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Atomic Structure Chemistry 112 History of the Atom: Time Line • Find the following scientists in your textbook. They are not given to you in the order they are found in the text. That would be too easy! Some may be found in other chapters besides 4 and 5!! • Write down the years they were alive or the year that their main discovery was made. • Write down their contributions to the development of today’s model of the atom. • Create a creative, colorful, and informative timeline! Scientist to Include on Timeline: • • • • • • Rutherford Millikan Dalton Marie and Pierre Curie Democritus Thompson • • • • • • Roentgen De Broglie Heisenberg Chadwick Bohr Aristotle Most of the following terms will be a review for you What is an Atom? • Atom – The smallest particle of an element that retains the properties of the element. • How big is an atom? VERY SMALL!!! • World Population 6,000,000,000 • Number of Atoms in a penny: 29, 000,000,000,000,000,000,000 • Scanning tunnel microscope allows us to see atoms Subatomic Particles found in the nucleus of an Atom: • Neutron: neutral (no charge) particle found inside the nucleus of an atom. • Proton: positive particle found in the nucleus of an atom. • Both the proton and neutron are approximately the same mass (1 amu) Subatomic particle outside of the nucleus: • Electrons are negatively charged particles found outside the nucleus of an atom. • The mass of an electron is negligible, only 1/1836th of an amu. • The mass of the electrons do not enter into calculating the atomic mass of an atom. Atomic Mass Unit • Abbreviated: amu • Chemists developed a method of measuring the mass of an atom without using very small numbers in scientific notation. • They chose an atomic standard: Carbon-12 • They agreed the mass of Carbon-12 was 12 amu. • Therefore, 1 amu is 1/12 the mass of a Carbon-12 atom. Atomic Number • The whole number next to each element that represents the number of protons in an atom. (Proton number can never be altered). • In a neutral atom the proton number will equal the electron number. However atoms can lose or gain electrons, even though they can’t lose or gain protons under normal circumstances. Atomic Mass • The atomic mass (when rounded to a whole number) represents the # of protons + the number of neutrons. • If you subtract the atomic number (# of protons) from the atomic mass (rounded to a whole #) you will find the number of neutrons in an atom. Isotopes • Atoms of the same element can be found in nature with a different atomic mass than others. This is because they can have different #’s of neutrons • An Isotope of an element is an atom of the same element with a different # of neutrons. Isotopes • Isotopes can be expressed two ways: – 1. C-14 (where 14 would be the whole number mass) – 2 146C (Where 14 is the mass and 6 is the atomic number) Why is the Atomic Mass not a Whole Number? • Atomic masses have decimals! • Atomic mass – weighted average mass of the isotopes of that element. • For Example: Mass of Cl is 35.453 • Isotopes: 75% Chlorine-35 25% Chlorine-37 Finding Protons, Neutrons, and Electrons • Example: K-39 – This is a potassium atom with a mass of 39 – Look on the periodic and see the the atomic number is 19. – Therefore # of protons = 19 (equal to atomic #) – Since it is a neutral atom (no charge) number of electrons = 19 – Number of neutrons = 39 – 19 = 20 Ions • Ions are charged atoms that have lost or gained electrons. • Two types of ions: – Cations: positive ions that have been formed by the loss of electrons. (Ca+2) – Anions: negative ions that have been formed by the gaining of electrons. (Cl-1) 4.4 Unstable Nuclei and Radioactive Decay • Late 19th century scientists noticed some elements spontaneously emitted energy and particles called radiation. • Elements that give off radiation are said to be radioactive. • Thus, atoms are not unchangeable as Dalton once thought. • Nuclear reactions – a reaction that involves a change in the nucleus of an atom. • Radioactive decay – when nuclei are unstable and gain stability by emitting radiation. • Fill out the Chart for Types of Radiation Types of Radiation Alpha Beta Gamma What is it made of? 2 protons and 2 neutrons (like the nucleus of a Helium atom) An electron Electromagnetic radiation Mass 4 amu 1/1836th of an amu No mass; energy Charge positive negative none Penetrating ability low medium high What stops it? paper Foil/metal lead Information from Chapter 25 • Transmutation – conversion of an atom of one element to an atom of another element by spontaneous emission of radiation. • Induced Transmutation – nuclear reactions produced artificially by striking a nucleus with a high-velocity charged particle. • Transuranium elements – all elements after Uranium on Periodic Table. – Produced in laboratory by