IB Chemistry Module 1.3: The Periodic Table - PDF

PC306.301 IB Chemistry I Module 1. The Atom as the Professor: Enrique Rodríguez, MSc. Fundamental Particle of Matter  Module 1.3 The Periodic Table –...

PC306.301 IB Chemistry I Module 1. The Atom as the Professor: Enrique Rodríguez, MSc. Fundamental Particle of Matter  Module 1.3 The Periodic Table – classification of elements Professor: Enrique Rodríguez, MSc. PC306.301 IB Chemistry I 1.3.1 A history of organizing matter Professor: Enrique Rodríguez, MSc. 7 It all started with Alchemy… The history of the periodic table is also the history of Chemistry and the discovery of elements. During the Middle Ages, where mysticism and Science were still the same thing, working with substances was known as Alchemy. Alchemists dedicated their work to study, purify and perfect materials and substances through esoteric, pseudoscientific methods. In his search for the philosopher’s stone in 1669 (which was supposed to turn every material into gold), Henning Brand distilled and purified human urine until he got an incandescent white liquid. This is the first ever recorded discovery of a new element: phosphorus (P). 2 The first approach More than 30 new element discoveries followed in the next 10 years, rising the need to enlist and classify them to facilitate the Alchemist’s job. a. Etienne-François Geoffroy (1718) He organized the elements in a tabular format, where he grouped them according to their affinity. His proposal is called the Affinity Table. His approach, while still considered Alchemical, first exhibited the systematic nature of modern Chemistry. 2 The birth of modern Chemistry b. Antoine Lavoisier (1789) He published the Elementary Treatise of Chemistry, the first textbook that gathered scientific notions of Alchemy, grouping it into a formal science he called Chemistry. Among the many contributions and revelations contained in the Treatise, Lavoisier presented a list of the known elements and classified into metals and nonmetals. Lavoisier’s list included Light and Energy, which were thought to be substances at that time. 2 Addition of Atomic Weights c. John Dalton (1810) Agreeing with Lavoisier’s conjectures, John Dalton defined a way to determine the atomic weights of elements and rearranged Lavoisier’s list in increasing order of mass. Dalton formally removed light and energy as elements. d. William Prout (1815) Noticed that all atomic weights seemed to be multiples of Hydrogen. Modified Dalton’s table by simplifying H weight to 1 unit. 2 Trends instead of lists e. Johann Wolfgang Döbereiner (1817) He was the first person to try to classify the elements (instead of just enlisting them). He classified elements in groups of three, called triads. The came from the observation that Strontium (Sr) had intermediate properties than those of Calcium (Ca) and Berilium (Be)… and also an intermediate weight. He then organized several elements in triads, where the middle element had intermediate weight, physical and chemical properties with respect to the other two. 2 Döbereiner’s triads still prevail 2 Trends instead of lists f. John Newlands (1864) Newlands expanded Döbereiner’s observations and proposed the Law of Octaves, in which he proposed that elements group according to their properties in eight groups of eight. Just like octaves in music, Newlands proposed that the properties within a group progressed following a trend along each group and within groups (periodicity!). Despite the genius of his ideas, he was ridiculed by colleagues:( 2 The final form g. Dimitri Mendeleev (1869) Mendeleev obsessively played “chemical solitaire”, trying to arrange the elements in a logical order of weight and properties. Newlands’ octaves did make sense but couldn’t explain properties for all known (until that point) 63 elements. Mendeleev finally came to the conclusion that the discrepancies existed not because the system was wrong, but because many elements were not yet discovered. Mendeleev then proposed a new table organized in columns (groups) and rows (periods), where the elements were ordered according to atomic mass and classified according to chemical and physical properties. 2  Monument to Dimitri Mendeleev The final form St. Petersburg Institute of Technology (RUS) g. Dimitri Mendeleev (1869) In this new table, Mendeleev left empty spaces in the groups for elements that were not yet discovered. Using the periodic trends, Mendeleev was able to successfully predict the masses and properties of the missing elements. Almost like magic, newly discovered elements started filling and fitting Mendeleev’s table, giving raise to our current classification system: the periodic table of elements. Mendeleev’s final table (1910) 2 Mendeleev’s organization (1910) 2 Modern organization 2 …only 118 elements? 2 1.3.2 Properties of the periodic table Professor: Enrique Rodríguez, MSc. 7 Structure of the modern periodic table Inherited from Mendeleev’s model, modern periodic tables are organized in 18 groups and 7 periods: 2 Metals, nonmetals and metalloids All the known elements can be classified according to their properties (and therefore, their position in the periodic table) into one of these categories: 2 Metals, nonmetals and metalloids All the known elements can be classified according to their properties (and therefore, their position in the periodic table) into one of three categories: a. The metals Located at the left side of the periodic table. Malleable Shiny High melting points Heat and electricity conductors Tend to lose electrons (to form cations). Gold (Au) Bismuth (Bi) 2 Metals, nonmetals and metalloids All the known elements can be classified according to their properties (and therefore, their position in the periodic table) into one of three categories: b. The nonmetals Located at the right side of the table. Typically brittle. Dull. Lower melting points. Low thermal and electrical conductivity. Tend to gain electrons (to form anions). Bromine (Br) Carbon (C) 2 Metals, nonmetals and metalloids All the known elements can be classified according to their properties (and therefore, their position in the periodic table) into one of three categories: c. The metalloids Located in the boundary between metals and nonmetals, showing intermediate properties. Lustrous or dull. High melting points. Intermediate thermal and electrical conductivity. Some disagreement prevails about the metalloid character of Po and At. Silicon (Si) Tellurium (Te) 2 The full picture Usually, there are two series of elements cut out from the periodic table and placed at the bottom. They are known as the Lanthanide (La†) and actinide (Ac‡) series. The separation of these series from the rest of the periodic table is related to their last oribitals filled in their electronic configuration: all La† and Ac‡ have their last electrons in the f orbital. 