Introduction to Chemistry Chapter 1

The Nature of Science and Chemistry Chemistry is essentially the study of matter and its interactions Includes both physical and chemical changes and the energy flow associated with them

Chemistry and the Scientific Method Identify the problem Make Observations Hypothesis Experiment Draw a Conclusion

Parts of an Experiment Experiment: a set of controlled observations that test a hypothesis Consists of: Independent variable - the variable you plan to change Dependent variable - the variable that changes in response to a change in the independent variable Control - a standard for comparison Data – information/ observation recorded 1) Qualitative: uses descriptive words Example- colors 2) Quantitative: uses numbers

Observation vs. Inference is an experience perceived through one or more of the senses What you feel, smell, see, hear, or taste Example of observation The temperature of a flask goes up as a reaction proceeds Inference an interpretation of an observation which goes one step beyond an observation. The “story” you create about what you see, feel, hear, smell or taste which is based upon a direct observation

Qualitative vs. Quantitative Examples 1. Smells like apples 6. red 2. 100 lbs 7. 40 mph 3. 105 degrees C 8. round 4. Tall 9. cold 5. Pink with purple dots

Scientific Method Theory & Law a broad and extensively tested explanation of “WHY” experiments give results. Supported by many experiments Considered successful if it can be used to make predictions that are true Natural Law a concise statement that summarizes the results of many observations and experiments

Matter Matter is anything that has mass and takes up space (by definition) Weight is the force of attraction that results from an object possessing mass They are not synonymous

Classification of Matter Pure substance Has a constant composition It’s identified by a formula or symbol Examples water- H2O – always has 2 hydrogens and 1 oxygen Helium – He Compound Composed of two or more different types of atoms that cannot be decomposed by physical means Chemically combined Sodium chloride- NaCl Carbon dioxide- CO2

Classification of Matter Element Simple substance that cannot be decomposed into simpler substances by any chemical change

Mixtures of Matter A mixture is a combination of two or more pure substances The individual pure substances do not change their chemical properties, however There are two types of mixtures: Heterogeneous mixtures do not have the same consistency throughout Homogeneous mixtures do have the same consistency

Examples of Mixtures Solution A on the left is an example of a heterogeneous mixture Solution B on the right is an example of a homogeneous mixture

Completed Classification Flowchart Matter Pure Substances Elements Compounds Mixtures Heterogeneous Homogeneous

Can be divided into two categories: Properties of Matter PHYSICAL Properties: Can be measured or observed without changing the chemical makeup (composition) of the object being studied Changes: Changes in physical appearance that do not result in a change in composition CHEMICAL Properties: Describe chemical reactivity (cannot be observed without changes in composition) Changes: Akin to chemical properties. Occur as one substance is converted into a completely different substance Can be divided into two categories: Extensive Properties: Depend on the amount of sample Intensive Properties: Depend on what the sample is; not amount

Measurements and Uncertainty Accuracy involves comparison with an accepted value (Poor accuracy results in high error) Precision involves comparison between measurements of the same quantity (Poor precision results in high uncertainty)

Rules for Determining Significant Figures All nonzero digits are significant. Ex: 678.2 (4 sig. figs.) Zeros between significant digits are significant. Ex: 7009 (4 sig. figs.) Zeros at the beginning of a number are never significant. Ex: 0.00045 (2 sig. figs.) Zeroes at the end of a number are significant IF the number contains a decimal. Ex: 0.9700 (4 sig. figs.)

Number of Significant Figures? 0.709 0.4600 12000 12001 12000.00 0.00390900 See Example 1.3 (Pg. 13) 3 4 2 5 7 6

Significant Figures in Calculations “The operator must report the proper number of significant figures; the calculator will not.” More rules regarding significant figures (why not?) Addition/Subtraction Answers must be reported to the fewest number of decimal places Multiplication/Division Answers must be reported to the fewest number of significant figures

Example #1 What is the volume of an object (expressed to the appropriate number of sig. figs.) with the following dimensions: length = 1.2 m, width = 1.45 m, height = 0.678 m? See Example 1.5 (Pg. 16) Formula: L x W x H (1.2 m)(1.45 m)(0.678 m) = 1.17972 m3 = 1.2 m3

Example #2 A gas at 25 °C fills a container whose volume is 1.05 x 103 cm3. The container plus gave have a mass of 837.6 g. The container when emptied of gas, has a mass of 836.2 g. What is the density of the gas at 25 °C? See Example 1.5 (Pg. 16) Density = mass/Volume D = (837.6 - 836.2)g / 1.05 x 103 cm3 = (1.4 g) / 1.05 x 103 cm3 = 0.0013 g/cm3 OR 1.3 x 10-3 cm3

Quantities That Are Not Limited by Significant Figures Counted numbers Represent exact quantities (as opposed to measured quantities) Ex: Number of students present in the classroom, number of eggs in a carton, etc. Defined numbers Typically refers to conversion factors Ex: 1 foot = 12 inches; 1 yard = 3 feet, 1 minute = 60 seconds, etc. The power of 10 Scientific notation Ex: The “10” in 6.62 x 1023

Examples There are 67 people in this classroom There is no range between 66.1 and 67.0 people, there are exactly 67 people The thermometer shown here reads approximately 27 C. This is a measured quantity and must therefore contain uncertainty.

Measurements and Units Every measurement is made up of two components: a quantity and a unit The modern scientific community uses the SI (Le Système International d’Unités) system of units to express measurements Each fundamental quantity has its own base unit Everything else is expressed as a combination of units or derived units Ex: cm3, g/mL, g/mol

SI Base Units Some of the base units have prefixes that can be used for convenience Base Units in italics are not required for this course and may be ignored Time Second (s) Length Meter (m) Mass Kilogram (kg) Temperature Kelvin (K) Amount of a substance Mole (mol) Electrical Current Ampere (A) Luminous Intensity Candela (cd)

SI Prefixes Used for sake of convenience

Temperature Conversion Factors There are three common temperature scales Fahrenheit ( °F ) Celsius ( °C ) Kelvin ( K ) The conversion between the Celsius and Kelvin scales is the most important: TK = TC + 273.15

Celsius and Fahrenheit Conversions Celsius to Fahrenheit Conversion: Fahrenheit to Celsius Conversion: