Unit 1

 Introduction  Matter  Physical/Chemical Properties and Changes  Extensive/Intensive Properties  Scientific Notation  Metric System  SI units  Conversions  Density  Significant Figures  Uncertainty

 The study of matter and the changes it undergoes.  Major divisions Inorganic  InorganicCompounds of elements other than carbon Organic  OrganicCompounds of carbon Biochemistry  BiochemistryCompounds of living matter Physical  PhysicalTheory and concepts Analytical  AnalyticalMethods of analysis

 We can explore the MACROSCOPIC world — what we can see —  to understand the PARTICULATE worlds we cannot see.  We write SYMBOLS to describe these worlds.

 Matter has mass and occupies space  What is not matter? ◦ Everything in the universe is either matter or energy

Matter Pure Substance Mixture ElementCompoundHomogeneous Heterogeneous IronCO 2 Juice Trail Mix

 Element ◦ Cannot be converted to a simpler form by a chemical reaction. ◦ Example ◦ Examplehydrogen and oxygen  Compound ◦ Combination of two or more elements in a definite, reproducible way. ◦ Example ◦ Examplewater - H 2 O

A combination of two or more pure substances. ◦ Homogeneous ◦ Homogeneous - Looks the same throughout ◦ Heterogeneous ◦ Heterogeneous - Does not look the same throughout  Which are homogeneous or heterogeneous? ◦ Blood Skittles “T-Bone” steak ◦ Orange Juice Vegetable Soup Salad Dressing

 Qualitative analysis  Qualitative analysis is data that is observed  Colors, textures, smells, tastes, appearance, etc  Quantitative analysis  Quantitative analysis is data that can be measured  Length, height, area, volume, mass, speed, time, temperature, humidity

 Physical properties  Physical properties can be measured or observed  colordensity  odormelting point  tasteboiling point  Chemical properties  Chemical properties describe matter’s ability to change into another substance  ability to burn  reactivity  ability to decompose

 Physical changes  Physical changes do not change identity of substance  tearingmelting  freezing boiling  grindingcutting  Chemical changes  Chemical changes change identity of substance  burning  reacting  combusting

 Extensive properties Depend on the quantity of sample measured. Example Example - mass and volume of a sample.  Intensive properties Independent of the sample size. Properties that are often characteristic of the substance being measured. Examples Examples - density, melting and boiling points.

Method to express really big or small numbers. ◦ Format isMantissa x Base Power We just move the decimal point around Decimal part of original number Decimals you moved

 If a number is larger than 1 The original decimal point is moved X places to the left. The resulting number is multiplied by 10 X. The exponent is the number of places you moved the decimal point = 1.23 x 10 8

 If a number is smaller than 1 The original decimal point is moved X places to the right. The resulting number is multiplied by 10 -X. The exponent is the number of places you moved the decimal point = 1.23 x 10 -7

 Most calculators use scientific notation when the numbers get very large or small.  How scientific notation is displayed can vary.  It may use x10 n  or may be displayed  using an E.  They usually have an Exp or EE ◦ This is to enter in the exponent.

 English units ◦ Still commonly used in the United States  Examples: pound, inch, foot, cup, pint  Why English system not used in chemistry ◦ Very confusing and difficult to keep track of the conversions needed ◦ Vary in size so you must memorize many conversion factors

 Changing the prefix alters the size of a unit. Prefix Symbol Factor megaM kilok hectoh dekada base decid centic millim microµ nanon picop

System International  SI - System International - systematic subset of the metric system  Physical QuantityNameAbbreviation Mass kilogramskg Lengthmetersm Timesecondss TemperatureKelvinK Amountmolemol Electric CurrentAmpereA Luminous Intensitycandelacd

 Give you ability to convert between units  Problem solving technique (factor label method) 1. Write what you know 2. Game plan 3. Set up units 4. Conversion factors 5. Solve

 Convert 26 gallons to cups ◦ Answer: 416 cups  Convert 18 miles to centimeters ◦ Answer: 2.9×10 6 cm

 Allow you to convert between metric prefixes  Write what you know  Set up units  Bigger unit gets the “1”  Smaller unit is 10 x where “x” is how many places apart the two units are  Example ◦ Convert.25 kg to mg.  Answer: 2.5x10 5 mg

 Ratio of mass to volume of matter  Common units are g / cm 3 or g / mL.  Example: what is the density of 5.00 mL of a fluid if it has a mass of 5.23 grams?  d = mass / volume  d = 5.23 g / 5.00 mL  d = 1.05 g / mL  Example 2: What would be the mass of 1.00 liters of this sample? Density = Mass Volume cm 3 = mL

 Mass  Mass - the quantity of matter in an object.  Weight  Weight - the effect of gravity on an object.  Since the Earth’s gravity is relatively constant, we can interconvert between weight and mass. kilogram (kg) gram (g)  The SI unit of mass is the kilogram (kg). However, in the lab, the gram (g) is more commonly used.

 Volume  Volume - the amount of space that an object occupies. liter (L) The base metric unit is the liter (L). milliliter (mL) The common unit used in the lab is the milliliter (mL). cm 3 One milliliter is exactly equal to one cm 3. SIm 3 The derived SI unit for volume is the m 3 which is too large for convenient use.

 Method used to express accuracy and precision.  You can’t report numbers better than the method used to measure them units = three significant figures

 The number of significant digits is independent of the decimal point These numbers All have three significant figures!

are not Leading zeros are not significant. Leading zero Captive zero Trailing zero three significant figures four significant figures five significant figures are Zeroes between non-zeros are significant. areONLY IF Trailing zeros are significant ONLY IF there is a decimal point in the number.

 How many significant figures are in the following? m  5 sig figs kg  4 sig figs 100,890 L  5 sig figs 3.29 x 10 3 s  3 sig figs cm  2 sig figs 3,200,000  2 sig figs

g g g g g 83.7 g  Addition and subtraction  Report your answer with the same number of digits to the right of the decimal point as the number having the fewest to start with.

 Multiplication and division. Report your answer with the same number of digits as the quantity have the smallest number of significant figures.  Example. Density of a rectangular solid kg / [ (18.5 m) ( m) (2.1m) ] = 2.8 kg / m 3 (2.1 m - only has two significant figures)

 After calculations, you may need to round off 5 or more up ◦ If the first insignificant digit is 5 or more, you round up 4 or less down ◦ If the first insignificant digit is 4 or less, you round down

 A properly written number in scientific notation always has the proper number of significant figures = 3.21 x Three Significant Figures Three Significant Figures

In science, all of our numbers are either measured or exact. Exact Exact - Infinite number of significant figures. 1 foot = exactly 12 inches Do not count toward significant figures Measured Measured - the tool used will tell you the level of significance and varies based on the tool.

 When using a measuring tool, record the numbers that are certain and add a guess number

 Systematic Errors in a single direction (high or low). Can be corrected by proper calibration or running controls and blanks.  Random Errors in any direction. Can’t be corrected. Can only be accounted for by using statistics.

 All measurements contain some uncertainty. We make errors Tools have limits Accuracy AccuracyHow close to the true value Precision PrecisionHow close to each other Neither accurate nor precise Precise but not accurate Precise AND accurate

◦ Instrument not ‘zeroed’ properly ◦ Reagents made at wrong concentration ◦ Temperature in room varies ‘wildly’ ◦ Person running test is not properly trained Random Systematic