induced transmutation – All are radioactive Radioactive Decay Rates • Half-life – the time required for one-half of a radioisotope’s nuclei to decay into its products. (see table 25-4) – Ex. Strontium-90 half life is 29 years • If today I have 10.0 g, in 29 years I would have 5.0g • In another 29 years (total of 58), I would have 2.5 g Nuclear Fission/Fusion • HUGE amounts of energy produced by Nuclear reactions. • Nuclear Fission – splitting of a nucleus into fragments to become more stable. • Used in nuclear power plants (controlled) • And nuclear bombs (uncontrolled) • Nuclear Fusion – Combination of nuclei to form a more stable nucleus. • Large energy released! • Sun powered by Fusion of hydrogen into helium. • Requires extremely high temperature to occur. • Scientist researching Cold Fusion Atomic Model Development… • Around the end of the 1700s Dalton believed that an atom was a solid indestructible mass and so did most scientists. • In the 1800s when JJ Thomson discovered the electron (first subatomic particle) that theory was shattered. Model Continued • The other subatomic particles were found and further disproved the idea of the atom being a solid mass. • Niels Bohr, who was a student of Rutherford, came up with the planetary model where electrons followed specific paths around the nucleus. He also stated that for an electron to move from one energy level to another it gained or lost a quantum of energy. Model cont… • Around 1925 Erwin Schrodinger came up with the quantum theory of electron placement which put electrons in clouds or regions not in specific paths. These areas of probability for electron location is based on a mathematical calculation. Identifying Where Electrons are Located in the Atom • The number and arrangement of electrons around the nucleus determines the atoms chemical properties. Specifically the outer electrons (valence electrons) Therefore, scientists need a shorthand way to write out electron arrangement. • Electron Configuration electrons in an atom. – Arrangement of • Principle Quantum Number – Energy level the electron occupies (n) • Examples: n= 1, 2, 3, 4, 5, 6, 7 The larger the value of n, the farther away the electrons are from the nucleus and the higher the energy of the electron. Electrons are lazy, so they want to be as close to the nucleus as they can so they can expend the LEAST amount of energy. Only a certain amount of electrons can fit in each energy level. The way to find the maximum in each level is: 2n2 •o •o •o •o Example: n =1 Can have 2(1)2 electrons = 2 electrons in energy level 1 You solve for n = 2 n=3 n=4 How many energy levels does hydrogen have? • Helium? • Lithium? • Sodium? • Calcium? • Xenon? Sublevels • Sublevels – further explain where the electrons are located in each energy level. The names of the sublevels we use are: s, p, d, and f Sublevels The number of sublevels in each energy level is equal to n •o n = 1 has one sublevel s •o n = 2 has two sublevels s, p •o n = 3 has three sublevels s, p, d •o n = 4 has four sublevels s, p, d, f •o n = 5 has five sublevels but the fifth one is not used in ground state elements. • Notice, energy levels must have an s sublevel before a p! Orbitals • Orbitals – Each sublevel has specific orbitals the electrons can be in: • Each orbital can hold 2 electrons. Therefore, the maximum number of electrons in the s orbital is 2. S has 1 orbital maximum electrons = 2 P has 3 orbitals maximum electrons = 6 D has 5 orbitals maximum electrons = 10 F has 7 orbitals maximum electrons = 14 Order of Filling • http://lectureonline.cl.msu.edu/~mmp/perio d/electron.htm • Three Main Rules Electrons Abide By: Aufbau principle: (“Electrons are Lazy”) each successive electron occupies the lowest energy orbital available. Exclusion Principle Pauli (“Two to Tango”): maximum of two electrons may occupy a single atomic orbital, but only if the electrons have opposite spins. • Hund’s Rule (“Why share if you don’t have to?”) : Single electrons with the same spin must occupy each equal-energy orbital before additional electrons with opposite spins can occupy the same orbital. • Orbital Diagrams Used to show how electrons are distributed within sublevels and to show the direction of spin. Boxes are used to represent orbitals. Arrows are used to represent electrons. The first electron in the orbital is represented by an arrow pointing up, , meaning clockwise spin. The second electron in the orbital is represented by an arrow pointing down, , meaning counterclockwise spin. An electron configuration notation is an abbreviated form of the orbital diagram. • Electron Dot Diagrams Valence Electrons – electrons in the atom’s outermost energy level. These electrons are involved in forming chemical bonds. They are represented visually by an electron dot structure. Also known as Lewis Dot Formulas. • Examples: