5 Early Atomic Theory and Structure

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1 5 Early Atomic Theory and Structure Chapter Outline Electric Charge A. Discovery of Ions 5.3 Subatomic Parts of the Atom Lightning occurs when electrons move to neutralize charge difference between the storm clouds and Earth. Foundations of College Chemistry, 14 th Ed. Morris Hein and Susan Arena 5.4 The Nuclear Atom Early Theories on the Structure of MaHer Early models of the atom were developed by the Greeks. Empedocles proposed matter was composed of four basic elements: earth, air, water and fire. Democritus proposed matter was composed of small, indivisible particles he called atoms. Atoms could combine in different ways, giving rise to the diversity of compounds we observe. Aristotle, an influential philosopher, supported Empedocles theory, so atomic theory was not fully accepted until 2000 years later. Dalton s theory of atoms, proposed in the early 1800s, states: 1. Elements are composed of small, indivisible particles called atoms. 2. Atoms of the same element are identical in mass and size. 3. Atoms of different elements differ in their mass and size. 4. Compounds are formed by combining two or more atoms of different elements. 5. Atoms combine to form compounds in simple whole number ratios. 6. Atoms of two elements may combine in different ratios, leading to formation of different compounds. Revisions to Dalton s Theory H 2 O H 2 O 2 1. Elements can be decomposed under certain conditions. a. Atoms are individual particles which are different for each element. b./c. Atoms combine in fixed ratios to form compounds. Two elements can combine in varying ratios to give different compounds. Most of Dalton s theory remains valid today. 2. Not all atoms of the same element have identical mass. These are called isotopes. 3. Atoms are not indivisible. Atoms are composed of subatomic particles. 1

2 Electric Charge Properties of Electric Charge 1. Charge may be either positive or negative. 2. Opposite charges (positive and negative) attract while like charges (i.e. negative and negative) repel. 3. Charge may be transferred from one object to another, by contact or induction. 4. The force of attraction between charges (F) is related to the distance between charges by: kq 1 q 2 F = r 2 where q 1 and q 2 are the charges, r is the distance between charges, and k is a constant. Discovery of Ions Michael Faraday: English scientist who discovered electrolytes (compounds that conduct electricity when dissolved in water). Faraday also discovered that some compounds decompose in water into their elements. These elements were attracted to either negatively or positively charged electrodes in the solution, meaning they were no longer neutral. These charged elements are called ions. A light bulb glows when ions are present in a saltwater solution when current is passed through it. The Nature of Ions Arrhenius extended Faraday s work. He proposed ions are atoms (or groups of atoms) that carry a positive or negative charge. Ex. NaCl in water dissociates into two ions, Na + and Cl. The Na + (cation) produced is attracted to the negatively charged electrode (cathode). The Cl (anion) produced is attracted to the positively charged electrode (anode). Based on Faraday s and Arrhenius work, Stoney proposed the electron was a fundamental unit of electricity associated with atoms. J. J. Thomson later experimentally confirmed the existence of electrons. Subatomic Parts of the Atom A single atom is tiny (diameter of 0.1 to 0.5 nm). Because atoms are so small, determining the presence of subatomic particles was very difficult. New instruments in the early 1900s permitted detection of these particles. A scanning tunneling microscope (STM) image shows an array of Cu atoms. Subatomic Parts of the Atom A Crooks tube permits generation of cathode rays, which are streams of electrons. Electrons and Protons Electrons (e ): A particle with negative electrical charge (assigned a relative charge of 1). Electrons have a very small mass (9.110 x g) and size (<10 12 cm). Protons (p): A Crooks (cathode) ray tube. The stream of electrons passes between the electrodes. The electron beam is deflected by both electric and magnetic fields, indicating it has charge. A particle with positive electrical charge (assigned a relative charge of +1). Protons have a much larger mass (~1837 times the mass of an electron). 2

