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Bell Ringer List 3 facts about atoms, which you’ve learned in past courses (do not talk with neighbors… come up with your own ideas!) CHAPTER 3 Protons, Neutrons, Electrons Atomic Mass, Ions and Isotopes Atom History Atom Models Atom • Greek for “Atomos” which means “indivisible” • Smallest unit that can exist alone or in combination with other atoms Structure of an Atom 1. Nucleus – contains protons (p+) and neutrons (n0) – dense, small 2. Electron Cloud – contains electrons (e-) – surrounds nucleus, mostly empty space – arranged in shells or energy levels Atom itself is neutral: Why? Subatomic Particles Phet Interactive • Watch for patterns: – What happens when a proton, electron or neutron is added to the charge? How about the mass? – What element has which particles in it? – Which atoms are stable? Which are not? • List the parts of an atom that have mass. • List the charges of each subatomic particle. Subatomic Particles Particle Symbol Charge Mass Number Proton p+ +1 1 Neutron n0 0 1 Electron e- -1 0 Atomic Number – number of p+ in the nucleus of an atom (always equal to number of e-) ‘Big Blue’ Atomic mass – the average mass of the isotopes ‘little red’ Mass Number – number of p+ and n0 in the nucleus of an atom The mass number is the atomic mass rounded off to a whole number!! Periodic Table How many protons, neutrons and electrons are in aluminum? p+ = Atomic Number (Big Blue) n0 = Mass Number – p+ Mass Number is the Atomic Mass (little red) rounded to a whole number Mass Number = p+ + n0 e- = p+ because atoms are electrically neutral p+ = 13 n0 = 27-13 = 14 e- = p+ = 13 Atomic Mass Atomic # Mass # # p+ # e# n0 Try Sodium (Na): Atomic Mass = Atomic # = Mass # = # p+ = # e- = # n0 = = = = = = = (Avg. mass of isotopes) (number of p+ or e-) (Atomic mass rounded) (same as Atomic #) (same as Atomic #) (Mass # - Atomic #) ________ ________ ________ ________ ________ ________ = = = = = = 32.065 16 32 16 p+ 16 e32 – 16 = 16 n0 Find the number of protons, neutrons and electrons for potassium. 1. 19 39.098 K 2. Atomic Number = 19 Atomic Weight = 39.098 Mass Number = 39 Proton = Neutrons = Electrons = Find the number of protons, neutrons and electrons. 1. 19 39.098 K 2. 84 36 Kr 3. Uranium -235 Protons Neutrons Electrons K 19 20 19 Kr 36 48 36 U 92 143 92 Bell Ringer Define: Atomic Mass (also known as atomic weight) Mass Number Atomic Number More on Structure Atom is very, very small overall • Measure in Angstroms (1 Å = 1 x 10-10 m) • • Nucleus is extremely tiny compared to the electron cloud How much smaller??? Nucleus is a marble in a football stadium. Nucleus “marble” Electron Cloud “football stadium” Particles in the Atom Nucleus – protons – neutrons Cloud – electrons mass: 1 amu mass: 1 amu mass: “0” amu (5.45 x 10-10 amu) 1 amu (atomic mass unit) = 1.673 x 10-24 g 0.000000000000000000000001673 g amu – special unit for the mass of an atom – 1/12 mass of carbon-12 atom Similar to: • 2000 lbs = 1 ton • 1 amu (atomic mass unit) = 1.673 x 10-24 g • Ion: When the # of p+ ≠ # e– This changes the charge of the atom • Isotope: When two atoms of one element have different masses. Mass Number v Atomic Weight Atomic weight: The average atomic mass of all the isotopes for an element – Measured in “amu” (atomic mass units) – AKA atomic mass Mass Number: The mass of a single atom for one isotope of an element - measured in amu Bell Ringer • What happens when you change the protons, neutrons, or electrons in an atom? Example of an Average Mass You have a bag of marbles. Half of the marbles have a mass of 5 g, 30% of the marbles have a mass of 2 g, and 20% of the marbles have a mass of 1 g. What is the average mass of a marble in your bag? Mass of marble % 5g .5 2g .3 1g .2 Avg. (5g .5) (2g .3) (1g .2) Avg. 2.5g 0.6g 0.2g 3.3g Example: The 4 Isotopes of Chromium (amu x %) Isotope Atomic Mass Units % Cr - 50 49.9461 amu .0435 2.17 amu Cr - 52 51.9405 amu .8379 43.52 amu Cr - 53 52.9407 amu .0950 