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  Topic 3: Periodicity 3.1 Periodic Table 3.1.1 The periodic table is arranged into four blocks associated with the four sub-levels  –  s, p, d, and f 3.1.2 The periodic table consists of groups (vertical columns) and periods (horizontal rows) 3.1.3 The period number (n) is the outer energy level that is occupied by electrons 3.1.4 The number of the principal energy level and the number of the valence electrons in an atom can be deduced from its position on the periodic table 3.1.5 The periodic table shows the position of metals, non-metals and metalloids 3.1.6 Deduction of the electron configuration of an atom from the element’s position on the periodic table, and vice versa   Periodic Table Arrangement    The periodic table is arranged into 4 blocks associated with their last orbital: s, p, d and f    Elements in the periodic table are arranged by increasing atomic number    Group:  A vertical column of elements o   Group number is the same as the number of valence electrons    Period:  A horizontal row of elements o   Period number is the same number of shells in the atom   o   All shells except the outer shell will be full      There is a division between metals and non-metals. Metals are on the left and non-metals are on the right    The farther sub-level occupied is found by the location of the element (which block it is in):      The location of the following families of the periodic table must be known: alkali metals, alkali earth metals, halogens, noble gases, transition metals, lanthanides and actinoids 3.2 Periodic trends 3.2.1 Vertical and horizontal trends in the periodic table exist for atomic radius, ionic radius 3.2.2 Trends in metallic and non-metallic behavior are due to the trends above 3.2.3 Oxides change from basic through amphoteric to acidic across a period 3.2.4 Prediction and explanation of the metallic and non-metallic behavior of an element based on its position in the periodic table 3.2.5 Discussion of the similarities and differences in the properties of elements in the same group, with reference to alkali metals and halogens    Trends in the Periodic Table    Periodicity refers to repeating trends or patterns of physical and chemical properties in elements  Atomic Radius    Atomic radius increases down a group  as the number of electron shells increases    Atomic radius decreases across a period o   This is because electrons are added to the same main energy level (n=3) the nuclear charge also increases o   The attraction between the nucleus and the outer electrons increases resulting in a smaller radius Ionic Radius    Ionic radius of positive ions decrease across a period  as the number of protons in the nucleus increases but the number of electrons remain the same      Ionic radius of negative ions decrease across a period  as the number of protons in the nucleus increases but the number of electrons remains the same   Ionization energy    Ionization energy increases across a period      The increase in nuclear charge across a period causes an increase in the attraction between the outer electrons and the nucleus makes the electrons more difficult to remove    Ionization energy decreases down a group      The electron being removed is from the energy level furthest from the nucleus so it gets easier to remove valence electrons as atomic radius increases down a group o   Valence Electrons: The outermost electrons of an atom   Electron affinity    Generally, metals have a low EA and non-metals have a higher EA      The greater the distance between the nucleus and the outer energy level, the weaker the electrostatic attraction and the less energy is released when an electron is added to the atom Electronegativity    Metals have low electronegativities because they lose electrons easily    Non-metals have high electronegativities as they gain electrons to complete their outer shell    Generally, electronegativity tends to increase across a period and decrease down a group Definitions Atomic Radius    –  The distance from the nucleus to the outermost electron Ionic Radius    –  The distance from the nucleus to the outermost electrons in an ion First ionization energy    –   The energy required to remove one mole of electrons from one mole of gaseous atoms. This is shown by:   ()  →  ()+  −   Electron affinity  –   The energy released when one mole of an electron is added to one mole gaseous atoms. This is shown by:   ()   − →  ()−  Electronegativity  –  A measure of the attraction an atom has for a shared pair of electrons in a covalent bond  Melting points    Melting points generally increase across a period until group 14, then they decrease      This depends on the type of bonding (covalent, ionic or metallic), structure (ionic lattice, molecular covalent, giant covalent, or metallic structures), and strength of metallic bond Metallic Character    Metallic character: How easily an atom can