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History of the Atom Scientists and Their Discoveries 1 Atoms - + + + + - - • Atoms- smallest possible unit into which matter can be divided, while still maintaining itsis the For example, what smallest possible unit properties. into which a long essay can be divided and still have some meaning? • Made up of: + - – protons – neutrons – electrons • Electrically they are 2 NEUTRAL!!!! Elements • Element- made up of one kind of atom that can’t be broken down into simpler substances by physical or chemical means • 90 occur naturally on Earth • 29 were synthesized (made) by scientists 3 The Periodic Table of Elements-Reference • Periodic table tiles contain a lot of information and to understand it, it’s necessary to know the parts of the atom and some terminology concerning them 8 O Oxygen 15.999 Atomic Number (number of protons) Element Symbol (a capital letter or a capital followed by a lower case letter) Element Name Atomic Mass (weighted average of all isotopes mass) 4 The Atom’s “Center” • Nucleons- particles in the nucleus of atoms – Protons – Neutrons Notice that the electrons are not a part of the nucleus - + - + + 5 Neutrons - ++ + + + + + + - - - - - • Neutrons- neutral particles; have no electric charge – Help make up the nucleus of the atom – Contribute to the atomic mass 1.67 X 10-24 g 6 Protons (+) • Protons- positively charged particles – In nucleus – They ID an atom— atomic number – Contribute to the mass of the atom 1.67 X 10-24 g – Charge = + 1.6 X 10-19 coulombs (same value as electron but positive) - ++ + + + + + + - - - - - + 7 Atomic Number • Atomic number - the number of protons in the nucleus of an atom…its ID #!!!! • Represented by “Z” - - + ++ - What would be the atomic number of the atom to the left? What element is it? 8 Electrons (-) • Electrons- negatively charged particles – Outside the nucleus of the atom in electron orbits/levels – Move rapidly and create an electron cloud – Mass is insignificant – Valence electrons- the outermost electrons involved in the formation of chemical bonds - ++ + + + + + + - - - - - 9 Neutral Atoms • Most atoms are neutral and the number of protons = the number of electrons • 1 +protons and 1 –electrons=0 (neutral) • Atomic number = protons = electrons 10 Electromagnetic Force • Electromagnetic force is the force that results from the + + repulsion of like charges and the attraction of opposites + AND NEUTRALIZE ONE ANOTHER - • This the force that holds Notice how the the electrons around the particles with the nucleus same charge move • Bill Nye Atoms apart and the Why are neutrons not pictured above? particles with different charges move together. 11 Atomic Models • Model – a familiar idea used to explain unfamiliar facts observed in nature • Theory- an explanation of observable facts and phenomena • To remain valid, models and theories must: • Explain all known facts • Enable scientists to make correct predictions 12 Democritus (460 BC – 370 BC) • Proposed the existence of atoms from Greek word “atomos” which means “not to cut” or “indivisible” • Thought you could cut matter in half until you got an indivisible (not dividable) particle Image taken from: https://reichchemistry.wikispaces.com/T.