5.03 amu Cr - 54 53.9389 amu .0236 1.27 amu 2.17amu 43.52amu 5.03amu 1.27amu 51.99 amu Aristotle – 350 BC • All matter composed of 4 elements: earth, air, fire, water • People believed for over a thousand years…into the middle ages Democritus – 400 BC • Greek Philosopher • The idea of the atom • Not scientific, based on philosophy not evidence • First to use ‘atomos’ the smallest particle of matter that cannot be divided Antoine Lavoisier - 1782 • With the recent invention of the balance - he did a lot of quantitative measurements • Leading him to the Law of Conservation of Mass – which states mass cannot be created nor destroyed • Mass of reactants will equal mass of products Joseph Proust - 1799 • Law of constant composition–a compound is composed of exact proportions of elements by mass regardless of how the compound was created NaCl by mass, is always: 39.34% Na and 60.66 % Cl John Dalton – 1803 • English Chemist/schoolteacher • Law of multiple proportions – if compounds are composed of the same elements, the masses of the elements can be expressed as ratios of whole numbers. Carbon monoxide CO C = 12 g O = 16 g C = 12 g O = 32 g Ratio 1 : 1 Carbon dioxide CO2 Ratio 1 : 2 Same elements but different ratios make different compounds Dalton’s Atomic Theory 1. All matter is composed of extremely small particles. 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 divided, created, nor destroyed. 4. Atoms combine in whole number ratios to form compounds. 5. Atoms are combined, separated, or rearranged in chemical reactions. What ideas did each of these people have about the atom? Democritus Aristotle Lavoisier Dalton Joseph John Thomson - 1877 • Discovered electrons had a negative charge • Cathode Ray Tube How did we figure out all this stuff about the atom??? Because the atom is so small !!!!! Joseph Thompson THE ELECTRON! Used a cathode ray tube to prove there were negative charged particles (now known as electrons) in an atom. This opened the way to the idea that an atom was not just a solid sphere not able to be broken down anymore. 1897 Robert Millikan - 1909 • Oil Drop Experiment • Figured out the charge of the electrons Ernest Rutherford - 1911 • Gold Foil Experiment • Nucleus, very small and very dense • Positive charge part was a proton Rutherford’s Model • There was a nucleus – very small, very dense • It is positively charged • Atoms are mostly empty space James Chadwick - 1932 • English physicist • Worked with Rutherford • Discovered the neutron Niels Bohr - 1913 • Danish physicist • Worked with Rutherford • Theory explained how the electrons moved around the nucleus • Electrons moving around nucleus can only follow certain orbits What are the four main models of the atom? Name each, and draw a picture to represent it The Evolving Atomic Model Dalton Model Thomson Model 1803 – John Dalton believed that an atom was an indestructible particle with no internal frame. 1897 – J.J. Thomson discovers the electron. He believed electrons were embedded in positive charge sphere. (Solid Ball Model) (Plum pudding Model or chocolate chip cookie dough) Rutherford Model Bohr Model 1911 – Ernest Rutherford discovers that there is a dense, positively charged nucleus. Electrons go around the nucleus. 1913 – Niels Bohr enhances Rutherford’s model by having electrons move in a circular orbit at fixed distances from the nucleus. (Nuclear Model) Bohr Model Nucleus Shells K L e- = 2 e- = 8 M N p+ n0 e- = 18 e- = 32 Fill shells inside to outside Bohr Model of the Atom – Similar to concert seating! Put the number of p+ , n0 , and e - in the diagram as shown… Lets diagram Sulfur: p+ = Atomic # = 16 e - = Atomic # = 16 n0 = Mass # - Atomic # = 32 – 16 = 16 Maximum electrons = 2 (seats) p+ = 16 n0 = 16 e- = 2 8 e- = 8 18 32 e- = 6 e- = Now try Bromine… p+ = n0 = e- = e- = e- = e- = Now try Phosphorus… p+ = n0 = e- = e- = e- = e- = Bell Ringer 1. What is the difference between an ion and an isotope? Please be checking side board for assignment updates