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Investigating Atoms and Atomic Theory Students should be able to: Describe the particle theory of matter. Use the Bohr model to differentiate among the three basic particles in the atom (proton, neutron, and electron) and their charges, relative masses, and locations. Compare the Bohr atomic model to the electron cloud model with respect to their ability to represent accurately the structure of the atom. Atomos: Not to Be Cut The History of Atomic Theory Atomic Models This model of the atom may look familiar to you. This is the Bohr model. In this model, the nucleus is orbited by electrons, which are in different energy levels. A model uses familiar ideas to explain unfamiliar facts observed in nature. A model can be changed as new information is collected. The atomic model has changed throughout the centuries, starting in 400 BC, when it looked like a billiard ball → Who are these men? In this lesson, we’ll learn about the men whose quests for knowledge about the fundamental nature of the universe helped define our views. Democritus This is the Greek philosopher Democritus who began the search for a description of matter more than 2400 years ago. He asked: Could matter be divided into smaller and smaller pieces forever, or was there a limit to the number of times a piece of matter could be divided? 400 BC Atomos His theory: Matter could not be divided into smaller and smaller pieces forever, eventually the smallest possible piece would be obtained. This piece would be indivisible. He named the smallest piece of matter “atomos,” meaning “not to be cut.” Atomos To Democritus, atoms were small, hard particles that were all made of the same material but were different shapes and sizes. Atoms were infinite in number, always moving and capable of joining together. This theory was ignored and forgotten for more than 2000 years! Why? The eminent philosophers of the time, Aristotle and Plato, had a more respected, (and ultimately wrong) theory. Aristotle and Plato favored the earth, fire, air and water approach to the nature of matter. Their ideas held sway because of their eminence as philosophers. The atomos idea was buried for approximately 2000 years. Foundations of Atomic Theory This was until the 18th century when scientists began to gather evidence favoring the atomic theory of matter. All chemists in the late 1700’s accepted that an element cannot be further broken down by ordinary chemical means, however, there was controversy as to whether elements always combine in the same ratio when forming a particular compound. Aided by new balances and methods, scientist began to accurately measure the masses of the elements and compounds—this led to the discovery of various laws: Law of Conservation of Mass (Lavoisier, 1777) : mass is neither created nor destroyed during ordinary chemical reactions or physical changes. Also found out that pure chemical compound is composed of a fixed proportion of elements. Law of Definite Proportion: (Proust, 1797) A chemical compound contains the same elements in the same proportions by mass regardless of the size of the sample or source of the compound. Law of Multiple Proportions: If two or more different compounds are composed of the same elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a small whole number ratio. Law of Definite Proportion Water Hydrogen + Oxygen 18.0 g 2.0 g 16.0 g Mass ratio = 16.0 g of O = 8 gives 8:1 ratio 2.0 g of H 8 Oxygen to 1 Hydrogen ratio Always results from the same combinations of atoms. Law of Multiple Proportions Also known as Dalton’s Law CO and CO2 CO = C to O ratio is 1:1 O= 16g C = 12g CO2 = C to O ratio is 1:2 O = 2(16g) = 32g C = 12g If CO and CO2 were each produced by the reaction of 12 g of carbon with oxygen, the resulting compounds would contain 16 g and 32 g of oxygen. The ratio of 16 g of O to 32g of O is ½ . This indicates that in CO there are half as many oxygen atoms for every carbon atom in CO2. Dalton’s Model In the early 1800s, the English Chemist John Dalton performed a number of experiments that eventually led to the acceptance of the idea of atoms. Dalton’s Theory He deduced that all elements are composed of atoms. Atoms are indivisible and indestructible particles. Atoms of the same element are exactly alike. Atoms of different elements are different. Compounds are formed by the joining of atoms of two or more elements. . This theory became one of the foundations of modern chemistry. Dalton’s Theory Not all of Dalton’s Atomic Theory have proven to be correct…we know that atoms are divisible into smaller particles, and know that for a given element can have atoms with different masses…the theory has been modified to explain the new observations. Important concept is: All matter is composed of atoms Atoms of any one element differ in properties from atoms of another element remain unchanged. Thomson’s Plum Pudding Model In 1897, the English scientist J.J. Thomson provided the first hint that an atom is made of even smaller particles. Thomson Model He proposed a model of the atom that is sometimes called the “Plum Pudding” model. Atoms were made from a positively charged substance with negatively charged electrons scattered about, like raisins in a pudding. Thomson Model Thomson studied the passage of an electric current through a gas. As the current passed through the gas, it gave off rays of negatively charged particles. A cathode ray can also be deflected by a magnet. Thomson Model This surprised Thomson, because the atoms of the gas were uncharged. Where had the negative charges come from? Where did they come from? Thomson concluded that the negative charges came from within the atom. A particle smaller than an atom had to exist. The atom was divisible! Thomson called the negatively charged “corpuscles,” today known as electrons. Since the gas was known to be neutral, having no charge, he reasoned that there must be positively charged particles in the atom. But he could never find them. The U.S. physicist Robert A. Millikan (1868–1953) carried out experiments to find the quantity of an electron’s charge. From his data, he found that the charge on each oil droplet was a multiple of 1.60 10–19 coulomb, meaning this must be the charge of an electron. •Using this charge and Thomson’s charge-to-mass ratio of an electron, Millikan calculated an electron’s mass. •Millikan’s values for electron charge and mass are similar to those accepted today. •An electron has one unit of negative charge, and its mass is 1/1840 the mass of a hydrogen atom. Protons and Neutrons If cathode rays are electrons given off by atoms, what remains of the atoms that have lost the electrons? For example, after a hydrogen atom (the lightest kind of atom) loses an electron, what is left? You can think through this problem using four simple ideas about matter and electric charges. 