+Glenn+ Time+Line+Project 13 Aristotle (384 BC – 322BC) • Rejected idea of the atom • Said matter could be cut continually • Aristotle was more influential than Democritus so atoms were forgotten about until late 1700’s 14 Antoine Lavoisier (1743 – 1794) “Father of Modern Chemistry” Generated a list of 33 elements Devised the metric system Discovered/proposed the Law of Conservation of Mass - matter can’t be created or destroyed, it just changes form (beginning mass = end mass) Image taken from: www.ldeo.columbia.edu/.../v1001/geo time2.html 15 John Dalton (1766 – 1844) In 1803 he proposed first experimentally based Atomic Theory that states atoms: o are building blocks of matter o are indivisible o of the same element are identical o of different elements are different o “Billiard Ball Model” 16 John Dalton (1766 – 1844) o Dalton’s Atomic Theory also explained The Law of Multiple Proportions -ratio of the masses of combined elements are WHOLE numbers which become subscripts for chemical formulas o Nitrogen and Oxygen combine to form NO or NO2, but not NO1.5 17 Law of Definite Composition • Law of Definite Composition (Proust’s Law)elements combine in a definite (constant) ratio by atomic mass – Water (H2O) is always 2 hydrogens for each oxygen – 16:2 oxygen:hydrogen mass ratio or 8:1 reduced – For Carbon dioxide (CO2) there is always a 32:12 oxygen to carbon mass ratio, or 8:3 reduced. 18 J.J. Thomson (1856 – 1940) Put electricity through a vacuum tube and produced a beam that was negatively charged Cathode Ray Tube Experiment 19 J.J. Thomson (1856 – 1940) • Credited with discovery of electron; a blow to Dalton’s indivisible atom idea – “Plum Pudding Model” – Also because atoms are neutral, the negative electrons must be embedded in a ball of positive charge 20 Millikan Oil Drop Experiment • Calculated the mass and quantified the charge of electrons! • Mass of electron = 9.1 X 10-28 grams (0.000000000000000000 00000000091 g) • Charge on the electron = -1.6 X 10-19 coulombs (unit of charge) - Millikan Animation and Interactive 21 Millikan Oil Drop Experiment-Reference 22 JJ Thomson and E. Goldstein • Realized if neutral atoms contain negative electrons, they must contain positive particles • Used vacuum tube similar to discovery of electron to discover protons – Mass = 1.67 X 10-24 grams (much heavier than electron!) – Charge = + 1.6 X 10-19 coulombs (same value as electron but positive) Ernest Rutherford (1871 – 1937) • Gold-foil experiment Positively charged alpha particles aimed at thin gold foil, but most passed through A few were deflected and some even bounced right back Gold foil experiment 24 Rutherford’s Gold Foil Experiment-Reference (1871 – 1937) 25 Ernest Rutherford (1871 – 1937) Conclusions: Disproved Thomson because showed most of atom is empty space Discovered dense, positively charged core, or nucleus, repels the + alpha particles Protons are surrounded by negatively charged electrons “Planetary Model” You’ll never see life the same way again 26 James Chadwick (1891 – 1974) Discovered the atomic mass of most elements was double the number of protons discovery of the neutron in 1932 Worked on the Manhattan Project Worked with Ernest Rutherford Like many others before him, he won a Nobel Prize 27 Review Early Atomic Theory video 5 min • Number your paper from 1-5 and answer the following questions. Two will be cumulative review! – 1. Which of these has 3 significant figures? • • • • a. 3340 b. 3.340 c. 0.001334 d. 334.00 28 Review • A • 2. Which of these is a homogenous mixture? • • • • a. salt b. iced tea c. pizza d. your computer 29 Review • B • 3. Which of these is true about subatomic particles? • a. electrons are negatively charged and in the nucleus • b. protons are negatively charged and in the nucleus • c. protons are positively charged and fly around the outside of the nucleus • d. neutrons are neutral and are in the nucleus 30 Review • D • 4. Which of these is true about the discovery of Millikan’s oil drop experiment – a. He discovered the electron – b. He discovered the mass of the neutron – c. He discovered the mass and the charge of the electron – d. He discovered the proton 31 Review • C • 5. Which of these is false? – – – – a. Neutrons are neutral b. Protons are positive and two will repel c. Electrons are negative and two will attract d. Protons are positive and they will attract negative electrons 32 Review • C 33 Mass of Sub-Atomic ParticlesReference (protons, neutrons, electrons) Neutron = 1.6749286 x10-24 g Proton = 1.6726231 x10-24 g Electron = 9.1093897 x10-28 g - - - - - - - - - - - - - - - - - - - - - - - - - - - - 1839 electrons = 1 neutron + 1836 electrons = 1 proton + How do you think the mass of a neutron compares to that of a proton? 1 neutron ≈ 1 proton ≈ 1.67 x10-24 g 34 Mass Number • Mass number – number of particles of significant mass in the atom • Represented by “A” protons + neutrons = mass number • Electrons are NOT included, their mass is zero • NOT found directly on the periodic table! Particle Charge Mass number Proton 1 Neutron 1 Electron 0 Location in atom 35 Let’s Do It!!! What would be the mass number of this atom? + - 3 4 + ++ 3 protons + 4 neutrons = a mass number of 7 Why did we not account for the electrons when calculating the mass number? - 36 Calculating the Actual Mass of 1 Atom • Actual mass of an atom is determined by the protons and neutrons (electrons have virtually no mass) • Each proton and neutron mass= ~1.67 x 10-24 g • Ex: a hydrogen atom has 1 electron and 1 proton: – – – – Proton = 1.67 x10-24 g No neutrons 0g Electron = + 0g Mass of the entire hydrogen atom = 1.67 x10-24 g 37 Calculating the Actual Mass of 1 Atom-Reference • The actual mass of an atom= (#protons + neutrons)(mass of p’s and n’s)= (#protons + neutrons) (1.67 x 10-24 g) = don’t count sd’s • Ex: What’s the actual mass of a Lithium atom with 3 protons and 4 neutrons? (#protons + neutrons)(mass of p’s and n’s)= (7)(1.67 x 10-24g) = 1.17 x 10-23 g 38 Relate Actual Mass to Mass Number • We can say the actual mass of an atom= (protons + neutrons)(mass of p’s and n’s)= OR (mass number) (1.67 x 10-24 g) = don’t count sd’s 39 Let’s Do It!!! • What’s the actual mass of a Carbon atom with 6 protons and 8 neutrons? 40 Let’s Do It!!! • (14)(1.67 x 10-24g) = 2.34 x 10-23 g 41 Calculating the Actual Mass of 1 Atom • Actual mass of an atom based on the idea that the whole atom is equal to the sum of the parts • Not exactly correct because binding energy is needed to hold the parts of an atom together • Some mass converted to this binding energy in a nuclear reaction so the calculation gives a value that is a little larger than reality • (E = mc2) Isotopes • What mass numbers do these atoms have? • What elements are these? • How do you know? 43 Isotopes • There mass numbers are 1, 2, and 3 but they all have one proton and therefore are all hydrogen! • So what is going on? They are isotopes. 44 Isotopes • Isotopes – different versions of atoms of an element that have same # of protons but different # of neutrons • This discovery disproved one of Dalton’s idea that atoms of the same element are exactly alike! 45 Isotopes – Hydrogen-1 (Protium): 1 proton, no neutrons and is most common – Hydrogen-2 (Deuterium): 1 proton and 1 neutron – Hydrogen-3 (Tritium): 1 proton and 2 neutrons – All versions of hydrogen! 