Chapter 1Chemical Foundations CHY 115: General Chemistry I
Chapter Outline • Overview of chemistry • Methods of science • Measurements • Unit systems • Quantities measured • Taking measurements • Accuracy and precision, types of error • Significant figures and calculations • Dimensional analysis • Temperature • Density • Classification of matter
Chemistry • Chemistry – study of the matter of the universe and the changes that this matter undergoes • Matter – anything that has mass and occupies space • Examples of matter and “non-matter”
Chemistry • What chemists study about the matter: • Composition • Structure • Properties • Chemical changes the matter will undergo • Relationship between the matter and: • The environment • Human health
The Scientific Method (1.2) • The scientific method describes a framework by which science is conducted. • Scientific method has been described as organized common sense. • Methodical approach to problem-solving.
Scientific Method • Observations questions • Develop hypotheses to explain observations • Test hypotheses • Make predictions, if hypothesis is true then…. • Test accuracy of the prediction • Repeat the process, test new hypotheses… • Analyze results and share findings and conclusions
Scientific Inquiry • Observations Questions • Observations must be recordable and repeatable • Observations may be: • Quantitative or • Qualitative
Theories • After significant research theories/models are developed to explain the observations • Theory – well-tested explanation of some part of nature that explains a broad range of observations • Supported by significant data • Subject to rigorous testing and revision when needed
Natural Law versus Theory • Natural Law • Statement/summary of observed behavior • Law of conservation of matter • Theory (also called a model) • Explanation of observed behavior based on significant data • Theories attempt to explain laws
Factors Impacting Scientific Inquiry • The direction of scientific research is impacted by many factors: • Theories and technology of the day • Money • Religion • Politics • World conditions
Units of Measure (1.3) • Much of chemistry is based on analysis of quantitative observations • A quantitative observation is obtained by measurement and includes a number and a unit.
Unit Systems • English system • Used in United States • Metric system • Used in science • International system (SI) • Based on the metric system UNITS MATTER, see page 9!
Unit Systems • English System • Used in U.S. • Little logic to the units • Examples
Unit Systems • Metric System • Developed in the late 1700’s and adopted after the French Revolution • A base (or fundamental) unit is defined for each quantity measured • The size of the base unit can be modified by adding a prefix
Unit Systems - metric Quantity Base unit Symbol
Unit Systems - meteric • Metric Prefixes, see page 10
Unit Systems • Using prefixes • Base unit = meter • Kilometer = km = _________ m
Unit Systems • International System (SI) • Adopted in 1960 • Internationally agreed upon set of units • Used in industry and science • See page 9
What chemists measure • Length – distance between 2 points • Metric base = _____________ • Mass – quantity of matter present • Base unit: __________ • Weight – measure of gravitational pull on an object • Base unit: ___________
Mass and Weight • Measure mass on a balance. • Measure weight on a scale.
Volume • Volume – amount of three dimensional space occupied by an object • SI base = meter3 • Metric base = Liter = dm3 • dm = ______ cm • dm3 = ___________ cm3 • Equivalent units: • mL = _______ = ________
Uncertainty in Measurement (1.4) • All measurements include some degree of uncertainty • A properly taken measurement includes all of the certain digits and one uncertain (estimated) digit
Taking measurements • When taking a measurement you record: • All known digits • those marked on the measuring device • One estimated digit • A multiple of 1/10 the smallest marked unit on the measuring device
Taking measurements • Graduated cylinder example Thermometer Example – on board
Accuracy and Precision • Accuracy – how close a measured value agrees with the true value • Ideally values will differ in only the estimated digit • Precision – how closely repeated measurements agree with each other • Ideally the values will differ in only the estimated digit
Types of Errors • Random error • Value has an equal probability of being high or low • Compensate for random errors by averaging multiple sets of data • Systematic error • Value recorded is consistently low or high • Compensate for systematic errors by…
Evaluating a measuring device • Good measuring devices are both accurate and precise • Readings taken with a precise, but not accurate measuring device can be corrected if the error is systematic.
Significant Figures and Calculations (1.5) • A measurement includes all the certain digits and one estimated (uncertain) digit • These digits are called the significant figures of a measurement. • All calculations based on measurements must reflect the uncertainty of the original measurements.
Significant Figures • Rules for counting significant figures • Rules for rounding off calculations based on significant figures
Counting Significant Figures • All nonzero integers are significant. • 35.76 g = _______ sig. fig.
Counting Significant Figures • Zeros • Leading zeros are NEVER significant • 0.0037 mL = _____ sig fig • Captive zeros are ALWAYS significant • 7.098 g = ______ sig. fig.
Counting Significant Figures • Zeros • Trailing zeros are significant only if the value includes a decimal place. • 3570 g = ______ sig. fig. • 7.500 kg = ______ sig. fig.
Counting Significant Figures • Exact Numbers have unlimited significant figures • Numbers obtained by counting • 23 students • Definitions • 1 foot = 12 inches • 1 inch = 2.54 cm (exactly)
Rules for Rounding Off • If the first digit to be removed is: • Less than 5, the preceding digit remains the same • 5 or greater, the preceding digit is increased by 1
Significant Figures and Calculations • Multiplication and Division • The answer is rounded to the same number of sig. fig. as the measurement with the fewest sig. fig. (3.50 x 102 mL) x 0.7030 g/mL = • How to enter #s in scientific notation on your calculator
Significant Figures and Calculations • Addition and Subtraction • The answer is rounded to the same number of decimal places as the measurement with the fewestdecimal places. 32.05 g + 5.3978 g + 6.30 g =
Dimensional Analysis (1.7) • Convert the number of minutes left in class to seconds. • Open to back cover of textbook. • PRACTICE!
Dimensional Analysis • The largest pumpkin at the 2012 Windsor fair weighed 1094 pounds. • Express the mass of this pumpkin in grams and in kg.
Dimensional Analysis • A block of wood has a volume of 2.50 ft3. • Express the volume of the wood block in cm3.
Dimensional Analysis • The world's oceans have a surface area of 361,100,000 square kilometers. • Express this surface area in square miles. Please put your final answer in scientific notation.
Dimensional Analysis • Water has a density of .998 g/mL at room temperature. • Express the density of water in pounds/gallon. • Game plan?
Temperature (1.8) • Temperature Scales (units) • Fahrenheit (0F) • Used in this country • Celsius (0C) • Used in the physical science • Kelvin (K) • SI unit for temperature • Used in gas law calculations
Temperature Conversions TK = TC + 273.15 TC = (TF - 32) OR… 5/9 (TF - 32) 1.8 TF = (1.8) TC + 32 …OR… 9/5x TC + 32
Temperature Conversions • Express 68.20 F in in 0C and in K • Express 78.50 C in 0F. • Express room temperature in K.
Density (1.9) • Density = mass of an object volume of object • Density is a physical property that is often used to identify an object (along with bp and mp)
Density • Density = mass of an object volume of object • Mass - is expressed in grams • Volume is expressed in: • mL or cm3 for solids and liquids • L for gases
Density • Density units: • Solids and Liquids • g/mL or g/cm3 • Gases • g/L • English System • Pounds/foot3