UNDERLYING STRUCTURE OF MATTER

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1 1 UNDERLYING STRUCTURE OF MATTER Chapter 4 Atomic Structure DEFINING THE ATOM Earlier theories of matter: A. Even though his hypothesis lacked evidence at the time, the Greek philosopher Democritus (460 B.C. 370 B.C.) believed that all substances were composed of small indivisible particles called atoms. B. John Dalton's ( ) Atomic Theory states: 1) All elements are composed of tiny indivisible particles called atoms. 2) Atoms of the same element are identical. The atoms of any one element are different from those of any other element. 3) Atoms of different elements can physically mix together or can chemically combine in simple whole number ratios to form compounds. 4) Chemical reactions occur when atoms are separated, joined or rearranged. Atoms of one element, however, are never changed into atoms of another element as a result of a chemical reaction. Despite their small size, individual atoms are observable with instruments such as scanning tunneling microscopes. This has fostered the development of atomic-scale (aka nanoscale) technology used for medicine, communications and space exploration. STRUCTURE OF THE NUCLEAR ATOM Discovering subatomic particles: C. J. J. Thomson ( ) discovered negatively charged particles in the atom and called them electrons. He passed an electric current through gases at low pressure within sealed glass tubes (cathode ray tubes). The resulting glowing beam consisted of tiny negatively charged particles moving at high speed. Thomson concluded that electrons must be part of every atom. D. Robert Millikan ( ) determined the charge and the mass of the electron. The electron carries exactly one unit of negative charge, and its mass is 1/1840 the mass of a hydrogen atom. The following ideas about matter and electrical charges allowed scientists to logically discover two other subatomic particles: 1) Atoms are neutral (do not have an overall electric charge). 2) Electric charges are carried by particles of matter. 3) Electric charges always exist in whole-number multiples of a single basic unit (that is, no fractions of a charge). 4) When a given number of negatively charged particles combine with an equal number of positively charged particles, an electrically neutral particle is formed. E. Eugen Goldstein ( ), while observing a cathode ray tube, found rays travelling in the direction opposite to that of the cathode ray. He concluded that they were composed of positive particles later called protons. Protons are described to have a mass 1840 times that of an electron. F. James Chadwick (1932) demonstrated that the atomic nucleus must contain heavy neutral particles, as well as the protons, called neutrons. Properties of subatomic particles are summarized in the following table: Particle Symbol Relative Charge Relative Mass Actual Mass (g) Electrons e - 1-1/ x Protons p x Neutrons N x 10-24

2 2 Theoretical Physicists believe that these subatomic particles are composed of yet smaller sub-nuclear particles called quarks. Thomson, and most scientists, originally thought that electrons were evenly distributed throughout the atom filled uniformly with positively charged material. In this theory (known as the plum-pudding model or raisin bun model), the electrons were stuck in a chunk of positive charge. G. During his studies at McGill University in Montreal, Ernest Rutherford ( ) hypothesized that atoms contained a tiny, positively charged core called the nucleus surrounded by electrons. Rutherford s atomic model, the nuclear atom, was also known as the bee hive model. The nucleus contains most of the atom s mass but only 1/1000 of the total size of the atom (Gold-Foil Experiment, page 107). This was an improvement upon Thomson s models but it needed to be revised in order to explain the chemical properties of the elements. DISTINGUSHING AMONG ATOMS The atomic number of an element is the number of protons in the nucleus of an atom of that element. Elements are different because they contain different numbers of protons. The mass number is the total number of protons and neutrons in an atom. Standard atomic notation An element can come in different varieties called isotopes. Isotopes are a form of an element in which the atoms have the same number of protons but different numbers of neutrons. The mass of isotopes will vary due to the uneven number of neutrons. Example: isotopes of carbon Carbon -14 does not last as long as Carbon -12 and eventually loses the two neutrons. This is called radioactive decay and takes up to 2000 years. Example: isotopes of hydrogen

