1	"Professor" Professor Abdul Muttaleb Jaber Chemistry Department Office: Room # 261F Tel: 2611 E-mail: amjaber@kfupm.edu.sa Office hours: S 9-10; 11-12 am U 9-10 am M 9-10; 11-12 am
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3	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
4	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
5	"Matter is composed of various..." 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 H₂to explode forming H₂O. One substance changes to another by reorganizing the way atoms attached to each other
6	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
7	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
8	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 X10³g 20 m g = 20 X10^-3g 6.63 ´ 10^-34 Joule seconds
9	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)
10	Prefixes used in SI units giga- G 1,000,000,000 10⁹ mega - M 1,000,000 10⁶ kilo - k 1,000 10³ deci- d 0.1 10^-1 centi- c 0.01 10^-2 milli- m 0.001 10^-3 micro- m 0.000001 10^-6 nano- n 0.000000001 10^-9
11	Volume measurement: Liter Liter is defined as the volume of 1 dm³ 1 dm³ = (10cm)³ = 1000 cm³ = 1000mL
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13	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
14	Electronic Analytical Balance
15	Errors in Measurement
16	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).
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18	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.
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20	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.
21	Precision and Accuracy
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23	"For multiplication and division" For multiplication and division the results are reported to the least number of significant figures 354.760 X 0.0004567 = 0.162018892 = 0.1620 For addition and subtraction The results are reported to the least number of decimal places 345.672 – 34.56720 = 311.1048 = 311.105
24	Rules for Counting Significant Figures Nonzero integers always count as significant figures. 3456 has 4 sig figs.
25	"Zeros" Zeros Leading zeros do not count as significant figures. (Zeros before the nonzero digit) 0.0486 has 3 sig figs. 0.0003 has one significant figure
26	"Zeros" Zeros Trailing zeros are significant only if the number contains a decimal point. 9.300 has 4 sig figs.
27	"Exact numbers have an infinite..." Exact numbers have an infinite number of significant figures. 1 inch = 2.54 cm, exactly
28	Scientific Notation
29	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!!
30	Multiple computations 2.54 X 0.0028 = 0.0105 X 0.060 1) 11.3 2) 11 3) 0.041 Continuous calculator operation = 2.54 x 0.0028 ¸ 0.0105 ¸ 0.060
31	"Here" Here, the mathematical operation requires that we apply the addition/ subtraction rule first, then apply the multiplication/division rule.
32	1.6 Dimensional Analysis Using the units to solve problems
33	Dimensional Analysis
34	How many minutes are in 2.5 hours? Initial unit 2.5 hr Conversion Final factor unit 2.5 hr x 60 min = 150 min 1 hr
35	"How many seconds are in..." How many seconds are in 1.4 days? Unit plan: days hr min seconds 1.4 days x 24 hr x ?? 1 day
36	Multiple units The speed limit is 65 mi/hr. What is this in m/s? 1 mile = 1760 yds 1 meter = 1.094 yds
37	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 8450 feet? Initial 8450 ft x 12 in. x 2.54 cm x 1 m 1 ft 1 in. 100 cm x 1 min x 60 sec = 2400 sec 65 m 1 min
38	Units to a Power How many m³is 1500 cm³?
39	1.7 Temperature Define the three temperature scales: Celsius , Fahrenheit and Kelvin Perform conversion from one to another.
40	Units of Temperature between Boiling and Freezing Fahrenheit Celsius Kelvin Water boils 212°F 100°C 373 K 180° 100°C 100K Water freezes 32°F 0°C 273 K
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42	1.8 Density Density is the mass of substance per unit volume of the substance:
43	Densities of Various Common Substances* at 20° C
44	Density Problem An empty container weighs 121.3 g. When filled with a liquid (density 1.53 g/cm³ ) the container weighs 283.2 g. What is the volume of the container?
45	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
46	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)
47	Organization of Matter
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49	Properties of Matter
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52	Separation of Mixtures
53	Distillation
54	Distillation is a physical change: No chemical change occurs when salt water is distilled.
55	Separation of Mixtures
56	Paper chromatography: A Line of the mixture to be separated is placed at one end of a sheet
57	The Paper Acts as a Wick to Draw up the Liquid
58	Component with the weakest attraction for the paper travels faster
59	Compounds and elements Element: A substance that cannot be decomposed into simpler substances by chemical means.