PDF Periodic Table Classification

## 3.2. Classification of the Elements in Periodic Table ### 3.1. General Features of the Periodic Table The periodic table is a tabular arrangement of elements in order of increasing atomic number. When elements are thus arranged, there is a recurring pattern in their chemical and physical proper...

## 3.2. Classification of the Elements in Periodic Table ### 3.1. General Features of the Periodic Table The periodic table is a tabular arrangement of elements in order of increasing atomic number. When elements are thus arranged, there is a recurring pattern in their chemical and physical properties. This periodic repetition of the elements' properties is called the periodic law. The properties of elements are periodic functions of their atomic numbers. The periodic table consists of 18 groups and 7 periods: * **Groups**: Vertical columns * **Periods**: Horizontal rows Elements in the same group have similar chemical and physical properties because they have the same number of valence electrons. Elements in the same period possess the same energy level, and they have properties that change progressively across the table. ### 3.2. Classification of the Elements in Periodic Table Elements in the periodic table can be classified according to their natures, electronic configurations, and properties, as follows: **Based on nature** * **Metals** * **Non-Metals** * **Metalloides** **Based on Electronic Configuration** * **s-block** * **p-block** * **d-block** * **f-block** **Based on Properties** * **Representative elements** * **Noble gases** * **Transition elements** * **Inner transition elements** ### Classification based on the element's nature: The elements of the periodic table can be divided into three main categories: * **Metals** * Good conductors of heat and electricity * Lose electrons (1, 3) * e.g., K and Fe * **Non-Metals (5, 7)** * Poor conductors of heat and electricity * Gain or share electrons * e.g., Cl and O * **Metalloids (Semi-metals)** ( * They have properties that are intermediate between those of metals and non-metals * e.g., Si and Ge | 1A (1) | 2A (2) | 3B (3) | 4B (4) | 5B (5) | 6B (6) | 7B (7) | 8B (8) | 1B (9) | 2B (10) | 3A (13) | 4A (14) | 5A (15) | 6A (16) | 7A (17) | 8A (18) | |---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---| | Li | Be | | | | | | | | | B | C | N | O | F | He | | Na | Mg | Sc | Ti | V | Cr | Mn | Fe | Co | Ni | Cu | Zn | Ga | Ge | As | Se | Br | Kr | | K | Ca | Y | Zr | Nb | Mo | Tc | Ru | Rh | Pd | Ag | Cd | In | Sn | Sb | Te | I | Xe | | Rb | Sr | | | | | | | | | | | | | | | | | Cs | Ba | La | Hf | Ta | W | Re | Os | Ir | Pt | Au | Hg | Tl | Pb | Bi | Po | At | Rn | | Fr | Ra | Ac | Rf | Db | Sg | Bh | Hs | Mt | Ds | Rg | 112 | 113 | 114 | 115 | 116 | 117 | 118 | | | | Ce | Pr | Nd | Pm | Sm | Eu | Gd | Tb | Dy | Ho | Er | Tm | Yb | Lu | | | | | | Th | Pa | U | Np | Pu | Am | Cm | Bk | Cf | Es | Fm | Md | No | Lr | | | *Classification of periodic table elements according to their nature.* ### Note: * Metals constitute the largest class of elements of the periodic table, and all of them are solids, except mercury (Hg) is a liquid at room temperature. Metals are divided into the following two categories: * **Light metals** * Soft, have a low density, are very reactive chemically, and are unsatisfactory as structural materials. * They are located at the far left of the table (Groups IA and IIA) * e.g. Sodium (Na) and Magnesium (Mg). * **Transition metals** * Hard, have a high density, readily react, and are useful structural materials. * They are located in the middle of the table (the B groups) * e.g. Iron (Fe) and Titanium (Ti). * **Non-metals** are gases e.g., Oxygen, Nitrogen, and Chlorine (Cl), are solids e.g., Sulphur and Iodine (S, I), one is a fuming dark liquid: Bromine (Br). * **Metalloids (Semi-metals)** are semiconductors of electricity and widely used in electrical components, e.g., Silicon (Si) and Germanium (Ge). ### Classification based on the element's electronic configuration Elements in the periodic table can be sorted into 4 different groups (blocks) based on their electron configurations. * **s-block elements**: These elements have the valence shell configuration of *ns*. * e.g., Sodium 11Na: 1s² 2s² 2p⁶ 3s¹ * **p-block elements**: These elements have the valence shell configuration of *ns²np¹-⁶*. * e.g., Phosphorus ¹⁵P: 1s² 2s² 2p⁶ 3s² 3p³ * **d-block elements**: All these elements are metals and are characterized by electrons being added to *d* orbitals. * e.g., Iron ²⁶Fe: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶4s² * **f-block elements**: All these elements are metals and are characterized by electrons being added to *f* orbitals. They are two series: Lanthanides and Actinides. * e.g., Promethium ⁶¹Pm: [Xe] 4f⁵ 6s² _Classification of periodic table elements according to their electronic configurations._ ### Note: 1. Helium (He) is an s-block element. 2. All the elements of the s-block (except Hydrogen and Helium), d-block, and f-block are metals. 