Slide 1:Chapter 02 Atomic Theory and Atomic Structure


Slide 2:Democritus 440 B.C. Aristotle Expressed the belief that all matter consists of very small, indivisible particles he named atomos (ucuttable or indivisible) Aristotle didn’t think so. He believed that everything is just a combination of Earth, Air, Fire and Water.


Slide 3:Antoine Lavoisier 1774 Lavoisier observed that in a chemical reaction, the mass of the reactants is equal to the mass of the products. Law of Conservation of Mass “Mass can neither be created nor destroyed in chemical reactions” 4.55 g + 2.02 g = 6.57 g 3.25 g + 3.32 g = 6.57 g


Slide 4:Joseph Proust 1799 Proust observed that water’s percentage by mass is 88.8% Oxygen and 11.2% Hydrogen from any source. Law of Constant Composition “Different samples of a pure chemical substance always contain the same proportion of elements by mass.”


Slide 5:Law of Multiple Proportions “Elements can combine in different ways to form different substances, whose mass ratios are small whole-number multiples of each other.”


Slide 6:John Dalton 1808 Dalton's Atomic Theory Elements are composed of extremely small particles, called atoms All atoms of a given element are identical, having the same size, mass, and chemical properties. The atoms of one element are different from the atoms of all other elements. Compounds are composed of atoms of more than one element. In any compound, the ratio of the numbers of atoms of any two of the elements present is either an integer or a simple fraction. A chemical reaction involves only the separation, combination, or rearrangement of atoms; it does not result in their creation or destruction. ISOTOPES Law of Conservation of Mass Law of Constant Composition Law of Multiple Proportions


Slide 7:Atom is the basic unit of an element that can enter into chemical combination


Slide 8:Around 1850s scientists started experimenting in the phenomena of radiation Radiation is the emission and transmission of energy through space in the form of waves


Slide 9:To study some aspects of radiation, scientists used a cathode ray tube. Which has a negatively charged plate (cathode) emits an invisible ray which is drawn to the positively charged plate (anode). And when it strikes a specially coated surface it produces a bright light (fluorescence).


Slide 10:Electromagnetic properties of Cathode Rays


Slide 11:Cathode rays are streams of negative particles of energy and Thomson identified them as electrons. With his knowledge of electromagnetic theory and the cathode ray tube, Thomson determined the ratio of electric charge to the mass of an individual electron ( -1.76 x 108 C/g) John Joseph Thomson 1897


Slide 13:R. A. Millikan eventually found the charge of an electron to be – 1.6022 x 10-19 C. And from these data he found that the calculated mass of an electron is 9.10 x 10-28 g.


Slide 15:Roentgen observed that cathode rays caused glass and metals emit very unusual rays. These rays penetrated matter, darkened covered photographic plates, and caused a variety of substances to fluoresce but are not affected by an electric field and a magnet. He called it X-rays. Wilhelm Roentgen 1895


Slide 16:Henri Antoine Becquerel 1896 Becquerel observed that uranium spontaneously emit rays, similar to X-rays, and darkened photographic plates without the simulation of cathode rays.


Slide 17:Marie Curie 1896 Curie, a student of Becquerel, suggested the name radioactivity. Radioactivity is the spontaneous emission of particles and/or radiation.


Slide 18:Rays emitted by radioactive elements


Slide 19:Two features of atoms had become clear: They contain electrons and are electrically neutral. Thomson proposed that an atom is a uniform, positive sphere of matter in which electrons are embedded. And he likened it to a traditional English desert plum-pudding.


Slide 20:In 1910, Ernest Rutherford started experimenting on alpha particles hitting thin gold foils


Slide 21:Rutherford proposed that the atom’s positive charges are concentrated in the nucleus, a dense central core within the atom. Protons are positively charged particles in the nucleus Ernest Rutherford 1910


Slide 22:If the size of an atom were expanded to that of a sports stadium, the size of the nucleus would be that of a marble.


Slide 23:Chadwick showed evidence of electrically neutral particles having a mass slightly greater than that of protons and named it neutrons. James Chadwick 1932


Slide 24:Mass and Charge of Subatomic Particles


Slide 25:Practice on the Development of the Atomic Theory


Slide 26:Practice on the Development of the Atomic Theory


Slide 27:Isotopes are atoms of the same element but have different masses. Isotopes have the same number of protons but different number of neutrons. Hydrogen has three naturally occurring isotopes protium, deuterium, and tritium.


