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History of the Atom A Look at Some Famous Scientists History and Experiments • • • Comes from the time around 400 B.C. Ancient Greek philosopher He performed no experiments during his time Conclusions or Discoveries Came up with the word “atomos” • He said that the atom was a small indestructible molecule • It was the smallest piece of matter • DEMOCRITUS History and Experiments Conclusions or Discoveries 384-322 B.C Greek Philosopher Used reason rather than experiments Stated that all matter was made up of fire, wind, earth or water Views were unchallenged for 2000+ years Aristotle History and Experiments Conclusions or Discoveries 1. • English school teacher who developed the Atomic Theory in 1803 • Studied the atom in 1803 • Studied the masses of elements and compounds and their ratios 3. • Used experimental data to support his theory 4. 2. 5. All matter is made of indivisible and indestructible atoms Atoms of the same element are identical Atoms of different elements are different Atoms of different elements combine in whole number ratios and form compounds Chemical reactions consist of the combination, separation, or rearrangement of atoms. John Dalton 1. Law of Conservation of Mass Mass (matter) is neither created or destroyed, just rearranged. Mass of reactants = mass of products C + O2 12 g 2(16) g CO2 44 g History and Experiments • • Researched the atom back in 1897 He used a cathode ray tube and observed the cathode ray is attracted to a positive (+) charge Conclusions or Discoveries Discovered electrons • Plum pudding model of the atom • (Chocolate Chip Cookie) • J.J.Thompson Observations Inferences When using a magnetic field, the rays bent away from the field The rays must have a charge When using electric fields the rays were attracted to the (+) field. Support for rays having (-) charge. When a small paddle wheel was placed in the path of the cathode rays the wheel was set in motion. Rays must consist of tiny particles that have mass (electrons). Cathode Ray Tube History and Experiments • • • Some of the most important discoveries related to the atom were completed by Rutherford His work was completed in 1908 Gold Foil Experiment Conclusions or Discoveries Discovered the nucleus • The nucleus is very small and dense • The atom is mainly empty space • Ernest Rutherford James Chadwick In 1932, James Chadwick discovered the neutron, a neutral subatomic particle with relatively the same mass as a proton. The masses of all atoms except hydrogen were known to be greater than the combined masses of their protons and electrons. Theorized that there must a 3rd particle Neutron – neutrally charged particle (no charge) Nucleus – made up of protons and neutrons History and Experiments • • Conclusions or Discoveries Bohr’s research on the atom was completed in 1913 Still highly relevant today • Electrons orbited the nucleus Niels Bohr History of the Atom - Timeline 1766 – 1844 Antoine Lavoisier Thomson makesJ.J. a substantial the number discovers of contributions electron and to the field of proposes the Chemistry Plum Pudding Model 1871 in 1897 – 1937 Niels Bohr proposes the Bohr Model in 1913 1887 – 1961 James Chadwick discovered the neutron in in 1932 1700s 1800s 1900s 460 – 370 BC 0 Democritus proposes the 1st atomic theory 1743 – 1794 Erwin John Dalton Ernest Rutherford Schrodinger proposes performs his the Gold Foil describes 1891 – 1974 atomic theory Experiment in in 1909 the electron 1803 cloud in 1926 1885 – 1962 Click on picture for more information 1856 – 1940 Definitions Atom ◦ The smallest part of an element that upholds the chemical properties of that element Element ◦ A substance that can’t be broken down into simpler chemical substances. ◦ What’s smaller: an atom or element? The 3 Subatomic Particles Symbol Charge Location Mass (amu) Electrons e- negative Electron Cloud Protons p+ positive Nucleus 1 Neutron n0 Nucleus 1 neutral 1/1836 (= zero) PERIODIC TABLE OF ELEMENTS • Just like the 26 letters of the alphabet combine to form all the words in the English language, the 100 or so elements combine to form everything that exists in the world • About 90 of the elements on the PT are found naturally in nature, the rest have been created in a laboratory • Any material made of one type of atom is classified as an ELEMENT PERIODIC TABLE OF ELEMENTS • The Periodic Table is organized into ROWS and COLUMNS • Each vertical column is called a group or family • Elements within the same group have similar properties • Au, Ag, Cu • Each horizontal row is called a period • Properties of the elements gradually change when you move through a period • Elements get smaller when you move from LEFT to RIGHT How to Read the Periodic Table Each element is designated by its atomic symbol ◦ Some symbols do not match actual element name because they were named under their Latin name The first letter of an atomic symbol is capitalized and the second letter is lowercase How to Read the Periodic Table The atomic number is the special I.D. of the element ◦ It can tell you how many protons are in the element Protons have an electric charge of +1 Electrons are equal to the number of protons in an element ◦ Electrons have an electric charge of -1 ◦ Atomic # = # Protons = # electrons How to Read the Periodic Table The atomic mass of an element is the number at the bottom of the square Atomic mass is the sum of the protons and neutrons in the atom ◦ Neutrons have an electric charge of zero ◦ Atomic mass = protons + neutrons The nucleus of an atom contains all the protons and neutrons Let’s practice ISOTOPE MASS # ATOMIC # 35 17 7 3 # PROTONS # NEUTRONS # ELECTRONS 17 35-17 18 17 3 7-3 4 3 Isotopes Atoms of the same element that contain different numbers of neutrons are called Isotopes. We identify isotopes by their mass number. The mass number is the total number of protons and neutrons it contains. Mass # = # Protons + # Neutrons Isotopes One way to express isotopes Mass Number ◦ OVER Atomic Number The total number of neutrons in an isotope is to subtract its atomic number from the mass number. Mass number -Atomic Number Number of Neutrons Examples How many Protons? How many Neutrons? ◦ 6 protons, 6 neutrons How many Protons? How many Neutrons? ◦ 6 protons, 8 neutrons Atomic Mass Atomic mass is the average atomic mass of an element and its various isotopes. Measured in Atomic Mass Units (amu) ◦ Example: Imagine 100 students take a test worth 100 points. If all students scored an 80 then the average would be an 80%. But what would happen if one student gets a 100%? It would raise the average slightly. ◦ Same thing happens with isotopes Are these elements the same, different or isotopes 1. Element X has 45 protons and 50 neutrons, while Element Y has 50 protons and 45 neutrons Same, Different, Isotope? Element A has 13 protons and 15 neutrons, while Element B has 13 protons and 14 neutrons Same, Different, Isotope? Electromagnetic Spectrum All types of electromagnetic radiation share common characteristics: ◦c =λxν c - Speed of light (in a vacuum) = 3.0 x 108 m/s λ – wavelength ν - frequency Atomic Line Spectrum Atomic Line/Emission spectrum-set of frequencies of the electromagnetic wave emitted by atoms of the element. Each element is unique. Flame gives off colors Quantum Theory The quantum theory describes the probability of finding an electron in a given region of space called an orbital. Principal Quantum Number (n) n = principal quantum number Major energy levels Larger n, the further away electron is from nucleus and the higher its energy Values of n are whole numbers from 1 infinity N= Periods- Horizontal rows Sublevel (l) l = sublevel within a major energy level Designates the shape of the orbital Values of l are whole numbers from l = 0 to n – 1 N=2 L= 0 n-1 L=0,1 Sublevel (l) values = orbital type L value = 0 = s “spherical” L value = 1 = p “dumbbell” Sublevel (l) = orbital type L value = 2 = d “Daisy” L value = 3 = f “Fantastic” Magnetic quantum number (ml) Designates the orientation of the orbital in space Values of m are whole numbers: -L 0 +L For example, if l = 1 calculate m -1 0 +1 Therefore, m = -1, 0, +1 Spin quantum number (ms) Describes the direction of an electron’s spin Electrons spin on their axis either clockwise or counterclockwise Only two possible values: +1/2 or -1/2 Quantum Theory Summary Quantum numbers (n, l, ml, and ms) attempt to locate where electrons in an atom “should” be Quantum numbers are analogous to the information give on a ticket stub: Stadium, Section, Row, and Seat Practice Problems n=4 ◦ l= 4-1=3 ◦ ml= -3,-2,-1,0,1,2,3 -2,-1,0,1,2 -1,0,1 0 ◦ ms= +1/2, -1/2 0,1,2,3 n=2 ◦ l= 2-1=1 ◦ ml= -1,0,1 0,1 0 ◦ ms= +1/2, -1/2 n=1 ◦ l= 1-1=0 ◦ ml= 0 ◦ ms= +1/2, -1/2 0 Electron Shells We identify elements based on how the electrons are arranged in shells. At least 7 shells and each shell can hold a certain amount of electrons. Electron Configuration For full configuration you must have all prior parts to it as well! Fluorine: 1s2 2s2 2p5 All adds up to 9! Francium 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s1 All adds up to 87! s-block Subshell Orbital blocks p-block d-block f-block Quick Review Quick Review Orbital Notation Orbital Notation is a way to show how many electrons are in an orbital for a given element. They can be shown with arrows. Take what we just learned with electron configuration but take it a step further. Each line can hold up to 2 electrons. S= ____ P= ____ ____ ____ D= ____ ____ ____ ____ ____ F=____ ____ ____ ____ ____ ____ ____ Pauli Exclusion Principle A max. of two electrons can occupy each orbital and these electrons must have opposite spin Pauli Exclusion Principle Aufbau Principle Electrons will occupy the lowest energy levels possible Aufbau (German) = “building up” Hund’s Rule Electrons enter orbital's of the same energy one at a time before pairing up It is more STABLE to have the max. number of unpaired electrons Hund’s Rule Mole Conversions 1 mole of any substance has 6.02 x 1023 atoms ◦ Avogadro’s number 1 mole of any substance is equal to its mass on the periodic table in grams. ◦ Molar mass Lewis Dot Structure Depending on what group they are in on the periodic table determines how many valence electrons they have. You cannot double up valence electrons until each side of the element has an electron ◦ Example: ◦ C C Lewis Dot Structure Practice Try creating Lewis Dot structures for the first 10 elements! H He Li B Be Lewis Dot Structure Practice C N O F Ne Answers