05 Periodic Properties of the Elements.pdf

Periodic Properties of the Elements Outline Development and Organization of the Periodic Table Trends in the Periodic Table Atomic Structure and Reactivity Development and Organization of the Periodic Table Dmitri Mendeleev (1834-1907) – a Russian chemist who arranged 65 elements then...

Periodic Properties of the Elements Outline Development and Organization of the Periodic Table Trends in the Periodic Table Atomic Structure and Reactivity Development and Organization of the Periodic Table Dmitri Mendeleev (1834-1907) – a Russian chemist who arranged 65 elements then known into a periodic table of rows and columns and summarized their behavior in the periodic law: When arranged by atomic mass, the elements exhibit a periodic recurrence of similar properties. Focused on chemical properties Given greater credit because he was able to predict the properties of several as-yet-undiscovered elements, for which he had left blank spaces in his table. Development and Organization of the Periodic Table Julius Lothar Meyer (1830-1895) – a German chemist who worked independently from Dmitri Mendeleev but arrived at virtually the same organization of elements almost simultaneously. Focused on physical properties Development and Organization of the Periodic Table Today’s periodic table resembles Mendeleev’s in most details, although it includes at least 50 elements that were unknown in 1870. The only substantive change is that the elements are no arranged in order of atomic number rather than atomic mass. Credited to English physicist Henry Gwyn Jeffreys Moseley, more popularly Henry Moseley. Trends in the Periodic Table The periodic table is an Periodic Law – refers to the arrangement of the elements, periodic recurrence of certain by atomic number, in which physical and chemical elements with similar physical properties when the elements and chemical properties are are considered in terms of grouped together in vertical increasing atomic number. columns. Trends in the Periodic Table Group – vertical column of elements in the periodic table. Members of a group have similar properties. Group numbers at the top Period – horizontal row of the periodic table. All members of a period have atoms with the same highest principal quantum number. Arranged in order of increasing atomic number from left to right Numbered at the extreme left Trends in the Periodic Table Trends in the Periodic Table The first two groups – the s-block – and the last six groups – the p-block – together constitute the main-group elements. Group 1 is also known as alkali metals Group 2 is also known as alkaline earth metals Group 16 is also known as chalcogens Group 17 is also known as halogens Group 18 is also known as noble gases Trends in the Periodic Table The d-block elements are also known as the transition elements, they come between the s-block and p-block The f-block elements are also known as the inner transition elements, would extend the table if incorporated in the main body of the table so they are extracted from the table and placed at the bottom. The 15 elements following barium (Z = 56) are called lanthanides. The 15 elements following radon (Z = 88) are called actinides. Trends in Atomic Size We commonly represent atoms as spheres in which the electrons spend 90% of their time. However, we often define atomic size in terms of how closely one atom lies next to another. In practice, we measure the distance between identical, adjacent atomic nuclei in a sample of an element and divide that distance in half. Because atoms do not have hard surfaces, the size of an atom in a compound depends somewhat on the atoms near it. Trends in Atomic Size In other words, atomic size varies slightly from substance to substance. The metallic radius is one-half the distance between nuclei of adjacent atoms in a crystal of the element; we typically use this definition for metals. For elements commonly occurring as molecules, mostly nonmetals, we define atomic size by the covalent radius, one-half the distance between nuclei of identical covalently bonded atoms. Trends in Atomic Size Atomic size varies within both a group and a period. These variations in atomic size are the result of two opposing influences: 1. Changes in n – As the principal quantum number (n) increases, the probability that the outer electrons will spend more time farther from the nucleus increases as well; thus, atomic size is larger 2. Changes in Zeff – As the effective nuclear charge (Zeff) – the positive charge “felt” by an electron – increases, outer electrons are pulled closer to the nucleus; thus, the atoms are smaller. Trends in Atomic Size The net effect of these influences depends on shielding of the increasing nuclear charge by inner electrons: 1. Down a group, n