Slide 28:Atomic number (Z) is the number of protons in the nucleus of an element *In a neutral atom (without charge), the number of protons is equal to the number of electrons Mass Number (A) is the total number of neutrons and protons present in the nucleus


Slide 29:Practice on Writing Nuclear Symbols Write the nuclear symbols of the following isotopes of hydrogen.


Slide 30:How nuclear symbols are read: Uranium-235 (“uranium two thirty-five”) is used in nuclear reactors and atomic bombs while Uranium-238 (“uranium two thirty-eight”) is not.


Slide 31:How to determine number of subatomic particles from nuclear symbols Uranium-235 has 92 protons 92 electrons 143 neutrons How about Uranium-238?


Slide 32:Practice on Determining the Number of Subatomic Particles with Nuclear Symbols How many protons, neutrons, and electrons are in the following isotope of copper 29 protons, 34 neutrons, and 29 electrons


Slide 33:For each of these species, determine the number of protons and the number of neutrons in the nucleus: Practice on Determining the Number of Subatomic Particles with Nuclear Symbols


Slide 34:2.4 The Periodic Table is a chart in which elements having similar chemical and physical properties are grouped together. The horizontal rows are called periods. The vertical columns are called groups. Some groups have special names.


Slide 35:The Periodic table can be subdivided into metals, nonmetals and metalloids Metals are good conductors of heat and electricity Nonmetals are poor conductors of heat and electricity Metalloids are intermediate between those of metals and nonmetals


Slide 36:Practice with the periodic table


Slide 37:Practice with the periodic table


Slide 38:Monatomic elements are elements that exists as single atoms in nature Such as the noble gases.


Slide 39:A Molecule is an aggregate of at least two atoms in a definite arrangement held together by chemical forces (chemical bonds) A diatomic molecule contains only two atoms A polyatomic molecule contains more than two atoms H2, N2, O2, Br2, HCl, CO O3, H2O, NH3, CH4


Slide 40:An ion is an atom, or group of atoms, that has a net positive or negative charge. cation – ion with a positive charge If a neutral atom loses one or more electrons it becomes a cation. anion – ion with a negative charge If a neutral atom gains one or more electrons it becomes an anion.


Slide 41:A monatomic ion contains only one atom A polyatomic ion contains more than one atom Na+, Cl-, Ca2+, O2-, Al3+, N3- OH-, CN-, NH4+, NO3-


Slide 42:Practice on Determining the Number of Subatomic Particles with Nuclear Symbols of Ions How many protons, neutrons, and electrons are in 13 protons, 14 neutrons, and 10 electrons


Slide 43:Practice on Determining the Number of Subatomic Particles with Nuclear Symbols of Ions How many protons, neutrons, and electrons are in 34 protons, 44 neutrons, and 36 electrons


Slide 44:Give the number of protons and electrons in each of the following ions: Practice on Determining the Number of Subatomic Particles with Nuclear Symbols of Ions


Slide 45:Common Monatomic ions


Slide 46:Practice on Nuclear Symbols


Slide 47:Practice on Nuclear Symbols


Slide 48:Practice on Nuclear Symbols


Slide 49:Practice on Nuclear Symbols


Slide 50:Practice on Nuclear Symbols


Slide 51:A molecular formula shows the exact number of atoms of each element in the smallest unit of a substance An empirical formula shows the simplest whole-number ratio of the atoms in a substance H2O C6H12O6 CH2O O3 O N2H4 NH2


Slide 52:An Allotrope is one of two or more distinct forms of an element.


Slide 53:Molecular Models


Slide 54:Ionic compounds consist of a combination of cations and an anions The formula is always the same as the empirical formula The sum of the charges on the cation(s) and anion(s) in each formula unit must equal to zero.


Slide 55:Al2O3 Al3+ O2- CaBr2 Ca2+ Br- Na2CO3 Na+ CO32- Formula of Ionic Compounds The subscript of the cation is numerically equal to the charge on the anion, and the subscript of the anion is numerically equal to the charge on the cation


Slide 56:Atomic mass is the mass of an atom in atomic mass units (amu) By definition: 1 atom 12C “weighs” 12 amu On this scale 1H = 1.008 amu 16O = 16.00 amu


Slide 57:Average Atomic mass is the weighted average of the masses of the naturally occurring isotopes of a particular element based on their abundance. Natural lithium is: 7.42% 6Li (6.015 amu) 92.58% 7Li (7.016 amu) = 6.941 amu


Slide 58:This is the mass that you see in the periodic table.


