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Chapter 4 Atomic Structure 1 Section 4.1 – Defining the Atom Objectives Describe Democritus’s ideas about atoms Explain Dalton’s atomic theory Identify what instrument is used to observe individual atoms 2 Trust? I am going to ask you to believe in something that you can not see with your unaided eye… We can not see the tiny fundamental particles that make up matter – yet all matter is composed of such particles called atoms 3 Atom Atom The smallest particle of an element that retains its identity in a chemical reaction Think of a coin the size of a penny and composed of pure copper (Cu) Imagine grinding the coin into fine dust … each speck still has the properties of Cu If you could continue to grind the dust into smaller and smaller Cu dust particles you would eventually come upon a particle of Cu that could no longer be divided and still have the chemical properties of Cu … this final particle is an atom 4 Early Models of the Atom Greek Philosopher Democritus (460 BC – 370 BC) Among the first to suggest the existence of atoms He believed that atoms were indivisible and indestructible Although his ideas agree with later scientific theory, they did not explain chemical behavior and lacked experimental support – not based on the scientific method – just philosophy 5 Early Models of the Atom Who was next? John Dalton (1766 – 1844) Experiment based – transformed Democritus’ ideas into scientific theory Results of his experimentation and observations are summarized in Dalton’s Atomic Theory Dalton studied the ratios in which elements combine in chemical reactions 6 Early Models of the Atom Dalton’s Atomic Theory 1. 2. 3. 4. All elements are composed of tiny indivisible particles called atoms Atoms of the same element are identical. The atoms of any other elements are different from those of any other element Atoms of different elements can physically mix together or can chemically combine in simple whole-number ratios to form compounds In chemical reactions, atoms are combined, separated or rearranged – but never changed into atoms of another element 7 Sizing up the atom Remember our copper (Cu) penny…We ground it into smaller and smaller particles until we came upon a Cu atom How big (or small) would that Cu atom be? Here’s a little perspective… a pure Cu coin the size of a penny contains 2.4 x 1022 atoms! 8 Sizing up the atom Earths population is only 6 x 109 people! If you could line up 1 x 108 Cu atoms side by side, they would produce a line only 1 cm long! The radii of most atoms is 5 x 10-11 to 2 x 10-10 m 9 Sizing up the atom Despite their small size individual atoms are observable with instruments such as scanning tunneling microscope The image below shows an image of iron atoms generated by a scanning tunneling microscope 10 Section 4.2 – Structure of the Nuclear Atom Objectives Identify three types of subatomic particles Describe the structure of atoms according to the Rutherford atomic model 11 Subatomic Particles One important change to Dalton’s Atomic Theory… Atoms are known to be divisible Atoms can be broken down into even smaller, more fundamental particles called subatomic particles Three kinds of subatomic particles are electrons, protons and neutrons How was this discovered? 12 Subatomic Particles Electrons Discovered in 1897 by J.J. Thomson Electrons are negatively charged subatomic particles Thomson performed experiments using a cathode ray tube to deduce the presence of negatively charged particles Cathode ray tubes pass electricity through a gas that is contained at a very low pressure 13 Thomson’s Experiment Voltage source - + Metal Disks 14 Thomson’s Experiment Voltage source + Passing an electric current makes a beam appear to move from the negative to the positive end 15 Thomson’s Experiment Voltage source + By adding an electric field he found that the moving pieces were negative 16 Thomson’s Experiment During experiment he used many different metals By the amount the beam bent he could find the ratio of charge to mass This ratio was the same for every material Thomson concluded it must be the same type of piece in every kind of atom…the electron 17 Mass of the Electron In 1961 Robert Millikan determined the mass of the electron to be 1/1840 the mass of a hydrogen atom The mass of an electron is 9.11 x 10-28 g 18 Conclusions from the study of the Electron 1. Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons. 2. Atoms are neutral, so there must be positive particles in the atom to balance the negative charge of the electrons 3. Electrons have so little mass that atoms must contain other particles that account for most of the mass 19 Conclusions from the study of the Electron 4. Eugen Goldstein in 1886 observed what is now called the “proton” - particles with a positive charge, and a relative mass of 1 (or 1840 times that of an electron) 5. 1932 – James Chadwick confirmed the existence of the “neutron” – a particle with no charge, but a mass nearly equal to a proton 20 Subatomic particles Name Relative Symbol Charge mass Actual mass (g) Electron e- -1 1/1840 9.11 x 10-28 Proton p+ +1 1 1.67 x 10-24 Neutron n0 0 1 1.67 x 10-24 21 The Atomic Nucleus Once subatomic particles were discovered scientists wondered how they were put together in as atom We will look at two models: Thomson’s Atomic Model (Plum Pudding Model) Rutherford’s Atomic Model 22 The Atomic Nucleus He believed the electrons were struck into a lump of positive charge Thomson’s model turned out to be short-lived due to groundbreaking work of Ernest Rutherford a former student of Thomson 23 The Atomic Nucleus Ernest Rutherford’s Gold Foil Experiment Ernest Rutherford decided to test the plum pudding model Wanted to determine how large the atoms were 24 The Atomic Nucleus Used Alpha particles Helium atoms that have lost their two electrons leaving a double positive charge Helium atom Alpha particle 25 The Atomic Nucleus Ernest Rutherford’s Gold Foil Experiment The alpha particles were fired at a thin sheet of gold surrounded by a detecting screen Particles that are hit on the detecting screen are recorded 26 The Atomic Nucleus What