Survey
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
Atomic Theory: History and Structure Atomos is Greek for ‘indivisible’ Ancient Greeks • Democritus (~460-370 BC) and Leucippus, his teacher, believed that everything is composed of "atoms“ that are physically indivisible, indestructible, and are continuously in motion with empty space between atoms; that there are an infinite number and kinds of atoms, which differ in shape and size. Aristotle (384-322 BC) The heavens were made of a fifth element, ‘aether’. • Aristotle believed a piece of matter could be divided an infinite number of times. • He also believed that matter was made up of 4 elements which had 4 associated qualities (e.g., fire was hot and dry) • The two forces of gravity (the tendency for earth and water to sink) and levity (the tendency for air and fire to rise) acted upon elements. Particulate Nature of Matter In the early 18th Century, evidence accumulated for the view that matter was made of particles, which led to new models: • Robert Boyle postulated that gases are composed of discrete particles separated by a void • In 1704 Sir Isaac Newton proposed a mechanical model of the universe, with small solid masses in motion. Chemical Laws of the 18th Century Near the end of the 18th Century, three laws emerged that governed chemical reactions. • Antoine Lavoisier formulated the law of conservation of mass in 1789 which states that the total mass in a chemical reaction remains constant (that is, the reactants have the same mass as the products). • The law of definite proportions was proven by the French chemist Joseph Louis Proust in 1799. This law states that if a compound is broken down into its constituent elements, then the masses of the constituents will always have the same proportions, no matter what the source was or the quantity of the original substance. John Dalton (1766 -1844) • Dalton, an English chemist, meteorologist, and physicist, combined these three laws (laws of conservation of mass, definite proportions, and multiple proportions) into a single explanation. Dalton turned Democritus’ idea of the atom into a scientific theory that could be tested by experiment by relating atoms to the measurable property of mass. Dalton’s Atomic Theory 1. All matter is composed of extremely small particles called atoms. 2. Atoms of a given element are identical in size, mass and other properties; atoms of different elements differ in size, mass, and other properties. 3. Atoms cannot be subdivided, created, or destroyed. Dalton’s Atomic Theory (cont.) 4. Atoms of different elements combine in simple whole-number ratios to form chemical compounds. 5. In chemical reactions, atoms are combined, separated, or rearranged. Dalton’s Contributions • In 1808 Dalton published a textbook A New System of Chemical Philosophy in which he explained his method for estimating relative atomic weights. His method was to compare the mass ratios in which they combined, with the hydrogen atom taken as unity. • Unfortunately, Dalton did not understand that some elements exist as diatomic molecules — such as oxygen which occurs as O2. He also mistakenly believed that the simplest compound between any two elements is always one atom of each (so he thought water was HO, not H2O). J.J. Thomson • Atoms were thought to be the smallest possible division of matter until 1897 when J.J. Thomson discovered the electron (which he called a ‘corpuscle’) through his work on cathode rays. Cathode-Ray Tubes • A cathode-ray tube (CRT) is a glass tube filled with some gas at low pressure through which electric current is passed. The surface of the tube directly opposite the cathode glowed, it was theorized, due to a stream of particles called a cathode ray which traveled from the cathode to the anode. CRT Experiments • Experiments with cathode-ray tubes showed that 1. Cathode rays are deflected by a magnetic field just as a wire carrying an electric current is (which they knew had a negative charge) 2. Cathode rays were deflected AWAY from a negatively charged object. • This led to the hypothesis that cathode rays were made of negatively charged particles. Thomson’s Discovery • Thomson’s experiments measured the ratio of the charge of the CRT particles compared to their mass. He learned that this ratio was the same, regardless of which metal he used to make the cathode or which gas filled the tube. • Thomson concluded that all cathode rays are made of the same negatively charged particles. Electron Charge and Mass • Furthermore, these negatively charged particles must be present in atoms of ALL elements and so the atom must be divisible. • Experiments by Robert Millikan in 1909 measured the charge of this newly discovered particle. Since Thomson had already found the charge-to-mass ratio, now the electron’s mass could be determined. • The mass of an electron is 9.109 x 10-31 kg which is 1/1837 the mass of a hydrogen atom. Thomson’s Atomic Model • Scientists needed to explain (1) how atoms are electrically neutral and (2) since the mass of an electron is so small compared to an atom, where was the rest of the mass of atom? • Thomson proposed that the electrons were set in a uniform sea of positive charge (although they could move about). This was known as the ‘plum pudding model.’ (small fruit pieces among the cake) Discovery of the Nucleus • Ernest Rutherford was a former student of Thomson who in 1909 performed an experiment with colleagues that showed that most of the mass and the positive charge of an atom is contained in a very small volume. Gold Foil Experiment • A beam of alpha particles (helium nuclei consisting of 2 protons and 2 neutrons) was shot at a thin piece of gold foil. If Thomson’s plum pudding model was correct, the alpha particles were expected to pass through the foil without being deflected, as their momentum would overcome both the small mass of the electrons and the repulsion of the uniformly Expected result of experiment distributed positive charge. based on ‘plum pudding’ model Unexpected Result • To their surprise, about 1 in 800 alpha particles was deflected at very large angles! • Rutherford’s famous quote about this is, it was “as if you had fired a 15inch [artillery] shell at a piece of tissue paper and it came back and hit you.” Rutherford’s Planetary Model of the Nucleus If the nucleus were the size of a marble, the size of the atom would be about the size of a football field. • Rutherford then proposed a planetary atomic model in which a cloud of electrons surrounded a small, compact nucleus of positive charge. Only such a concentration of charge could produce an electric field strong enough to cause the deflection of the alpha particle beam. Atoms are Divisible! • Nowadays we define an atom as the smallest particle of an element that retains the chemical properties of that element. • We view atoms as consisting of two regions: the nucleus is very small and located at the center of an atom, consisting of one or more protons and usually neutrons, while electrons occupy a very large region surrounding the nucleus. • Protons, neutrons and electrons are referred to as SUBATOMIC PARTICLES. To be continued… • Meanwhile, back at the ranch . . . . So, what’s in the nucleus? • We know now that, except for the simplest type of hydrogen atom, all atomic nuclei are made of two kinds of particles, PROTONS (p+) and NEUTRONS (n°). • However, neutrons were not discovered until 1932, 21 years after Rutherford’s discovery of the nucleus! • For the purposes of chemistry, other subatomic particles besides protons, neutrons, and electrons (e-) don’t play a part in chemical properties, so we’re going to ignore them. Composition of the Nucleus Just to complete the set… ELECTRONS e- or −𝟏𝟎𝒆 • Each electron has a negative charge • Each electron has a mass equal to 9.109 x 10-31 kg Atoms are Electrically Neutral • This means that every atom has equal numbers of protons and electrons (neutrons have no charge). • The nuclei of atoms of different elements differ in the number of protons they have and therefore the amount of positive charge they possess. • The number of protons in an atom determines an element’s identity and is known as its ATOMIC NUMBER. How big are atoms? • It is convenient to think of electrons moving rapidly outside the nucleus in energy regions that can be imagined as ‘clouds’ or shells of negative charge. • The radius of an atom is the distance from the center of the nucleus to the outer portion of this electron cloud. • Atomic radii range from about 40 to 270 pm (that’s picometers or 10-12 m) while nuclei of atoms are much smaller, about 0.001 pm. Isotopes • Isotopes are atoms of an element that have the same number of protons but different numbers of neutrons. Because these atoms have the same number of electrons as all other atoms of that element, they behave similarly in chemical reactions. • Almost all elements have more than one isotope. Xenon has 9 and tin has 10. Fluorine, however, has only one. Isotropes vs. Allotropes • Isotopes are NOT different forms of an element. For instance, carbon occurs as graphite, diamond, carbon nanotubes, and buckminsterfullerene which all have very different properties. These are ALLOTROPES of carbon, not isotopes Hydrogen’s Isotopes MASS NUMBERS OF HYDROGEN ISOTOPES Atomic number (# of protons) # of neutrons Mass number (protons + neutrons) % of hydrogen atoms Protium 1 0 1+0=1 99.9885 Deuterium 1 1 1+1=2 0.0115 Tritium 1 2 1+2=3 Very rare Tritium is radioactive. It exists in very small amounts in nature, but it can be prepared artificially. Designating Isotopes • The isotopes of hydrogen have distinct names, although isotopes are usually identified by their mass number. • There are two methods for specifying isotopes that DON’T have names: 1. The mass number is written with a hyphen after the name of the element (tritium is hydrogen-3). This is HYPHEN NOTATION. Carbon’s Isotopes • Carbon has 3 