1. Atoms have no net electric charge; they are electrically neutral. 2. Electric charges are carried by particles of matter. 3. Electric charges always exist in whole-number multiples of a single basic unit; that is, there are no fractions of charges. 4. When a given number of negatively charged particles combines with an equal number of positively charged particles, an electrically neutral particle is formed. It follows that a particle with one unit of positive charge should remain when a typical hydrogen atom loses an electron. In 1886, Eugen Goldstein (1850–1930) observed a cathode-ray tube and found rays traveling in the direction opposite to that of the cathode rays. • He concluded that they were composed of positive particles. • Such positively charged subatomic particles are called protons. The Neutron!!!! In 1932, the English physicist James Chadwick (1891–1974) confirmed the existence of yet another subatomic particle: the neutron. Neutrons are subatomic particles with no charge but with a mass nearly equal to that of a proton. Rutherford’s Gold Foil Experiment In 1908, the English physicist Ernest Rutherford was hard at work on an experiment that seemed to have little to do with unraveling the mysteries of the atomic structure. Rutherford’s experiment Involved firing a stream of tiny positively charged particles at a thin sheet of gold foil (2000 atoms thick) Gold Foil Experiment Most of the positively charged “bullets” passed right through the gold atoms in the sheet of gold foil without changing course at all. Some of the positively charged “bullets,” however, did bounce away from the gold sheet as if they had hit something solid. He knew that positive charges repel positive charges. This could only mean that the gold atoms in the sheet were mostly open space. Atoms were not a pudding filled with a positively charged material. Rutherford concluded that an atom had a small, dense, positively charged center that repelled his positively charged “bullets.” He called the center of the atom the “nucleus” The nucleus is tiny compared to the atom as a whole. Rutherford Rutherford reasoned that all of an atom’s positively charged particles were contained in the nucleus. The negatively charged particles were scattered outside the nucleus around the atom’s edge. Bohr Model In 1913, the Danish scientist Niels Bohr proposed an improvement. In his model, he placed each electron in a specific energy level. Bohr Model According to Bohr’s atomic model, electrons move in definite orbits around the nucleus, much like planets circle the sun. These orbits, or energy levels, are located at certain distances from the nucleus. Wave Model The Wave Model Today’s atomic model is based on the principles of wave mechanics. According to the theory of wave mechanics, electrons do not move about an atom in a definite path, like the planets around the sun. The Wave Model In fact, it is impossible to determine the exact location of an electron. The probable location of an electron is based on how much energy the electron has. According to the modern atomic model, at atom has a small positively charged nucleus surrounded by a large region in which there are enough electrons to make an atom neutral. Electron Cloud: A space in which electrons are likely to be found. Electrons whirl about the nucleus billions of times in one second They are not moving around in random patterns. Location of electrons depends upon how much energy the electron has. Electron Cloud: Depending on their energy they are locked into a certain area in the cloud. Electrons with the lowest energy are found in the energy level closest to the nucleus Electrons with the highest energy are found in the outermost energy levels, farther from the nucleus. Indivisible Electron Greek X Dalton X Nucleus Thomson X Rutherford X X Bohr X X Wave X X Orbit Electron Cloud X X The table below summarizes the properties of these subatomic particles. Properties of Subatomic Particles Particle Symbol Relative charge Relative mass (mass of proton = 1) Actual mass (g) Electron e– 1– 1/1840 9.11 10–28 Proton p+ 1+ 1 1.67 10–24 Neutron n0 0 1 1.67 10–24 Size of atoms An atom is small, but the particles associated with the atom are much smaller! Atomic Numbers and Mass Numbers The atomic number of an element is the number of protons in the nucleus of each of the elements. In a neutral atom, the atomic number also designates how many electrons the atom contains. The atomic number identifies an element. For each element listed in the table below, the number of protons equals the number of electrons Mass Number The total number of protons and neutrons in an atom is called the mass number. If you know the atomic number and mass number of an atom of any element, you can determine the atom’s composition. The number of neutrons in an atom is the difference between the mass number and atomic number. Mass Number The composition of any atom can be represented in shorthand notation using atomic number and mass number. • The atomic number is the subscript. • The mass number is the superscript. Au is the chemical symbol for gold. ISOTOPES