1. E-mail: amjaber@kfupm.edu.sa Office hours: S 9-11 am U 10-11 am
Professor Abdul Muttaleb Jaber
Chemistry Department
Office: Room # 261F
Tel: 2611
Office hours: S am
U am
M 9-11am

2. Chapter 1 Chemical Foundations
An overview
The scientific method
Units of measurements
Uncertainty in measurements
Significant figures and
calculations
Dimensional analysis
Temperature
Density
Classification of matter

3. 1.1 Chemistry: an overview
Matter is composed of atoms
Atoms are found as individuals or molecules
Atoms and molecules are connected by electrons
The challenge of chemistry is to think of the material the atomic level
100 different types of atoms form all substances in the world

4. Matter is composed of various types of atoms or molecules.
Water is composed of O and H; H2O
An electric spark causes a mixture of O2 and H2 to explode forming H2O.
One substance changes to another by
reorganizing the way atoms attached to each other

5. Scientific method It is a way of solving problems
It consists of the following steps:
Observation- what is seen or measured
Hypothesis- guess of why things behave the way they do. (possible explanation for an observation)
Experiment- designed to test hypothesis
These steps would lead to new observations, and the cycle goes on
Once a set of hypotheses agree with
observations, they are grouped into a theory

6. Scientific method
Thery is a set of tested hypothesis that gives an overall explanation for a natural phenomenon
Laws are summaries of observations
Often mathematical relationship

7. Examples: 20 grams 20 k g = 20 X103 g 20 m g = 20 X10-3 g
1.3 Units of measurements
Every measurement has two parts
Number
Scale (called a unit)
SI system (le Systeme International in French) based on the metric system
Examples:
20 grams
20 k g = 20 X103 g
20 m g = 20 X10-3 g
6.63   Joule seconds
Prefix

8. Metric System Fundamental SI base Units Mass - kilogram (kg)
Length- meter (m)
Time - second (s)
Temperature- Kelvin (K)
Electric current- ampere (amp, A)
Amount of substance- mole (mol)

9. Prefixes used in SI units
giga- G 1,000,000,
mega - M 1,000,
kilo - k 1,
deci- d
centi- c
milli- m
micro- m
nano- n

10. Volume measurement: Liter
Liter is defined as the volume of 1 dm3
1 dm3 =
(10cm)3 =
1000 cm3 =
1000mL

11. Pipet
Graduated
Cylinder
Buret
Volumetric
Flask

12. Mass and Weight Mass is measure of resistance to change in motion
Weight is force of gravity.
Sometimes used interchangeably
Mass can’t change, weight can

13. Electronic Analytical Balance

14. Errors in Measurement
All scientific measurements are subject to error.
These errors are reflected in the observation that two successive measurements of the same quantity are different.

15. Types of Errors
Random Error (Indeterminate Error) - measurement has an equal probability of being high or low.
Systematic Error (Determinate Error) - Occurs in the same direction each time (high or low), often resulting from poor technique of measurement or bad equipment.
You can have precision without accuracy
You can’t have accuracy without precision (unless you’re really lucky).

16. Uncertainty in Measurements

17. Uncertainty in Measurement
A measurement always has
some degree of uncertainty.
Uncertainty has to be indicated in any measurement.
Any measurement has certain digits and one uncertain digit.
A digit that must be estimated is
called uncertain.
The number of certain digits + the uncertain digit is called number of significant figures.

19. Precision and Accuracy
Accuracy: Agreement of a particular value to the true value (degree of correctness)
Measurements that are close to the “correct” value are accurate.
Precision: The degree of agreement among several measurements of the same quantity (degree of repeatability).
Measurements that are close to each other are precise.

20. Precision and Accuracy

21. 1.5 Significant figures and calculations
The number of digits reported in a measurement reflect the accuracy of the measurement and the precision of the measuring device.
# Sig figs. =?
All the figures known with certainty plus one extra figure are called significant figures.
# Sig figs. =?

22. Significant figures in computation
For multiplication and division
the results are reported to the least number of significant figures
X = =
For addition and subtraction
The results are reported to the least number of decimal places
– = =

23. Rules for Counting Significant Figures
Nonzero integers always count as significant figures.
3456 has
4 sig figs.

24. 0.0486 has 3 sig figs. 0.0003 has one significant figure
Zeros
Leading zeros do not count as significant figures.
(Zeros before the nonzero digit)
has
3 sig figs.
has
one significant figure
Captive zeros always count as significant figures.
16.07 has
4 sig figs.

25. 9.300 has 4 sig figs. if the number contains a decimal point. Zeros
Trailing zeros are significant only
if the number contains a decimal point.
9.300 has
4 sig figs.

26. Exact numbers have an infinite number of significant figures.
1 inch = cm, exactly
Zeros at the end of a number before a
decimal place are ambiguous
10,300 g has: 3 or 4 or 5??

