Chapter 1 Tools of Chemistry

1.1 Study of Chemistry O _____________ R _____________ I _____________ Observation Representation Interpretation Macroscopic Observations Microscopic Interpretations

Law Modify Observations Theory (Model) Hypothesis Experiment 1.1 Scientific Method Observations Theory (Model) Hypothesis Prediction Modify Experiment Experiment Law COMMUNICATE

1.2 Classification of Matter

1.2 Matter: Definitions Substance Mixture Solution Homogenous mixture Heterogeneous mixture Elements Compounds: Law of Definite Proportions (definite composition)

1.2 Elements If a pure substance cannot be decomposed into something else, then the substance is an element. There are 114 elements known. Each element is given a unique chemical symbol (one or two letters). Elements are building blocks of matter. The earth’s crust consists of 5 main elements. (O, Si, Al, Fe, Ca) The human body consists mostly of 3 main elements. (O, C, H)

1.2 Elements Chemical symbols with one letter have that letter capitalized (e.g., H, B, C, N, etc.) Chemical symbols with two letters have only the first letter capitalized (e.g., He, Be). C Cu Na U

1.2 Pure Substances and Mixtures Atoms consist of only one type of element. Molecules can consist of more than one type of element. Molecules can have only one type of atom (an element). Molecules can have more than one type of atom (a compound). If more than one atom, element, or compound are found together, then the substance is a mixture.

1.2 Pure Substances and Mixtures If matter is not uniform throughout, then it is a heterogeneous mixture. If matter is uniform throughout, it is homogeneous. If homogeneous matter can be separated by physical means, then the matter is a mixture. If homogeneous matter cannot be separated by physical means, then the matter is a pure substance. If a pure substance can be decomposed into something else, then the substance is a compound.

1.2 Compounds Most elements interact to form compounds. The proportions of elements in compounds are the same irrespective of how the compound was formed. Law of Constant Composition (or Law of Definite Proportions): The composition of a pure compound is always the same.

1.2 Substances If water is decomposed, then there will always be twice as much hydrogen gas formed as oxygen gas. Pure substances that cannot be decomposed are elements.

1.2 Mixtures Heterogeneous mixtures are not uniform throughout. Homogeneous mixtures are uniform throughout. Homogeneous mixtures are called solutions.

1.3 Separation of Mixtures Mixtures can be separated if their physical properties are different. Filtration: Solids can be separated from liquids by means of filtration. The solid is collected in filter paper, and the solution, called the filtrate, passes through the filter paper and is collected in a flask. Evaporation Sieving Magnetic

1.3 Separation of Mixtures Distillation: Distillation requires the different liquids to have different boiling points. In essence, each component of the mixture is boiled and collected. The lowest boiling fraction is collected first.

1.3 Separation of Mixtures: Distillation

1.3 Separation of Mixtures: Chromatography Chromatography can be used to separate mixtures that have different abilities to adhere to solid surfaces. (based on polarity) The greater the affinity the component has for the surface (paper) the slower it moves. The greater affinity the component has for the liquid, the faster it moves. Chromatography can be used to separate the different colors of inks in a pen.

1.3 Chromatography Separation

1.3 CHROMATOGRAPHY

1.3 Principles of Chromatography The order of peaks depends on few factors: Molar mass Attractive forces between the compound and substrate (adhesive forces) Attractive forces between the molecules of the compound (cohesive force) Solubility of the compound in the driving liquid 9/7

1.4 Periodic Table Periods: how many? How do we number them? Groups (families): how many? How do we number them? Metals? Which groups? How many in PT? Nonmetals? Which groups? How many? Metalloids? Noble gases? Lanthanides? Actinides? Relationship between quantum numbers and PT?

1.5 States of Matter Matter can be a gas, a liquid, or a solid. Gases have no fixed shape or volume. Gases can be compressed to form liquids. Liquids have no shape, but they do have a volume. Solids are rigid and have a definite shape and volume.

1.5 The Three States of Matter  

1.5 Pure Substances and Mixtures

1.5 States of Matter States of matter differ by ______________ (Distance of separation between particles (what about density?) Melting Boiling Condensation Evaporation Sublimation Deposition Freezing Mp, bp, fusion, condensation point

1.6 Physical and Chemical Properties of Matter Physical property can be ___________ Examples: Chemical property can be __________ Physical change __________________ Chemical change__________________ Differences between physical and chemical properties:_________________

1.6 Properties of Matter Extensive property_________ Examples: Intensive property _________

1.6 Physical and Chemical Properties Intensive physical properties do not depend on how much of the substance is present. Examples: density, temperature, and melting point. Extensive physical properties depend on the amount of substance present. Examples: mass, volume, pressure.

