Periodicity Objectives 1 Explain the reason behind the

Periodicity

Objectives: 1. Explain the reason behind the arrangement of elements of the periodic table. 2. Show the link between atomic structure, oxidation state and the position of elements. 3. Predic the properties of unknown elements. 4. Identify trends of Group II, VII. 5. Identify trends of Group 3, specifically metallic and non-metallic characteristics.

The Periodic table: There are 118 elements of the Periodic table which are arranged by increasing atomic number. These elements are presented as vertical columns or Groups and horizontal columns or rows.

Groups contain the same number of outer shell electrons Be 2, 2 Mg 2, 8, 2 Ca 2, 8, 8, 2 Sr 2, 8, 18, 8, 2 Ba 2, 8, 18, 8, 2 All elements of the same Group (vertical column) have the same outer shell configuration.

Groups contain the same number of outer shell electrons Elements with the same number of outer shell electrons have similar properties and behaviour. E. g all elements of group 1 are highly reactive metals. Elements of group 7 are highly reactive non-metals.

Elements of the same group have similar properties Lithium Sodium Potassium Rubidium All of these elements are highly reactive metals.

Trends of Group IIA: 1. Atomic radius 2. Ionization energy 3. Melting Point 4. Boiling Point 5. Relative density

Group 2 trends: 1. Atomic Radius Atoms have electrons arranged as rings or levels at fixed distances from the nucleus. Each ring or shell holds a specific number of electrons and the greater the numbers of electrons, the more occupied shells there and the larger the radius of the atom. -Atomic radius increases down the group. And decreases left to right across the period.

2. Ionization energy: Decreases down the group Atoms consist of a nucleus whose positive charge holds negatively charged electrons at specific levels or shells around the nucleus much like the moon orbiting the earth. The farther from the nucleus an electron is the weaker the attractive force exerted by the nucleus upon that electron and the less work or energy it would take to liberate that electron from the atom. The amount of energy it takes to remove an electron from an atom is called ionization energy. As the atoms become bigger lower down the group their outer electrons move farther from the nucleus and the ionization energy of that element decreases.

3. Melting Point & Boiling point: Increases Up the group To melt a substance the forces pulling one molecule or atom from another Should overcome the attractive forces (bonds) holding the molecules together. As atoms become larger they possess increasing amounts of Protons and the attraction between the nucleus of one atom and the shells of electrons around neighbouring atoms becomes stronger and so elements with high atomic numbers are usually liquids and solids at room temperature. This also applies to boiling. -Melting and Boiling point increases Up the group.

3. Melting Point & Boiling point: decreases down the group

4. Density: Increases down the group Be Ra ` *electrons are not shown. Moving from top to bottom along group 2 represents and increase of atomic number and mass of the elements, the atom of each element having a bigger nucleus with more protons and neutrons than the element above it. The amount of matter that each atom contains increases going down the group. Materials made up elements lower down the group will contain more mass pe unit volume than members of the group higher up, as each of their atoms contains Why does an increase of more protons and neutrons. E. g cubic electrons not 1 count meter of Beryllium weighs tonnes towards 1. 85 the density of anwhile element? 1 cubic metre of Radon weighs 5. 5 tonnes.

Introduction to Group VIIA: The halogens Group Fluorine Chlorine Bromine Iodine The members of group 7 are a highly reactive class Members of group 7 lower down the of group solids, the while. Greek those such as Bromine near elements called halogens, are from “salt former”. the middle are liquid and Fluorine & Chlorine Due to their reactivity these areatnever found nearelements the top are gases room temperature. Explain this as graded change of statewith from naturally as pure deposits but exist compounds solids to gas? other elements or as ions. Members higher up on the group form Diatomic gases at room temperature, while Bromine near the middle is liquid and Iodine near to the bottom is a solid as room temperature.

Trends of Group VII: 1. Atomic radius 2. Ionization energy 3. Melting Point 4. Boiling Point 5. Density

1. Atomic radius Similar to that of Group 2, atomic radius increases down the group. Fluorine is the smallest member of Group 7 with Iodine and Astatine being the largest.

