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 E-mail: amjaber@kfupm.edu.sa Office hours: S 9-11 am U 10-11 am M 9-11am

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

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

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

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

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

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

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)

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

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

Pipet Graduated Cylinder Buret Volumetric Flask

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

Electronic Analytical Balance

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.

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).

Uncertainty in Measurements

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.

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.

Precision and Accuracy

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. =?

Significant figures in computation 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

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

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) 0.0486 has 3 sig figs. 0.0003 has one significant figure Captive zeros always count as significant figures. 16.07 has 4 sig figs.

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.

Exact numbers have an infinite number of significant figures. 1 inch = 2.54 cm, exactly Zeros at the end of a number before a decimal place are ambiguous 10,300 g has: 3 or 4 or 5??

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)

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

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

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

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

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

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

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 1 hr 1 min = 1.2 x 105 sec

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

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

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

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

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

K = oC + 273

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

Densities of Various Common Substances* at 20° C

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?

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

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)

Organization of Matter

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 + O2 pure water. In the flask containing water, there is no oxygen or hydrogen left over.

Physical and Chemical Changes Properties of Matter Physical and Chemical Changes

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.

Filtration

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.

Distillation

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

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.

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

The Paper Acts as a Wick to Draw up the Liquid

Component with the weakest attraction for the paper travels faster

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.