Basic Concepts of Matter Chapter 1 Basic Concepts of Matter

Classifying Matter Matter Mass vs. Weight Kinetic-Molecular Theory Anything that has mass and occupies space Mass vs. Weight Kinetic-Molecular Theory All matter consists of extremely tiny particles in constant motion

States of Matter Solid Liquid Gas -Closely packed together with a definite ridged shape -Vibrate back and forth in a confined space -the particles are not able to move past one another Liquid -arranged randomly with a definite volume -“fluid” -the particles are not confined in space and can move past one another Gas -no definite shape or volume -the particles are far apart and move very rapidly colliding with other particles and the container walls

Categorizing Matter Elements Pure Substance -cannot be decomposed into simpler form via chemical reactions -found on periodic chart -atoms are the smallest particle that retains the characteristic properties of the elements Pure Substance -consists of all the same substance (pure gold, distilled water, etc) -have a set of unique properties that identifies it

Categorizing Matter Chemical Compounds -two or more elements in a definite ratio by mass with unique properties that separate them from the individual elements -can be decomposed into the constituent elements by chemical reactions -chemical compounds are held together by a chemical bound Water– hydrogen and oxygen Carbon dioxide – carbon and oxygen

Categorizing Matter Mixtures homogeneous mixtures (solution) two or more pure substances in the same container homogeneous mixtures (solution) -uniform composition throughout -single phase -cannot be separated easily heterogeneous mixtures -nonuniform composition thoughout -easily separated

Physical and Chemical Changes Physical changes changes in physical properties -melting, boiling, and cutting Chemical changes changing one or more substances into one or more different substances (chemical reaction) 2H2 + O2 -> 2H2O

Chemical and Physical Properties Chemical Properties observed during a chemical reaction (change in chemical composition) -rusting, oxidation, burning… -chemical reactions Physical Properties observed without changing the substance’s composition -allow for identification and classification -density, color, solubility, melting point…

Classification of Physical Properties Extensive Properties depend on the amount of substance present -mass or volume Intensive Properties do not depend on the amount of substance present -melting point, boiling point, density…

Density Describes how compact a substance is Who “discovered” density? Density = mass/volume or D = m/V

Density Example: Calculate the density of a substance if 742 grams of it occupies 97.3 cm3. 1cm3 = 1mL => 97.3cm3 = 97.3mL

Density Example: You need 125 g of a corrosive liquid for a reaction. What volume do you need? liquid’s density = 1.32 g/mL

Units of Measure Qualitative measures Quantitative measures Nonnumerical experimental observations describing the identity of a substance in a sample Quantitative measures Numerical experimental observations describing how much of a particular substance is in a sample System International d’Unites (SI) measurement system used in the sciences based on the metric system

Math Review and Measurements We make measurements to understand our environment: Human senses: sight, taste, smell, hearing… Our senses have limits and are biased Instruments: an extension of our senses meter sticks, thermometers, balances These are more accurate and precise All measurements have units METRIC SYSTEM vs. British System

SI units Quantity Unit Symbol length meter m mass kilogram kg time second s current ampere A temperature Kelvin K amt. substance mole mol

Measurements in Chemistry Name Symbol Multiplier mega- M 106 (1,000,000) kilo- k 103 (1,000) deka- da 10 deci- d 10-1 (0.1) centi- c 10-2 (0.01)

Measurements in Chemistry Name Symbol Multiplier Milli- m 10-3(0.001) Micro-  10-6(0.000001) Nano- n 10-9 Pico- p 10-12 Femto- f 10-15

Units of Measurement Length Measure of space in any direction -derived unit cm -standard length is a meter (m)

Units of Measurement Volume Amount of space occupied by matter -derived unit: mL or cm3 (cc) -liter (L) is the standard unit

Units of Measurement Time (t) Mass (m) Weight (W) Interval or duration of forward events -standard unit is the second (s) Mass (m) measure of the quantity of matter in a body Weight (W) measure of the gravitational attraction (g) for a body (w=m x g) 1 kg = 1000g 1 kg = 2.2 lbs 1 g = 1000mg

Heat and Temperature Heat (q) vs. Temperature (T) 3 common temperature scales: all use water as a reference -Fahrenheit (F) -Celsius (C) -Kelvin (K)

Temperature Reference Points Melting Point Boiling Point of water of water 32 oF 212 oF 0.0 oC 100 cC 273 K 373 K Body temperature 37.0 oC or 98.6 oF 37.2 oC and greater—sick 41 oC and greater, convulsions <28.5 oC hypothermia Fahrenheit Celsius Kelvin

Temperature Scales

Fahrenheit to Centigrade Relationships Temperature Scales Fahrenheit to Centigrade Relationships Example: Convert 211 oF to degrees Celsius. Example: Express 548 K in Celsius degrees.