2. Ionization energy decreases going down the group from Fluorine to Iodine and Astatine. The trend of ionization energy is also similar to that of Group II. However the ionization energies of Group 7 are much larger than those of group II. This is because ionization energy is a measure of the amount of energy it would take to remove electrons from an atom. However members of group 7 have outer shells with 7 electrons it is far easier to lose 1 or 2 electrons than it is to lose 7 and so more energy or work is done to remove an electron from a group 7 element. However it is still easier to remove the electrons from larger atoms than from smaller ones and so the general trend remains the same: Ionization energy decreases going down the group from Fluorine to Iodine and Astatine. Element: Ionization Energy: Fluorine 1690 Chlorine 1260 Bromine 1150 Iodine 1020

Melting and Boiling point: Group 7 shows identical trends to group 2 with melting and boiling points increasing going down the group. Boiling point and melting point increase down the group Members of group 7 lower down the group are solids, while those such as Bromine near the middle are liquid and Fluorine & Chlorine near the top are gases at room temperature. Explain this graded change of state from solid to gas?

Density: Decreases down the group Identical to that of Group 2 Density of materials made of group 7 elements increases with elements lower down the group than those above. • Density increases down the group.

Introduction to Period 3 Trends of period 3: Period 3 contains the three different types of elements: metals, metalloids and non-metals and refers to the elements left to right from Sodium to argon.

Trends of period 3: 1. Atomic radius 2. Electronegativity 3. Metallic properties 4. Reactivity 5. Bonding 6. Melting point 7. Boiling point 8. Density

1. Atomic radius: Each element of period 3 has one more proton and one more electron than the one before it. Each extra electron going to the outer shell of the element while the proton resides at the nucleus. As one goes across the period the nucleus experiences an increase of positive charge while it’s electrons are still at the same fixed distance away. This increase of nuclear positive charge without an increase of distance between the nucleus and electrons causes the atomic radius to shrink as one goes across the period. T’s as if the nucleus is pulling the atom together harder because of the extra “pulling” force of the protons upon the orbiting electrons. So the elements at the beginning of period 3 are the largest while those at the end are the smallest.

2. Electronegativity: Electronegativity is the ability of an element to attract electrons to itself and is a balance between 2 factors: 1. The size of the nucleus and 2. The number of occupied electron shells. The positive charge from the protons of the nucleus of each element emanates outwards to attract negatively charged bodies. This force however is decreased by 1. The distance over which it acts; the longer the distance between the nucleus and the electron the weaker the attraction. 2. The Number of occupied shells that the atom has. Each electron orbiting the atom is capable of absorbing the positive attractive force produced by the nucleus. Much of the attraction is able to propagate to the outside of the atom but where ever an electron intersects with the positive rays from the nucleus the positive attractive force is absorbed. . An effect referred to as “shielding” as the occupied shells serve to shield outside electrons from the pull of the nucleus. Very large atoms with large distance between electrons outside the nucleus and many occupied shells of electrons absorbing positive force from the nucleus are so well shielded that they often will not accept an electron unless forced. Additional electrons do not affect the positive forces coming from the nucleus as much as the addition of entire shells.

2. Electronegativity: electron Both Chlorine and Sodium have the same numbers of shells, smaller atomic radius. However the Chlorine atom is more electronegative element due to 1. Larger Nucleus with more protons. Electronegativity is a balance between the pull of positively charged Protons, Shielding effect of electrons and the distance between the nucleus and electrons outside.

2. Electronegativity: Periodic table showing electronegativity values of all elements(Pauling scale). Electronegativity increases up the group and left to right across the period. Fluorine is the most electronegative element.

3. Metallic properties All metals react by losing electrons to produce positive ions, the more readily an ion loses electrons, the more metallic it's nature. The first 3 elements of period 3 consist of sodium, magnesium and Aluminum, which have 1, 2 and 3 valence electrons respectively. it is easier to lose 1, 2 or 3 electrons than to gain 7, 6 or 5 electrons so these elements react by losing electrons. This behaviour means that these elements are true metals. Silicon which has 4 valence electrons ad neither loses nor gains but shares electrons as covalent bonds. Phosphorous , Sulphur, Silicon and Chlorine readily gain electrons instead of losing them and are the least metallic elements of period 3. Metallic nature decreases from right to left along the period 3.

4. Reactivity The reactivity of an element is defined as how readily an element will partake of chemical reactions. Chemical reactions take place when electrons are lost, gained or exchanged. Metals react by losing electrons, so metals which have a very weak hold on the outer electrons react readily. Nonmetals react by gaining electrons, so non-metals whose nuclei have a very strong attraction to electrons are very reactive. Several factors affect the reactivity of an element 1. Electronegativity 2. Atomic radius 3. Chemical nature: whether metal or non metal. Metals with 1 valence electron lose their electrons to chemical reactions more easily than those with 2 or 3 electrons, so reactivity of metals increases from right to left. Reactivity of non-metals also varies based on the electronegativity of the element, the more however the reactivity increase from left to right instead.