3 Since the masses of individual atoms are so small and impractical to work with, we compare the relative mass of atoms using a reference isotope as a standard. The isotope chosen is carbon-12. This isotope of carbon was assigned a mass of exactly 12 atomic mass units, 12 amu. So in summary 1 amu = 1/12 the mass of a carbon-12 atom The atomic mass of an element is a weighted average mass of the atoms in a naturally occurring sample of the element. A weighted average mass reflects both the mass and the relative abundance of the isotopes as they occur in nature. To calculate the atomic mass of elements, multiply the mass of each isotope by the natural abundance, expressed as a decimal, and then add the products (see page 117 in text for an example). Chapter 5 Electrons in Atoms MODELS OF THE ATOM Rutherford s planetary model of the atom (the bee hive model) could not explain the chemical properties of elements. For example it could not explain why metals and compounds of metals give off characteristic colors when heated in a flame or why objects, when heated to higher and higher temperatures first glow dull red, then yellow, and then white. H. Niels Bohr ( ) changed Rutherford s model to include how the energy of an electron changes when it absorbs or emits light. Bohr proposed that an electron is found in only specific circular path, or orbits, around the nucleus. Each orbit has a fixed energy called energy levels (like the unequal rungs of the ladder on the right). To move from one energy level to another, electrons must gain or lose just the right amount of energy. They can t exist between levels. Quantum is considered the amount of energy required to move an electron from one energy level to another energy level. The amount of energy an electron gains or loses in an atom is not always the same. As mentioned above, and as seen in the image on the right, the energy levels in an atom are not evenly spaced. The higher energy levels are closer together. It takes less energy to move between the higher levels. I. In 1926, Erwin Schrodinger developed a mathematical wave equation that indicated the probability of finding an electron in any given space. Some areas have higher probability than others. It is impossible to locate the exact location of electrons due to Heisenberg s Uncertainty Principle (see p. 145) (e.g., rotating propeller blade). All the points of high probability can be connected and forms a 3-D fuzzy cloud called an orbital. The size and shape of the orbital depends on the energy of the electron. The four orbitals discussed here are s orbital spherical p orbital dumbbell and have three different orientations. d orbital cloverleaf and have five different orientations. f orbital complicated shapes, and have seven different orientations. 3

4 4 There are four specific quantum numbers: 1. Principle Quantum Number n Indicates the energy level. The levels are numbered 1, 2, 3 n. Describes the size of the fuzzy cloud. The greatest # of electrons in an energy level is calculated by the formula 2n 2 : n = 1 2(1) 2 = 2 electrons n = 2 2(2) 2 = 8 electrons n = 3 2(3) 2 = 18 electrons 2. Second Quantum Number - l Indicates specific sublevel in the energy level. Describes the shape of the fuzzy cloud. The number of sublevels in each energy level varies. It depends on the size of the level. The sublevels are named in order of energy (s, p, d, f). For example n = 1 1 sublevel s n = 2 2 sublevels s,p n = 3 3 sublevels s,p,d n = 4 4 sublevels s,p,d,f n = 5 4 sublevels s,p,d,f Overlapping of Energy levels: Notice that 4s has lower energy that the 3d sublevel. The lower the energy the more stable the sublevel is. 3. Third Quantum Number - m Indicates the number of orbitals (the space occupied by two electrons). This describes the orientation of the cloud. Sublevel s contains 1 pair of electrons Sublevel p contains 3 pairs of electrons Sublevel d contains 5 pairs of electrons Sublevel f contains 7 pairs of electrons 4. Fourth Quantum Number s Indicated the spin of the electron [clockwise (1/2) or counterclockwise (-1/2)]. The two electrons in the same orbital have opposite spin. ELECTRON ARRANGEMENT OF ATOMS An orbital diagram is a 1-D energy graph showing the relative energies of orbitals in each energy level and sublevel. According to the Aufbau principle, electrons occupy the orbitals of lowest energy first. According to Hund s Rule, no electron pairing takes place in sublevels until each has one electron first. According to the Pauli Exclusion Principle, no two electron in an atom can have the same set of four quantum numbers. The way in which electrons are arranged in various orbitals around the nucleus is called electron configuration. The following are steps to writing electron configurations: 1. Determine the total number of electrons in the atom or ion. 2. Start assigning electrons in increasing order of energy level and sublevels. 3. Continue this assignment using the diagonal rule here on the right (helps with overlapping).

