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
DEMOCRITUS DALTON THOMSON RUTHERFORD BOHR DeBROGLIE The Greeks - Democritus • 450 B.C. • Believes there is a fundamental particle that cannot be broken. ( won’t subdivide to smaller components) • The particle is invisible / indivisible • Coins the term “atomos” John Dalton • 1810 • Believes atoms are invisible and indivisible. • All atoms of the same element are alike in every respect – especially mass. • Atoms of different elements are different in every respect – especially mass. • -- combine in whole number ratios • --Billiard Ball Theory Early symbols representing compounds in various simple ratios. Dalton J.J. Thomson • 1897 • An atom consists of a positively charged globule. • Negative particles, called electrons, are embedded in the globule. • These electrons move about in the globule. • Jellied salad or raisin bun theory. (Cathode Ray Tube) JJ Thomson’s model of the atom Positive charged mass with embedded electron particles electrons Positive mass RUTHERFORD EXPT Ernest Rutherford / Nuclear model • 1908 • Performs his famous scattering experiment HYPOTHETICAL Observations - Most alpha particles pass through the foil as if there was nothing there - A few veer off as they are slightly deflected - A very few bounce back at the alpha gun Conclusions - Atoms are mainly empty space (over 99.9 %) -The mass of the atom is concentrated in a core or nucleus which is positively charged. - Electrons, which are negative move about in a large region of space outside the nucleus Rutherford Model Negative Electron Positive Nucleus Quantum Mechanical Model OR the Nuclear Atom I. General Information 1. Atoms have a very dense core called the nucleus.The nucleus contains two sub-atomic particles: A) Protons which have a positive charge B) Neutrons which do not carry a charge (neutral) 2. The electrons, which carry a negative charge, are in a state of constant motion in a relatively large region of space away from the nucleus. 3. Atoms are mainly empty space. (over 99.9 %) NUCLEAR MODEL RUTHERFORD - DISCOVERED NUCLEUS -POSITIVE CHARGED PROTON -ELECTRONS ORBIT OUTSIDE Described as - - SOLAR MODEL Neils Bohr Energy Levels Electrons are located at different energy levels due to energy of the electrons Quantum model Orbitals s p d f s – 2 electrons p – 6 electrons d – 10 electrons F – 14 electrons [ Electron configuration ] Neils Bohr ENERGY LEVELS - 7 Levels for 7 periods Electron orbits / electron configuration s p d f orbits / electron location ELECTRON LOCATION • Electrons are located in specific energy levels • These energy levels can be thought of as globe shaped regions of space surrounding the nucleus. • Electrons in the inner level have the lowest amount of energy and energy increases as the levels move outward. • Electrons enter and fill these levels in order and each level has a limit to how many electrons it can hold. • Level 1 holds 2 electrons • Level 2 holds 8 electrons • Level 3 holds 8 electrons • Level 4 holds 18 electrons • Level 5 hold 18 electrons • Level 6 holds 32 electrons • Level 7 holds 32 electrons Shows the 3 orbitals and how they overlap DUST CLOUD MODEL De Broglie - Model Probability of finding electrons at different levels Electrons are particles / bundles of Energy. Dual principle theory – wave mechanics / particle theory PROBABILITY OF ELECTRON LOCATION (DOTS) II. Summary of Atomic Particles Particle Symbol Relative Relative Location Mass charge Proton H or p+ 1 +1 Nucleus neutron n or n 1 0 Nucleus electron e or e-1 0 -1 Constant motion outside nucleus Relative => means with respect to each other About 1864 ( 2000) electrons = 1 proton III. Atomic Number • This is used to identify the element(name and/or symbol) • This number ALWAYS tells us the how many protons are in the nucleus of atoms of that element. • NO