naturally occurring isotopes: Carbon-12 (6 protons, 6 neutrons) Carbon-13 (6 protons, 7 neutrons) Carbon-14 (6 protons, 8 neutrons) • Carbon-14 is unstable and undergoes radioactive decay. This rate of decay is known and is used to find the ages of fossils, rocks and minerals. • Other elements have unstable isotopes which are also used to ‘date’ fossils, rocks and minerals like potassium and uranium. Designating Isotopes (cont.) ISOTOPES OF HYDROGEN AND HELIUM Isotope Nuclear Symbol # protons # electrons #neutrons Hydrogen-1 (protium) 1 1 0 Hydrogen-2 (deuterium) 1 1 1 Hydrogen-3 (tritium) 1 1 2 Helium-3 2 2 1 Helium-4 2 2 2 Relative Atomic Masses • Because the masses of atoms expressed in grams are very small, it is more convenient to use relative atomic masses. • Scientists have chosen the carbon-12 atom to be the standard atom for this purpose, and it has been assigned a mass of EXACTLY 12 atomic mass units or amu. One amu is 1/12 of the mass of a carbon-12 atom. Average Atomic Masses • Most elements occur naturally as a mixture of isotopes with the percentage of each isotope being nearly always the same, regardless of where on Earth the element is found. • AVERAGE ATOMIC MASS is the weighted average of the atomic masses of the naturally-occurring isotopes of an element. Calculating a Weighted Average Example problem Problems with Rutherford’s Planetary Model • Remember Rutherford’s planetary model? (a cloud of electrons, like planets, orbiting a small, compact nucleus of positive charge). Unfortunately, when electrical charge is involved, it gets more complicated. • First, planets do not have electric charge, and orbiting charged particles would lose energy and spiral into the nucleus in less than a second. • Secondly, the planetary model had no way of explaining the highly peaked emission and absorption spectra that were observed for atoms. Wavelength (nm) Quantum Theory to the Rescue • At the beginning of the 20th Century, Max Planck and Albert Einstein postulated that light energy is emitted and absorbed in discrete (distinct) amounts called quanta which we know as photons. • Quantum mechanics is important for interactions on the scale (size) of atoms. The Bohr Model • In 1913 Niels Bohr, with Rutherford, incorporated the idea of quanta into his model of the atom, in which an electron could only orbit the nucleus in particular circular orbits which depended on the electron’s energy and angular momentum. This meant an atom had fixed energy levels for its electrons. Bohr Model Photons have energies according to this formula, E=hf, where h is Planck’s constant (6.626 × 10-34 m2-kg/s) and f is the frequency of the light in hertz. Bohr Model • In this model an electron could not spiral into the nucleus because it could not lose energy in a continuous manner. Electrons could only make instantaneous "quantum leaps" between fixed energy levels, and when this happened, light was emitted or absorbed at a frequency that depended upon this change in energy. • This model successfully explained the observed emission and absorption spectra of light. More Modern Theories • Bohr’s model could only partially explain the hydrogen atom, but it is still useful in explaining quantized energy levels. • However, the modern model of the atom describes the positions of electrons in an atom in terms of probabilities. An electron can potentially be found at any distance from the nucleus, but, depending on its energy level, exists more frequently in certain regions around the nucleus than others. • This pattern is referred to as its atomic orbital. Orbitals come in a variety of shapes—sphere, dumbbell, torus, etc.—with the nucleus in the middle. Atoms have the same number of protons and electrons and so are electrically neutral. The electron clouds are made up of different energy levels. Each energy level can only hold a certain number of electrons. The electrons in the outermost level can interact with other atoms’ electrons and form chemical compounds. How many electrons in each level? • Electrons fill the energy levels from lowest to highest. –The lowest level (Period 1) can hold only 2 e–The 2nd level (Period 2) can hold 8 e–The 3rd level (Period 3) can hold 18 e–The 4th level (Period 4) can hold 32 e- Electron orbital ring notation Boron-10 Oxygen-16 Which element is this? Orbitals and Quantum Numbers • In order to completely describe orbits, scientists use quantum numbers which specify the properties of atomic orbitals and the properties of electrons in those orbitals. • The principal quantum number n indicates the main energy level occupied by the electron. Values of n are positive integers only. The lowest energy level is n = 1 and is closest to the nucleus. Each higher value of n is successively further out. Orbital Shapes • Except at the first main energy level, orbitals of different shapes—known as sublevels—are defined by the second quantum number, the angular momentum quantum number l. ORBITAL LETTER DESIGNATIONS