The Foundations of Chemistry CHAPTER ONE The Foundations of Chemistry

Chapter Outline Matter and Energy States of Matter Chemical and Physical Properties Chemical and Physical Changes Mixtures, Substances, Compounds, and Elements Measurements in Chemistry Units of Measurement

Chapter Outline Use of Numbers The Unit Factor Method (Dimensional Analysis) Percentage Density and Specific Gravity Heat and Temperature Heat Transfer and the Measurement of Heat

Matter and Energy - Vocabulary Chemistry Science that describes matter – its properties, the changes it undergoes, and the energy changes that accompany those processes Matter Anything that has mass and occupies space. Energy The capacity to do work or transfer heat. Scientific (natural) law A general statement based the observed behavior of matter to which no exceptions are known.

Natural Laws Law of Conservation of Mass Law of Conservation of Energy Law of Conservation of Mass-Energy Einstein’s Relativity E=mc2

Scientific Method Observation Hypothesis Observation or experiment Theory Law

States of Matter Solids

States of Matter Solids Liquids

States of Matter Solids Liquids Gases

States of Matter Change States heating cooling

States of Matter Illustration of changes in state requires energy

Chemical and Physical Properties Chemical Properties - chemical changes rusting or oxidation chemical reactions Physical Properties - physical changes changes of state density, color, solubility Extensive Properties - depend on quantity Intensive Properties - do not depend on quantity

Mixtures, Substances, Compounds, and Elements matter in which all samples have identical composition and properties Elements substances that cannot be decomposed into simpler substances via chemical reactions Elemental symbols found on periodic chart

Mixtures, Substances, Compounds, and Elements

Mixtures, Substances, Compounds, and Elements substances composed of two or more elements in a definite ratio by mass can be decomposed into the constituent elements Water is a compound that can be decomposed into simpler substances – hydrogen and oxygen

Mixtures, Substances, Compounds, and Elements

Mixtures, Substances, Compounds, and Elements composed of two or more substances homogeneous mixtures heterogeneous mixtures

Measurements in Chemistry 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 Metric Prefixes Name Symbol Multiplier mega M 106 kilo k 103 deka da 10 deci d 10-1 centi c 10-2

Measurements in Chemistry Metric Prefixes Name Symbol Multiplier milli m 10-3 micro  10-6 nano n 10-9 pico p 10-12 femto f 10-15

Units of Measurement Definitions Mass Weight measure of the quantity of matter in a body Weight measure of the gravitational attraction for a body

Units of Measurement Common Conversion Factors Length Volume 1 m = 39.37 inches 2.54 cm = 1 inch Volume 1 liter = 1.06 qt 1 qt = 0.946 liter See Table 1-7 for more conversion factors

Use of Numbers Exact numbers Accuracy Precision 1 dozen = 12 things for example Accuracy how closely measured values agree with the correct value Precision how closely individual measurements agree with each other

Use of Numbers Significant figures digits believed to be correct by the person making the measurement Measure a mile with a 6 inch ruler vs. surveying equipment Exact numbers have an infinite number of significant figures 12.000000000000000 = 1 dozen because it is an exact number

Significant Figures - Rules Use of Numbers Significant Figures - Rules Leading zeroes are never significant 0.000357 has three significant figures Trailing zeroes may be significant must specify significance by how the number is written 1300 nails - counted or weighed? Use scientific notation to remove doubt 2.40 x 103 has ? significant figures

Use of Numbers Scientific notation for logarithms take the log of 2.40 x 103 log(2.40 x 103) = 3.380 How many significant figures? Imbedded zeroes are always significant 3.0604 has five significant figures

Use of Numbers Piece of Black Paper – with rulers beside the edges

Use of Numbers Piece of Paper Side B – enlarged How long is the paper to the best of your ability to measure it?

Use of Numbers Piece of Paper Side A – enlarged How wide is the paper to the best of your ability to measure it?

Use of Numbers Determine the area of the piece of black paper using your measured values. Compare your answer with your classmates. Where do your answers differ in the numbers? Significant figures rules for multiplication and division must help us determine where answers would differ.

Use of Numbers Multiplication & Division rule Easier of the two rules Product has the smallest number of significant figures of multipliers

Use of Numbers Multiplication & Division rule Easier of the two rules Product has the smallest number of significant figures of multipliers

Use of Numbers Multiplication & Division rule Easier of the two rules Product has the smallest number of significant figures of multipliers

Use of Numbers Determine the perimeter of the piece of black paper using your measured values. Compare your answer with your classmates. Where do your answers differ in the numbers? Significant figures rules for addition and subtraction must help us determine where answers would differ.

