1. Chemistry 100(02) Fall 2001 Dr. Upali Siriwardane
CTH 311 Phone
Office Hours: 8:00-9:00, 11:00-12:00 M, W
Tu, Th, F 10:00-12:00 a.m.
Test 1 : Chapters 1, 2: September 26
Test 2: Chapters 3, 4: October 31
Test 3: Chapters 5, 6: November 14
Make-up, Comprehensive, November 15

2. Chemistry 100(04) Fall 2001 Dr. Upali Siriwardane
CTH 311 Phone
Office Hours: M, Tu, W, Th, F 9:00-11:00 a.m.
Test 1 : Chapters 1, 2: October 2.
Test 2: Chapters 3, 4: October 30
Test 3: Chapters 5, 6: November 13
Make-up, Comprehensive, November 15

3. Chemistry 100(05) Fall 2001 Dr. Upali Siriwardane
CTH 311 Phone
Office Hours: M, Tu, W, Th, F 9:00-11:00 a.m.
Test 1 : Chapters 1, 2: October 2.
Test 2: Chapters 3, 4: October 30
Test 3: Chapters 5, 6: November 13
Make-up, Comprehensive, November 15

4. KEY CONCEPTS What is chemistry? Physical & chemical changes.
Physical & chemical properties.
Categories of matter Separating Mixtures. Scientific Method Scientific Measurement Observation Uncertainty.
Significant figure Precision Accuracy Significant figures in calculations Unit Conversions Temperature Conversions Unit conversion method.
Density Calculations.

5. What is chemistry?
Chemistry deals with non-reversible changes of matter.
Chemistry explains using atoms and molecules.
Chemical Concepts and Models improve your problem solving skills
Chemistry is a Central Science

7. Chemistry “The study of matter and the changes it undergoes.”
Major divisions
Inorganic Compounds of elements other than carbon
Organic Compounds of carbon
Biochemistry Compounds of living matter
Physical Theory and concepts
Analytical Methods of analysis

8. What is Matter Matter: Anything that has a mass and volume. Energy:
Manifestations of matter.
Matter and Energy is intertwined.

9. Classification of matter
Pure Substance
Mixture
Element
Compound
Homogeneous
Heterogeneous
Iron
Hemoglobin
Plasma
Blood

10. Hierarchy of Matter Mixtures Heterogeneous Homogenous Pure Substances
Compounds
Elements
Atoms
Nucleus
Electrons
Neutrons
Protons

11. Mixtures A combination of two or more pure substances.
Homogeneous - Uniform composition
Heterogeneous - Non-uniform composition
Which are homogeneous or heterogeneous?
Blood Urine “T-Bone” steak
Gasoline Twinkie Salad Dressing

12. How do you Separate Mixtures?
Flotation: based on density
Filtration: Solid- liquid
Distillation- Liquid-liquid
Magnetic Separation- Magnetic-
Chromatography:
1) Paper 2) Column 3) Gas

13. Pure substances Element Compound
Cannot be converted to a simpler form by a chemical reaction.
Example hydrogen and oxygen
Compound
Combination of two or more elements in a definite, reproducible way.
Example water - H2O
Both elements and compounds have characteristic properties such as color, boiling point and reactivity

14. Pure substances
The properties of a compound and the elements it is made of can differ greatly.
Formula BP density Other
Hydrogen H Flammable
Oxygen O Supports combustion
Water H2O Not flammable

15. Properties of Substances
Physical properties:
Physical properties are descriptions of matter such as color, density, viscosity, boiling point, and melting point.
Chemical properties:
Chemical properties relates to the changes of substances making up the matter. For example, corrosiveness, Flammability

16. Extensive and intensive properties
Extensive properties
Depend on the quantity of sample measured.
Example - mass and volume of a sample.
Intensive properties
Independent of the sample size.
Properties that are often characteristic of the substance being measured.
Examples - density, melting and boiling points.

17. Physical properties
Properties that do not involve substances changing into another substance.
Examples
color density
odor melting point
taste boiling point
feel compressibility

18. Chemical properties
Properties that involve substances changing into another substance.
Chemical Reaction - one or more substances are changed into other substances.
Example A chemical property of wood is it’s ability to burn - combustion.
wood + oxygen carbon dioxide + water + heat
Reactants Products
The reactants and products are very different.

19. Example Which are chemical or physical changes? Mulching leaves
Milk turning sour
Making wine
Making ice water
Beer going flat
Leaves changing color

20. Type of Changes Physical change:
A change in the state of matter. It does not involve a change in the substances. E.g. melting of wax and water.
Chemical change:
A change involving at least one of the substances making the matter. E.g. Electrolysis of water, formation of rust: reaction of iron and oxygen to from iron oxide.

