Atoms, Molecules and Chemical Reactions CHM 101 PDF

Atoms, Molecules and Chemical Reactions Week 2 CHEMISTRY The language used in chemistry is seen and heard in many disciplines, ranging from medicine to engineering to forensics to art. The language of chemistry includes its own vocabulary as well as its own form of shorthand. C...

Atoms, Molecules and Chemical Reactions Week 2 CHEMISTRY The language used in chemistry is seen and heard in many disciplines, ranging from medicine to engineering to forensics to art. The language of chemistry includes its own vocabulary as well as its own form of shorthand. Chemical symbols are used to represent atoms and elements. Chemical formulas depict molecules as well as the composition of compounds. Chemical equations provide information about the quality and quantity of the changes associated with chemical reactions.  ATOMS These are the smallest constituent unit of matter which posses properties of the chemical element. – Atoms are defined as the basic building blocks of matter. - Atoms do not exist independently, rather, they form molecules and ions which then combine in large numbers to form matter that we are able to see, feel and touch. - The human eye is incapable of seeing an atom, instead, experiments are carried out to find out their structure, and behavior e.g. a scanning electron microscope (SEM)  ATOMS All atoms of the same element are identical and different elements have different types of atoms. All the elements listed in the periodic table are made of atoms. An example is with Aluminum on the periodic table. A rearrangement of atoms is what leads to a chemical reaction. Atoms are the building blocks of everything we see around us yet we cannot see an atom or even a billion atoms with the naked eye  2 main parts of an atom: Nucleus-99.9% of the atom’s mass Electron cloud or energy rings Atomic Theory through the Nineteenth Century The earliest recorded discussion of the basic structure of matter comes from ancient Greek philosophers, the scientists of their day. In the fifth century BC, Leucippus and Democritus argued that all matter was composed of small, finite particles that they called atomos, a term derived from the Greek word for “indivisible”. They thought of atoms as moving particles that differed in shape and size, and which could join together. Later, Aristotle and others came to the conclusion that matter consisted of various combinations of the four “elements”—fire, earth, air, and water—and could be infinitely divided.  Interestingly, these philosophers thought about atoms and “elements” as philosophical concepts, but apparently never considered performing experiments to test their ideas. The Aristotelian view of the composition of matter held sway for over two thousand years, until English schoolteacher John Dalton helped to revolutionize chemistry with his hypothesis that the behavior of matter could be explained using an atomic theory. First published in 1807, many of Dalton’s hypotheses about the microscopic features of matter are still valid in modern atomic theory Postulates of Dalton’s atomic theory can be found on the next slide: Daltons Atomic Theory Daltons Atomic Theory Dalton’s atomic theory provides a microscopic explanation of the many macroscopic properties of matter that you’ve learned about. For example, if an element such as copper consists of only one kind of atom, then it cannot be broken down into simpler substances, that is, into substances composed of fewer types of atoms. And if atoms are neither created nor destroyed during a chemical change, then the total mass of matter present when matter changes from one type to another will remain constant (the law of conservation of matter). Daltons Atomic Theory Dalton knew of the experiments of French chemist Joseph Proust, who demonstrated that all samples of a pure compound contain the same elements in the same proportion by mass. This statement is known as the law of definite proportions or the law of constant composition. The suggestion that the numbers of atoms of the elements in a given compound always exist in the same ratio is consistent with these observations. Daltons Atomic Theory Dalton also used data from Proust, as well as results from his own experiments, to formulate another interesting law. The law of multiple proportions states that when two elements react to form more than one compound, a fixed mass of one element will react with masses of the other element in a ratio of small, whole numbers. For example, copper and chlorine can form a green, crystalline solid with a mass ratio of 0.558 g chlorine to 1 g copper, as well as a brown crystalline solid with a mass ratio of 1.116 g chlorine to 1 g copper. These ratios by themselves may not seem particularly interesting or informative; however, if we take a ratio of these ratios, we obtain a useful and possibly surprising result: a