Use of Numbers Addition & Subtraction rule More subtle than the multiplication rule Answer contains smallest decimal place of the addends.

Use of Numbers Addition & Subtraction rule More subtle than the multiplication rule Answer contains smallest decimal place of the addends.

Use of Numbers Addition & Subtraction rule More subtle than the multiplication rule Answer contains smallest decimal place of the addends.

The Unit Factor Method Simple but important method to get correct answers in word problems. Method to change from one set of units to another. Visual illustration of the idea.

The Unit Factor Method Change from a to a by obeying the following rules.

The Unit Factor Method Change from a to a by obeying the following rules. Must use colored fractions.

The Unit Factor Method Change from a to a by obeying the following rules. Must use colored fractions. The box on top of the fraction must be the same color as the next fraction’s bottom box.

The Unit Factor Method Fractions to choose from R R O B O B B O R O B

The Unit Factor Method Fractions to choose from O R R R O B O B B O R

The Unit Factor Method Fractions to choose from O B R R O R O B O B B

The Unit Factor Method Fractions to choose from O B B R B R O B R O B

The Unit Factor Method Fractions to choose from O B B R B R O B R O B

The Unit Factor Method Fractions to choose from O B B R B R O B R O B

The Unit Factor Method Fractions to choose from O B B R B R O B R O B

The Unit Factor Method colored fractions represent unit factors 1 ft = 12 in becomes or Example 1-1: Express 9.32 yards in millimeters.

The Unit Factor Method

The Unit Factor Method

The Unit Factor Method

The Unit Factor Method

The Unit Factor Method O B B T R T R O B B

The Unit Factor Method Example 1-2: Express 627 milliliters in gallons. You do it!

The Unit Factor Method Example 1-2. Express 627 milliliters in gallons.

The Unit Factor Method Example 1-3: Express 3.50 (short) tons in grams. You do it!

The Unit Factor Method Example 1-3: Express 3.50 (short) tons in grams.

The Unit Factor Method Area is two dimensional thus units must be in squared terms. Example 1-4: Express 2.61 x 104 cm2 in ft2.

The Unit Factor Method Area is two dimensional thus units must be in squared terms. Example 1-4: Express 2.61 x 104 cm2 in ft2. common mistake

The Unit Factor Method Area is two dimensional thus units must be in squared terms. Example 1-4: Express 2.61 x 104 cm2 in ft2. O R P

The Unit Factor Method Area is two dimensional thus units must be in squared terms. Example 1-4: Express 2.61 x 104 cm2 in ft2. O R R

The Unit Factor Method Area is two dimensional thus units must be in squared terms. Example 1-4: Express 2.61 x 104 cm2 in ft2.

The Unit Factor Method Area is two dimensional thus units must be in squared terms. Example 1-4: Express 2.61 x 104 cm2 in ft2.

The Unit Factor Method Volume is three dimensional thus units must be in cubic terms. Example 1-5: Express 2.61 ft3 in cm3. You do it!

The Unit Factor Method Volume is three dimensional thus units must be in cubic terms. Example 1-5: Express 2.61 ft3 in cm3.

Percentage Percentage is the parts per hundred of a sample. Example 1-6: A 335 g sample of ore yields 29.5 g of iron. What is the percent of iron in the ore? You do it!

Percentage Percentage is the parts per hundred of a sample. Example 1-6: A 335 g sample of ore yields 29.5 g of iron. What is the percent of iron in the ore?

Density and Specific Gravity density = mass/volume What is density? Why does ice float in liquid water?

Density and Specific Gravity density = mass/volume What is density? Why does ice float in liquid water? H2O(s) H2O(l)

Density and Specific Gravity Example 1-7: Calculate the density of a substance if 742 grams of it occupies 97.3 cm3.

Density and Specific Gravity Example 1-7: Calculate the density of a substance if 742 grams of it occupies 97.3 cm3.

Density and Specific Gravity Example 1-8 Suppose you need 125 g of a corrosive liquid for a reaction. What volume do you need? liquid’s density = 1.32 g/mL You do it!

