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1 ATOMIC STRUCTURE A. ATOMIC THEORY Due to contributions of numerous scientists 1. John Dalton: basic unit of matter - tiny particle called the atom. Dalton’s Theory A. All matter is composed of extremely small particles called atoms, which cannot be subdivided, created, or destroyed. B. All atoms of the same element are identical C. Atoms of different elements are different D. In chemical reactions, atoms can be combined to form compounds. E. In chemical reactions, atoms can be combined, separated, or rearranged, but never created, destroyed or changed. Summary: Most of Dalton’s Theory was accurate, but some was incorrect to some degree. 2. J.J. Thomson: atom was made up of EVEN SMALLER PARTS Summary: So: If electrons are present in atoms, then what makes up the rest of the atom? 3. Ernest Rutherford: Sought to answer this question. -Rutherford’s Gold-Foil Experiment (KNOW!!!) What did Rutherford’s experiment show? (KNOW!!) 1. 2. 4. Neils Bohr : Proposed a model of the atom showing a central dense nucleus with electrons found in surrounding orbits. 5. The present day model of an atom –-> The WAVE MECHANICAL MODEL -Central Nucleus, with clouds of electrons around the nucleus. -Electrons are found in surrounding orbitals -Orbital: *Modern model of the atom has evolved over a long period of time through the work of many scientists. Some other scientists: 1) Planck –said that atoms absorb energy in discrete amounts called quanta 2) Chadwick –Discovered neutrons in the nucleus of an atom 3) Moseley –Coined the term Atomic Number 2 B. THE ATOM Protons, neutrons, and electrons are all inside the atom, called: ___________________ SUBATOMIC PARTICLE COMPARISON (MEMORIZE!) LOCATION CHARGE SYMBOL MASS PROTON NEUTRON ELECTRON C. ATOMIC NUMBER: ATOMIC # = _______________________ (in the nucleus of an atom) -In a neutral atom the # of protons = the # of electrons **Atomic # is indicated in bold in the bottom left corner of each elements’ box on The Periodic Table** The atomic # for Carbon (C) is ____________ How many protons does a neutral C atom have? _________ How many electrons? ___________ Atomic # Ex. 2: The element Lithium (Li) has an atomic # of _______ How many protons does a neutral Li atom have? _________ How many electrons? _________ D. MASS NUMBER: MASS # = ___________________ + ______________________ # OF NEUTRONS = _______________ - ___________________ Standard Example: Mass # 40 Element Symbol X 20 Atomic # # of neutrons = ___________ 3 Additional Examples: 1. An atom of Lithium (Li) has a mass # of 7. Find the following: Atomic #: _______________ # of protons: ____________ # of neutrons: ____________ # of electrons: ___________ 2. Find the following with regard to 39K: Mass #: __________________ Atomic #: ________________ # of protons: _____________ # of neutrons: ____________ 3. Find the following with regard to Nitrogen-14: Mass #: _______________ Atomic #: ______________ # of protons: ____________ # of neutrons: ____________ # of electrons: _____________ 4. An atom has 15 protons and 16 neutrons. What is its atomic #? _______________ What is its mass # ? ________________ What element is this? _______________ # of electrons: ____________ E. ISOTOPES AND ATOMIC MASS ISOTOPE - atoms of same element that have SAME _________ but DIFFERENT __________ Isotopes of atoms can be stable or unstable. Stability of an isotope is based on the # of protons and neutrons in its nucleus. If the nucleus is unstable, the isotope may spontaneously decay and emit RADIATION. EXAMPLES: Hydrogen has 3 isotopes: Carbon has 2 isotopes: ATOMIC MASS: WEIGHTED AVERAGE MASS of all the NATURALLY OCCURING ISOTOPES of that element. o Represented by a decimal # at the top left hand corner of each element box. Atomic Mass 4 F. CALCULATING AVERAGE ATOMIC MASS Example: Chlorine has two naturally occurring isotopes – Cl-35 and Cl-37. 