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The Structure of the Atom Section 4.1 Early Theories of Matter Early Greeks believed in 4 fundamental substances: Early Greeks believed matter could continuously be divided into smaller and smaller particles Philosophers And Early Scientists fire, earth, water, and air. Democritus (460-370 B.C.) First person to propose the idea that matter was not infinitely divisible. Matter is composed of empty space that small particles move through Matter made up of tiny individual particles – atomos (Greek for atom) o Atoms are solid; cannot be created, destroyed, or further divided o Various sizes and shapes of atoms o Different properties of matter are due to size, shape, and movement of atoms – different atom for each substance Changes in matter result from changes in the groupings of atoms, not changes in the atoms themselves General ideas, not science – not supported experimentally 1 Aristotle (384-322 B.C.) Rejected Democritus’ atomic theory (didn’t agree with his own ideas on nature) Didn’t believe in the ―nothingness of space‖ Because Democritus could not experimentally support his theory and prove what held atoms together, Aristotle’s reputation killed the idea of the atom for 1600 years! Next 2000 years of chemistry history is dominated by alchemy – pseudoscience that was obsessed with turning cheap metals into gold. John Dalton (1766-1844) Beginning of development of modern atomic theory Revised Democritus’ ideas based upon scientific research Dalton’s Atomic Theory Isaac Newton (1642 – 1727) & Robert Boyle (1627 – 1691) Renewed support of Democritus Still no experiments or predictions Boyle First ―chemist‖ to perform quantitative experiments Measured relationship between pressure and volume of air Substance is an element if it can’t be broken down into 2+ simpler substances All matter is composed of extremely small particles called atoms All atoms of a given element are identical – same size, mass, chemical properties Atoms of one element are different from atoms of another element Atoms cannot be created, divided into smaller particles or destroyed Different atoms combine in wholenumber ratios to form compounds In a chemical reaction, atoms are separated, combined, or rearranged The Structure of the Atom Dalton and the Law of Conservation of Mass Chemical reactions merely separate, combine, or rearrange atoms Experimental evidence and explanation of the composition of compounds (with the conservation of mass) led to the acceptance of Dalton’s atomic theory Atom (definition): The smallest particle of an element that retains all the properties of that element 2 Problems with Dalton’s atomic theory Atoms are indivisible – we now know that there are subatomic particles that make up atoms All atoms of the same element are identical – some atoms of an element may have slightly different masses (isotopes) Atoms are incredibly small, but can be seen with a scanning tunneling microscope (STM). Using the STM, scientists can: Move individual atoms around to form shapes, patterns and simple machines Nanotechnology – field of manipulating atoms – to manufacture machines the size of molecules. Section 4.2: Subatomic Particles and the Nuclear Atom History of discovery Due to the presence of static electricity, scientists in the 1800s began to look for a connection between matter and electric charge o How would electricity behave in the absence of matter (in a vacuum with the air/matter removed)? William Crookes’ experiment with a cathode tube – he saw a flash of light produced by some form of radiation strike a coating at the opposite end of the tube. o Because of the experiment, these became known as cathode rays. http://micro.magnet.fsu.edu/electroma g/java/crookestube/ Impact of Thomson’s work Particle smaller than an atom o Atoms are divisible into smaller subatomic particles Dalton’s theory was wrong o Not believed initially, but further experiments proved Thomson right. Lessons from the cathode ray tubes Cathode rays are streams of charged particles Particles have a negative charge, but the exact value of the charge is unknown Found in all forms of matter o Now known as electrons J.J. Thomson Because experiments to measure the particle’s mass directly had failed, Thomson worked to compare the ratio of the cathode ray charge to its mass Determined that the mass of the charged particle was less than a hydrogen atom (the lightest atom known) Thomson discovered the electron in the late 1890s! The Structure of the Atom 3 Next step… Questions about the atom: Robert Millikan determined the charge of an electron (1909). 1. How can the atom be electrically neutral, if it is known to have a negative charge? 2. Mass of an electron is miniscule, what accounts for the rest of the mass in a typical atom? Equivalent to a single unit of negative charge Mass of an electron = 9.1 x 10-28 g (1/1840 of the mass of a hydrogen atom) Thomson’s atomic model ―Plum pudding model‖ (for us, it would be known as the chocolate chip cookie dough model). o Spherical model o Chocolate chips are the electrons o Cookie dough is the uniformly distributed positive charge Plum pudding model didn’t last long! Ernest Rutherford (1911) Completing the atom 1920 Rutherford refined his concept of the nucleus o Nucleus contains positively charged particles called protons Proton subatomic particle with a charge equal to but opposite of electron +1 charge 1932 Chadwick showed that the nucleus contains another subatomic particle, a neutral particle called the neutron o Mass of neutron is nearly equal to that of the proton o No electrical charge Gold foil experiment Plum pudding model is incorrect Atom is mostly empty space Center of the atom is a dense region (he called it the nucleus), that contains most of the atom’s mass and all of its positive charge Beginning of the modern model of the atom Rutherford’s new nuclear atomic model Electrons move rapidly through empty space and are held by the attraction of a positively-charged nucleus Atom: Electrically neutral Spherical Protons (+) = Electrons (―) Nucleus (+ charged) Neutrons = no charge Electrons (―) surround nuc. http://www.mhhe.com/physsci/chemist ry/essentialchemistry/flash/ruther14.sw f The Structure of the Atom 4 Modern Atomic Theory: Atoms are composed of a nucleus (containing positively charged protons and neutrally charged neutrons) surrounded by negatively charged electrons) Protons and electrons are identical in number in a neutral atom. Nucleus of the atom: o Positively charged o Composed of protons and neutrons o 99.97% of the mass of the atom Section 4.3: How Atoms Differ Henry Moseley (1887 – 1915) Discovered that atoms of each element contain a unique positive charge in their nuclei o Number of protons in an atom identifies that atom as being a particular element o # proton = atomic number o Atomic number determines the element’s position on the periodic table o Moseley reordered Mendeleev’s periodic table o Periodic table is organized left-to-right and top-to-bottom by increasing atomic number o All atoms are neutral, so the # of protons equals the #s of electrons chemical name 1 Atomic number H Chemical symbol 1.008 Average atomic mass Isotopes and mass number Dalton’s errors: atom is indivisible, all atoms of an element are identical All atoms of an element have the same # of protons (and electrons) # of neutrons in nuclei may differ o Potassium (20 neutrons, 21 neutrons, 22 neutrons) Solving for number of neutrons: Atoms with the same number of protons but different numbers of neutrons are called isotopes o Most elements in nature exist as a mixture of isotopes Mass number is the sum of the number of protons and neutrons in an element # N0 (# of neutrons) = atomic mass # — atomic # Example: K-39 = 39-19 = 20 N Hydrogen o All react the same chemically and physically 0 Isotopes differ in mass: more neutrons = more mass Identify isotopes by their mass number, ex. potassium-39, potassium-40, potassium-41. The Structure of the Atom 5 Naming and writing isotopes Mass of individual atoms Chemists often write out isotopes using a shortened type of notation using only the chemical symbol, atomic number, and mass number. Mass of neutrons and protons are ~ 1.67 x 10-24g 40 41 K K 19 19 Mass number is written as a superscript to the left of the chemical symbol Atomic number is written as a subscript to the left of the chemical symbol Atomic mass Atomic mass is the weighted average of the mass of all the isotopes of an element Most of the mass is in the nucleus Mass of electron: ~ 9.1 x 10-28g Small numbers are difficult to work with. A new measuring unit: mass of an atom is relative to the mass of a atomic standard, which is carbon12. Carbon-12 atoms have an atomic mass of exactly 12 atomic mass units. Calculating atomic mass 1. Mass of each isotope multiplied by the percentage of the time it occurs naturally 1 atomic mass unit (amu) = 1/12 the mass of a carbon-12 atom 1 amu is approximately equal to the mass of a single proton or neutron 2. Add the products derived for each isotope Analyzing atomic mass: Example: Looking at the atomic mass of an element will help you determine what the most abundant isotope may be. Cl-35 = 76% (.76 x 35) = 26.6 Cl-37 = 24% (.24 x 37) = 8.9 Atomic mass of Cl 35.5 amu Ex. Exception to this logic: Bromine has an atomic mass of 79.904 amu, leading you to expect that the most abundant isotope is bromine-80. Not so. Bromine has 2 isotopes (bromine-79 (50.69%) and bromine-81 (49.31%), giving a weighted atomic mass of approximately 80 amu. Fluorine (F) has an atomic mass close to 19 (18.998) amu. Fluorine has an atomic number of 9, so it has 10 neutrons. If fluorine had several fairly abundant isotopes, it would be unlikely that its atomic mass would be so close to a whole number. The Structure of the Atom 6 Section 4.4: Unstable Nuclei and Radioactive Decay Radioactivity Types of radiation Chemical reaction: involves the change of one or more substances into new substances. Alpha (α) radiation Atoms may be rearranged, but their identities do not change Chemical reactions only involve an atom’s electrons – the nucleus is not changed Each alpha particle contains 2 protons and 2 neutrons (2 P+, 2 N0) Has a 2+ charge Equivalent to a helium-4 nucleus Represented by Reactions that change an atom of one element into another element Nuclear reactions: reactions that involve a change in an atom’s nucleus These reactions spontaneously emit radiation – (radioactivity) Changes the identity of the element o Atoms of one element change into atoms of a different element Spontaneously emit radiation because their nuclei are unstable o Unstable systems gain stability by losing energy o Unstable nuclei lose energy by emitting radiation – called radioactive decay o Unstable nuclei will undergo radioactive decay until they form a stable nonradioactive atom, usually of a different element Beta (β) radiation Nuclear Equation Equations that account for the changing of atomic number and atomic masses during radioactive decay. During alpha decay, mass number decreases by 4 and atomic number decreases by 2 Since opposites attract, the positively charged alpha particle is attracted to the negatively charged plate shown below Radiation is deflected toward the positive plate (see below) Fast moving electrons, so has a 1— charge Represented by Gamma (γ) radiation or gamma rays High-energy radiation that possesses no mass Usually accompanies the other 2 types of radiation Accounts for the most energy lost during the radioactive decay process Does not result in the formation of a new atom Represented by The Structure of the Atom 7 Nuclear Equations Representation of radioactive decay Sample problems (pages 106-107) 1. 2. 3. Nuclear Stability Primary factor in determining nuclear stability is the ratio of neutrons to protons Nuclei with too many or too few neutrons are unstable Unstable nuclei lose energy through radioactive decay and become stable with the appropriate ratio of neutrons to protons Few radioactive atoms are found in nature – most have already decayed into stable atoms. Strong force Force that holds N0 and P+ together in the nucleus Only effective over very short distances (≈10-3 pm) 100,000 times stronger than gravity