AP Chemistry Chapter 1: Chemical Foundations. The only thing that matters is Matter!

AP Chemistry Chapter 1: Chemical Foundations The only thing that matters is Matter!

The Scientific Method 1. Observations (collecting data) -quantitative or qualitative 2. Formulating hypothesis - possible explanation for the observation 3. Performing experiments - gathering new information to decide whether the hypothesis is valid

Outcomes Over the Long-Term Theory (Model) - A set of tested hypotheses that give an overall explanation of some natural phenomenon. Natural Law - The same observation applies to many different systems - ex: Law of conservation of mass

Law vs. Theory A natural law summarizes what happens Law of gravity A theory (model) is an attempt to explain why it happens. Einstein's theory of gravity describes gravitational forces in terms of the curvature of spacetime caused by the presence of mass

Nature of Measurement A quantitative observation is a measurement. A measurement consists of both a number and a unit. Examples: 20 grams 6.63 x 10-34 J s

International System of Units The International System of Units, symbolized SI, is the modern version of the metric system Seven SI Units: Quantity Unit Symbol length meter m mass kilogram kg time second s electric current ampere A temperature kelvin K amount of substance mole mol luminous intensity candela cd

SI Units

SI Prefixes Common in Chemistry

1 Liter = 1dm 3 = (10 cm) 3 = 1000cm 3 Since 1 cm 3 = 1 ml, 1L= 1000ml

Precision and Accuracy Accuracy refers to the agreement of a particular value with the true value. Precision refers to the degree of agreement among several measurements made in the same manner. Neither accurate nor precise Precise but not accurate Precise AND accurate

Types of Error Random Error (Indeterminate Error) measurement has an equal probability of being high or low. Systematic Error (Determinate Error) Occurs in the same direction each time (high or low), often resulting from poor technique or incorrect calibration. This can result in measurements that are precise, but not accurate.

Uncertainty in Measurement The volume of a buret is read at the bottom of the liquid curve (meniscus). Meniscus of the liquid occurs at about 20.15 ml. Certain digits: 20.15 Uncertain digit: 20.15

Uncertainty in Measurement A digit that must be estimated is called uncertain. A measurement always has some degree of uncertainty. Measurements are performed with instruments No instrument can read to an infinite number of decimal places

Significant Figures Significant Figures: all the digits that can be known precisely in a measurement, plus one last estimated (uncertain) digit To determine if a figure is significant, you need to follow the rules! 4 5 9 2 7 9 8 4 0 3

1. All non-zero integers are always significant 341 = 3 Sig Figs 2. All trapped zeros are always significant 7003 = 4 Sig Figs 3. Leading zeros are NEVER significant. 0.0071 = 2 Sig Figs 4. Trailing zeros are ONLY significant when there is a DECIMAL in the number. 43.00 = 4 Sig Figs

5. Zeros that are placeholders at the end of a number are NOT significant. 6. Unlimited number of sig. figs: 1. Counted objects Ex: 36 students in the class. 2. Defined or Exact quantities Ex: 60 minutes = 1 hour 300 = 1 Sig. Figs

Sig Fig Practice #1 How many significant figures in each of the following? 1.0070 m 5 sig figs 17.10 kg 4 sig figs 100,890 L 5 sig figs 3.29 x 10 3 s 3 sig figs 0.0054 cm 2 sig figs 3,200,000 mol 2 sig figs

Significant Figures in Mathematical Operations Multiplication and Division: The answer is rounded to the same number of significant figures as the measurement with the least significant figures. 6.38 x 2.0 = 12.76 13 (2.0 ONLY has 2 sig figs; therefore, round final answer to 2 sig figs)

Sig Fig Practice Calculation Calculator says: Answer 3.24 m x 7.0 m 22.68 m 2 23 m 2 100.0 g 23.7 cm 3 4.219409283 g/cm 3 4.22 g/cm 3 0.02 cm x 2.371 cm 0.04742 cm 2 0.05 cm 2 710 m 3.0 s 236.6666667 m/s 240 m/s 1818.2 lb x 3.23 ft 5872.786 lb ft 5870 lb ft 1.030 g 2.87 ml 2.9561 g/ml 2.96 g/ml

