1. Chemistry 100(02) Fall 2014 Instructor: Dr. Upali Siriwardane
Office: CTH 311
Phone
Office Hours: M,W, 8:00-9:30 & 11:30-12:30 a.m
Tu,Th,F 8: :00 a.m.   Or by appointment
Test Dates:
September 29, 2014 (Test 1): Chapter 1 & 2
October 20, (Test 2): Chapter 3 & 4
November 12, (Test 3) Chapter 5 & 6
November 13, (Make-up test) comprehensive: Chapters :30-10:45:15 AM, CTH 328

2. Text Book & Resources REQUIRED :
Textbook:  Principles of Chemistry: A Molecular Approach, 2nd Edition-Nivaldo J. Tro - Pearson Prentice Hall and also purchase the Mastering Chemistry
Group Homework, Slides and Exam review guides and sample exam questions are available online: and follow the course information links.
OPTIONAL :
Study Guide: Chemistry: A Molecular Approach, 2nd Edition-Nivaldo J. Tro 2nd Edition
Student Solutions Manual: Chemistry: A Molecular Approach, 2nd Edition-Nivaldo J. Tro 2nd

3. Chapter 1. Matter, Measurement, and Problem Solving
1. 1 Atoms and Molecules…………………………………
1 .2 The Scientific Approach to Knowledge……………
1 .3 The Classification of Matter……………………………
1 .4 Physical and Chemical Changes and Physical and Chemical Properties……………………………………
1 .5 Energy: A Fundamental Part of Physical and Chemical Change……………………………………………………
1 .6 The Units of Measurement……………………………
1 .7 The Reliability of a Measurement………………………
1 .8 Solving Chemical Problems……………………………

4. Chapter 1. KEY CONCEPTS What is chemistry? Scientific Method.
Properties of the three states of matter
Physical changes and properties.
Chemical change and properties.
Categories of matter.
Elements and Compounds
Atomic symbols
Chemical Elements and properties
Chemical Symbolism
Separating Mixtures.
Scientific Measurement
Prefixes of SI units
Macro, micro and nano-scales
Conversion factors.
Factor label method.
Uncertainty and significant figures
Temperature Conversions.
Density Calculations.
Three chemical Laws
Dalton's atomic theory
Interpreting chemical formulas and chemical reaction.
Concept of mole
Gram to mole conversion

5. Units

6. SI UNITS OF MEASUREMENT
Five Basic Units
Length meter (m)
Mass kilogram (kg)
Time second (s)
Temperature kelvin (K)
Amount mole (mol)
6.02 x 1023 units

7. The Units of Measurement
1) Give the name and abbreviation of the SI Unit for:
Length b) Mass c) Time
d) Amount of substance e) Temperature

8. Metric System How to change measurements to reasonable numbers Prefix
Symbol
Decimal
Equivalent
Power of 10
mega- M
1,000,000
Base x 106
kilo- k
1,000
Base x 103
deci- d
0.1
Base x 10−1
centi- c
0.01
Base x 10−2
milli- m
0.001
Base x 10−3
micro-
m or mc
Base x 10−6
nano- n
Base x 10−9
pico p
Base x 10−12

9. Macroscale, Microscale, and Nanoscale

10. 2) Give the abbreviation for the following units and describe what they are used to measure:
cubic centimeter b) micrometer
c) nanoseconds d) millimole

11. 3) Give the name and the abbreviation (without looking in the book) of the SI or metric prefix for:
b) c) d) 10-2
e) f) g) i) 10-6

12. Uncertainty in a Measurement
types of errors, random and systematic.
Careless measurements
Low resolution instruments
Calibration errors
The last digit is an estimate. .
Measurement ≈ cm

13. Uncertainty and Significant Figures
All measurements involve some uncertainty.
Scientists write down all the digits that have no uncertainty plus one additional uncertain digit.
If an object is reported to have a mass = g, the last digit (“2”) is uncertain ( it is probably close to 2, but may be 4, 1 … etc).
There are five significant figures in this number. All the digits are meaningful.