Isotopes- Atoms having different mass numbers but the same atomic numbers. Nuclide is the term for any isotope of any element. EX; Hydrogen has 3 isotopesHydrogen, Deuterium, Tritium There are three different kinds of neon atoms. How do these atoms differ? • All have the same number of protons (10). • All have the same number of electrons (10). • But they each have different numbers of neutrons. Because isotopes of an element have different numbers of neutrons, they also have different mass numbers. Despite these differences, isotopes are chemically alike because they have identical numbers of protons and electrons, which are the subatomic particles responsible for chemical behavior. Analogy to Dogs… • Their color or size doesn’t change the fact that they are all dogs. • Similarly, the number of neutrons in isotopes of an element does not change which element it is because the atomic number does not change. Relative Atomic Masses Masses of atoms measured in grams are very small. In order to set up a relative scale of atomic mass, one atom is arbitrarily chosen as the standard and assigned a relative mass value. The masses of all the other atoms are then expressed in relation to this defined standard. The average atomic masses are important to chemist because they indicate relative mass relationships in chemical reactions. • It is more useful to compare the relative masses of atoms using a reference isotope as a standard. • The reference isotope chosen is carbon12. This isotope of carbon has been assigned a mass of exactly 12 atomic mass units. An atomic mass unit (amu) is defined as onetwelfth of the mass of a carbon-12 atom. A carbon-12 atom has six protons and six neutrons in its nucleus, and its mass is set at 12 amu. • The six protons and six neutrons account for nearly all of this mass. • Therefore, the mass of a single proton or a single neutron is about one-twelfth of 12 amu, or about 1 amu. Average Atomic Masses Most elements in nature occur naturally as mixtures of isotopes. The percentage of each isotope in the naturally occurring element is nearly always the same, no matter where the element is found. Atomic masses for the elements are averages for these naturally occurring mixtures of the isotopes. In nature, most elements occur as a mixture of two or more isotopes. Each isotope of an element has a fixed mass and a natural percent abundance. Chlorine occurs as two isotopes: chlorine-35 and chlorine-37. If you calculate the arithmetic mean of these two masses (34.968 amu + 36.966 amu)/2), you get an average atomic mass of 35.986. • However, this value is higher than the actual value of 35.453. To explain this difference, you need to know the natural percent abundance of the isotopes of chlorine. • Chlorine-35 accounts for 75 percent of the naturally occurring chlorine atoms; chlorine-37 accounts for only 24 percent. Because there is more chlorine-35 than chlorine-37 in nature, the atomic mass of chlorine, 35.453 amu, is closer to 35 than to 37. The atomic mass of an element is a weighted average mass of the atoms in a naturally occurring sample of the element. A weighted average mass reflects both the mass and the relative abundance of the isotopes as they occur in nature. Understanding Relative Abundance of Isotopes The atomic mass of copper is 63.546 amu. Which of copper’s two isotopes is more abundant: copper-63 or copper-65? Compare the atomic mass to the mass of each isotope. The atomic mass of 63.546 amu is closer to 63 than it is to 65. Determine the most abundant isotope based on which isotope’s mass is closest to the atomic mass. Because the atomic mass is a weighted average of the isotopes, copper-63 must be more abundant than copper-65. You can determine atomic mass based on relative abundance. To do this, you must know three things: the number of stable isotopes of the element, the mass of each isotope, and the natural percent abundance of each isotope. To calculate the atomic mass of an element, multiply the mass of each isotope by its natural abundance, expressed as a decimal, and then add the products. Carbon has two stable isotopes: carbon-12, which has a natural abundance of 98.89 percent, and carbon-13, which has a natural abundance of 1.11 percent. • The mass of carbon-12 is 12.000 amu; the mass of carbon-13 is 13.003 amu. • The atomic mass of carbon is calculated as follows: Atomic mass of carbon = (12.000 amu x 0.9889) + 13.003 amu x 0.0111) = (11.867 amu) + (0.144 amu) = 12.011 amu Let’s try one! Element X has two naturally occurring isotopes. The isotope with a mass of 10.012 amu (10X) has a relative abundance of 19.91 percent. The isotope with a mass of 11.009 amu (11X) has a relative abundance of 80.09 percent. Calculate the atomic mass of element X. The mass each isotope contributes to the element’s atomic mass can be calculated by multiplying the isotope’s mass by its relative abundance. The atomic mass of the element is the sum of these products. • Isotope 10X: mass = 10.012 amu relative abundance = 19.91% = 0.1991 • Isotope 11X: mass = 11.009 amu relative abundance = 80.09% = 0.8009 Use the atomic mass and the decimal form of the percent abundance to find the mass contributed by each isotope. for 10X: 10.012 amu x 0.1991 = 1.993 amu for 11X: 11.009 amu x 0.8009 = 8.817 amu Add the atomic mass contributions for all the isotopes. For element X, atomic mass = 1.953 amu + 8.817 amu = 10.810 amu Why is the atomic mass of an element usually not a whole number? The atomic mass of an element is usually not a whole number because it is a weighted average of the masses of the naturally occurring isotopes of the element. Subatomic Particles of the atom Each basic element has a certain number of electrons and protons, which distinguishes each element from all other basic elements. In most elements, the number of electrons is equal to the number of protons. This maintains an electrical balance in the structure of atoms since protons and electrons have equal, but opposite electrostatic fields.