4. Reactivity: • Reactivity of metals increase from right to left across period 3, sodium being the most reactive. • Reactivity of non-metals increases from right to left across period 3, Chlorine being the most reactive. Reactivity of metals Reactivity of Non-metals

4. Reactivity: Question: Sodium explodes on contact with air and is stored under mineral (paraffin) oil, while magnesium and aluminium can be stored safely using an ordinary reagent bottle. Why? Question: Fluorine is so reactive that when stored with glass containers it begins to react with the silicon of the vessel walls. explain Fluorine's highly reactive nature.

5. Bonding: Na 2, 8, 1 Mg 2, 8, 2 Al 2, 8, 3 Si 2, 8, 4 P 2, 8, 5 S 2, 8, 6 F 2, 8, 7 Ar 2, 8, 8 Metals react by losing electrons and are said to be Electropositive. Metals of group 1 and 2 lose electrons to form ions with a charge of +1 or +2. When an element bonds with metals of group 1 or 2 it accepts the electrons which the metal loses and it itself becomes negatively charged. The oxidation state of an element is equal to it’s charge so elements which are bonded to group 1 or 2 metals usually have negative oxidation states. When non-metals of group 6 or 7 bond react they gain electrons and become negatively charged taking on oxidation states of -2 or -1 respectively. When 2 non-metals bond the more electronegative partner carries the electrons and has a negative oxidation state, while the less electronegative partner has a positive oxidation state. The exception is Fluorine which always has a negative oxidation state of -1. Oxygen also always has a negative oxidation state of -2 except when bonded to fluorine. Elements of Group 4 and 5 may have positive or negative oxidation states according to what it they are bonded to. When bonded to elements less electronegative than themselves, e. g metals, they have negative oxidation states as they accept the electrons from the metals. When bonded to more electronegative elements, e. g group 6 and 7 non-metals, they lose electrons and have a positive oxidation state.

5. Bonding: Question: No matter which element is bonded to fluorine, fluorine will always have a negative oxidation state, while it’s partner will take on a positive oxidation state. Why?

6. Melting Point & 7. Boiling Point: Na Mg Al Si P S Cl Ar Melting Point 98 650 660 1423 44 117 -101 -189 Boiling Point 890 1120 2447 2680 280 445 -34 -186 Both the melting and Boiling point increase to a maximum at silicon and then decrease with the exception of Sulphur. Silicon has 4 valence electrons, why is it the element with the highest melting and boiling point?

8: Density (gcm-3) Na Mg Al 0. 97 1. 74 2. 70 Si P S 2. 33 1. 82 1. 96 Cl Ar 2. 99 x 10 -3 1. 66 x 10 -3 Density follows the same trend as melting and boiling points. Increasing up to silicon and then decreasing to argon, with the exception of sulphur.

Physical Properties of period 3 elements: Element Atomic Number Description Sodium (Na) 11 Soft metal; low density; very/reactive Magnesium (Mg) Aluminum (Al) Silicon (Si) 12 14 Harder metal than sodium; low density; less reactive than sodium Metal as hard as magnesium; low density less reactive than magnesium Brittle non metal; not very reactive Phosphorus (P) Sulphur (S) 15 Non-metal; low melting point; white solid; reactive 16 Non-metal; low melting point; yellow solid; moderately reactive Chlorine (Cl) 17 Non-metal; green gas; extremely reactive. Argon (Ar) 18 Non-metal; colorless gas, unreactive 13

Oxides of Group 3 Elements: When dissolved with water the oxides of group 1 and group 2 metals produce Bases, Aluminium oxide is Amphoteric, and when dissolved reacts with both acids and bases. Silicon oxide reacts with bases at high temperature. The oxides of Phosphorous are acidic and react with bases to form silicates. All sulphur oxides are acidic , Sulphuric acid is a solution of Sulphur trioxide. Chlorine Oxides are very rare. Sodium Oxide - Basic Magnesium Oxide - Basic Aluminium Oxide -Amphoteric (reacts with both acids and bases) Silicon Oxide - Basic ( reacts only at high temperature) Phosphorus Oxide - Acidic Sulphur Oxide - Definitively acidic

Chlorides of some group 3 elements: Sodium Chloride and Magnesium Chloride are ionic compounds, when molten or dissolved with water they conduct electricity. All other Chlorides are Co-valent.