AND SHAPES Value of l Letter designation Shape of orbital 0 s Spherical 1 p Dumbbell 2 d Dumbbell, torus, etc. 3 f Various s, p, d and f orbital shapes How many electrons can the f orbitals hold? QUANTUM NUMBER RELATIONSHIPS IN ATOMIC STRUCTURE Principal quantum number Sublevels in main energy level Number of electrons per sublevel Number of electrons per main energy level 1 s 2 2 2 s p 2 6 8 3 s p d 2 6 10 18 4 s p d f 2 6 10 14 32 Electron Orbitals Gallery http://en.wikipedia.org/wiki/Atomic_orbital_model Wait a minute… • Did the pattern in the last slide remind you of anything else? • Which orbital shape corresponds to the blue ones? • Which orbital shape corresponds to the yellow ones? • Which orbital shape corresponds to the red ones? • Which orbital shape corresponds to the green ones? Sublevel Blocks of the Periodic Table Electron Configurations • The quantum model improves on the Bohr model because it describes the arrangements of electrons in atoms other than hydrogen. The arrangement of electrons in an atom is known as the atom’s electron configuration. • Because atoms of different elements have different numbers of electrons, a unique electron configuration exists for the atoms of each element. • Like all systems in nature, electrons in atoms tend to assume arrangements that have the lowest possible energies. • The lowest energy arrangement of electrons for each element is called the element’s ground-state electron configuration. • To build up electron configurations for the ground state of a particular atom, the energy levels of the orbitals are found, then electrons are added to the orbitals one by one according to three basic rules. Rule #1: Aufbau Principle • This rule shows the order in which electrons occupy orbitals. An electron occupies the lowestenergy orbital that can receive it. • The orbital with the lowest energy is the 1s orbital. The 2s orbital is the next highest in energy, then the 2p orbitals. • Beginning with the third main energy level (n = 3), the energies of the sublevels in different main energy levels begin to overlap. Sublevels in Order of Increasing Energy Notice that the 4s energy sublevel is lower than the 3d energy sublevel, which is lower than the 4p sublevel. The same pattern holds for the 5s, 4d and 5p sublevels. Things are even more complicated for higher energy sublevels. Notice that the 6s sublevel is lower than the 4f sublevel, which is lower than the 5d sublevel, which is lower than the 6p sublevel. Rule #2: The Pauli Exclusion Principle No two electrons in the same atom can have the same set of 4 quantum numbers. • In addition to the two quantum numbers already mentioned, there are two more: the magnetic quantum number (which indicates the orientation of an orbital around the nucleus) and the spin quantum number (which explains why a single orbital can hold a maximum of two electrons with opposite spin states). Pauli Exclusion Principle 1s orbital for helium This is referred to as an electron “pair”. Rule #3: Hund’s Rule • As many unpaired electrons as possible are placed in separate orbitals in the same sublevel. Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin state. Hund’s Rule CORRECT 2s 2p INCORRECT 2s 2p Hund’s Rule CORRECT 2s 2p INCORRECT 2s 2p Orbital Notation • In orbital notation, an unoccupied orbital is represented by a line, ____, with the orbital’s name written underneath the line. • An orbital containing one electron is written with one arrow, shown as _ ↑ _. • An orbital containing two electrons is written as _↑↓ _, showing the electrons paired and with opposite spin states. • The lines are labeled with the element’s symbol, the principal quantum number and sublevel letter. H ↑_ He _↑↓ _ 1s 1s Electron Configuration Notation No lines, no arrows—instead the number of electrons in a sublevel is shown by adding a superscript to the sublevel designation. Hydrogen’s configuration is 1s1. The superscript shows that one electron is present in hydrogen’s 1s orbital. Helium’s configuration is 1s2. The superscript shows that there are two electrons in helium’s 1s orbital. Sample Problem • The electron configuration of boron is 1s22s22p1. How many electrons are present in an atom of boron? What is the atomic number of boron? Write the orbital notation for boron. First draw the lines representing orbitals: Next, add arrows showing the electron locations, starting with the lowest energy levels: Your turn • The electron configuration of nitrogen is 1s22s22p3. How many electrons are present in an atom of nitrogen? What is the atomic number of nitrogen? Write the orbital notation for nitrogen. 7 electrons, 7, _↑↓ _ _↑↓ _ _↑_ _↑_ _↑_ 1s 2s 2p ASPEN HILTON WS Order of Fill of Orbitals • Would you like a simplerto-write-and-remember guide to filling orbitals than the energy diagram you colored above? • This should help! Start at the end of the top arrow, continue down the arrow to its tip then drop down a level to the end of the next arrow, etc.