14. Uncertainty and Significant Figures
To find the number of significant figures:
Read a number from left to right and count all digits, starting with the first non-zero digit.
All digits are significant except those zeros that are used to position a decimal point (“placeholders”).
5 sig. figs.
Scientific Notation ( x 10-4)
significant
placeholders
significant

15. Significant Figure Primer
All non-zero digits are significant.
Example: 536 has three significant figures.
Zeros between non-zero digits are also significant.
Example: 6703 has four significant figures.
Place holder zeros:
Zeros to the left of a non-zero digit are NOT significant.
Example: has two significant figures.
Example: has three significant figures.
Zeros to the right or after a non-zero digit to the decimal point are NOT significant.
Example: 7000 has only one significant figure.
Example: has four significant figures.
Example: 50.0 has three significant figures; all the zeros are significant.

16. Uncertainty and Significant Figures
Examples
Number Sig. figs. Comment
2.12 3
The zeros are not placeholders. They are significant.
The zeros are placeholders (not significant).
Only the last two zeros are significant.
, 2, 3 ? Ambiguous. If a number lacks a decimal point the zeros may be placeholders or may be significant.
Adding a decimal point is one way to show that the zeros are significant.
5.0 x No ambiguity.

17. The Reliability of a Measurement
4) Express the following numbers in scientific notation with the appropriate number of significant figures: b)
c) d) x 107

18. Rounding is > 5, round the last retained digit up by 1.
Look at the 1st non-significant digit (the digit after the last one retained). If it:
is > 5, round the last retained digit up by 1.
is < 5, make no change.
equals 5, and the 2nd non-significant digit is:
absent, round the last retained digit up by 1.
odd, round the last retained digit up by 1.
even, make no change.
Consider rounding to 3 significant figures.
last retained digit
2nd non-significant digit
1st non-significant digit
It rounds up to 37.7

19. Rounding
Examples
Round the following numbers to 3 significant figures:
1st non-sig. 2nd non-sig. Rounded
Number digit digit Number

20. Can you define accuracy vs. precision?
Precision = the degree of exactness of a measurement that is repeatedly recorded.
Accuracy = the extent to which a measured value agrees with a standard value
Which set is more precise?
18.2 , 18.4 , 18.35
17.9 , 18.3 , 18.85
16.8 , 17.2 , 19.44

21. Can you hit the bull's-eye?
Three targets with three arrows each to shoot.
Both accurate and precise
Precise but not accurate
Neither accurate nor precise
How do they compare?

22. Unit Conversion Calculation
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).

23. Using Conversion Factors 5) Convert 78
Using Conversion Factors 5) Convert inches into: a) feet b) meters 6) Convert meters into: a) kilometers b) micrometers

24. Significant Figures and Scientific Notation
Example: 301 in scientific notation is 3.01 x NOTE: The decimal point was moved two places to the left. Example: in scientific notation is 3.01 x NOTE: The decimal point was moved two places to the right. Both of these value indicate THREE significant figures.

25. Significant Figures in Add/Sub
The answer you report in a problem should only include significant digits.
Addition and subtraction
Find the number of digits after the decimal point (adp) in each number.
answer adp = smallest input adp.
Example
Add:
17.245
adp = 3
adp = 4
Rounds to: (adp = 3)

26. Addition and Subtractions Examples:
Examples: in in
in but you would report in cm cm
cm but you would report cm

27. Significant Figures Add/Sub
Example
Subtract 6.72 x 10-1 from 5.00 x 101
Write the numbers down with the same power of 10:
x 101
– x 101
x 101
Rounds to: 4.93 x adp = 2
adp = 2
adp = 4

28. Significant Figures Mul/Div
Multiplication and Division
Find the number of significant figures (sig. fig.) in each number. Answer has sig. fig = smallest input sig. fig.
Example
Multiply and
17.245 x
Rounds to: sig. fig. = 4
Multiply x 12.1 x
Rounds to: x 104 (3 sig. fig.)
sig. fig. = 5
sig. fig. = 4
= 14,

29. 7) Perform the following calculations and give the answer in the correct number of significant figures.
Answer: Scientific notation:
b) Answer: Scientific notation:
c) x Answer: Scientific notation:

30. 7) Perform the following calculations and give the answer in the correct number of significant figures.
2 adp 2 adp adp
Answer: Scientific notation: x 101
2 adp 5 adp adp
b) Answer: Scientific notation: x 101
0 adp 8 adp adp
c) x Answer: Scientific notation: x 102