EXCEPTIONS • Examples: Carbon atoms (atomic # of 6) => 6 protons Sodium atoms ( atomic # 11 ) => 11 protons Osmium atoms ( ? ) => ? Protons ?????? Atoms (atomic # 47) => 47 protons IV. Nuclear Charge • This is the total positive charge present on the nucleus. • It is always positive because the protons are the only charged particle in the nucleus. • Therefore, the nuclear charge is equal to the number of protons in the atom or ion. (equal to the atomic number) • Examples: The nuclear charge on K is + 19 The nuclear charge on Ni + 28 V. ATOMIC MASS • The mass number is the total mass in grams for a constant and specific number of atoms of each element. • This number is very large (because atoms are very small) and is known as the mole number. • 1.0 mole is equal to 6.02 x 1023 particles. Examples: The mass of one mole of Hydrogen atoms is 1.01 g. Therefore 1.01 g/mol is the mass number for Hydrogen. ATOMIC MASS OF Zr 91.22 g / mol MASS NUMBER - ATOMIC MASS ROUNDED TO A WHOLE NUMBER. ATOMIC MASS ROUNDED MASS NUMBER CARBON - ATOMIC MASS OF 12.01 => 12 amu SILVER - ATOMIC MASS OF 107.87 => 108 amu ANTIMONY - ATOMIC MASS OF 121.75 => 122 amu THE NUMBER OF NEUTRONS IS CALCULATED AS: MASS NUMBER - ATOMIC NUMBER = NEUTRONS VI. Number of Neutrons • Since only protons and neutrons contribute to the mass of an atom, the sum of these two particles must be equal to the mass number. • To calculate the number of neutrons in a given atom: MASS NUMBER – # OF PROTONS = # OF NEUTRONS Examples:How many neutrons are in each of the following? 1. A carbon atom? 12 – 6 = 6 2. A zinc atom? 65 – 30 = 35 3. 44.96 Sc 21 Scandium has 45 – 21 = 24 neutrons VII. ISOTOPES All atoms of the same element are not identical. Isotopes - 2 or more atoms of the same element that have (same atomic number and same number of protons) Different number of neutrons ===> Different mass Hydrogen 1 1 1 H 1 proton 0 neutrons Protium Occur naturally Hydrogen 2 2 1 H 1 proton 1 neutron Deuterium Hydrogen 3 3 1 H 1 proton 2 neutrons Tritium Man-made / synthetic Please do these in your NOTE BOOK !! A carbon isotope has 8 neutrons. What is its mass number? Answer: 14 An isotope of zinc has a mass number of 62. How many neutrons does it have? Answer: 32 A given isotope has a mass number of 52 and contains 26 neutrons. What element is it? Answer: Iron (Fe) VIII. WRITING ISOTOPES 2 methods to represent the different isotopes of an element 1.) X(Y) where ‘X’ is the element symbol and ‘Y’ is the mass number of that particular isotope. Examples: H(1) H(2) H(3) C(12) C(14) U(238) Zn(64) A Z X where ‘X’ is the element symbol, ‘A’ is the 2.) mass number of the isotope and ‘Z’ is the atomic number Examples: 14 6 65 30 C , Zn , 238 92 90 38 U , Sr etc. IX. CALCULATING AVERAGE MASS • An example: A theoretical element • Element J has 4 known isotopes. Their mass distribution in nature is as follows: 1) J(40) = 30.0% 2) J(43) = 40.0% ( largest % in nature) 3) J(44) = 10.0% ( smallest % in nature) 4) J(50) = 20.0% Calculate the average mass of this element ( J ) Take the % of each mass and add the results together to get a total. Each Isotope may occur in different % in nature 30/100 x 40g = 12.0 g 40/100 x 43g = 17.2 g 10/100 x 44g = 4.4g 20/100 x 50.0g = 10.0g 43.6 g The average mass of element J is 43.6 g / mol Remember – the total percent of any element must be 100%. If you are given numbers like 20% and 40% for 2 of the 3 isotopes… 100 – ( 20 + 50) = 30 % for the third isotope. IMPORTANT ****when completing tables and answering questions on atomic theory, ALWAYS USE THE GIVEN DATA FROM THE QUESTION. The values from the periodic table are used if no other information is available. X. NUMBER OF ELECTRONS • In any neutral element , the number of electrons is equal to the number of protons. (equal to the atomic number) • That is: [ p+ = e-] XI. IONS Species where the number of protons is NOT equal to the number of electrons. That is: [ p+ = e- ] These are charged particles known as ions Type: II It is also possible to remove one or more electrons from an atom or ion. This will result in the formation of a positive ion. Energy must be supplied in order to remove each electron. REACTANTS Na + energy → Al + energy → PRODUCTS Na+ + eAl 3+ + Mg + energy → Mg 2+ + OIL <-> RIG 3e- 2e- This kind of reaction is known as an OXIDATION reaction **All ions are formed by the addition or removal of electrons. NEVER CHANGE THE NUMBER OF PROTONS.** Type I: It is possible for an atom to gain one or more electrons. This results in the formation of a negative ion. Energy is released as each electron is accepted. Examples: F + e- → F- + energy S + 2e- → S2- + energy N + 3e- → N3- + energy This kind of reaction is known as a REDUCTION reaction. To help remember oxidation and reduction !! (1 WAY) OIL OXIDATION IS LOSS OF ELECTRONS RIG REDUCTION IS GAIN OF ELECTRONS 2 ND WAY My name is LEO the lion says GER Oxidation - Reduction Left side is reactants Right side is products Na + Na becomes what type of ion ? ( + or - ) That indicates that it lost electrons….which are the products Losing electrons means gaining Energy Na + Energy == > Na + + Energy is a reactant 1e- Sodium plus energy == > sodium ion and 1 electron Predicting Ion Formation • Stability in ions is represented by the total number of electrons present. • The stable numbers are those found in atoms of the inert gases. (2, 10, 18, 36, 54, 86) These are unreactive – that is they are stable. • All other atoms will gain (become reduced) or lose (undergo oxidation) in order to achieve one of these numbers. Predict the ion that each of the following atoms would form in order to become stable. Write out the oxidation or reduction reaction. • • • • • • • • • K Cl O P Ca Ga Cs N Br ANSWERS K + energy → K+ + eCl + e- → Cl- + energy O + 2e- → O2- + energy P + 3e- → P3- + energy Ca + energy → Ca2+ + 2eGa + energy → Ga3+ + 3eCs + energy → Cs+ + eN + 3e- → N3- + energy Br + e- → Br- + energy Dmitri Mendeleyev given credit for the first working periodic table 1. The periodic table is arranged in horizontal rows and vertical columns. 2. Each horizontal row is called a PERIOD or SERIES. 3. Each vertical column is known as a GROUP or FAMILY. ( same chemical characteristics) Label the following on your blank table: - Families (show numbers & Roman numerals) - Periods (show number of each period) - Alkali metals ------ - -- - ( group 1A ) - Alkali earth metals - - - ( group 2A) - Transition elements - - - ( group B ) - Halogens ------ -- -- -- -- -- - - - (group 7A ) - Inert gases (noble gases)- - (group 8A ) VALENCE ELECTRONS • All of the electrons located in the outermost (highest) energy level are called valence electrons. • For all ‘A’ elements, the group number equals the number of valence electrons. Group I A to VIII A • Carbon is in group IVA and therefore has 4 valence electrons. • All transition (B family) elements have 2 valence electrons. Metallic Properties Metallic properties decrease as we move left to right across the table. Metallic properties increase as we move from top to bottom in the table. The most metallic element is in the bottom – left of the table and the least metallic element is in the top right. ATOMIC SIZE IN ANY FAMILY As atomic number increases – Atomic size increases This is due to more energy levels holding electrons. Eg. Cs atoms are larger than Na atoms because Na atoms have electrons in 3 energy levels holding electrons and Cs has 6 energy levels holding electrons IN ANY PERIOD As atomic number increases – Atomic size decreases This is due to increased nuclear attraction on electrons in the same energy level. Eg. A Br atom is smaller than a Ca atom. 35 protons