Chapter 1 Introduction: Matter and Measurement Chapter 1: Matter and Measurement Matter Classification of Matter Properties Measurement Units (SI, metric) Significant Figures Unit Conversions What is Matter? Matter is anything that has mass and occupies space. Classification of Matter Why do we classify matter? Why do we classify anything? Classification of Matter Pure Substances Elements Compounds Molecules Mixtures Homogeneous Heterogeneous Elements Elements are composed of atoms of only one type of element. Examples: Iron metal (Fe) Sulfur (S or S8) Oxygen (O2) Bromine (Br2) Elements Know the element names and symbols for the first 36 elements. We will add a few more to this list in chapters 2 and 3. Compounds Compounds are combinations of atoms of more than one type of element. Examples: Iron(III) sulfide (Fe2S3) Methane (CH4) Water (H2O) Hydrogen peroxide (H2O2) Pure Substances Elements and compounds have definite compositions, and each has a set of properties that are unique. Molecules Molecules are groups of atoms bonded chemically. Molecules can be elements or compounds H2 CO2 O2 CH3CH2OH Elements and Compounds Molecules Molecules can be Elements or Compounds mixtures elements compounds Classify: A C B D Classification of Matter Mixtures 2 or more pure substances that exist together but are not combined chemically. Components of a mixture can be separated by physical processes. Distillation separates water (H2O) from salt (NaCl) Mixtures Homogeneous Mixtures (solutions) Constant composition throughout air, salt water, alloys Heterogeneous Mixtures non-uniform composition potting soil, oil and water, iron and sulfur States of Matter (Phases) Substances can exist in the solid, liquid, or gas state. Microscopic View of the 3 States of Water Phase Changes (Changes in Physical State) Boiling, Evaporating Condensing Melting Freezing These are physical changes. liquid solid gas Sublimation Chemical Change Physical Change No change in identity of substances eg. H2O(l) H2O(g) Chemical Change (Chemical Reaction) A substance is changed into a different substance eg. 2H2O 2H2 + O2 Chemical Changes Rusting 2Na + Cl2 2NaCl VD02_007.mov 2H2 + O2 2H2O Properties Types of Properties Chemical Physical Intensive Extensive Properties Identify the following as chemical or physical properties. Melting point of aluminum Ability of nitric acid to dissolve copper Density of mercury Boiling point of water Instability of hydrogen in a balloon Properties Identify the following as chemical or physical properties. Melting Point of aluminum Ability of nitric acid to dissolve copper Density of mercury Boiling Point of Water Instability of a hydrogen in a balloon Intensive and Extensive Properties Extensive Properties depend upon the amount of substance present Amount of heat needed to melt gold Volume of water Mass of oxygen gas Intensive Properties are independent upon the amount of substance. (Useful for identification) Melting Point of aluminum Boiling point of ethanol Density of mercury Measurements - are used to communicate information. A measurement consists of a number and an appropriate unit. Examples: 125C; 10.0 kg; 10.0 mg 1.5 yards; 1.004 meters 0.936 g/mL; 10cm3 Measurement In Chemistry What is the size of an atom? of Earth? We use exponential notation to express such numbers conveniently. Powers of Ten Review of Exponential Notation Appendix A.1 Exponential Notation Metric Units - SI Units Metric units are based on powers of 10. SI units are a set of metric units used preferentially in scientific measurements (International System). Base Units (T 1.4) Prefixes (T 1.5) SI Units (Metric Units) Base Units (T 1.4) Base Units (Table 1.4) Mass: kilogram (kg) Length: meter (m) Time: second (s, sec) Temperature: Kelvins (K) Electric Current: ampere (A) Luminous Intensity: Candela (cd) Amount of Substance: mole (mol) SI Units (Metric Units) Prefixes Prefixes Example: Length measurements Giga - G 109 1Gm=109 m Mega- M 106 1Mm=106 m Kilo- k 103 1km=103 m Centi- c 10-2 1cm=10-2 m Milli- m 10-3 1mm=10-3 m Micro- 10-6 1m = 10-6 m Nano- n 10-9 1nm=10-9 m Pico- p 10-12 1pm=10-12 m SI Units (Metric Units) Prefixes Relationships between prefixes: 1cm = 10-2 m or 102 cm = 1 m 1mm = 10-3 m or 103 mm = 1 m 1m = 10-6 m or 106 m = 1 m 1nm = 10-9 m or 109 nm = 1 m 1pm = 10-12 m or 1012 pm = 1 m SI Units (Metric Units) Temperature The Kelvin (K) temperature scale is an absolute scale (its zero value is the lowest possible temperature). How are these temperature scales related? SI Units (Metric Units) Temperature Temperature Relationships Kelvin(K) = ºC + 273.15 ºC = 5/9 (ºF - 32) Common Temperatures you should know in K and ºC: Normal Freezing Point of Water - Normal Boiling Point of Water - Lowest Possible Temperature - SI Units (Metric Units) Derived Units SI-Derived Units Volume: (length x length x length) 1 mL = 1 cm3 Volumetric Glassware SI Units (Metric Units) Derived Units Density Units: (mass/Volume) kg/m3 g/cm3 g/mL g/L Pb Which has a greater density, lead or aluminum? Al SI Units (Metric Units) Derived Units Density Units: (mass/Volume) kg/m3 g/cm3 g/mL g/L Pb Which has a greater density, lead or aluminum? Al 11.34 g/cm3 2.70 g/cm3 Carbon dioxide gas is more dense than helium gas. Why? CO2 He Density = Mass / Volume Water is more dense than ice. Why? ICE WATER Density = Mass / Volume The number of significant digits describes the exactness of the measurement. To record a measurement to the proper # sig. figs, estimate between the tick marks. The last digit in your measurement should have some uncertainty. Significant Figures Rules for Determining # Significant Figures - p.22 Each of the following have three significant figures. Explain why. 435 g 405 g 40.5 g 5.00 g 0.151 g; 0.00405 g ** If in doubt, write the number in scientific notation Significant Figures How many significant figures in 1800 mg? Significant Figures and Mathematical Operations Adding and Subtracting The final answer is expressed with the same number of decimal places as the measurement with the fewest decimal places. Example: 89.332 cm +1.1 cm 90.432 cm Round to 1 decimal place: 90.4 cm Significant Figures and Mathematical Operations Multiplying and Dividing The final answer is expressed with the same number of significant figures as the measurement with the fewest significant figures. Example: 4.4643 cm x 2.8 cm 12.50004 cm2 Round to 2 sig. figs: 13 cm2 Significant Figures and Mathematical Operations Conversion factors are often exact numbers and are not measurements. Examples: 12 inches = 1 foot 1000 mm = 1 meter 60 sec = 1 minute Example: (4.55 feet)x(12 in./ft) =54.6 inches How can you use units to solve problems? Ratios Dimensional Analysis (Factor Label Method) Unit Conversions Example: Convert 45.0 seconds to minutes Convert 416 nm to meters Unit Conversions Example: Convert 416 nm to meters Answer: 4.16x10-7 m Unit Conversions Convert 65 miles per hour to meters per second. (1 mile = 1.6093 km) Unit Conversions: Units and SF Practice! You do not need to memorize any of the English to metric conversion factors, such as 2.54 cm/inch. You do need to know the metric (SI) conversion factors in T 1.5. You do need to know how to convert between Celsius and Kelvin. Unit Conversions