3. The p-block is the only block in the periodic table that contains metals, non-metals, and metalloids. For example, Tin (Sn) and Lead (Pb) are metals; Carbon, Oxygen, and Nitrogen (C, O, N), and Chlorine (Cl) are non-metals, while Silicon (Si) and germanium (Ge) are metalloids. ### Classification based on the element's properties: 1. Representative Elements 2. Noble Gases 3. Transition Elements 4. Inner Transition Elements ### I. Representative Elements (Main Group Elements): * They are the A group elements in the periodic table (also called main group elements). They are designated in Groups 1A to 7A. The number assigned to a specific group represents the number of valence electrons of the elements present in this group. For example, group 1A elements have one valence electron, while those of 7A possess seven valence electrons. * They have partly occupied the highest energy level. The last electron was added to *n*s or *n*p orbital. * They can be metals, nonmetals, or metalloids. * They can be solids, gases, or liquids. | 1A | 2A | 3A | 4A | 5A | 6A | 7A | 8A | |---|---|---|---|---|---|---|---| | A | A | A | A | A | A | A | A | _Elements in the 1A-7A groups are called the representative elements._ #### Some groups of elements & their family names <start_of_image>table: | Group | Family Name | Characteristics| |---|---|---| | 1A | Alkali Metals | *All end in ns¹* <br/> *All group 1A are metals except hydrogen.* <br/> *Very reactive metals.* <br/> *They easily lose one valence electron to form monovalent cations (M+).* <br/> *They react vigorously with water to form alkalis.* | | 2A | Alkaline Earth Metals | *All end in ns²* <br/> *Metals* <br/> *Very Reactive metals* <br/> *They lose two valence electrons to form divalent cations (M²+) * <br/> *They react less vigorously with water than group 1A.* | | 7A | Halogens | *All end in ns² np⁵* <br/> *Halogens means "salt-forming".* <br/> *Very reactive nonmetals.* <br/> *They easily gain an electron to form a monovalent anion (Χ)*. <br/> *They are not found as free elements in nature. Instead, they are found as halide ions in various minerals and in seawater.* | | | | |---|---| | Noble gases | **ns² np⁶** | | *8 Valence electrons* | #### II. Noble Gases: * **Group 8A** elements (also called group 18 or 0) * Their electron configuration: 1s² for helium and ns² np⁶ for others. * Their outer *s* and *p* electron configurations are filled. * Unreactive inert gases: because of their stable nature as their valence shell is completely filled with electrons. | | | |---|---| | He | helium | | Ne | neon | | Ar | argon | | Kr | krypton | | Xe | xenon | | Rn | radon | #### III. Inner Transition Metals: * Inner Transition Metals are known as f-block elements. * They are also called: rare-earth elements. * They are located below the main body of the table, in two horizontal rows. * Two series of *f* block: * a) Lanthanides: 4*f* subshell is being filled * b) Actinides: 5*f* subshell is being filled | | | |---|---| | Lanthanides | La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu | | Actinides | Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr | These are called the inner transition elements, and they belong here. ## 3.3. Atomic Properties ### 3.3.1. Sizes of Atoms and Ions (Atomic and Ionic Radii): The atomic radius (r) of an atom can be defined as half the distance (d) between the two nuclei of a diatomic molecule. #### Group Trends in Atomic Size: Due to the increase in the number of energy levels (shells), the atomic size (atomic radius) increases from top to down through a group, so that the atoms get bigger. #### Period Trends in Atomic Size: As the atomic number (Z) increases, electrons are added at the same energy level, and the nuclear charge increases. So, the attractive forces between the nucleus and the outermost electrons (valence electrons) will increase and the electrons are pulled closer to the nucleus. Therefore, atomic size decreases from left to right through a period, so that the atom gets smaller. #### Ionic Size The ionic radius is the size of an ion compared to its neutral atom. * **Cations (+)**: * Cations are smaller in size (ionic radius) than the atoms they came from. * Because the attraction force between electrons and the nucleus increases (as the nuclear charge increases), thus the size of the cation decreases. * **Anions (-)**: * Anions are bigger than the atoms they came from. * Because the number of electrons increases, the repulsion forces between electrons are increased and electrons spread out more, thus the size of the anion increases. ### 3.3.2. Ionization Energy (IE): Ionization energy (IE) is the minimum amount of energy required to remove the most loosely bound electron from an isolated gaseous atom. **A (g) → A⁺ (g) + e⁻** Below are some important points concerning IE: * Ionization energy is always endothermic, that is energy is added to the atom to remove the electron. * IE measured in kilojoules, (kJ) 1\. The energy required to remove only the first electron is called the **first ionization energy**. 2\. The **second ionization energy** is the energy required to remove the second electron. 