Precision and Accuracy how closely measured values agree with the correct value Precision how closely individual measurements agree with each other Precise Accurate Both Neither

Mathematics in Chemistry Exact numbers (counted numbers) 1 dozen = 12 things Measured Numbers Use rules for significant figures Use scientific notation when possible Significant figures digits in a measured quantity that reflect the accuracy of the measurement -in other words, digits believed to be correct by the person making the measurement Exact numbers have an infinite number of significant figures 12.000000000000000 = 1 dozen

Significant figures (numbers/digits) Why use significant numbers? -Calculators give 8+ numbers -People estimate numbers differently -Dictated by the precision (graduation) on your measuring device -In the lab, the last significant digit is the digit you (the scientist) estimate Scientists have develop rules to help determine which digits are “significant”

Rules for Significant Figures 1. All Nonzero numbers are significant!!! 2. Leading zeroes are never significant 0.000357 3. Imbedded zeroes are always significant 3.0604 4. Trailing zeroes may be significant - You must specify significance by how the number is determined or even written 1300 nails - counted or weighed? 1.30000 –How many significant figures?

Significant Figures Multiplication & Division rule: The product retains the number of significant figures that corresponds to the multiplier with the smallest number of significant figure (sig. fig.)

Significant Figures Addition & Subtraction rule: Answer retains the smallest decimal place value of the addends.

Locating the decimal and deciding when to count the zeros!!! Scientific Notation Express answers as powers of 10 by moving the decimal place right (-) or left (+) Use of scientific notation is to remove doubt in the Significant Figures: 2000  2 x 103 15000  1.5 x 10? 0.004  4 x 10-3 0.000053  __.__ x 10? In scientific notation, zeros are given if they are significant!!! 1.000 x 103 has 4 significant figures 2.40 x 103 has ? significant figures Key to Sig. Figs… Locating the decimal and deciding when to count the zeros!!!

Review #2 Units of Measure Heat vs. Temperature Precision vs. Accuracy -length -volume -time -mass -weight Heat vs. Temperature -three temperature scales -temperature conversions Precision vs. Accuracy Significant Figures Scientific Notation

Conversion Factors Length Volume See Text for more conversion factors 1 m = 39.37 inches 2.54 cm = 1 inch Volume 1 liter = 1.06 qt 1 qt = 0.946 liter See Text for more conversion factors

Conversion Factors Why do conversions? -Scientists often must convert between units Conversion factors can be made for any relationship of units -Use known equivalence to make a fraction that can be used to “convert” from one unit to the other

Dimensional Analysis 1 inch = 2.54 cm Use the ratio to perform a calculation so the units will “divide out” Example: Convert 60 inches to centimeters

Dimensional Analysis Example: Express 9.32 yards in millimeters. 3 ft = 1 yard 1 ft = 12 in or 1 in = 2.54 cm 100 cm = 1 m 1000 mm= 1 m

Dimensional Analysis Example: Express 627 milliliters in gallons. 1 liter = 1.06 qt 1 qt = 0.946 liter

Practice on your Own 1kg = 2.20 lbs Convert 25 g to lbs Convert 1 mL to Liters Convert 20 meters to cm

Dimensional Analysis Area = length x width Area is two dimensional thus units must be in squared terms: Express: 2.61 x 104 cm2 in ft2

Volume =length x width x height Dimensional Analysis Volume =length x width x height Volume is three dimensional thus units must be in cubic terms Express: 2.61 ft3 in cm3 this volume is used in medical measurements--cc

Percentage Percentage is parts per hundred of a sample % = x100 Example: A 335 g sample of ore yields 29.5 g of iron. What is the percent of iron in the ore? g of substance total g of sample