5 The noble gases are unique because their electron configurations involve completely filled energy levels. You can abbreviate large electron configurations by starting with the noble gas that precedes it. This is called the core (or kernel) short cut. Ex. Cr [Ar]4s 2 3d 4 There are some exceptions to the rules. Some electron configurations differ from the rules because halffilled sublevels are not as stable as filled sublevels BUT they are more stable than other configurations. As an example, consider groups 6 and 11: Cr 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 instead of 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 WHY? Our primary concern when dealing with electrons involves the electrons in the outer energy level. These are called valence electrons. When drawing electron dot diagrams, the following rules apply: 1. The first two electrons are in the s orbital. 2. Other electrons (six more) fill the p orbitals one at a time (Hund s rule). 3. There will never be more than 8 electrons Chapter 6 The Periodic Table The periodic table allows you to easily compare the properties of one element (or group of elements) to another element (or groups of elements). Groups (columns) have similar chemical and physical properties. The properties of the elements within periods (rows) vary as you move across from element to element. This pattern of properties then repeats as you move to the next period. Electrons play a key role in determining the properties of elements. Based on their electron configurations, atoms can be sorted into four categories (see figure 6.12 on page166): 5 Groups Representative elements (Groups 1 and 2) Representative elements (Groups 13 to 17) Transition metals (Groups 3 to 12) Inner transition metals (the lanthanides and actinides are the two periods at the bottom) Electron Configuration The s sublevels are being filled one at a time. The p sublevels are being filled one at a time. The d sublevel is being filled one at a time. The f sublevel is being filled one at a time. PERIODIC TRENDS 1) The atomic radius (which is a reflection of the size of the atom) is defined as one half of the distance between

6 the nuclei of two atoms of the same element when the atoms are joined. There are two trends to consider with respect to atomic radius: a) In general, atomic size INCREASES from top to bottom within a GROUP. As atomic number increases within a group, the charge of the nucleus increases and the number of energy levels increases. The increase of positive charge draws the electrons closer to the nucleus BUT the increase of occupied orbitals shield the higher energy levels from the attraction of protons in the nucleus. The shield effect wins out so atomic size increases as you head down the group b) In general, atomic size DECREASES from left to right across a PERIOD. As the atomic number increases the electrons are going into the same energy level and the shielding affect is constant. At the same time the nucleus is growing and the positive charge increases pulling the electrons in the energy levels closer causing the atomic size to decrease as you move across the period. 2) Ions are atoms, or groups of atoms, that have a charge because of a transfer of electrons. Metals tend to lose electrons becoming positively charged (cations) and nonmetals tend to gain electrons becoming negatively charged (anions). When there is enough energy absorbed by an atom to overcome the attraction of the protons in the nucleus, electrons will move to higher energy levels. The energy required to actually REMOVE an electron from an atom is called the ionization energy. It is measured when an element is in its gaseous state. There are two trends to consider with respect to ionization energy: a) As the atomic size increases, nuclear charge has a smaller effect on the electrons in the highest occupied energy level. As a result, less energy is required to remove an electron from this energy level. b) Nuclear charge increases from left to right but the shielding effect remains constant. So, there is an increase in the attraction of the nucleus for an electron, therefore it takes more energy to remove an electron from an atom. In summary, the first ionization energy tends to decrease from top to bottom within a group and increase from left to right across a period. 3) When considering trends in the size of ions, cations are always smaller than the atoms from which they form and anions are always larger than the atoms from which they form. When metals lose an electron, the attraction between the remaining electrons and the nucleus is increased. The electrons are draw closer to the nucleus. Representative metals tend to lose all of their outermost electrons during ionization therefore the ion has one fewer occupied energy level. When nonmetals, like halogens, gain electrons the attraction of the nucleus for these electron decreases. The ion is much larger than the atom. 4) Electronegativity is the ability of an atom of an element to attract electrons when the atom is in a compound. It is helpful to consider the connection that electronegativity has with the radius of an atom (recall trend #1). In general, electronegativity values decrease from top to bottom within a group. For representative elements, the values tend to increase from left to right across a period. Note that the transition metals do not have consistent electronegativity values. 6

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