Slide 59:Practice on Average Atomic Mass The atomic masses of chlorine-35 (75.53 %) and chlorine-37 (24.47 %) are 34.968 amu and 36.956 amu, respectively. Calculate the average atomic mass of chlorine . The percentages in parentheses denote the relative abundances.


Slide 61:Quantum Numbers are required to describe the distribution of electrons in atoms.


Slide 62:Principal Quantum Number (n) Relates to the average distance of the electron from the nucleus in a particular orbital.


Slide 63:Angular Momentum Quantum Number (l) Shape of the “volume” of space that the e- occupies l has possible integral values from 0 to (n-1). l values are generally designated by letters n = 1, l = 0 n = 2, l = 0 or 1 n = 3, l = 0, 1, or 2 l = 0 s orbital l = 1 p orbital l = 2 d orbital l = 3 f orbital


Slide 64:Shapes of the Orbitals


Slide 66:Magnetic Quantum Number (ml) Describes the orientation of the orbital in space. The allowed values for ml are –l to +l for a given value of l ml = -l, …., 0, …. +l if l = 1 (p orbital), ml = -1, 0, or 1 if l = 2 (d orbital), ml = -2, -1, 0, 1, or 2


Slide 67:ml = -1 ml = 0 ml = 1 ml = -2 ml = -1 ml = 0 ml = 1 ml = 2 Orientations of orbitals in space


Slide 70:Experimental Arrangement for Demonstrating the Spinning Motion of Electrons


Slide 77:Electron Spin Quantum Number (mS) Describes the spin of the electron (clockwise or counterclockwise) The allowed values are +1/2 and -1/2 ms = -½ ms = +½


Slide 78:Allowed Combinations of Quantum Numbers n, l and ml for the First Four Shells


Slide 79:Example on Quantum Numbers Give the values of the quantum numbers associated with the orbitals in the 3p subshell. n = 3 l =1 ml = -1,0,1


Slide 80:Example on Quantum Numbers What is the total number of orbitals associated with the principal quantum number n = 4? 16 orbitals


Slide 81:Practice with possible quantum numbers


Slide 82:Practice with quantum numbers


Slide 83:Practice with quantum numbers


Slide 84:1s1 Electron configuration is how the electrons are distributed among the various atomic orbitals in an atom.


Slide 85:Pauli Exclusion Principle “No two electrons in an atom can have the same four quantum numbers.”


Slide 87:H 1 electron H 1s1 He 2 electrons He 1s2 Li 3 electrons Li 1s22s1 Be 4 electrons Be 1s22s2 B 5 electrons B 1s22s22p1 C 6 electrons Aufbau Building Up Principle “A maximum of two electrons are put into orbitals in the order of increasing orbital energy: the lowest-energy orbitals are filled before electrons are placed in higher-energy orbitals.”


Slide 88:As a consequence of the Aufbau Building Up Principle, this is the order by which the orbitals are filled up with electrons in an atom.


Slide 89:Hund’s Rule of Multiplicity “The most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins“ C 6 electrons C 1s22s22p2 N 7 electrons N 1s22s22p3 O 8 electrons O 1s22s22p4 F 9 electrons F 1s22s22p5 Ne 10 electrons Ne 1s22s22p6


Slide 94:Practice with electronic configuration


Slide 95:Practice with electronic configuration


Slide 96:Practice with electronic configuration


Slide 97:Practice with electronic configuration


Slide 98:Practice with electronic configuration


Slide 99:Practice with electronic configuration


Slide 100:Practice with electronic configuration


Slide 101:Paramagnetic unpaired electrons Diamagnetic all electrons paired Paramagnetic substances are those that contain net unpaired spins and are attracted by a magnet. Diamagnetic substances do not contain net unpaired spins and are slightly repelled by a magnet.


