Chemistry of Our Environment Notes. Chapter 1

Transcription

1 Chemistry of Our Environment Notes Introduction: Purpose of Course: 1. Meet Lib Ed Requirement 2. Fire Science & Dental Hygiene How we will meet both needs. Exams & Grading Homework Assignments Chapter 1 Reason for studying chemistry: Essentially effects all aspects of our lives. Study helps us to understand this and make rational decisions for today and for the future. We will learn some basic chemistry principles so that we can talk with knowledge later in the semester about relevant topics for our work and our lives. Basic or theoretical Science: The concepts that form the foundation of our scientific knowledge. This is essentially discovered through pure research, many times without a specific end goal in mind. Applied Science: Research towards the solution of a particular problem. In industry today, most research is applied research. Technology: Using basic science and/or applied science to produce products of practical application, frequently for mass production. Scientific Method: 1. Observations 2. Hypotheses 3. Experimental Testing 4. Theories And Laws Chemistry is the study of the structure and interaction of matter. Matter is anything that occupies space and has mass. Although sometimes considered the same as weight, mass does not change with location. Weight depends on gravity. Chemists study matter in quantities large enough to be seen and handled in an ordinary manner (which we call macroscopic), or so small to be only seen through a microscope (which we call microscopic). Finally, chemistry did not become a modern science until chemists realized there were particles that could not even be seen through microscopes (submicroscopic), and studied their effects on how matter behaves. At the macroscopic level: 3 Types of matter:

2 2 1. Elements Basic substances in nature. Cannot be decomposed into simpler substances. 2. Compounds 2 or more elements, chemically combined together in a specific way, and always the same way. Can only be chemically decomposed into other smaller compounds or elements. Compounds have completely different properties than the elements that they contain. a) Pure Substance (your book simply uses the term substance) Any pure element or compound 3. Mixtures 2 or more pure substances physically mixed together in any way and not always the same way. Frequently easy to separate into its components. The components retain some or all of their own properties. At the submicroscopic level: 2 Particles: (for now) 1. Atoms Smallest particle identifiable as an element. 2. Molecules 2 or more atoms combined together chemically. Smallest particle that can be identified as a compound. Mixtures can further be divided into 2 types: 1. Homogeneous All parts of the mixture are identical, even at the microscopic level. All homogeneous mixtures are also called solutions. Some common examples are clean air, sugar water, brass, soda water, Scotch. a) Also, all elements and compounds, by definition, are homogeneous. 2. Heterogeneous Different parts have different composition. In many cases these different parts can be obviously seen ( cinnamon & sugar, basket of different fruit, metal ores), and sometimes only under special conditions (dusty air, milk, blood) There are 3 forms or phases of Matter: 1. Solid Holds its shape and size no matter what container it is in. 2. Liquid Holds its size but takes shape of its container. 3. Gas Takes shape & size of its container. Any substance can be (under the proper conditions) a solid, liquid or gas. Think of ice, water and steam. They are all different states for the substance water.

3 3 Chemists study matter by studying its identity and behavior. In doing so, they look at the properties of matter. 2 Main Classifications: 1. Physical properties (or characteristics) that can be studied without changing the substance into a new substance or trying to do so. Examples are color, physical state (solid, liquid or gas) at specific temperatures, and something called density. Density defined as the mass (or weight) in a specific volume of space for that substance. For example, water has a mass of 1.0 grams in every 1.0 milliliter. This is a much better way to identify a substance than using mass or volume separately. We will look at this in one of our experiments. Usually reported as 1.0 grams/ milliliter. When we study physical properties we frequently do change the substance, but only its appearance not its identity. For example, the temperature a substance freezes at is called its Freezing Point or FP. To measure FP, we simply have to freeze a liquid sample of the substance and measure the temperature this occurs at. We have only changed the state from liquid to solid, but the substance is the same. (Think of ice and water). These types of changes are called Physical Changes. 2. Chemical: Properties that can only be studied by changing or trying to change the substance into a new, different substance. For example to determine whether a substance is flammable, we take a small sample and burn it or at least try to burn it. If it burns, a chemical reaction occurs forming totally different substances. If it does not burn we have determined that it is not flammable. This is still a chemical property. In studying chemical properties we perform Chemical Changes on the substances. Reactants Products for example: Hydrogen + Oxygen water All Chemical & some physical changes also involve energy changes. Energy is the ability to do work (move objects). 3 Types for now: 1. Potential stored energy for later use (from position or chemical make-up) 2. Kinetic Energy of motion 3. Heat Energy is frequently converted from one type to another. Examples: Burning oil to produce heat (chemical potential heat energy) or turning on a light switch (electrical light energy). Temperature is the measure of intensity (how hot or cold) of heat.