2 Blocks in the periodic table Elements are divided in blocks according to which orbital shape (quantum number l) is occupied by the outermost electrons: 2 Writing electron configurations… Blocks can be used to write the electron configurations of elements: Electron configuration for Cl: Electron configuration for Mo: 2 Learning check Assign each configuration to a place in the periodic table: 2 Periodic trends a. Effective Nuclear Charge This refers to the net positive charge (attraction) experienced by valence electrons of an atom. Is the combination of the attractive force of the protons and the repulsive forces of core electrons. Higher effective nuclear charge = higher forces of attraction between nucleus and electrons. Effective nuclear charge:  Increases as number of protons increases.  Increases as number of valence electrons increases. 2 Periodic trends b. Atomic Radius The Radius of a sphere refers to the distance between its center and the outer edge. However, for quantum mechanics, the “edge” is not well defined, as valence electrons follow variable trajectories. Therefore, atomic radius is measured as the Wan der Waals distance: the distance between the nuclei of two adjacent atoms. Why do we care? For many applications, atoms can be modeled as spheres. Therefore, it is important to know the radii of these spheres. Radius for most atoms lies between 30 and 300 pm (one trillionth of a meter), aka 3 Å (ångströms). 1 Å = 1x10-10 m. 2 Periodic trends * Down a group: b. Atomic Radius * Across a period: 2 Atomic Radius: overall trends 2 Periodic trends * Ionic Radius The number of electrons spinning around the nucleus determines ionic radius. Therefore, ionization of elements changes their ionic radii. a. Positive ions Electrons removed from the parent atom: = higher effective nuclear charge. = increased attractions between nucleus and electrons. = reduced atomic radius. 2 Periodic trends * Ionic Radius The number of electrons spinning around the nucleus determines ionic radius. Therefore, ionization of elements changes their ionic radii. b. Negative ions Electrons added to the parent atom: = lower effective nuclear charge. = Decreased attractions between nucleus and electrons. = Increased atomic radius. 2 Periodic trends * Isoelectronic species For isoelectronic species (atoms and ions with the same number of electrons), atomic radius is defined in terms of effective nuclear charge. Higher ENC = smaller radius. 2 Learning check Assign each isoelectronic species to a place in the table. Respect the color schemes. 2 Periodic trends e. Ionization energy The energy needed to remove 1 electron from an atom’s outermost orbital (Valence electrons). In terms of spontaneous processes, this term is specific to metals, who tend to lose electrons. 2  ? Relationship: atomic radius and ionization energy Group trend of atomic radius 2 Periodic trends c. Electron affinity Is the opposite to ionization energy. Electron affinity refers to the energy released when an electron is added to an atom to form an anion. This energy is released as a result of the force of attraction needed to “secure" the electron in place. 2  ? Periodic trends Group trend of c. Electron affinity atomic radius 2 Periodic trends d. Electronegativity Electronegativity refers to the capacity of an atom to attract an electron pair while forming a covalent bond with another atom. 2 Periodic trends d. Electronegativity The periodicity is as follows: 2 Periodic trends e. Metallic and non-metallic character * Metallic character Is the tendency of an atom to lose electrons and form a cation. Lower ionization energy = higher metallic character. * Non-metallic character Is the tendency of an atom to gain electrons and form an anion. Greater electron affinity = higher non-metallic character. 2 Periodic trends 2 1.3.3 Group 1 and 17 elements Professor: Enrique Rodríguez, MSc. 7 Group 1: the Alkali metals The word “alkali” comes from the Arabic al-qalïy, which means “burnt ashes”. This is because potash (K2CO3), an alkali, was first discovered in burnt plant ashes. These metals are often also called the “Lithium series”.Lithium (Li) Properties of Group 1 metals: They are very soft metals. Sodium (Na) They have one valence electron in an s Alkali metals orbital (s1). They are highly reactive – eager to lose their electron and become cations. Potassium (K) They are especially reactive with water and halogens, in consequence, they are never found in nature as pure solids, rather as their resulting compounds: salts or oxides. 2 Reactivity of Group 1 with water Eager to lose electrons, all alkali metals are highly reactive with water: Reactivity trends: Reactivity increases as ionization energy decreases. Lower ionization energies means it’s easier for electrons to be removed. The reaction is highly exothermic and violent. Because of this reactivity, pure alkalis are stored submerged in oils, as there is water in air. 2 Reactivity of Group 1 with water 2 Group 17: the Halogens The word Halogen comes from the Greek roots hals and gene, which mean “salt maker”. These elements are also highly reactive, and they form salts with metals. Properties of Group 17 nonmetals: These elements are highly electronegative. They also have high electron affinity. Because of this, they are very reactive and can be harmful to life. They exist in all three states of matter: F and Cl are gases; Br is a liquid and I,At and Ts are solids (at NaCl (table salt) standard P and T conditions). Chlorine gas (Cl2) 2