lose electrons    Metallic character is displayed by metals, which are all on the left-hand side of the periodic table including alkali metals, alkalie earth metals, tarnsition metals the lanthanide and actinides, and the basic metals    Metallic character increases down a group    Metallic character decreases across a period   Trends Across: Group 1  –  The Alkali Metals    Atomic/Ionic radius increases  down the group as there are more electron shells    First ionization energy decreases  down the group as the valence electron is further from the nucleus so its easier to remove    Electronegativity decreases  because of increased distance and shielding    Melting points decrease  as atoms become larger and therefore metallic bonds becomes weaker    Reactivity increases down the group  as the valence electron is easier to lose, due to shielding Trends Across: Group 7  –  Halogens    Same trends as alkali metals except melting points increase  as Van der Waal forces becomes greater with more electrons      Reactivity decreases down group as with each consecutive element the outer shell gets further from the nucleus. So the attraction between the nucleus and electrons gets weaker, so an electron is less easily gained.   Trends Across: period 3    In period 3, sodium, magnesium and aluminum are metals located on the left of period 3 o   Metals are shiny solids that are excellent thermal and electrical conductors. They are ductile and malleable o   Metals are reducing agents and form cations. Their oxides and hydroxides behave like bases and neutralize acids    Aluminum oxide, however is insoluble in water and amphoteric    Phosphorus, sulfur, chlorine and argon are non-metals located on the right of period 3 o   Non-metals can be solids, liquids are gas o   Non-metals tend to be oxidizing agents and form anions. Their oxides tend to be acidic and are neutralized by bases    All period three oxides will react with water to form either an acidic or alkali (basic) solution: Group 1 2 13 14 15 16 17 18 Element Na Mg Al Si P S Cl Ar Structure of element Giant Metallic Giant Covalent Simple Covalent Structure of Oxide Formula of oxide Na 2 O MgO Al 2 O 3 SiO 2  P 4 O 10  SO 3  Cl 2 O 7  P 4 O 6  SO 2  Cl 2 O Acid-Base character of oxide Basic Amphoteric Acidic No oxide    Transition metals    Transition metals: Elements whose atoms have   incomplete d-orbitals  or can form positive ions with an incomplete d sub-level o    Zinc is not considered a transition element as it does not have an incomplete d orbital    Transition metals have variable oxidation states, display catalytic and magnetic properties,  form complex ions with ligands and have colored compounds (13.2)   Oxidation states    All the d-block elements (except scandium and zinc) can exist in more than one oxidation state    There is no increase in successive ionization energies in transition metals. There is only a slightly higher effective nuclear charge so there are no great jumps in successive ionization energies. So, ions can lose varying numbers of electrons with the same relative ease, depending on the oxidative nature of the surrounding environment.    All transition elements can form ions with an oxidation number of +2 (by losing the 2e -  from the 4s shell). So, transition elements show an oxidation state of +2 when the s-electrons are removed    In addition, each element can form a number of ions with different oxidation numbers    In order to form ions: o   Electrons in the 4s subshell have less energy than the 3d subshell o   Ions are formed as the electrons are lost fr om the 4s and 3d subshell’s   o   Electrons are removed from the valence shell (4s) before the 3d subshell Properties of transition metals 13.1 First row d-block elements 13.1.1   Transition elements have variable oxidation states, form complex ions with ligands, have colored compounds, and display catalytic and magnetic properties   13.1.2 Zn is not considered to be a transition element as it does not form ions with incomplete d-orbitals 13.1.3 Transition elements show an oxidation state of +2 when the s-electrons are removed 13.1.4 Explanation of the ability of transition metals to form variable oxidation states from successive ionization energies 13.1.5 Explanation of the nature of the coordinate bond within a complex ion 13.1.6 Deduction of the total charge given the formula of the ion and ligands present 13.1.7 Explanation of the magnetic properties in transition metals in terms of unpaired electrons Physical properties   Chemical properties      High electrical and thermal conductivity    High melting point    High tensile strength    Malleable (Can be beaten into shape)    Ductile (Can be drawn into wires)    Variable oxidation states    Formation of complex ions    Colored compounds    Catalytic behavior o   Fe in the Haber process o   V 2 O 5  in the Contact process o   Fe 2+  
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