1.6 Physical and Chemical Changes When a substance undergoes a physical change, its physical appearance changes. Ice melts: a solid is converted into a liquid. Physical changes do not result in a change of composition. When a substance changes its composition, it undergoes a chemical change: When pure hydrogen and pure oxygen react completely, they form pure water. In the flask containing water, there is no oxygen or hydrogen left over.

1.6 Physical and Chemical Changes

1.7 Measurements Calculate the diameter of water molecule (assume that water molecule is spherical). How would you do it? Assumptions for calculations? Macroscopic (direct) measurements Microscopic (indirect) measurements

1.7 Measurements: SI Units There are two types of units: fundamental (or base) units; derived units. There are 7 base units in the SI system.

1.7 Measurements: Fundamental Base Quantity Name of Unit Symbol Length Mass Time Electrical Current Temperature Amount of Substance Luminous Intensity

1.7 Measurements Derived Units: _____________________ Examples:________________________ Density Volume Area Weight Pressure Velocity

Derived Unit: Volume

1.7 Metric System: SI Units

1.7 SI Units Note the SI unit for length is the meter (m) whereas the SI unit for mass is the kilogram (kg). 1 kg weighs 2.2046 lb. Temperature There are three temperature scales: Kelvin Scale Used in science. Same temperature increment as Celsius scale. Lowest temperature possible (absolute zero) is zero Kelvin. Absolute zero: 0 K = -273.15 oC.

1.7 Temperature

1.7 Examples: An object with a mass of 17.95 g occupies a volume of 11.8 mL. What is its density? A sample with a density of 3.75 g cm-3 occupies a volume of 10.44 cm3. What is the mass of the sample? A graduated cylinder is filled with 15.0 cm3 of water. An object with a mass of 29.66 g causes the total volume to increase to 23.4 mL. What is the density of the sample?

1.8 Handling Numbers Significant Figures: Accuracy Precision Exact numbers: counting and given conversion factors Measured numbers: uncertainty due to instruments Interpret the meaning of the following measurements (uncertainties?): 12.0 m 0.001 mg 1.2 x 10-2 cm 1.43 x 103 g

1.8 Uncertainty in Measurement All scientific measures are subject to error. These errors are reflected in the number of figures reported for the measurement. These errors are also reflected in the observation that two successive measures of the same quantity are different. Precision and Accuracy Measurements that are close to the “correct” value are accurate. Measurements that are close to each other are precise.

1.8 Precision and Accuracy

1.8 Significant Figures The number of digits reported in a measurement reflect the accuracy of the measurement and the precision of the measuring device. All the figures known with certainty plus one extra figure (estimated digit) are called significant figures. In any calculation, the results are reported to the fewest significant figures (for multiplication and division) or fewest decimal places (addition and subtraction).

1.8 Significant Figures: Rules Non-zero numbers are always significant. Zeros between non-zero numbers are always significant. Zeros before the first non-zero digit are not significant. (Example: 0.0003 has one significant figure.) Zeros at the end of the number after a decimal place are significant. Zeros at the end of a number before a decimal place are ambiguous (e.g. 10,300 g).

1.8 Look for a Decimal Point Pacific Atlantic Present Absent Moving from the proper side, skip all zeros until you get to the first non-zero digit. All remaining digits, including zeros, are significant.

1.7 Significant Figures How many SF in the following? 1.23 x 10-5 m 0.0123 cm 7.801 x 10-4 mg 1.2 x 100 g/cm3 123,000 km 1.2 x 10-12 m 513 pencils

1.7 Significant Figures How many SF in the answers? 5.456 m + 0.6 m + 14.33 m 3.70 g – 2.99 g 3.45 cm x 12.345 cm (2.02 x 106 m) + ( 4.56 x 107 m)

1.8 Handling Numbers Rules for: Determining # of SF? ________________ Adding and subtracting SF? ___________ Multiplying and dividing with SF? _______ Rounding? _________________________ Complex calculations? ________________ For multi-step problems using the calculator, round at the END of the calculations.

1.8 Handling Numbers Scientific Notations: Without using calculators solve the following problems. Report your answers in scientific notations: 14.45 + (1.3 x 10-1) 88,000 / (2.2 x 102) 1.23 x 104 + 1.45 x 103 = 1.5 x 10-1 + 0.012 = 2.3 x 104 x 3.2 x 10-8 =

1.9 Dimensional Analysis Method of calculation utilizing a knowledge of units. Given units can be multiplied or divided to give the desired units. Conversion factors are used to manipulate units: Desired unit = given unit  (conversion factor) The conversion factors are simple ratios:

1.9 Factor Label Method Convert 6.6 m to miles Convert 12.5 cm3 to m3; to ft3 The density of iron is 7.8 g/cm3. Convert it to lb/in3