21. Chemical verses Physical change
Sodium reacting Iodine changing
with chlorine from a solid to a gas

22. Scientific method
All scientific studies follow the same approach to examining a problem.
The scientific method requires that we:
Make observations
Apply logical, organized reasoning to observations made.
Form a hypothesis.
Reject or confirm that hypothesis through experiments.

23. Scientific Method. A method common to all sciences has
Four Basic Steps:
a) Experiment
b) Data or Results
c) Hypothesis
d) Further experiments to test hypothesis

24. Scientific method Try new tests No Make observations Organize
Make hypothesis
Do experiments
Try
new
tests No
Did hypothesis work?
Yes
Develop a theory
Do more experiments

26. Measurement Measurements or observations are made
using our physical senses or using scientific instruments.
1) Qualitative measurements.
Changes that cannot be expressed in terms of a number.
2) Quantitative measurements.
expressed in terms of a number and an unit.

29. Units are important 45 000 has little meaning, just a number
$45,000 has some meaning - money
$45,000/yr more meaning - person’s salary

30. SI units SI - System International Only uses certain metric units.
Systematic subset of the metric system.
Only uses certain metric units.
Mass kilograms
Length meters
Time seconds
Temperature kelvin
Amount mole
Other SI units are derived from SI base units.

31. Metric prefixes Changing the prefix alters the size of a unit.
Prefix Symbol Factor
mega M
kilo k
hecto h
deka da
base
deci d
centi c
milli m

33. Other Units
Derived Units. Units consisting of more than one one base unit. E.g. g/cm3
English units.
Still commonly used in the United States. Weight ounce, pound, ton Length inch, foot, yard, mile Volume cup, pint, quart, gallon
Not often used in scientific work.
Very confusing and difficult to keep track of the conversions needed.

34. Converting units Factor label method
Regardless of conversion, keeping track of units makes thing come out right
Must use conversion factors
- The relationship between two units
Canceling out units is a way of checking that your calculation is set up right!
Other names used Unit Conversion Method dimensional(Unit) Analysis

35. Example. Metric conversion
How many milligrams are in a kilogram?
1 kg = 1000 g
1 g = 1000 mg
1 kg x x 1000
= mg kg g mg g

36. ( ) ( ) Example  = 10-6 = micro
Creatinine is a substance found in blood. If an analysis of blood serum sample detected 0.58 mg of creatinine, how many micrograms were present?
 = 10-6 = micro
10-3 g
1 mg
( )
1 g
10-6 g
( )
0.580 mg
= 580 g

37. Common conversion factors
English Factor
1 gallon = 4 quarts qt/gal
1 mile = 5280 feet ft/mile
1 ton = 2000 pounds lb/ton
Common English to Metric conversions Factor
1 liter = quarts qt/L
1 kilogram = 2.2 pounds 2.2 lb/kg
1 meter = yards yd/m
1 inch = 2.54 cm cm/inch

38. .Speed of light is 3.00 x 108 m s-1 . Convert the speed of light to miles per year (1 mile = 1.61 km).

39. Measurement Number Part Exact Measurements
Uncertainty in Measurement
Significant Figures
Exact Measurements
Extensive and Intensive Properties
Density
Measuring Temperature and Volume

41. Uncertainty in Measurement
All measurements contain some uncertainty.
We make errors
Tools have limits
Uncertainty is measured with
Accuracy How close to the true value
Precision How close to each other

42. Accuracy How close our values agree with the true value. Here the
average value
would give a
good number
but the numbers
don’t agree.
Large random error

43. Precision How well our values agree with each other. Here the numbers
are close together
so we have good
precision.
Poor accuracy.
Large systematic
error.

44. Accuracy and precision
Predict the effect on accuracy and precision.
Instrument not ‘zeroed’ properly
Reagents made at wrong concentration
Temperature in room varies ‘wildly’
Person running test is not properly trained

45. Types of errors Systematic Random Instrument not ‘zeroed’ properly
Reagents made at wrong concentration
Temperature in room varies ‘wildly’
Person running test is not properly trained

46. Significant figures Method used to express accuracy and precision.
You can’t report numbers better than the method used to measure them.
67.2 units = three significant figures
Certain
Digits
Uncertain
Digit

47. Significant Not significant Non-zero digits are always significant.
Any zeros between two significant digits
Trailing zeros in the decimal portion
Not significant
Leading zeros
Trailing zeros in whole numbers (use scientific notion to avoid confusion.
Exact numbers: unit definition has an unlimited number of sig. figs. 1 ft = 12 in