46 Hyphen Notation-Reference • We use mass numbers to distinguish between isotopes because they differ in their number of neutrons • Hydrogen-1 =1 proton + 0 neutron=mass # 1 • Hydrogen-2 = 1 proton + 1 neutron =mass # 2 • Hydrogen-3 = 1 proton + 2 neutrons=mass # 3 • This is written in hyphen notation 47 Nuclear Symbols-Reference • Nuclear symbols are a way to write atoms using the mass number and atomic number • Format: • Hydrogen-1 1 Hydrogen-2 2 H 1 Hydrogen-3 3 H 1 H 1 48 Let’s Do It!!! • Naturally occurring carbon consists of three isotopes, Carbon-12, Carbon-13, and Carbon14. State the number of protons, neutrons, and electrons in each of these carbon atoms • Think-Pair-Share 12C 13C 14C 6 6 6 #p _______ _______ _______ #n _______ _______ _______ #e _______ _______ _______ 49 Review Atomic Number video 9 min • Number your paper from 1-5 and answer the following questions. Two will be cumulative review! – 1. Which of these is proper scientific notation? • a. 4.56 X 106 • b. 4.5 X 32 • c. 45.6 X 1010 • d. 456 X 106 50 Review • A • 2. Which of these would be the correct answer with the proper number of significant digits if you multiplied 0.05 X 3.01 – – – – a. 0.1505 b. 0.150 c. 0.15 d. 0.2 51 Review • D Remember beginning zeros are never significant, so 0.05 has only 1! • 3. What does this mean? 14C 6 – a. this atom has 6 neutrons and 20 electrons – b. this atom has 6 protons and 8 neutrons for a combined mass number of 14 – c. this atom has 6 protons and 14 electons – d. this atoms has 6 electrons and 14 neutrons 52 Review • B • 4. Which of these is a correct definition of an isotope? – a. different versions of an element that have a different number of neutrons – b. atoms of the same element with the same atomic number but different mass number – c. different version of an element that have the different number of electrons – d. A and B 53 – e. B and C Review • D • 5. How do you calculate the actual mass of an atom – a. add up all the protons, electrons, and neutrons – b. add up all the protons and neutrons and divide by two – c. add up all the protons and neutrons and multiply by 1.67 x 10-24 g – d. add up all the protons and neutrons 54 Review • C 55 Actual Mass of 1 AtomReference • The actual mass of an atom is a super small number and is cumbersome in calculations, so scientists assigned a relative scale to the mass of these particles and created a new unit called “atomic mass unit”, or amu () • amu () = atomic mass unit 56 Finding Atomic Mass of a Single Atom – To convert from actual mass to this new amu, scientists set Carbon as the standard and the value of the amu unit is defined by the actual mass of Carbon-12: Actual mass of C-12= (12)(1.67x10-24g) = 2.00x10-23g and we use this as a standard to create a conversion factor: 2.00x10-23g = 12 (amu) 57 Reference-Atomic Mass of a Single Atom • Find the atomic mass of an Oxygen-18 atom • Step 1: calculate the actual mass of an Oxygen-18 atom 18 x 1.67x10-24 g = 3.01 x10-23 g • Step 2: use dimensional analysis to convert to amu with our Carbon standard conversion factor • 3.01 x10-23 g x 12 = 18.1 amu 2.00 x10-23 g • **Don’t use 12 when figuring SDs because it58 is a standard, not a measurement Let’s Do It!!! • What’s the atomic mass of a 7 Li atom? 59 Let’s Do It!!! • What’s the atomic mass of a 7 Li atom? • Step 1: calculate the actual mass of this lithium atom 7 x 1.67x10-24 g = 1.17 x10-23 g • Step 2: use dimensional analysis to convert to amu with our Carbon standard conversion • 1.17 x10-23 g x 12 = 7.02 amu 2.00 x10-23 g 60 Isotopes • What elements are these? • What are their mass numbers? ++ + + + 61 Isotopes • How would we write the hyphen notation of these isotopes? Nuclear notation? • Students on board ++ + + + 62 Isotopes • If we pick up one of the trillions of boron atoms in the world, it could be either of these 2 types because they are both present ++ + + + • Boron-10 Boron-11 63 Isotopes • As it turns out, any mass of Boron, and all Boron