76% of the isotopes of Cl are Cl-35, and 24% of its isotopes are Cl-37. Calculate the average atomic mass of chlorine. Step Calculation Multiply each % by the appropriate mass # 76 x 35 = 2660 24 x 37 = 888 2660 + 888 = 3548 3548 / 100 = 35.48 Add these answers together Divide by 100 Examples: Data in various chemical handbooks indicate the relative percentages of these naturally occurring isotopes of Magnesium (Mg): 24Mg = 79.3% , 25Mg = 10.1%, 26Mg = 10.6%. Given this data, calculate the average atomic mass for Magnesium. G. PRINCIPAL ENERGY LEVELS Show how far away the electrons are from the nucleus There are up to 7 PEL’s that surround the nucleus Electrons in _______________________________ have __________________________ Low PEL # = ___________________________ = ____________________________ High PEL # = ___________________________ = ____________________________ Diagram of Atom and PEL’s surrounding nucleus: How many electrons can each energy level hold? Principle Max # of electrons Energy Level (n) level can hold (2n2) 1 _____electrons 2 _____ electrons 3 _____ electrons 4 _____electrons 5 H. ELECTRON CONFIGURATIONS AND VALENCE ELECTRONS ELECTRON CONFIGURATION: o Can be found below the atomic # in the bottom left hand corner of each element box. Electron Configuration o Examples: 1. Magnesium (Mg): Atomic #: ___________ 2. Zinc (Zn): Atomic #: ____________ # of electrons: ________ # of electrons: __________ Electron Config: ____________ Electron Config: ______________ # of PEL’s occupied: ______________ # of PEL’s occupied: __________ VALENCE ELECTRONS: o participate in chemical reactions with other atoms o last # in electron configuration electron configuration on periodic table shows # of valence electrons a particular element has o Examples: 1. How many valence electrons does Mg have? ______ 2. How many valence electrons does Zn have? ______ NON-VALENCE ELECTRONS: All the other inner electrons. I. ELECTRON DOT STRUCTURES (AKA. LEWIS DOT STRUCTURES) A diagram that depicts valence electrons ________ around the symbol of the element Method: X Draw the electron dot structures for the following elements: 1. Helium 3. Sulfur 5. Boron 7. Nitrogen 9. Neon 2. Lithium 4. Calcium 6. Silicon 8. Chlorine 10. Iodine 6 J. GROUND STATE VS. EXCITED STATE GROUND STATE: low-energy electron configuration, e- fill in from inner PELs to outer PELs o Periodic table shows ground state e- config EXCITED STATE: An atom can become excited when its electrons move to a higher energy level without fully filling up the lower energy level. o How do they do this? ___________________________________________________ o Excited electrons move back down to ground state by ___________________ energy. o Energy released called: ____________________ Examples: • • • • Fluorine - ground state e- config (Periodic Table) 2-7 Possible excited state e- configs 1-7-1 2-6-1 1-8 An inner level is left partially filled b/c an e- has jumped to a higher PEL TOTAL # OF e- REMAINS THE SAME! 1. Which is the electron configuration of an atom in the excited state? A) 2 – 8 – 2 B) 2 – 7 - 1 C) 2 – 8 – 1 2. Which electron configuration represents an atom in the excited state? A) 1 – 2 B) 2 – 3 C) 2 – 1 D) 2 – 8 – 3 D) 2 - 7 BRIGHT LINE SPECTRUM: Ex: The bright-line spectrum of lithium looks like this: So: If you have an unknown element and you want to see if it is lithium, compare the bright-line spectrum of the sample you have with the bright line spectrum of lithium. FLAME TESTS: Which sample is lithium? 7 K. IONS IONS ARE _____________________________ Atoms are called ions when they are no longer neutral atoms Atoms become ions by: __________________________________________ If an atom GAINS ELECTRONS its charge becomes ___________________ If an atom LOSES ELECTRONS its charge becomes ___________________ The charge of an ion is indicated as a superscript! (If there is no superscript, it is a neutral atom.) How many protons does N+3 have? _________ How many neutrons does N+3 have? ________ How many electrons does N+3 have? ________