Slide 102:Practice with paramagnetism and diamagnetism


Slide 103:Practice with paramagnetism and diamagnetism


Slide 105:3d 4d 5d 6d 1s 2s 3s 4s 5s 6s 7s 4f 5f 1s 2p 4p 5p 6p 3p 7p Outermost subshell being filled with electrons


Slide 107:1s 2s 3s 4s 5s 6s 7s


Slide 108:1s 2p 4p 5p 6p 3p 7p


Slide 109:3d 4d 5d 6d


Slide 110:4f 5f


Slide 111:3d 4d 5d 1s 2s 3s 4s 5s 6s 7s 4f 5f 1s 2p 4p 5p 6p 3p 6d 7p


Slide 112:Practice with electronic configuration


Slide 114:Newlands noticed that when elements were arranged according to their atomic masses, every eight element had similar properties – Law of Octaves. John Newlands 1864


Slide 115:Lothar Meyer 1869 Meyer and Mendeleev independently proposed a more extensive tabulation of the elements Dmitri Mendeleev 1869


Slide 116:Mendeleev provided spaces for still undiscovered elements


Slide 117:Comparison of Predicted and Observed Properties for Gallium (eka-Aluminum) and Germanium (eka-Silicon)


Slide 118: This suggested that there is a better basis for the arrangement of the elements in the periodic table than the average atomic mass .


Slide 119:Henry Moseley 1913 Moseley suggested to arrange the periodic table according to their atomic numbers.


Slide 123:Valence Electrons are the outer electrons of an atom which are involved in chemical bonding


Slide 124:Valence Electrons of the representative elements are in the ns2np6


Slide 125:Transition metals have incompletely filled d orbitals


Slide 126:Lanthanides and Actinides have incompletely filled f orbitals


Slide 127:Ground State Electron Configurations of the Elements


Slide 128:Blocks of the periodic table, corresponding to filling the different kinds of orbitals.


Slide 129:Classification of the Elements.


Slide 130:Na [Ne]3s1 Na+ [Ne] Ca [Ar]4s2 Ca2+ [Ar] Al [Ne]3s23p1 Al3+ [Ne] Atoms lose electrons so that cation has a noble-gas outer electron configuration. H 1s1 H- 1s2 or [He] F 1s22s22p5 F- 1s22s22p6 or [Ne] O 1s22s22p4 O2- 1s22s22p6 or [Ne] N 1s22s22p3 N3- 1s22s22p6 or [Ne] Atoms gain electrons so that anion has a noble-gas outer electron configuration. Electron Configurations of Cations and Anions Of Representative Elements


Slide 131:Cations and Anions Of Representative Elements


Slide 132:Na+: [Ne] Al3+: [Ne] F-: 1s22s22p6 or [Ne] O2-: 1s22s22p6 or [Ne] N3-: 1s22s22p6 or [Ne] Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne Ions, or atoms and ions, that have the same number of electrons, and hence the same ground-state electron configuration are said to be isoelectronic.


Slide 133:Electron Configurations of Cations of Transition Metals When a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital and then from the (n – 1)d orbitals. Fe: [Ar]4s23d6 Fe2+: [Ar]4s03d6 or [Ar]3d6 Fe3+: [Ar]4s03d5 or [Ar]3d5 Mn: [Ar]4s23d5 Mn2+: [Ar]4s03d5 or [Ar]3d5


Slide 148:Atomic Radius is one half the distance between the two nuclei in two adjacent atoms


Slide 149:Atomic Radii (in picometers) of representative elements according to their positions in the periodic table.


Slide 150:Ionic Radius is the radius of a cation or an anion. Cation is always smaller than atom from which it is formed. Anion is always larger than atom from which it is formed.


Slide 151:The radii (in picometers) of ions of familiar elements arranged according to the elements’ positions in the periodic table.


Slide 152:I1 first ionization energy I2 second ionization energy I3 third ionization energy I1 < I2 < I3 Ionization energy is the minimum energy (kJ/mol) required to remove an electron from a gaseous atom in its ground state.


Slide 153:General Trend in First Ionization Energies


Slide 154:Electron affinity is the negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion. DH = -328 kJ/mol EA = +328 kJ/mol DH = -141 kJ/mol EA = +141 kJ/mol


Slide 164:Group 1A Elements (ns1, n  2)


Slide 165:Group 2A Elements (ns2, n  2)


Slide 166:Group 3A Elements (ns2np1, n  2)


Slide 167:Group 4A Elements (ns2np2, n  2)


Slide 168:Group 5A Elements (ns2np3, n  2)


Slide 169:Group 6A Elements (ns2np4, n  2)


Slide 170:Group 7A Elements (ns2np5, n  2)


Slide 171:Group 8A Elements (ns2np6, n  2)