small, whole-number ratio. This 2-to-1 ratio means that the brown compound has twice the amount of chlorine per amount of copper as the green compound. Atomic Theory after the Nineteenth Century If matter were composed of atoms, what were atoms composed of? Were they the smallest particles, or was there something smaller? In the late 1800s, a number of scientists interested in questions like these investigated the electrical discharges that could be produced in low-pressure gases, with the most significant discovery made by English physicist J. J. Thomson using a cathode ray tube. Fig: Cathode ray tube Atomic Theory after the Nineteenth Century Based on his observations, here is what Thomson proposed and why: The particles are attracted by positive (+) charges and repelled by negative (−) charges, so they must be negatively charged (like charges repel and unlike charges attract). They are less massive than atoms and indistinguishable, regardless of the source material, so they must be fundamental, subatomic constituents of all atoms. Although controversial at the time, Thomson’s idea was gradually accepted, and his cathode ray particle is what we now call an electron, a negatively charged, subatomic particle with a mass more than one thousand-times less that of an atom. The term “electron” was coined in 1891 by Irish physicist George Stoney, from “electric ion.” The nucleus contains the majority of an atom’s mass because protons and neutrons are much heavier than electrons, whereas electrons occupy almost all of an atom’s volume. The diameter of an atom is on the order of 10−10 m, whereas the diameter of the nucleus is roughly 10−15 m—about 100,000 times smaller. For a perspective about their relative sizes, consider this: If the nucleus were the size of a blueberry, the atom would be about the size of a football stadium  The number of protons in the nucleus of an atom is equal to the atomic number (Z). The number of electrons in a neutral atom is equal to the number of protons. The mass number of the atom (M) is equal to the sum of the number of protons and neutrons in an atom The number of neutrons is equal to the difference between the mass number of the atom (M) and the atomic number (Z)  QUESTIONS The existence of isotopes violates one of the original ideas of Dalton’s atomic theory which one? How are electrons and protons similar? How are they different? How are protons and neutrons similar? How are they different? Answers Dalton originally thought that all atoms of a particular element had identical properties, including mass. Thus, the concept of isotopes, in which an element has different masses, was a violation of the original idea. To account for the existence of isotopes, the second postulate of his atomic theory was modified to state that atoms of the same element must have identical chemical properties. Both are subatomic particles that reside in an atom’s nucleus. Both have approximately the same mass. Protons are positively charged, whereas electrons are negatively charged. Both are subatomic particles that reside in an atom’s nucleus. Both have approximately the same mass. Protons are positively charged, whereas neutrons are uncharged. MOLECULES Molecules consist of one or more atoms bound together by covalent (chemical) bonds. Atoms may be depicted by circle shapes, each of which has a nucleus at the center (containing protons and neutrons), surrounded by one or more concentric circles representing the ‘shells’ or ‘levels’ in which the electrons surrounding the nucleus of the atom are located and markings indicating the electron at each level. A molecule is the smallest thing a substance can be divided into while remaining the same substance. It is made up of two or more atoms that are bound together by chemical bonding. MOLECULES A molecule is a collection of two or more atoms that make up the smallest recognizable unit into which a pure material may be split while maintaining its makeup and chemical characteristics. Some examples of molecules are H2O (water) N2 (nitrogen) O3 (ozone) The simplest forces between atoms are those which arise as a result of electron transfer. A simple example is that of say sodium fluoride. The sodium atom has a nuclear charge of +11, with 2 electrons in the K shell, 8 in the L shell and 1 in the M shell. The fluorine atom has a nuclear charge of 9 with 2 electrons in the K shell and 7 in the L shell. The outermost electron in the sodium atom may transfer readily to the fluorine atom; both atoms then have a complete shell but the sodium now has a net charge of +1 and the fluorine a net charge of -1. These ions, therefore, attract one another by direct coulombic interaction. Attraction or repulsion of particles or objects because of their electric charge. CHEMICAL REACTIONS CHEMICAL REACTIONS TYPES OF CHEMICAL REACTIONS CHEMICAL FORMULAS ANSWERS QUESTIONS ANSWERS QUESTIONS ANSWERS