pulling on outer electrons vs 20 protons pulling on outer electrons. THE ELECTRON AFFINITY ADDITION OF AN ELECTRON TO AN ATOM ELECTRON ADDED / ENERGY IS LOST MOVIE 1 MOVIE 2 IONIZATION ENERGY This is the amount of energy required to remove an electron(s) from an atom or ion. (See formation of a positive ion..) IN ANY FAMILY As atomic number increases – ionization energy decreases This is a function of size. The outermost electrons are further and further from the nucleus and are easier to remove. IN ANY PERIOD As atomic number increases – ionization energy increases This is due to increased nuclear attraction on electrons in the same energy level. BACKGROUND / Electron Configuration 1] Electrons are located in specific energy levels surrounding the nucleus. These are numbered consecutively from 1 – 7, starting from the nucleus and working outward. 2] Each energy level has a specific maximum capacity for holding electrons. 3] Energy levels can be sub-divided into sub-levels (or subshells). These are identified with letters: s, p, d, f…etc. 4] These sub-levels also have maximum capacities. 5] Each sub-level are, in turn, divided into orbitals that hold 2 electrons each. 6] Levels and sub-levels fill systematically with electrons starting from the inside (lowest energy) and working outward (highest energy). ELECTRON CONFIGURATION • That is: level one completely fills before level 2 starts filling…level 2 before level 3….etc. • Note: Not all energy levels contain all sub-levels. • Capacities: s – 2 electrons – 1 orbital 1 (2electrons) p – 6 electrons – 3 orbitals 3 (2electrons) d – 10 electrons – 5 orbitals 5 (2electrons) f – 14 electrons – 7 orbitals 7 (2electrons) 1s2 / 2s22p6 / 3s23p6 / 4s23d104p6 / 5s24d105p6 / 6s24f145d106p6 / 7s25f146d107p6 1 2 3 4 5 6 7 32 32 ENERGY LEVELS 2 8 8 18 18 Number of electrons in each energy level Notation Energy levels 1 2 3 1 s 2 2 s 2 2 p 6 3s 2 3p 6 electrons Orbitals / sublevel Prime quantum number/ energy level (refer to period on periodic table) Writing Electron Configurations / Review H: 1s1 He: 1s2 Li: 1s2 2s1 Ca: 1s2 2s2 2p6 3s2 3p6 4s2 Ag: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d9 U: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f4 Your turn --- make my day N: 1s2 2s2 2p3 Na: 1s2 2s22p6 3s1 Br: 1s2 2s22p6 3s23p6 4s23d104p5 VALENCE ELECTRONS Electrons in the last energy level / max 8 Bonding electrons / octets - only 8 F - 1s2 2s2 2p5 look at the last energy level F – 1s2 2s2 2p5 Filled up [ count only the s and p electrons 2 electrons + 5 electrons = Cl - has 7 valence electrons Cl - will gain 1 electron to become stable Cl becomes Cl 1- [ valence of 1- ] 7 Other Topics • • • • Last electron added (valence electrons) How to remember the sequence (two methods) Grouping quantum numbers Configurations of ions Short Cut Version / Noble Gas Method (Only use when instructions allow it) Write the symbol for the last complete inert gas and then continue the configuration from the next energy level. Example: Write the Noble Gas configuration for Ra. Ra: 86Rn – [ 7s2 ] Using this format, write configurations for: – [ 6s2 4f14 5d9 ] Au: 54Xe Ga: 2 3d10 4p1 ] Ar – [ 4s 18 Cf: Sc: 86Rn – [ 7s2 5f10 ] 2 3d1 ] Ar – [ 4s 18 Questions for Periodic Table Review Which Element? • • • • • • • • • • • • • Is an alkali metal? Is chemically inert? Is in group IV A? Is in period 2? Has the largest mass number? Is a transition element? Forms a 3 – ion? Forms a 2+ ion? Tends to lose one electron? Tends to gain 2 electrons? Is Lithium? Has a nuclear charge of + 20? has 2 valence electrons? • Are chemically similar? • Has the largest radius? • Has the lowest ionization energy? • Has valence e- in period 3? • Last e- added was d1? • Is the most metallic? • Configuration ends in 4p2? • Has 22 protons in nucleus? • Has e- in 3 energy levels only?