in the world, is ~20% Boron-10, and ~80% Boron-11—their relative abundances ++ + + + • Boron-10 20% Boron-11 80% 64 Isotopes • Many elements are like this • All chlorine in the world is 75% Chlorine35 and 25% Chlorine-37 • The majority are Chlorine-35 65 Isotopes-Reference • In nature there are always mixtures of isotopes and this can pose difficulties when we do calculations. Why? • Pencil lead (carbon) has some carbon-12, carbon-13 and carbon-14 mixed • The mass of one Carbon-12 atom with 12 protons and neutrons would be this: (12)(mass of protons and neutrons) = (12)(1.67 x10-24 g) = 12 amus () but that is just for Carbon-12!! 66 Isotopes and Atomic Mass of an Element • STOP! “Mass” confusion and Reference Chart • In calculations we need a mass value that represents the whole element mixture (carbon-12, carbon-13 and carbon-14 mixed), not just one isotope, like Carbon-12 How? 67 Atomic Mass of an ElementReference • Atomic mass- WEIGHTED average of the atomic masses of all the element’s isotopes as they are found in nature (don’t confuse it with mass number which is just p + n) (abundance isotope #1) (atomic mass isotope #1 ) + (abundance isotope #2) (atomic mass isotope #2 ) + continue for all isotopes 68 Isotopes and Atomic Mass of an Element • The atomic masses of each element (the weighted average of all its isotopes) is found in the periodic table ++ + + + • Boron-10 20% Boron-11 80% 69 Atomic Mass of an Element – Reference Use the following data to calculate the atomic mass for the element Magnesium Isotope Atomic Mass of Isotope Abundance Mg - 24 23.982628 78.600 % Mg - 25 24.963745 10.11 % Mg - 26 25.960802 11.29 % (.78600) (23.982628 ) + (.1011) (24.963745 ) + (.1129) (25.960802 ) = 18.850 + 2.52 + 2.931 = 24.305 You do SD’s for every individual calculation 70 Let’s Do It!!!! • The element copper has naturally occurring isotopes with mass numbers of 63 and 65 • The relative abundance and atomic masses are 69.2% for a mass of 62.93amu and 30.8% for a mass of 64.93amu. Calculate the atomic mass of the element copper 71 Let’s Do It!!!! • Divide the percentages by 100 to convert to decimals (.692) (62.93amu) + (.308)(64.93amu) = 43.5 amu + 20.0 amu = 63.5 amu This shows that the majority of the isotopes found in nature are 62.93 amu (closer to 63 than 65) 72 Let’s Do It!!! • There are two isotopes of silicon. The atomic mass of the element silicon found on the periodic table is 28.086amu. • Which of these is not possible to be one of the atomic masses of the individual isotopes? – – – – A. 26.065 B. 29.543 C. 28.086 D. 27.439 73 Atomic Mass of an Element • If an element has only 1 isotope, then the atomic mass of that isotope IS the atomic mass of the element 74 Review • Number your paper from 1-5 and answer the following questions. Two will be cumulative review! • 1. Which of these is the symbol for the metric unit micro (and also for amu)? – A. M – B. m – C. ρ – D. μ 75 Review • D • 2. Which of these is the amount of protons in an atom and therefore, the ID. – a. atomic mass – b. atomic number – c. mass number – d. isotope number 76 Review • B • 3. What is the bottom number on this periodic table tile? – a. atomic mass of the element Boron – b. atomic number – c. mass number – d. atomic mass of a Boron atom 77 Review • A • 4. The atomic mass of the element sulfur is 32.1 g. There are four common isotopes of sulfur, S-32 (32.2344 μ), S-33 (33.45676 μ), S-34 (34.15643 μ), S-36 (36.44321 μ) Which of these isotopes is the most abundant? – – – – a. S-32 b. S-33 c. S-34 d. S-35 78 Review • A • 5. How do you calculate the atomic mass of an element? – a. l x w x h – b. D = m/V – c. (abundance