4 4 There are 112 known elements (20 of which are so called man-made elements (we will explain this later in the semester hopefully). Each element has a 1 or 2 letter symbol to represent it. A chemist needs to know the names that go with the symbols and vice-versa. We will learn some of them in this course. Compounds also have symbols based on the symbols of the elements composing the compound. Examples: NaCl, NH 4 NO 3, Mg(OH) 2. These are called molecular formulas. Later in the semester we will learn about other ways to represent the formulas of compounds. Besides studying the identity and behavior of matter in a descriptive sense (qualitative analyses), chemistry is also very much of a quantitative science, involving many measurements and calculations based on these measurements. For measurements all scientists use the Metric System (actually the International System (SI))rather than the English system. For our purposes we will consider the classical Metric System. The differences between Metric and SI are small and not of great importance at this level. There are many advantages of Metric over English. A few of these are: 1. Simpler - Many fewer words and meanings to learn 2. More logical Like our money system, the Metric system is based on powers of ten 3. Easier calculations Frequently only need to move a decimal point. BASIC UNITS Name Abbrev Type of Approximate Measurement Eng Equivalent meter m length inches (about 1 yard) liter L volume 1.06 quarts gram g mass oz (about 1/30 oz)

5 5 METRIC SYSTEM Prefix Abbreviation Value giga G 1 billion times (1 x 10 9 ) mega M 1 million times (1 x 10 6 ) kilo k 1 thousand times (1 x 10 3 ) deci d one-tenth (1/10) centi c one-hundredth ( 1/100 ) milli m one-thousandth ( 1/ 1000 ) micro µ one-millionth ( 1 / 10 6 ) nano n one-billionth ( 1 / 10 9 ) pico p one-trillionth ( 1 / ) Finally, time is measured the same way in all systems. The basic unit is the second (abbreviated as s) Let s practice some conversions: 1. If converting from a large unit to a smaller unit, move the decimal point to the right. Fill in any empty spaces with zeroes. The number of decimal places to move equals the number of zeroes in the prefix of the new unit. 2. If converting from a smaller unit to a larger unit, move the decimal point to the left and continue as above.

6 6 Practice: 1. Convert 5.36 m into cm 2. Convert 3468 nm into m 3. Convert 3.67 kg to mg Frequently scientist have very large or very small numbers to deal with. To make it easier they use scientific notation g becomes 1.23 x g, while mm becomes 5.38 X mm. For large # s, the positive exponent indicates the # of decimal positions after the first digit. For very small # s, the negative exponent indicates the number of decimal positions (less 1) filled by zeros between the decimal point and the first digit. A physical property that is easily measured and frequently used to help identify a substance is Density. It is defined as Mass divided volume of a sample. Mathematically that is: D = M/V Besides helping identify substances, density tells us whether something will float or sink in a liquid. Chapter 2 Ancient Greek and Roman concepts of Nature: 1. Continuists Aristotle 2. Atomists Leucippus, Democritus & Lucretious Dark & Middle Ages Little change