3\. The **third ionization energy** is the energy required to remove a third electron. **Mg (g) → Mg⁺ (g) + e⁻, IE 1st = +738 kJ.** **Mg⁺ (g) → Mg²⁺ (g) + e⁻, IE 2nd = +1451 kJ.** **Mg²⁺ (g) → Mg³⁺ (g) + e⁻, IE 3rd = +7731 kJ.** **Give reasons:** Second IE is always greater than the first IE and third IE is always greater than the first or second IE. #### **Factors Affecting Ionization Energy:** * **a) Nuclear Charge:** * Nuclear charge increases * The attraction of the valence electrons increases * IE increases * **b) Shielding Effect:** * Greater distance from the nucleus decreases IE. * As more shells are added to atoms, the inner layers of electrons shield the outer electrons from the nucleus. #### **Group Trends in Ionization Energy:** As you go down a group, the first IE decreases because the electron is further away from the attraction of the nucleus, and there is more shielding. #### **Period Trends in Ionization Energy:** All the atoms in the same period have the same energy level, so same shielding. But, with increasing nuclear charge, Thus IE generally increases from left to right. **Note:** IE is inversely proportional to atomic size. ### 3.3.3. Electron Affinity (EA): Electron affinity (EA) is the energy change that occurs when an atom gains an electron. * EA is measured in kJ. * Electron affinity is usually **exothermic**, but not always. **F (g) + e⁻ → F⁻ (g) EA = -328 kJ** **F (1s²2s²2p⁵) + e⁻ → F⁻ (1s²2s²2p⁶)** #### Group Trends in Electron Affinity: As you go down a group, the EA generally decreases. The greater the distance, the less the attraction (due to the shielding effect) and so the less energy is released. #### Period Trends in Electron Affinity: EA generally increases from left to right in periods, as the nuclear charge increases and the filling of the valence shell of the atom becomes closer to the nucleus. **Note:** There are many exceptions ### 3.3.4. Electronegativity (EN): The electronegativity (EN) of an element is a measure of the relative tendency of an atom to attract electrons to itself when it is chemically combined with another element. An element with a high electronegativity means it pulls the electron towards itself strongly when it is chemically combined with another element. #### Group Trend in Electronegativity: From top to down in a group, the EN decreases. Because the atomic size increases, the electron is further away from the attraction of the nucleus and there is more shielding. #### Period Trend in Electronegativity: The electronegativity increases across a horizontal row of the periodic table from left to right as nuclear charge increases so the attraction of electrons becomes stronger. The table below illustrates the **Pauling scale**, which shows the electronegativities of the periodic table elements. | 1A | 2A | 3A | 4A | 5A | 6A | 7A | 8A | 3B | 4B | 5B | 6B | 7B | 1B | 2B | |---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---| | H 2.1 | Li 1.0 | Be 1.5 | | | | | | B 2.0 | C 2.5 | N 3.0 | O 3.5 | F 4.0 | | | | Na 0.9 | Mg 1.2 | | Sc 1.3 | Ti 1.5 | V 1.6 | Cr 1.6 | Mn 1.5 | Fe 1.8 | Co 1.8 | Ni 1.8 | Cu 1.9 | Zn 1.6 | Ga 1.6 | Ge 1.8 | | K 0.8 | Ca 1.0 | | Y 1.2 | Zr 1.4 | Nb 1.6 | Mo 1.8 | Tc 1.9 | Ru 2.2 | Rh 2.2 | Pd 2.2 | Ag 1.9 | Cd 1.7 | In 1.7 | Sn 1.8 | | Rb 0.8 | Sr 1.0 | | La 1.1 | Hf 1.3 | Ta 1.5 | W 1.7 | Re 1.9 | Os 2.2 | Ir 2.2 | Pt 2.2 | Au 2.4 | Hg 1.9 | Tl 1.8 | Pb 1.9 | | Cs 0.7 | Ba 0.9 | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | * The electronegativity of Fluorine (4.0) is higher than that of any other element. * Oxygen (3.5) is the second most electronegative element. * Nitrogen and Chlorine (3.0) are the third most electronegative elements. #### Summary of Periodic Properties & Trends: | Property | Trend | |---|---| |Atomic Radii | Increase down a group, decrease across a period | |Ionic Radii (Cations) | Decrease down a group, decrease across a period | |Ionic Radii (Anions) | Increase down a group, increase across a period | |Electron Affinity | Generally decreases down a group, increases across a period | |Ionization Energy | Generally increases across a period, decreases down a group | |Electronegativity | Generally increases across a period, decrease down a group | ## Practice Questions: 1. The correct order of radius of the following is ... * a) Li < Rb <K<Na * b) Li<Na < Rb< K * c) Na<K<Rb<Li * d) Li<Na<K<Rb 2. Which of the shown groups is commonly known as noble gases? * a) Group 1A * b) Group 2A * c) Group 7A * d) Group 8A 3. Which of the following is a halogen? * a) S * b) Ca * c) Fe * d) Cl 4. What is the trend of atomic size from top to down in a group of the periodic table? * a) Increases * b) Decreases * c) Remains constant * d) Randomly changes 5. Which elements have electron configurations that end in *ns² np⁶*? * a) Noble gases * b) Transition metals * c) Halogens * d) Inner transition metals What is true about all the elements in the same vertical column in the periodic table? * a) They possess the same number of energy levels * b) They have the same number of valence electrons * c) They have the same atomic size * d) They have the same electronegativity