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Early Atomic Theory 5.1 Dalton’s Model of the Atom 5.2 Electric Charge 5.3 Subatomic Parts of the Atom 5.4 The Nuclear Atom 5.5 Isotopes of the Elements 5.6 Atomic Mass Early Theories on the Structure of Matter 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 Model of the Atom Elements are composed of small, indivisible particles called atoms. • Atoms of the same element are identical in mass and size. • Atoms of different elements differ in their mass and size. • Compounds are formed by combining two or more atoms of different elements. • Atoms combine to form compounds in simple whole number ratios. • Atoms of two elements may combine in different ratios, leading to formation of different compounds. Dalton’s Model of the Atom Atoms are individual particles which are different for each element. ! Atoms combine in fixed ratios to form compounds. Two elements can combine in varying ratios to give different compounds. H 2O H2O2 Most of Dalton’s theory remains valid today. Dalton’s Model of the Atom Revisions to Dalton’s Theory 1. Elements can be decomposed under certain conditions. 2. Not all atoms of the same element have identical mass. Atoms of the same element with different atomic mass are called isotopes. 3. Atoms are not indivisible. They are composed of subatomic particles. 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: F= kq1q2 r2 where q1 and q2 are the charges, r is the distance between charges, and k is a constant. Subatomic Parts of the Atom Because atoms are so small, determining the presence of subatomic particles was very difficult. A single atom is tiny (diameter of 0.1 to 0.5 nm). 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. 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. 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 10–28 g) and size (<10–12 cm). Protons (p): 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). The Effect of Subatomic Particles 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. + charged sphere Electrons Thomson’s “plum p model of the atom. The Effect of Subatomic Particles Atoms can become ions by gaining or losing electrons from this sphere. Electrons are lost from atoms to give cations. Electrons are gained from atoms to give anions. Neutrons The last subatomic particle was discovered by Chadwick in 1932. Neutrons (n): A particle with no electrical charge. Neutrons have a mass similar to that of a proton. Summary of Subatomic Particles Atoms are composed of three smaller, subatomic particles: electrons, protons and neutrons. 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. 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). Nuclear Model of the Atom Because most of the particles were not deflected, this suggested most of the atom is empty space. 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. Atomic Number Atomic Number: 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. 27 Co Atomic Number Isotopes of the Elements 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. Isotopes: atoms of an element with the same atomic number but different numbers of neutrons. Isotopes of the Elements Example: Isotopes of Hydrogen Protium Deuterium 1 proton 0 neutrons 1 proton 1 neutron Tritium 1 proton 2 neutrons Standard Isotopic Notation Mass Number Atomic Number A E Z Element Symbol Mass number: Total number of protons and neutrons for an element. Isotopes of the Elements Practice: How many protons, neutrons, and electrons are found in each of the following isotopes? 64 Cu 29 Atomic Number: 29 protons (therefore 29 electrons) # Neutrons = Mass Number – Atomic Number 64 – 29 = 35 neutrons Let’s Practice! Which isotope corresponds to an element with 24 protons and 28 neutrons? a. 28 Cr 52 Solution: b. 52 Cr 24 # protons = Atomic Number = 24 Element: Cr c. 52 Ni 28 Mass Number = protons + neutrons 128 Te d. 52 e. 24 Cr 52 ! = 24 + 28 = 52 Atomic Mass Because the mass of a single atom is so small, it is inconvenient to use this as a mass unit. Instead, relative atomic mass units (amu) are used. Using carbon-12, as a standard, 1 atomic mass unit is equal to 1/12th the mass of a carbon-12 atom. 1 amu = 1.6606 x 10-24 g All periodic tables use atomic masses based on the carbon-12 isotope. Atomic Mass and Isotope Distribution 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 63.546 amu. Cu exists as 2 major isotopes, Cu-63 and Cu-65. Cu-63 is more abundant IS more abundant than Cu-65 The average atomic mass is equal to the sum of the atomic mass of each isotope multiplied by its % abundance. Atomic Mass and Isotope Distribution Average atomic mass of Cu: (62.9298) x (0.6909) + (64.9278) x (0.3091) = 63.55 amu Atomic Mass % Abundance Atomic Mass % Abundance Atomic Mass Practice Silver exists as two isotopes with atomic masses of 106.9041 and 108.9047 amu. Determine the average atomic mass for silver if the % abundance for each isotope is 51.82 and 48.18%, respectively. Average atomic mass of Ag: (106.9041) x (.5182) + (108.9047) x (0.4818) = 107.8680 amu Atomic Mass % Abundance Atomic Mass % Abundance Let’s Practice! Chlorine exists as two isotopes, Cl-37 (36.96590 amu) and Cl-35. If the percent abundance of each isotope is 24.47 % and 75.53 %, what is the atomic mass of Cl-35 if the average atomic mass is 35.46 amu? Solve for a: (36.96590) x (.2447) + (a) x (0.7553) = 35.46 amu a. b. c. d. e. 36.95690 34.96885 36.57823 33.56438 35.64544 Solution: 9.046 + (a) x (0.7553) = 35.46 amu (a) x (0.7553) = 26.41 amu a = 34.97 amu