isotope #1)(atomic mass #1) + (abundance isotope #2) (atomic mass #2) etc. – d. mass number – atomic number 79 Review • C 80 The Mole • The word “mole” in Chemistry is a term used to describe a certain amount of something • For example: – 1 dozen = 12 – 1 baker’s dozen = 13 – 1 gross = 144 – 1 mole = 6.02 X 10 23 which is 602,000,000,000,000,000,000,000 or 602 hextrillion ………of whatever 81 The Mole-Reference • A mole is the amount of a substance that contains 6.02 x 1023 of something – 1 mole of pencils = 6.02 x 1023 pencils – 1 mole of eggs = 6.02 x 1023 eggs – 1 mole of carbon = 6.02 x 1023 carbon atoms • 6.02 x 1023 is called Avogadro’s number • It helps us count atoms, which we can’t see, using the measurement of mass 82 Let’s Do It!!! • How many atoms of oxygen are in a mole of oxygen? • How many atoms of magnesium are in a mole of magnesium? • Would they have the same mass? 83 Atomic Mass of an Element = Mole-Reference • Each element has a unique atomic mass and this is a standard for each element • Atomic mass = mass of 1 mole of an element • The atomic masses are from the periodic table and we use grams 1 mole O = 6.02 x 1023 atoms O = 15.999 g O 1 mole Mg= 6.02 x 1023 atoms Mg= 24.305 g Mg 84 Molar Mass • Molar mass –mass of 1 mole of a pure substance • Molar mass of element (g/mol) – the atomic mass of an element • Molar mass of a compound (g/mol)- the sum of the molar masses of the elements making up the compound • What’s the molar mass of Cl? Don’t forget units! We are going to round to 3 SDs for the atomic mass 85 • What’s the molar mass of H2O? Moles to Atoms-Reference • 1 mole = 6.02x1023 atoms = atomic mass (g) • These can be used as a conversion factors 6.02 x 1023 atoms 1 mole atomic mass atomic mass 6.02 x 1023 atoms 1 mole • If you have 1.5 moles of potassium, how many atoms is this? • Question mark format for DA ? atoms of K= 6.02 x 1023 atoms K = 9.03 x 1023 1.50 mole K * --------------------atoms K 86 1 mole K Moles to Grams-Reference • 1 mole = 6.02x1023 atoms = Atomic Mass (g) • If you have 1.25 moles of potassium, how many grams is this? • Question mark format for DA ? grams of K = 39.1 g K 1.25 moles K * ----------- = 48.9 g K 1 mole K 87 Grams to Atoms-Reference • 1 mole = 6.02x1023 atoms = Atomic Mass (g) • If you have 12.5 grams of phosphorus, how many atoms is this? • Question mark format for DA ? atoms of P= 23 atoms P 6.02 x 10 12.5 g P = x --------------------31.0 g P 2.43 x 1023 atoms P 88 Let’s Do It!! • 1 mole = 6.02x1023 atoms = Atomic Mass (g) • If you have 2.93 x 1024 atoms of sodium, how many moles is this? • Question mark format for DA ? moles of Na= 2.93 x 1024 atoms Na = 4.87 moles Na 89 Let’s Do It!!! • 1 mole = 6.02x1023 atoms = Atomic Mass (g) • If you have 1.25 grams of sodium, how many moles is this? • Question mark format for DA ? moles of Na = 1 mole Na 1.25 g Na = 0.0543 mole Na * -------------23.0 g Na 90 Let’s Do It!!! • 1 mole = 6.02x1023 atoms = Atomic Mass (g) • If you have 3.52 x 1022 atoms of carbon, how many grams is this? ? grams of C = 1022 3.52 x atoms C 12.0 g C = .702 g C x ---------------6.02 x 1023 atoms C 91 Mole vs. Molecule • Don’t confuse mole and molecule – A mole is an amount – A molecule is a thing – We can have a mole of molecules • A mole of water molecules = 6.02 x 1023 molecules of water • If I have an amino acid molecule tryptophan, C11H12N2O2, how many molecules of tryptophan are in a mole? 92 Review • Number your paper from 1-5 and answer the following questions. Two will be cumulative review! – 1. Which of these is proper scientific notation? • • • • a. 5.56 X 106 b. 3.5 X 32 c. 95.6 X 1010 d. 256 X 