Chapter 1 Matter & Measurement

Chemistry is… …the study of the composition, structure, and properties of matter and the changes it undergoes C2H5OH + 3 O2  2 CO2 + 3 H2O + Energy Reactants  Products

Matter Mass A measure of the amount of matter Anything that has mass and occupies space Mass A measure of the amount of matter

Atom Element The smallest unit of an element that maintains the properties of that element Element A pure substance made of only one kind of atom

Properties of Matter Extensive properties depend on the amount of matter that is present. Extensive properties Volume Mass Energy Content (think Calories!) do not depend on the amount of matter present. Intensive properties Melting point Boiling point Density

Physical Change A change in a substance that does not involve a change in the identity of the substance. Example: Phase Changes

Phase Differences Solid – definite volume and shape; particles packed in fixed positions. Liquid – definite volume but indefinite shape; particles close together but not in fixed positions Gas – neither definite volume nor definite shape; particles are at great distances from one another Plasma – high temperature, ionized phase of matter as found on the sun.

Three Phases

Copper Phases - Solid

Copper Phases - Liquid

Copper Phases – Vapor (gas)

Chemical Change A change in which one or more substances are converted into different substances. Heat and light are often evidence of a chemical change.

Separation of a Mixture The constituents of the mixture retain their identity and may be separated by physical means.

Separation of a Mixture The components of dyes such as ink may be separated by paper chromatography.

Filtration:

Separation of a Mixture Distillation

Separation of a Compound The Electrolysis of water Compounds must be separated by chemical means. With the application of electricity, water can be separated into its elements Reactant  Products Water  Hydrogen + Oxygen H2O  H2 + O2

Measurement

Nature of Measurement Part 1 - number Part 2 - scale (unit) Examples: Measurement - quantitative observation consisting of 2 parts Part 1 - number Part 2 - scale (unit) Examples: 20 grams 6.63 x 10-34 Joule seconds

The Fundamental SI Units (le Système International, SI)

SI Prefixes Common to Chemistry Unit Abbr. Exponent Kilo k 103 Deci d 10-1 Centi c 10-2 Milli m 10-3 Micro  10-6

Temperature Scales

The Thermometer Determine the temperature by reading the scale on the thermometer at eye level. Read the temperature by using all certain digits and one uncertain digit. Certain digits are determined from the calibration marks on the thermometer. The uncertain digit (the last digit of the reading) is estimated. On most thermometers encountered in a general chemistry lab, the tenths place is the uncertain digit.

Do not allow the tip to touch the walls or the bottom of the flask. If the thermometer bulb touches the flask, the temperature of the glass will be measured instead of the temperature of the solution. Readings may be incorrect, particularly if the flask is on a hotplate or in an ice bath.

Reading the Thermometer Determine the readings as shown below on Celsius thermometers: _ _ . _ C 8 7 4 _ _ . _ C 3 5

Volume Instruments

Reading the Meniscus Always read volume from the bottom of the meniscus. The meniscus is the curved surface of a liquid in a narrow cylindrical container.

Try to avoid parallax errors. Parallax errors arise when a meniscus or needle is viewed from an angle rather than from straight-on at eye level. Correct: Viewing the meniscus at eye level Incorrect: viewing the meniscus from an angle

Measuring Volume Determine the volume contained in a graduated cylinder by reading the bottom of the meniscus at eye level. Read the volume using all certain digits and one uncertain digit. Certain digits are determined from the calibration marks on the cylinder. The uncertain digit (the last digit of the reading) is estimated.

Use the graduations to find all certain digits There are two unlabeled graduations below the meniscus, and each graduation represents 1 mL, so the certain digits of the reading are… 52 mL.

Estimate the uncertain digit and take a reading The meniscus is about eight tenths of the way to the next graduation, so the final digit in the reading is . 0.8 mL The volume in the graduated cylinder is 52.8 mL.

10 mL Graduate What is the volume of liquid in the graduate? 6 _ . _ _ mL 6 2

Uncertainty in Measurement A digit that must be estimated is called uncertain. A measurement always has some degree of uncertainty.

Why Is there Uncertainty? Measurements are performed with instruments No instrument can read to an infinite number of decimal places Which of these balances has the greatest uncertainty in measurement?

Precision and Accuracy Accuracy refers to the agreement of a particular value with the true value. Precision refers to the degree of agreement among several measurements made in the same manner. Neither accurate nor precise Precise but not accurate Precise AND accurate

Significant Digits: Atlantic/ Pacific Rule When the decimal is ABSENT, go to the Atlantic side of the number, start counting digits when you reach a non-zero number. Record

Practice 45,000 hrs 78,700 kilometers 3,000 liters two three one

Pacific - when the decimal is PRESENT, go to the Pacific side of the number, start counting digits when you reach a non-zero number

Practice .009999 grams 560.03 mL 100.0 meters .00506 4 5 none

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

Rules for Counting Significant Figures - Details Zeros - Captive zeros always count as significant figures. 16.07 has 4 sig figs.

More practice 56,000,000 seconds 33,000 candles 60900 milligrams .899000 centimeters .6700 meters 2 0 * 3 6 4

Sig Fig Practice #1 1.0070 m  5 sig figs 17.10 kg  4 sig figs How many significant figures in each of the following? 1.0070 m  5 sig figs 17.10 kg  4 sig figs 100,890 L  5 sig figs 3.29 x 103 s  3 sig figs 0.0054 cm  2 sig figs 3,200,000  2 sig figs

Density – a derived measure value is found through mathematical computation D=Mass/ Volume

Scientific Notation In science, we deal with some very LARGE numbers: 1 mole = 602000000000000000000000 In science, we deal with some very SMALL numbers: Mass of an electron = 0.000000000000000000000000000000091 kg

. 2 500 000 000 9 8 7 6 5 4 3 2 1 Step #1: Insert an understood decimal point Step #2: Decide where the decimal must end up so that one number is to its left Step #3: Count how many places you bounce the decimal point Step #4: Re-write in the form M x 10n

2.5 x 109 The exponent is the number of places we moved the decimal.

0.0000579 1 2 3 4 5 Step #2: Decide where the decimal must end up so that one number is to its left Step #3: Count how many places you bounce the decimal point Step #4: Re-write in the form M x 10n

5.79 x 10-5 The exponent is negative because the number we started with was less than 1.