Rutherford expected… Because evenly distributed + charges of the plum pudding model were not enough to stop the ++ of the alpha particles to pass through without 27 What he found … 28 The Atomic Nucleus What Rutherford’s results meant: The atom is mostly empty space. The electrons are distributed around the nucleus, and occupy most of the volume All the positive charge, and almost all the mass is concentrated in a small area in the center. This is why most alpha particles passed through without deflection (no + charge to deflect the ++ alpha particle) This is why some alpha particles were deflected. The + charge and mass of the center was enough to deflect the ++alpha particle. The center is composed of protons and neutrons. He called this a “nucleus” His model was called the nuclear atom 29 The Atomic Nucleus 30 Section 4.3 – Distinguishing Among Atoms Objectives Explain what makes elements and isotopes different from each other Calculate the number of neutrons in an atom Calculate the atomic mass of an element Explain why chemists use the periodic table 31 Atomic Number All atoms are composed of protons, neutrons and electrons How then are atoms of one element different than atoms of another element? Why is an atom of C different from an atom of H? Elements are different because they contain different numbers of protons The atomic number of an element is the number of protons in the nucleus of an atom of that element # of protons in an atom = the number of electrons Atoms are electrically neutral 32 Atomic Number Atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element Element # of Protons Atomic Number Carbon 6 6 Phosphorus 15 15 Gold 79 79 33 Mass Number The majority of mass of an atom is concentrated in the nucleus The mass of an atom depends on the number of protons and neutrons Total number of protons and neutrons in an atom is called the mass number The number of neutrons in an atom is the difference between the mass number and atomic number Number of neutrons = mass number – atomic number 34 Complete Symbols We began looking at symbols for some of the elements, however: Complete symbols contain the symbol of the element, the mass number and the atomic number Mass Superscript → Number Atomic Subscript → Number X 35 Complete Symbols Au 197 79 Find each of the following: Number of protons 79 Number of neutrons 118 Number of electrons 79 Atomic number 79 Mass number 197 Atoms can also be referred to by using the mass and the name of the element – gold – 197 36 Complete Symbols If an element has an atomic number of 34 and a mass number of 78, what is the: Number of protons? Number of neutrons? Number of electrons? Name of element? Complete symbol? 34 44 34 Selenium 78 Se 34 37 Complete Symbols If an atom has 91 protons and 140 neutrons what is the: Atomic number? Mass number? Number of electrons? Complete symbol? Name of the element? 91 91+140 = 231 91 231 Pa 91 Protactinium 38 Isotopes So, if two atoms have different number of protons, they are different elements. What if two atoms have different number of neutrons? Atoms with different number of neutrons are isotopes of the same element They have different mass numbers Despite these differences, isotopes are chemically alike because they have identical numbers of protons and electrons Protons and electrons are responsible for chemical behavior 39 Isotopes Another way to say it: Isotopes are atoms of the same element having different masses, due to varying numbers of neutrons There are three known isotopes of hydrogen Isotope Protons Electrons Neutrons Hydrogen–1 (protium) 1 1 0 Hydrogen-2 (deuterium) 1 Hydrogen-3 (tritium) 1 1 1 1 2 Nucleus 40 Isotopes • Elements occur in nature as mixtures of isotopes 41 Atomic Mass How heavy is an atom of oxygen? It depends, because there are different kinds of oxygen atoms (different isotopes of Oxygen) We are more concerned with the average atomic mass. This is based on the abundance (percentage) of each isotope of that element in nature. 42 Determining Atomic Mass We don’t use grams for this mass because the numbers would be too small Instead of grams, the unit we use is the Atomic Mass Unit (amu) It is defined as one-twelfth the mass of a carbon-12 atom Carbon-12 was chosen because of its isotope purity. Each isotope has its own atomic mass, thus we determine the average from percent abundance 43 Calculating Atomic Mass To calculate the average atomic mass of an element we need to know 3 things: What are the stable isotopes of the element The mass of each isotope of the element The natural percent abundance of each isotope of the element 44 Calculating Atomic Mass Sample Steps for An Element with 3 Naturally Occurring Isotopes massisotope1 x natural abundanceisotope1 = avg. massisotope1 massisotope2 x natural abundanceisotope2 = avg. massisotope2 massisotope3 x natural abundanceisotope3 = avg. massisotope3 avg. massisotope1 + avg. massisotope2 + avg. massisotope3 = avg mass of atom for given element The resulting sum is the weighted average mass of the atoms of the element that occur in nature 45 Atomic Mass Practice Problem The element copper has naturally occurring isotopes with mass numbers of 63 and 65. The relative abundance and atomic masses are 69.2% for mass = 62.93 amu, and 30.8% for mass = 64.93 amu. Calculate the average atomic mass of copper 46 Periodic Table – A Preview The periodic table is an arrangement of elements in which the elements are separated into groups based on a set of repeating properties The periodic table allows you to easily compare the properties of one element to another 47 Periodic Table – A Preview Each element is identified with its symbol placed in a square The elements are listed in order of increasing atomic number from left to right and top to bottom 48 Periodic Table – A Preview Each horizontal row (there are 7 of them) is called a period The properties of the elements vary as you move across it from element to element The pattern then repeats as you move to the next period Each vertical column is called a group or family Elements in a group have similar chemical and physical properties Identified with a number and either an “A” or a “B” 49