3 The Effect of Subatomic ParOcles Thomson s work demonstrated the atom is composed of smaller, charged particles. Dalton s theory of the atom then had to be revised. Thomson s Model of the Atom Electrons are negatively charged particles which are embedded in a positively charged atomic sphere. The Effect of Subatomic ParOcles Atoms can become ions by gaining or losing electrons from this sphere. Electrons are lost from atoms to give cations. + charged sphere Electrons Thomson s plum pudding model of the atom. Electrons are gained from atoms to give anions. Neutrons Summary of Subatomic ParOcles The last subatomic particle was discovered by Chadwick in Atoms are composed of three smaller, subatomic particles: electrons, protons and neutrons. Neutrons (n) A particle with no electrical charge. Neutrons have a mass similar to that of a proton. Chemical properties of atoms can be described based on the electrons, protons and neutrons. Though other subatomic particles are now known, the theories of atomic structure are based only on these 3 subatomic particles. Nuclear Model of the Atom In 1911, Ernest Rutherford established the nuclear model of the atom by bombarding gold atoms with α particles. Nuclear Model of the Atom Because most of the particles were not deflected, this suggested most of the atom is empty space. Most of the particles passed through the gold foil, but some were deflected and some even bounced back! This suggested the gold atoms must have a densely, positively charged nucleus to affect the path of an α particle (a positively charged He atom). Protons and neutrons are located in the nucleus. Electrons are dispersed throughout the remainder of the atom (mainly open space). Neutral atoms contain the same number of protons and neutrons to maintain charge balance. 3

4 : Number of protons in the nucleus of an atom. The atomic number determines the identity of the atom. Atomic numbers for every element are above the element s symbol in the periodic table. After discovery of the nuclear model of the atom, the mass of almost all atoms was found to be larger than expected, based on the number of protons and electrons. This led to the discovery of neutrons. Though all atoms of the same element have the same number of protons, atoms of the same element may have different numbers of neutrons. 27 Co Isotopes: atoms of an element with the same atomic number but different numbers of neutrons. Example: Isotopes of Hydrogen Protium Deuterium Tritium 0 neutrons 1 neutron 2 neutrons Standard Isotopic Notation Mass Number A E Z Element Symbol Mass number: Total number of protons and neutrons for an element. Practice: How many protons, neutrons, and electrons are found in each of the following isotopes? 64 Cu 29 : 29 protons (therefore 29 electrons) # Neutrons = Mass Number = 35 neutrons Let s PracOce! Which isotope corresponds to an element with 24 protons and 28 neutrons? Because the mass of a single atom is so small, it is inconvenient to use this as a mass unit. a. b. c. 28 Cr Cr Ni d. 52 Te e Cr Solution: # protons = = 24 Element: Cr Mass Number = protons + neutrons = = 52 Instead, relative atomic mass units (amu) are used. 12 Using carbon-12, 6 C, as a standard, 1 atomic mass unit is equal to 1/12th the mass of a carbon-12 atom. 1 amu = x g All periodic tables use atomic masses based on the carbon-12 isotope. 4

5 and Isotope DistribuOon and Isotope DistribuOon Since most elements are a mixture of isotopes, the atomic mass for an element is the weighted average of all naturally occurring isotopes of the element. Example: The atomic mass of Cu is amu. Cu exists as 2 major isotopes, Cu-63 and Cu-65. Cu-63 is more abundant, as the atomic mass is very close to 63 amu. Calculating average atomic mass: Sum of the atomic mass of each isotope multiplied by its % abundance. Average atomic mass of Cu: ( ) x (0.6909) + ( ) x (0.3091) = amu % Abundance % Abundance Measuring Cu isotope abundances by using mass spectrometry. Silver exists as two isotopes with atomic masses of and amu. Determine the average atomic mass for silver if the % abundance for each isotope is and 48.18%, respectively. Average atomic mass of Ag: ( ) x (.5182) + ( ) x (0.4818) = amu PracOce % Abundance % Abundance Let s PracOce! Chlorine exists as two isotopes, Cl-37 ( amu) and Cl-35. If the percent abundance of each isotope is % and %, what is the atomic mass of Cl-35 if the average atomic mass is amu? a d b e c Solution: Solve for a: ( ) x (.2447) + (a) x (0.7553) = amu (a) x (0.7553) = amu (a) x (0.7553) = amu a = amu Learning ObjecOves Learning ObjecOves 5.1 Describe Dalton s model of the atom and compare it to the earlier concepts of matter. 5.2 Electric Charge Use Coulomb s Law to calculate the force between particles and distinguish between a cation and anion. 5.3 Subatomic Parts of the Atom Describe the three basic subatomic particles and how they changed Dalton s model of the atom. 5.4 The Nuclear Atom Explain how the nuclear model of the atom differs from the Dalton and Thomson models. 5.5 Define the terms atomic number, mass number and isotope. 5.6 Define the relationship between the atomic mass of an element and the masses of its isotopes. 5