Mixtures, Substances, Compounds, and Elements
6
Substances and Mixtures
Matter
Physical process
Mixtures
Pure Substances
Homogeneous Mixture
Chemical Reaction
Heterogeneous Mixture
Compounds
Elements
7
Fig. 1-7, p. 10
8
States of Matter
  • Change States
  • heating
  • cooling

Vaporization Evaporation Boiling Freezing Soli
dification Crystallization Melting Fusion
9
States of Matter
Solid
Liquid
Gas
heat
heat
cool
cool
Less attractive force and more disordered.
10
States of Matter
  • Illustration of changes in state
  • requires energy

11
Mixtures, Substances, Compounds, and Elements
12
Elements you Need to Know
13
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.

14
Types of Solutions
15
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.

16
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

17
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.

18
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.

19
Law of Definite Proportions
11.1 H and 88.9 O by mass, 25.0 g sample of
water
20
Use of Numbers
  • Exact numbers
  • 1 dozen 12 things for example

21
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

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
23
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.
24
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.
25
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

26
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

27
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

28
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
29
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
30
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

31
Use of Numbers
  • Multiplication Division rule
  • Easier of the two rules
  • Product has the smallest number of significant
    figures of multipliers

32
Use of Numbers
  • Multiplication Division rule
  • Easier of the two rules
  • Product has the smallest number of significant
    figures of multipliers

33
Use of Numbers
  • Multiplication Division rule
  • Easier of the two rules
  • Product has the smallest number of significant
    figures of multipliers

34
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.

35
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

36
Use of Numbers
  • Addition Subtraction rule
  • More subtle than the multiplication rule
  • Answer contains smallest decimal place of the
    addends

37
Use of Numbers
  • Addition Subtraction rule
  • More subtle than the multiplication rule
  • Answer contains smallest decimal place of the
    addends

38
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.

39
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)

40
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)

41
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

42
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

43
Measurements in ChemistryMetric Prefixes
  • Name Symbol Multiplier
  • mega M 106
  • kilo k 103
  • deka da 10
  • deci d 10-1
  • centi c 10-2

44
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

45
Metric Conversions
  • 1 km 103 m
  • 1 dL 10-1 L
  • 1 msec 10-3 sec
  • 1 ?m 10-6 m

46
Fig. 1-20, p. 24
47
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)

48
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

49
Conversion Factors
  • Example 1-1 Express 9.32 yards in millimeters.

50
Conversion Factors
51
Conversion Factors
52
Conversion Factors
53
Conversion Factors
54
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.

55
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

56
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.

57
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.

58
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.

59
The Unit Factor Method
  • Volume is three dimensional thus units must be in
    cubic terms.
  • Example 1-4 Express 2.61 ft3 in cm3.

60
The Unit Factor Method
  • Example 1-2. Express 627 milliliters in gallons.

61
Conversions of Double Units
62
Density and Specific Gravity
  • density mass/volume
  • What is density?
  • Why does ice float in liquid water?

63
Density and Specific Gravity
  • Example 1-6 Calculate the density of a substance
    if 742 grams of it occupies 97.3 cm3.

64
Density and Specific Gravity
  • Example 1-6 Calculate the density of a substance
    if 742 grams of it occupies 97.3 cm3.

65
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

66
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.

67
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?

68
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?

69
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)
70
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?

71
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?

72
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

73
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

74
Heat and Temperature
  • MP water BP water
  • Fahrenheit 32 oF 212 oF
  • Celsius 0.0 oC 100 oC
  • Kelvin 273 K 373 K

75
Relationships of the Three Temperature Scales
76
Relationships of the Three Temperature Scales
77
Relationships of the Three Temperature Scales
78
Heat and Temperature
  • Example 1-10 Convert 211oF to degrees Celsius.

99.4
79
Heat and Temperature
  • Example 1-11 Express 548 K in Celsius degrees.

80
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.

81
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

82
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.

83
Heat Transfer and the Measurement of Heat
84
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
85
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

86
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
About PowerShow.com