48. Examples 0.00341........3 sig. digs. 1.0040.........5 sig. digs.
… 2 sig. digs. 6.5 x 104
sig. digs.
sig. digs x 105

49. Significant figures: Rules for zeros
Leading zeros are not significant.
three significant figures
Leading zero
Captive zeros are significant.
four significant figures
Captive zero
Trailing zeros are significant.
five significant figures
Trailing zero

50. Significant figures Zeros are what will give you a headache!
They are used/misused all of the time.
Example
The press might report that the federal deficit is three trillion dollars. What did they mean?
$3 x 1012 or
$3,000,000,000,000.00

51. Significant figures
In science, all of our numbers are either measured or exact.
Exact - Infinite number of significant figures.
Measured - the tool used will tell you the level of significance. Varies based on the tool.
Example
Ruler with lines at 1/16” intervals.
A balance might be able to measure to the nearest 0.1 grams.

52. Scientific notation
Most calculators use scientific notation when the numbers get very large or small.
How scientific notation is displayed can vary.
It may use x10n
or may be displayed
using an E.
They usually have an Exp or EE
This is to enter in the exponent.
E-2

53. Examples
3.78 x 10 5
8931.5
x 10 3
5.93 x
4 x

54. Significant figures and calculations
An answer can’t have more significant figures than the quantities used to produce it.
Example
How fast did you run if you
went 1.0 km in 3.0 minutes?
speed = 1.0 km / 3.0 min
= 0.33 km / min

55. Significant figure in Calculations
Different rules apply in each case
Addition and subtraction
In multiplication and division
The root or power of a measurement
Exact Numbers: Numbers coming from definitions such as 12 in = 1 foot. They are not considered in Sig. Fig. Calculations.

56. Significant figures and calculations
Addition and subtraction
Report your answer with the same number of digits to the right of the decimal point as the number having the fewest to start with. g g g
805.4 g g
83.7 g

57. Significant figures and calculations
Multiplication and division.
Report your answer with the same number of digits as the quantity have the smallest number of significant figures.
Example. Density of a rectangular solid.
25.12 kg / [ (18.5 m) ( m) (2.1m) ]
= 2.8 kg / m3
(2.1 m - only has two significant figures)

58. Rounding off numbers After calculations, you may need to round off.
If the first insignificant digit is 5 or more,
- you round up
If the first insignificant digit is 4 or less,
- you round down.
If the digit 5 exactly rounded off to a even

60. Temperature Conversions
oF -- > oC ; C = 5/9 (F - 32)
oC -- > oF ; F =9/5 C + 32
oC -- > K ; K = C
Human body temperature is 98.6 oF. Convert this temperature to oC and K scale
oC = 5/9 ( ) = 5/9 (66.6) = 37.0
oC--> K = 37.0 oC = K

61. Example. oF to K If the temperature is 75.0 oF, what is it in K?
First convert to oC
Then convert to K 5
oC = (75.0oF - 32) 9
= 23.9
K = 23.9oC
= 297

62. Measuring volume Volume - the amount of space that an object occupies.
The base metric unit is the liter (L).
The common unit used in the lab is the milliliter (mL).
One milliliter is exactly equal to one cm3.
The derived SI unit for volume is the m3 which is too large for convenient use.

63. Density
Density is an intensive property of a substance based on two extensive properties.
Common units are g / cm3 or g / mL.
g / cm3 g / cm3
Air Bone
Water Urine
Gold Gasoline
Density =
Mass
Volume
cm3 = mL

64. Example. Density calculation
What is the density of 5.00 mL of a fluid if it
has a mass of 5.23 grams?
d = mass / volume
d = g / 5.00 mL
d = g / mL
What would be the mass of 1.00 liters of this
sample?

65. Density Calculations Equation method: Density = mass ÷ volume; d = m/v
Factor Label method:14.2 g -- > ? cm3
conversion factor?
2.70 g 1 cm3 or
1 cm g
14.2 g x 1 cm3
= cm3
2.70 g

66. Example. Density calculation
What would be the mass of 1.00 liters of the fluid sample?
The density was 1.05 g/mL.
density = mass / volume
so mass = volume x density
mass = 1.00 L x x 1.05
= 1.05 x 103 g ml L g mL

67. Specific gravity
The density of a substance compared to a reference substance.
Specific Gravity =
Specific Gravity is unitless.
Reference is commonly water at 4oC.
At 4oC, density = specific gravity.
Commonly used to test urine.
density of substance
density of reference

68. Specific gravity measurement
Hydrometer
Float height will
be based on
Specific Gravity.