dominates – Atomic size generally increases in a group from top to bottom 2. Across a period, Zeff dominates – Atomic size generally decreases in a period from left to right These trends hold well for the main-group elements but not as consistently for the transition elements. Trends in Atomic Size Trends in Atomic Size Using only the periodic table, rank each set of main-group elements in order of decreasing atomic size: a) Ca, Mg, Sr b) K, Ga, Ca c) Br, Rb, Kr d) Sr, Ca, Rb Trends in Ionization Energy Ionization energy (IE) – energy (in kJ) required for the complete removal of 1 mol of electrons from 1 mol of gaseous atoms or ions. Pulling an electron away from a nucleus requires energy to overcome the attraction thus, IE is usually positive. First ionization energy (IE1) – energy to remove an outermost (highest energy sublevel) from the gaseous atom. Atoms with low IE1 often form cations Atoms with high IE1 (except noble gas) form anions Second ionization energy (IE2) – removes a second electron. IE2 > IE1 Ionization energies decrease as atomic size increase. Trends in Ionization Energy Using only the periodic table, rank each set of main-group elements in order of decreasing ionization energy: a) Kr, He, Ar b) Sb, Te, Sn c) K, Ca, Rb d) I, Xe, Cs Trends in Electron Affinity Electron affinity – is the energy change (in kJ) accompanying the addition of 1 mol of electrons to 1 mol of gaseous atoms or ions. First electron affinity (EA1) – energy to form 1 mol of monovalent (1-) gaseous anions In most cases, energy is released when the first electron is added thus, EA1 is usually negative. Second electron affinity (EA2) – adds a second electron. EA2 is always positive because energy must be absorbed to overcome electrostatic repulsions and add another electron to a negative ion. Trends in Electron Affinity It is more difficult to make generalizations about EAs than about IEs. The smaller atoms to the right of the periodic table (Group 17) tend to have large, negative electron affinities. EAs tend to become less negative in progressing toward the bottom of a group, with the notable exception of the second-period members of Groups 15, 16, and 17 (N, O, and F). Some atoms have no tendency to gain an electron, such as the noble gases where electrons have to enter the next shell, and Groups 2 and 12 where the electrons have to enter an empty p subshell, etc. Trends in Electron Affinity Trends in Electronegativity Electronegativity (EN) – is a measure of the ability of an atom to attract electrons towards itself in the context of a chemical bond. 1. Across a period: electronegativity increases in a period from left to right Due to the stronger attraction that the atoms obtain as the nuclear charge increases. 2. Down a group: electronegativity decreases in a group from top to bottom Due to the longer distance between the nucleus and the valence electron shell, thereby decreasing the attraction. Electronegativity decrease as atomic size increase. Trends in Electronegativity Trends in Metallic Character Metal – an element whose atoms have small numbers of electrons in the outermost electronic shell. Removal of an electron(s) from a metal atom occurs without great difficulty, producing a positive ion (cation). Metals generally have a lustrous appearance, are malleable and ductile, and are able to conduct heat and electricity. Nonmetal – an element whose atoms tend to gain small numbers of electrons to form negative ions (anions) with the electron configuration of a noble gas. Nonmetal atoms may also alter their electron configurations by sharing electrons. Nonmetals are mostly gases, liquid (bromine), or low melting point solids and are very poor conductors of heat and electricity Metalloid – an element that may display both metallic and nonmetallic properties under the appropriate conditions. Trends in Metallic Character Metallic property 1. Across a period: metallic property decreases in a period from left to right Due to the increase in number of valence electrons, as well as decrease in atomic radius 2. Down a group: metallic property increases in a group from top to bottom Due to the increase in the number of shells and atomic radius. Metallic property increase as atomic size increase. Magnetic Properties An important property related to the electron configurations of atoms and ions is their behavior in a magnetic field. A spinning electron is an electric charge in motion, which induces a magnetic field. A diamagnetic substance has all its electrons paired and the individual magnetic effects cancel out. It is slightly repelled by a magnetic field. A paramagnetic substance has one or more unpaired electrons in its atoms or molecules and the individual magnetic effects do not cancel out. It is attracted into a magnetic field. Any questions?