31. 7) Perform the following calculations and give the answer in the correct number of significant figures.
d) Answer: Scientific notation:
e) 1.43 x Answer: Scientific notation:
f) (7.601x107) x (8.09x10-4) Answer
Scientific notation:

32. 7) Perform the following calculations and give the answer in the correct number of significant figures.
2 adp 5 adp adp
d) Answer: Scientific notation: x 101
Multiplication
3 sfg 3 sfg sfg
e) 1.43 x Answer: Scientific notation: : x 100 = 1.21
4 sfg sfg sfg
f) (7.601x107) x (8.09x10-4) Answer =615
Scientific notation: x 102

33. Temperature Scales: Comparison

34. Temperature Conversions
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
Shift the scale to zero
Convert the scale 100/180
Shift the scale to zero K

35. Temperature Convesions
oF -- > oC ; C = 5/9 (F - 32)
oC -- > oF ; F =9/5 C + 32
oC -- > K ; K = C
8) Convert 98.6 °F into:
a) °C b) K.

36. Problem Solving by Factor Label Method
State question in mathematical form
Set equal to piece of data specific to the problem
Use conversion factors to convert units of data specific to problem to units sought in answer
Other names used Unit Conversion Method or dimensional (Unit) Analysis

37. Common Conversion Factors
Length 1 kilometer = 1000 m = mile
1 inch = 2.54 cm (exactly)
1 angstrom (Å) = 1 x m
Volume 1 liter (L) = 1 x 10-3 m3
= 1000 cm3 = 1000 mL
= quarts
1 gallon = 4 quarts = 8 pints
Mass 1 amu = x g
1 pound = g = 16 ounces
1 ton (metric) = 1000 kg
1 ton (US) = 2000 pounds

38. Solving Chemical Problems
9) Convert 75 miles per hour into: m s-1.

39. Solving Chemical Problems
10) Convert 100 m2 into cm2
11) Convert 1 m3 into cm3 .

40. Significant Figures and Mathematical Operations
Mathematical operations dictate the reporting of significant figures in an answer.
Multiplication and Division
The least precise measured value determines the number of significant figures in the reported answers.
Addition and Subtraction
The value with the smallest decimal measurement determines the answer’s significant figure.

41. Solving Chemical Problems 13) Perform the following mathematical operations and give the answer with the correct number of significant figures : a)
(2.481 x 12.74) =
2.69 b)
(4.73 x 10-4) - (72.85)
(872.3) - (0.305)

42. Density Density = mass (g) volume (cm3) An INTENSIVE physical property
The physical property does not depend on amount of substance.
The physical properties of mass and volume that determine a substance’s density are EXTENSIVE.
Extensive physical properties are dependent on amount.
Densities of liquid and gases are affected by temperature.

43. Density Calculations PROBLEM:
Mercury (Hg) has a density of 13.6 g/cm3. What is the mass of 95 mL of Hg in grams? In pounds?
Strategy:
1. Convert mL to cm3.
2. Solve for mass (in grams) using density relationship.
3. Convert grams to pounds.

44. A Density Calculation PROBLEM:
Mercury (Hg) has a density of 13.6 g/cm3. What is the mass of 95 mL of Hg in grams? In pounds?
Density = mass (g)
volume (cm3)
Step 1: 95 mL x (1cm3/1mL) = 95 cm3
Step 2: g/cm3 = x / 95 cm3
x = 1.29 x 103 g, but report 1.3 x 103 g
Step 3: 1.3 x 103 g x (1 lb/454 g) = 2.9 lb

45. Solving Chemical Problems
12) Aluminum block weighs 14.2 g and has a density of
2.70 g cm-3. Calculate the volume of the block.

46. Problem:
The density of octane (C8H18) is 7.00 lb/gal.
a) What is density in mg/cm3? b) What is the mass in grams of 1.25 liters of octane?
Strategy:
1. Convert 7.00 lbs to mg.
2. Change gallons to cm3.
3. Determine the density of octane in mg/cm3.
4. Convert 1.25 L to mL.
5. Determine the mass of octane in 1.25 L using density.