7 7 Modern Atomic Theory: John Dalton All matter made up from tiny, indestructible, indivisible spheres called atoms. 2. Atoms of an element are all identical. Atoms of different elements are different. 3. Compounds are formed when atoms combine together in small, whole number ratios to form molecules. This ratio is always the same for any compound. 4. A chemical reaction involves a rearrangement of atoms. No atoms are created, destroyed or broken apart in a chemical reaction Explained 2 previously discovered concepts (called laws, because statement of belief without explanation of why or how): 1. Law of Conservation of Mass Talked about earlier Discovered by Antoine Lavoisier 2. Law of Definite Proportions Discovered by Joseph Proust in 1799 All samples of a compound always has the same ratio of elements no matter where it is obtained from. By now the modern concept of elements was well accepted and many attempts had been made to arrange the known elements in some kind of coherent manner. All attempts had fatal flaws until a Russian chemist, Dmitri Mendeleev, in 1869 devised an approach which is still basically the same that is used today. His Periodic Table was based on 2 ideas: 1. Arrange elements in order of increasing atomic mass 2. Have groups or families of elements Predicted properties of missing elements for example, germanium: Property Predicted Observed Atomic mass Density (g/cm3) Color Dirty gray Grayish white

8 8 Chapter 3 By the end of the 19 th Century scientists had discovered that the atom was not solid and indivisible as Dalton had proposed. Several important experiments led to these discoveries. Let s look at 2 of them: 1. J.J. Thompson and the Crookes Tube Experiment

9 9 a) Led to discovery of electrons as basic unit of negative charge, found randomly inside atom. b) Plum-Pudding Model of the atom 2. Ernest Rutherford and discovery of the nuclear atom a) Natural radioactivity 3 types alpha particles (α), beta particles (β) and gamma radiation (γ)

10 10 b) Gold foil Experiment c) Nuclear Model of the atom Summary of particles inside atom: PARTICLE SYMBOL LOCATION Relative Charge Relative Mass (Mass #) Electron - e Outside = 0 nucleus Proton p + Inside nucleus Neutron n 0 Inside nucleus A tomic Number The number of protons in an atom. Also the number of electrons in a neutral atom. Unique for each element. Identifies the element. M ass Number The number of protons + the number of neutrons in an atom. This number is not unique for each element. All but one element have atoms with different numbers of neutrons and therefore different mass numbers. For example, there are 3 different atoms of C found in Nature, all having 6 protons, but one having 6 neutrons,

11 11 one 7 neutrons and one with 8 neutrons. What are their Mass Numbers? These 3 atoms are called isotopes of C. Isotopes are atoms of the same element with different Mass Numbers. Isotopic Symbols C A tomic Weight or better Atomic Mass(AW) The relative weight of an average atom of one element compared to that of another atom chosen to be the standard. The modern standard is the C-12 isotope (arbitrarily assigned a value of amu (atomic mass unit). Actually the AW of an element is the average (taking into account the relative amount of each) mass of all the naturally found isotopes. B ohr Model Also called Solar System Model. Nucleus with electrons revolving around in orbits of specific energies (called Energy levels) Each successive orbit has a maximum # of electrons that can exist (2, 8, 18, 32, 50 etc) We will only worry about the first 20 elements. For these the maximum # of electrons can be considered to be 2 for the first and 8 for the rest. For every atom the electrons fill up the orbits in order beginning with the lowest energy. The closer to the nucleus, the lower the energy. In other words Li has an atomic # of 3, therefore 3 electrons (2 in the first level and 1 in the next). See Table 3.2 on page 54. The electrons in the outermost energy level are called Valence electrons and that level is called the Valence Shell.

12 12 Dimitri Mendeleev and the Periodic Table Modern Version 1. Names of Families 2. Metals, Non-metals and metalloids 3. Representative elements and transition metals M odern Periodic Table: 1. Metals Left side - a) Properties Shiny, malleable, ductile, conduct electricity & heat 2. N on-metals Right Side b) Properties Brittle as solids, non-conductors 3. M etalloids In between c) Properties Some of both metals & non-metals 4. N oble Gases Far Right Column d) Properties All are gases even at very low temperatures, non-conductors, very little tendency to chemically react

13 13 R epresentative Elements Those in columns headed by an A (IA, IIA etc). We will spend most of our time dealing with these elements. T ransition Metals Those in columns headed by a B. In ner Transition Elements Bottom 2 rows