Rules for Significant Figures in Mathematical Operations Addition and Subtraction: Display the final answer with the same number of decimal places as the least precise measurement used in the calculation. 6.8 + 11.934 = 18.734 18.7 (6.8 goes to the tenth; therefore, round to the tenth: 3 sig figs)

Sig Fig Practice #3 Calculation Calculator says: Answer 3.24 m + 7.0 m 10.24 m 10.2 m 100.0 g - 23.73 g 76.27 g 76.3 g 0.02 cm + 2.371 cm 2.391 cm 2.39 cm 713.1 L - 3.872 L 709.228 L 709.2 L 1818.2 lb + 3.37 lb 1821.57 lb 1821.6 lb 2.030 ml - 1.870 ml 0.16 ml 0.160 ml

Concept Check You have water in each graduated cylinder shown. You then add both samples to a beaker (assume that all of the liquid is transferred). How would you write the number describing the total volume? 2.8 + 0.28 = 3.08 = 3.1 ml What limits the precision of the total volume?

Units of Temperature: Kelvin, Celsius, Fahrenheit The lowest theoretical temperature possible where all motion of particles stop is O K or absolute zero. K = o C + 273.15

Density Density: The ratio of the mass of an object to its volume. Common units are g/cm 3 or g/ml Density = Mass Volume

Example: A student determines that a piece of metal has a volume of 285 ml and a mass of 612 g. Is the shiny piece of metal Aluminum, which has a density of 2.70 g/ml? NO D= m v = 612 g 285 ml = 2.147 = 2.15 g/ml

Water has a density of 1.00 g/cm 3 at 4 0 C Materials that have a density lower than 1 g/cm 3 will float in water. Materials that have a density greater than 1 g/cm 3 will sink in water

Density generally decreases as its temperature increases. Water is an exception Ice has a density of.917 g/cm 3 at 0 0 C. Water has a density of 1.00 g/cm 3. That s why ice floats!

DA Covered while reviewing summer worksheets.

Square and Cubic units Use the conversion factors you already know, but when you square or cube the unit, don t forget to cube the number also! Best way: Square or cube the ENITRE conversion factor Example: Convert 4.3 cm 3 to mm 3 ( ) = 4.3 cm 3 10 3 mm 3 4.3 cm 3 10 mm 3 1cm = 4300 mm 3 1 3 cm 3

Classification of Matter Anything occupying space and having mass. Matter exists in three main states. Solid definite volume & shape Liquid definite volume, indefinite shape Gas indefinite volume & shape

Volume Properties of Matter Extensive properties depend on the amount of matter that is present. Mass Energy Content (think Calories!) Intensive properties do not depend on the amount of matter present. Melting point Boiling point Density

Matter consists of atoms and molecules in motion. Kinetic Nature of Matter

OTHER STATES OF MATTER PLASMA an electrically charged gas; Example: the sun or any other star BOSE-EINSTEIN CONDENSATE a condensate that forms near absolute zero that has superconductive properties; Example: supercooled Rb gas

Mixtures Have variable composition. Homogeneous Mixture Having visibly indistinguishable parts; solution. Heterogeneous Mixture Having visibly distinguishable parts.

Homogeneous Mixtures

Homogeneous vs. Heterogeneous Mixtures

Compound vs. Mixture

Physical Change Change in the form of a substance, not in its chemical composition. A physical change will not break up compounds Example: boiling or freezing water Distillation Filtration Chromatography

Separation of Mixtures Physical means can be used to separate a mixture into its pure components. Ex: dyes such as ink may be separated by paper chromatography.

magnet distillation 1.4

Chemical Change A given substance becomes a new substance or substances with different properties and different composition. Example: Bunsen burner (methane reacts with oxygen to form carbon dioxide and water)

Separation of a Compound The Electrolysis of water Compounds must be separated by chemical means. With the application of electricity, water can be separated into its elements Reactant Products Water Hydrogen + Oxygen 2 H 2 O 2 H 2 + O 2

Physical vs. Chemical Change Physical changes do not result in new substances. Chemical changes result in NEW substances Mrs. Kalmer is the one on the left...