Density and Specific Gravity Example 1-8 Suppose you need 125 g of a corrosive liquid for a reaction. What volume do you need? liquid’s density = 1.32 g/mL

Density and Specific Gravity Example 1-8 Suppose you need 125 g of a corrosive liquid for a reaction. What volume do you need? liquid’s density = 1.32 g/mL

Density and Specific Gravity Water’s density is essentially 1.00 at room T. Thus the specific gravity of a substance is very nearly equal to its density. Specific gravity has no units.

Density and Specific Gravity Example 1-9: A 31.10 gram piece of chromium is dipped into a graduated cylinder that contains 5.00 mL of water. The water level rises to 9.32 mL. What is the specific gravity of chromium? You do it

Density and Specific Gravity Example1-9: A 31.10 gram piece of chromium is dipped into a graduated cylinder that contains 5.00 mL of water. The water level rises to 9.32 mL. What is the specific gravity of chromium?

Density and Specific Gravity Example1-9: A 31.10 gram piece of chromium is dipped into a graduated cylinder that contains 5.00 mL of water. The water level rises to 9.32 mL. What is the specific gravity of chromium?

Density and Specific Gravity Example 1-10: A concentrated hydrochloric acid solution is 36.31% HCl and 63.69% water by mass. The specific gravity of the solution is 1.185. What mass of pure HCl is contained in 175 mL of this solution? You do it!

Density and Specific Gravity

Density and Specific Gravity

Density and Specific Gravity

Heat and Temperature Heat and Temperature are not the same thing T is a measure of the intensity of heat in a body 3 common temperature scales - all use water as a reference

Heat and Temperature Heat and Temperature are not the same thing T is a measure of the intensity of heat in a body 3 common temperature scales - all use water as a reference

Heat and Temperature MP water BP water Fahrenheit 32 oF 212 oF Celsius 0.0 oC 100 cC Kelvin 273 K 373 K

Relationships of the Three Temperature Scales

Relationships of the Three Temperature Scales

Relationships of the Three Temperature Scales

Relationships of the Three Temperature Scales Easy method to remember how to convert from Centigrade to Fahrenheit. Double the Centigrade temperature. Subtract 10% of the doubled number. Add 32.

Heat and Temperature Example 1-11: Convert 211oF to degrees Celsius.

Heat and Temperature Example 1-12: Express 548 K in Celsius degrees.

Heat Transfer and The Measurement of Heat SI unit J (Joule) calorie 1 calorie = 4.184 J English unit = BTU Specific Heat amount of heat required to raise the T of 1g of a substance by 1oC units = J/goC

Heat Transfer and the Measurement of Heat Heat capacity amount of heat required to raise the T of 1 mole of a substance by 1oC units = J/mol oC Example 1-13: Calculate the amt. of heat to raise T of 200 g of water from 10.0oC to 55.0oC

Heat Transfer and the Measurement of Heat Heat transfer equation necessary to calculate amounts of heat amount of heat = amount of substance x specific heat xDT

Heat Transfer and the Measurement of Heat Heat transfer equation necessary to calculate amounts of heat amount of heat = amount substance x specific heat xDT

Heat Transfer and the Measurement of Heat Heat transfer equation necessary to calculate amounts of heat amount of heat = amount substance x specific heat xT

Heat Transfer and the Measurement of Heat Example 1-14: Calculate the amount of heat to raise the temperature of 200.0 grams of mercury from 10.0oC to 55oC. Specific heat for Hg is 0.138 J/g oC. You do it!

Heat Transfer and the Measurement of Heat Example 1-14: Calculate the amount of heat to raise T of 200 g of Hg from 10.0oC to 55oC. Specific heat for Hg is 0.138 J/g oC. Requires 30.3 times more heat for water 4.184 is 30.3 times greater than 0.138

Synthesis Question It has been estimated that 1.0 g of seawater contains 4.0 pg of Au. The total mass of seawater in the oceans is 1.6x1012 Tg, If all of the gold in the oceans were extracted and spread evenly across the state of Georgia, which has a land area of 58,910 mile2, how tall, in feet, would the pile of Au be? Density of Au is 19.3 g/cm3. 1.0 Tg = 1012g.

Synthesis Question

Synthesis Question

Synthesis Question

Group Activity On a typical day, a hurricane expends the energy equivalent to the explosion of two thermonuclear weapons. A thermonuclear weapon has the explosive power of 1.0 Mton of nitroglycerin. Nitroglycerin generates 7.3 kJ of explosive power per gram of nitroglycerin. The hurricane’s energy comes from the evaporation of water that requires 2.3 kJ per gram of water evaporated. How many gallons of water does a hurricane evaporate per day?

End of Chapter 1