27. Addition and Subtraction (6.6 x 10-8) + (4.0 x 10-9) =
Scientific Notation
Addition and Subtraction
(6.6 x 10-8) + (4.0 x 10-9) =
(3.42 x 10-5) – (2.5 x 10-6) =
7 x 10-8
3.17 x 10-5
(Note that these answers have been expressed in standard form)

28. Rules for Rounding Off To get the correct number of significant digits
you need to round numbers
In a series of calculations get one extra digit then
round
If the digit to be removed
is less than 5, the preceding digit stays the same
is equal to or greater than 5, the preceding digit is increased by 1
Don’t forget to add place-holding zeros if necessary to keep value the same!! 13 13

29. Multiple computations
2.54 X =
X 0.060
1) ) )
Continuous calculator operation =
x  

30. Here, the mathematical operation requires that we apply the addition/ subtraction rule first, then apply the multiplication/division rule.
= 12

31. Using the units to solve problems
1.6 Dimensional Analysis
Using the units to solve problems

32. Dimensional Analysis
Method of calculation utilizing a knowledge of units.
Conversion factors are used to manipulate units:
The conversion factors are simple ratios:

33. How many minutes are in 2.5 hours?
Initial unit
2.5 hr
Conversion Final
factor unit
2.5 hr x min = min
1 hr

34. How many seconds are in 1.4 days? Unit plan: days hr min seconds
1.4 days x 24 hr x ??
1 day
Exact numbers
1.4 day x 24 hr x 60 min x 60 sec
1 day hr min
= 1.2 x 105 sec

35. Multiple units The speed limit is 65 mi/hr. What is this in m/s? 65 mi
1 mile = 1760 yds
1 meter = yds
65 mi hr
1760 yd
1 m
1 hr
1 min
1 mi
1.094 yd
60 min
60 s

36. If you are running at a speed of 65 meters per minute, how many seconds will it take for you to walk a distance of feet?
Initial
8450 ft x 12 in. x cm x 1 m
1 ft in cm
x 1 min x 60 sec = sec
65 m 1 min

37. Units to a Power
How many m3 is 1500 cm3?
1 m
100 cm 3
1500 cm3

38. 1.7 Temperature
Define the three temperature scales: Celsius , Fahrenheit and Kelvin
Perform conversion from one to another.

39. Units of Temperature between Boiling and Freezing
Fahrenheit Celsius Kelvin
Water boils °F °C K
180° °C K
Water freezes 32°F °C K

40. K = oC + 273

41. 1.8 Density Density is the mass of substance
per unit volume of the substance:

42. Densities of Various Common Substances* at 20° C

43. Density Problem
An empty container weighs g. When filled with a liquid (density 1.53 g/cm3 ) the container weighs g. What is the volume of the container?

44. 1.9 Classification of Matter
Matter: Anything occupying space and having mass.
Three States of Matter:
Solid: rigid - fixed volume and shape
Liquid: definite volume but assumes the shape of its container
Gas: no fixed volume or shape - assumes the shape of its container

45. Types of Mixtures Matter could be pure (one component only) or mixture
Mixtures have variable composition (more than one component)
A homogeneous mixture has visibly indistinguishable components usually called a solution (for example, tap water)
A heterogeneous mixture has visibly distinguishable components, clearly not uniform (for example milk)

46. Organization of Matter

47. Physical and Chemical Changes
A physical change, is associated with a change in the physical appearance but not in chemical composition
Ice melts: a solid is converted into a liquid.
When a substance changes its chemical composition, it undergoes a chemical change:
H2 + O pure water.
In the flask containing water, there is no oxygen or hydrogen left over.

48. Physical and Chemical Changes
Properties of Matter
Physical and Chemical Changes

49. Separation of Mixtures
Mixtures can be separated if their physical properties are different.
Solids can be separated from liquids by means of filtration.
The solid is collected in filter paper, and the solution, called the filtrate, passes through the filter paper and is collected in a flask.

50. Filtration

51. Separation of Mixtures
Homogeneous liquid mixtures can be separated by distillation.
Distillation requires the different liquids to have different boiling points.
Thus, each component of the mixture is boiled and collected.
The lowest boiling fraction is collected first.

52. Distillation

53. Distillation is a physical change: No chemical change occurs when salt water is distilled.

54. Separation of Mixtures
Chromatography can be used to separate mixtures that have different abilities to adhere to solid surfaces.
The greater the affinity the component has for the surface (paper) the slower it moves.
The greater affinity the component has for the liquid, the faster it moves.
Chromatography can be used to separate the different colors of inks in a pen.

55. Paper chromatography: A Line of the mixture to be separated is placed at one end of a sheet

56. The Paper Acts as a Wick to Draw up the Liquid

57. Component with the weakest attraction for the paper travels faster

58. Compounds and elements
Compound: A substance with a constant composition that can be broken down into elements by chemical processes.
Element: A substance that cannot be decomposed into simpler substances by chemical means.