Chapter 1 The Study of Chemistry

Topics Introduction Scientific Method Classifications of Matter Properties of Matter Units of Measurement – Metric system Temperature Conversion Metric Conversion (Prefixes) Accuracy vs. Precision Significant Figures Density

Scientific Method 1. Observation 2. Hypothesis – an educated guess 3. Experimentation 4. Results 5. Conclusion

States of Matter Solid Liquid Gas Plasma

ATOM Is the simplest unit of matter.

Definitions Elements – can’t be decomposed further into simpler substances - 118 elements presently - Lv (element 116 - livermorium) - Fl (element 114 - flerovium) Compound – combination of 2 or more elements Pure substance – has distinct properties and composition; does not vary from sample to sample (ex. Water, NaCl)

Definitions Mixtures – combinations of 2 or more substances (ex. sugar in water) 2 Types of Mixtures 1. Homogenous Mixtures (solutions) = 1 phase 2. Heterogeneous Mixtures = > 2 phases

SOLUTIONS Homogeneous mixtures are called SOLUTIONS.

Solution Solution – homogenous mixture A solution is not necessarily a liquid. Can be gas or solid.

Physical vs. Chemical Properties Physical properties – can be measured w/o changing identity and composition of substance (ex. Boiling pt.,freezing pt., color, odor, density, hardness) Chemical properties – describe how substance reacts or changes to form other compounds (ex. Flammability, toxicity)

Changes of State and Properties Physical changes – does not change composition of compound Chemical changes – converts to a different chemical substance Intensive Properties – independent of amt. (ex. Density, Temperature, Melting Pt) Extensive Properties – dependent on amt. (ex. Mass, Volume)

Units of Measurement Use the Metric System!

Units of Measurement Mass – grams; kilogram Length – centimeter; meter Volume – milliliter or cubic centimeter (cm3) Temperature – Celcius; Kelvin

Precision vs Accuracy Accuracy – when acquired value agrees with true value Precision – when acquired values exhibit reproducibility

Significant Figures More significant figures = more certainty Helps in determining how to round measured values and still precise

SIGNIFICANT FIGURES In counting and definitions, there are an infinite number of sig figs In measurements, the number of sig figs consists of all certain and the first uncertain digits Unit conversions do not determine # of sig. figs.

Rules of Significant Figures 1. Non-zero integers always count. Ex. 1234.5 grams = 5 Sig. Figs. 2. Captive zeros are always significant. Ex. 100.3 grams = 4 Sig. Figs.

Rules of Significant Figures 3. Leading zeros are NEVER significant. Ex. 0.6780 grams = 4 Sig. Figs. 4. Trailing zeroes are significant ONLY if there is a decimal point Ex. 12.0 grams = 3 Sig. Figs 120 grams = 2 Sig. Figs

Rules of Significant Figures 5. Exact numbers (obtained by counting) are infinite and do not determine the number of significant figures. Example: 4 cows = ?

Determine the # of Sig. Fig. 200.0 1050 3003 0.0006 10,000 0.5

Rules of Significant Figures Multiplication/Division Answer will have the same # of sig figs as the value with the least # of sig figs Ex: 3.8 x 200.0 = 2 Sig. Figs.

Rules of Significant Figures Addition/Subtraction Answer has the same # of decimal places as the number with the least # of decimal places Ex. 3.1 + 2.500 + 5.76 = 11.4

Order of Operations Parenthesis Multiplication/division Addition/subtraction

Rounding Look only to the right of the number you are rounding to: - If 5 or more, round up - If less than 5, round down

General Rule Carry ALL figures through to the end of a problem. Round the final answer to the correct number of significant figures

Problem Indicate the number of sig. figs. in each of the following measured quantities: A. 358 kg B. 0.054 s C. 6.3050 cm D. 0.0105 L E. 7.0500 x 10-3 m3

Problem Round each of the following numbers to 4 sig. figs. And express the result in standard exponential or scientific notation. A. 102. 53070 B. 656, 980 C. 0.008543210 D. 0.000257870 E. - 0. 0357202

Problem Carry out the following operations and express the answer with the appropriate number of sig. figs. A. 12.0550 + 9.05 B. 257.2 – 19.789 C. (6.21 x 103)(1.1050) D. 0.0577 / 0.753

Prefixes in Metric System Mega - million Kilo - 1,000 Hecto - 100 Deka - 10 ----- - 1 (liter, gram, meter) Deci - 1/10 or 0.1 Centi - 1/100 or 0.01 Milli - 1/1000 or 0.001

Trick to Unit Conversions Arrange the values with units such that you can cancel what you do not want. You should only end up with the desired unit.

Tricks to Conversion Big to small - multiply Small to Big - divide going down Small to Big - divide Going up

Temperature Conversions 0 oC = 273.15 K oF = 1.8 oC + 32

Things to Remember! 1 milliliter = 1 cc 1000 milliliter = 1 liter 0 oC = 32 oF = 273.15 K

Density Is the amount of mass in a unit volume of the substance Is affected by Temperature. The higher the temp., the lower the density. D = mass of substance = grams volume of substance mL or cm3

Density Density = mass volume = gram mL

Different ways of calculating volume I. For solids with regular shapes: A. For a cube: Vcube = s3 B. For a rectangular solid, V = L x W x H C. For a cylinder: V= pr2h D. For a sphere: V = 4/3 pr3

Different ways of calculating volume II. For an Irregular Solid Water displacement

Different ways of calculating volume III. For a liquid Use of graduated cylinder, beaker, pipet or buret.

Problem A cube of osmium metal 1.500 cm on a side has a mass of 76.31 grams at 25 oC. What is its density in g/cm3 at this temperature?

Problem The density of titanium metal is 4.51 g/cm3 at 25 oC. What mass of titanium displaces 65.8 mL of water at 25 oC?

Problem The density of benzene at 15 oC is 0.8787 g/mL. Calculate the mass of 0.1500 L of benzene at this temperature.