106 93 Review • A • 2. Which of the numbers in the nitrogen periodic table tile is the atomic number? • • • • a. N b. 14 c. 14.007 d. 7 94 Review • D • 3. Which of these is NOT a correct conversion factor for Carbon? • a. 1 mole C 6.02 x 1023 g C • b. 1 mole C 6.02 x 1023 atoms C • c. 1 mole C 12.011 g C • d. 12.011 g C 6.02 x 1023 atoms C 95 Review • A • 4. If I have 14.007 grams of nitrogen, how many atoms do I have? • • • • a. 1.67 x10-24 b. 1.17 x10-23 c. 2.24 x10-23 d. 6.02x1023 96 Review • D • 5. How many atoms are in 10 moles of carbon? • • • • a. 10 atoms b. 6 atoms c. 6.02x1023 atoms d. 6.02x1024 atoms 97 Review • D • 1 mole of ANYTHING = 6.02x1023 atoms • ? Atoms = 10 moles C x 6.02x1023 atoms C = 1 mole C 10 x 6.02x1023 atoms C= 6.02x1024 atoms 98 Radius of Nucleus of an Atom-Reference • Radius = 3 A (1.4 x 10-13 cm) • Where A = mass number of the atom 3 A = # of protons + neutrons and = the cube root of the mass number and 1.4 x 10-13 cm = a constant that describes the effective range of force 99 Radius of Nucleus of an Atom-Reference • Calculate the radius of the nucleus of a 56Fe atom 3 • Radius = A (1.4 x 10-13 cm) • Radius = 3 56 (1.4 X 10-13 cm) • Radius = 5.4 X 10-13 cm only round ONCE ! ! ! 100 Volume of a NucleusReference • Calculate the volume of the nucleus of a 56Fe atom (we model as a sphere) • Volume of a sphere = 4 r 3 3 • Volume = 4 (5.4 *10 13 cm) 3 3 • Volume = 6.6 * 10-37 cm3 101 Size of the Nucleus Atom Nuclear Radius Nuclear Volume 4 A (1.4 x 10-13 cm) r 3 Nuclear Mass 3 3 23 11 D m V Na 197 79 4 2 A(1.67*10-24g ) Nuclear Density Au He 64 29 Cu 80 35 Br 102 Size of the Nucleus Atom Nuclear Radius Nuclear Volume 4 A (1.4 x 10-13 cm) r 3 Nuclear Mass 3 3 23 11 Na 197 79 4 2 Au He m V 8.1 x 10-13 cm 2.2 x 10-13 cm 5.6 x 10-13 cm 80 35 6.0 x 10-13 cm Br D 4.0 x 10-13 cm 64 29 Cu A(1.67*10-24g ) Nuclear Density 103 Size of the Nucleus Atom Nuclear Radius Nuclear Volume 4 A (1.4 x 10-13 cm) r 3 Nuclear Mass 3 4.0 x 10-13 cm 3 2.7 x 10-37 cm3 8.1 x 10-13 cm 2.2 x 10-36 cm3 2.2 x 10-13 cm 4.5 x 10-38 cm3 64 29 5.6 x 10-13 cm 7.4 x 10-37 cm3 80 35 6.0 x 10-13 cm 9.0 x 10-37 cm3 23 11 Na 197 79 4 2 Au He Cu Br A(1.67*10-24g ) Nuclear Density D m V 104 Size of the Nucleus Atom Nuclear Radius Nuclear Volume 4 A (1.4 x 10-13 cm) r 3 Nuclear Mass 3 3 A(1.67*10-24g ) 4.0 x 10-13 cm 2.7 x 10-37 cm3 3.84 x 10-23g 8.1 x 10-13 cm 2.2 x 10-36 cm3 3.29 x 10-22g 2.2 x 10-13 cm 4.5 x 10-38 cm3 6.68 x 10-24g 64 29 5.6 x 10-13 cm 7.4 x 10-37 cm3 1.07 x 10-22g 80 35 6.0 x 10-13 cm 9.0 x 10-37 cm3 1.34 x 10-22g 23 11 Na 197 79 4 2 Au He Cu Br Nuclear Density D m V 105 Size of the Nucleus-Reference Atom Nuclear Radius Nuclear Volume 4 A (1.4 x 10-13 cm) r 3 Nuclear Mass 3 3 A(1.67*10-24g ) Nuclear Density D m V 4.0 x 10-13 cm 2.7 x 10-37 cm3 3.84 x 10-23g 1.4 x 1014 g/cm3 8.1 x 10-13 cm 2.2 x 10-36 cm3 3.29 x 10-22g 1.5 x 1014 g/cm3 2.2 x 10-13 cm 4.5 x 10-38 cm3 6.68 x 10-24g 1.5 x 1014 g/cm3 64 29 5.6 x 10-13 cm 7.4 x 10-37 cm3 1.07 x 10-22g 1.5 x 1014 g/cm3 80 35 6.0 x 10-13 cm 9.0 x 10-37 cm3 1.34 x 10-22g 1.5 x 1014 g/cm3106 23 11 Na 197 79 4 2 Au He Cu Br Densities of Atoms • Why are these density values nearly the same? • Because all nuclei are made up of the same material (protons & neutrons). The SAME material always has the SAME density! 107 Densities of Atoms • If we look up the density of the element sodium (it’s on some periodic tables) we see that its density is .971 g/cm3. • Why is this different from the value we calculated? • We calculated values for the nucleus only and the periodic table’s value is for the whole atom – including the space that the electrons occupy 108