P Block Elements for STD-XI Chemistry 2024-25 (Goa Board)

# The P Block Elements for STD-XI-Chemistry 2024-25 (As per Goa board syllabus 2024-25) ## UNIT: P-BLOCK ELEMENTS **Group -13 to Group 17 Elements** (10 Periods) General Introduction: Electronic configuration and general trends in physical and chemical properties of elements down the groups (ox...

# The P Block Elements for STD-XI-Chemistry 2024-25 (As per Goa board syllabus 2024-25) ## UNIT: P-BLOCK ELEMENTS **Group -13 to Group 17 Elements** (10 Periods) General Introduction: Electronic configuration and general trends in physical and chemical properties of elements down the groups (oxidation states, atomic/ionic radii, ionization enthalpy, electronegativity, electron gain enthalpy, reactivity towards oxygen, hydrogen, halogen and metals), anomalous behavior of the first element in each group. Group 18 elements: Electronic configuration, Physical properties and chemical properties (reasons for inertness to chemical reactivity) ## The p-Block Elements Elements belonging to groups 13 to 18 of the periodic table are called p-block elements. General electronic configuration: $ns^2np^{1-6}$ (except for He) | Group | 13 | 14 | 15 | 16 | 17 | 18 | |---|---|---|---|---|---|---| | General electronic configuration | $ns^2np^1$ | $ns^2np^2$ | $ns^2np^3$ | $ns^2np^4$ | $ns^2np^5$ | $ns^2np^6$ ($1s^2$ for He) | | First member of the group | B | C | N | O | F | He | | Group oxidation state | +3 | +4 | +5 | +6 | +7 | +8 | | Other oxidation states | +1 | +2, -4 | +3, -3 | +4, +2,-2 | +5, +3, +1, -1 | +6, +4, +2 | ## Group 13 Elements-The Boron Family Group 13 is sometimes referred to as the boron group, named for the first element in the family. These elements are--not surprisingly--located in column 13 of the periodic table. This group includes boron, aluminum, gallium, indium, thallium, and ununtrium (B, Al, Ga, In, Tl, and Uut, respectively). | Element | Symbol | Atomic No. | Electronic Configuration | Abundance in Earth's Crest (in ppm) | |---|---|---|---|---| | Boron | B | 5 | [He] $2s^2 2p^1$ | 8 | | Aluminium | Al | 13 | [Ne] $3s^2 3p^1$ | 81,300 | | Galium | Ga | 31 | [Ar] $3d^{10} 4s^2 4p^1$ | 15 | | Indium | In | 49 | [Kr] $4d^{10} 5s^2 5p^1$ | 1 | | Thallium | Tl | 81 | [Xe] $4f^{14} 5d^{10} 6s^2 6p^1$ | 0.3 | ### Outer Electronic Configuration: $ns^2np^1$ **Atomic Radii:** The atomic and ionic radii of group 13 elements are smaller than the corresponding elements of alkali and alkaline earth metals. **Reason:** On moving from left to right in a period the effective nuclear charge increases and the outer electrons are pulled more strongly towards the nucleus. This results in decrease in atomic size. On moving down the group, both atomic and ionic radii expected to increase due to the addition of a new electron shell with each succeeding element. **Exception:** Atomic radius of Ga is less than that of Al due to the presence of poor shedding 10d-electrons in gallium. **Ionisation enthalpies:** First ionisation enthalpies of the elements of group-13 are less than those of the elements present in group-2 in the same period. **Reason:** The removal of p-electron is much easier than the s-electron and therefore, the first ionisation enthalpies ($\Delta_1 H_1$) of the elements of group 13 are lower as compared to the corresponding elements of group 2. On moving down the group 13 from B to Al the first-ionization enthalpies ($\Delta_1 H_1$) decrease due to an increase in atomic size and screening effect which outweigh the effect of increased nuclear charge. There is discontinuity expected in the ionisation enthalpy values between Al and Ga and between In and Tl due to enability of d- and f-electrons which have low screening effect to compensate the increase in nuclear charge. **Electronegativity:** Down the group, electronegativity first decreases from B to Al and then increases. This is due to discrepancies in the atomic size of the elements. ### Physical Properties (i) Due to strong crystalline lattice boron has high melting point. Rest of the members of this family have low melting point. (ii) Boron is extremely hard and black coloured solid and non-metallic in nature. (iii) Other members of this family are soft metals with low melting point and high electrical conductivity. **Oxidation states:** The first two elements boron and aluminium show only +3 oxidation state ~ in the compounds but the other elements of this group gallium, indium and thalium also exhibit +1 oxidation state in addition to +3 oxidation state i.e., they show variable oxidation states. As we move down the group, the stability of +3 oxidation state decreases while that of +1 oxidation state progressively increases. ### Chemical Properties 1. **Group 13 reactivity towards oxygen** All group 13 elements at high temperatures react and form trioxides, $M_2O_3$. $4M (s) + O_2 (g) → 2 M_2O_3(s)$ The reaction of Ti with oxygen produces $Ti_2O$. Along with that another compound, $Ti_2O_3$, is also produced. On moving down the group, elements of group 13 start reacting more vigorously with oxygen. The boron in crystalline form does not react towards oxygen. Finely split amorphous boron combines with oxygen to generate $B_2O_3$when heated. Aluminium should react with air thermodynamically, but it does not. In this case, it is because $Al_2O_3$forms a protective layer on the surface of the metal, rendering it inert. The elements in group 13 generally do not react directly with hydrogen: **Boron** The first element in the group, boron is generally unreactive with other elements, except at high temperatures. Boron can form many compounds with hydrogen, called boranes, including diborane ($B_2H_6$) and $B_{10}H_{14}$. **Aluminum and gallium** These elements form fewer stable hydrides, but $AlH_3$ and $GaH_3$ do exist. $AlH_3$ is an insoluble, polymeric solid that breaks down quickly in water, and $GaH_3$ is unstable at room temperature. **Indium** Indium is not known to form many hydrides, except in complex compounds. **Thallium** No stable compound of thallium and hydrogen has been synthesized in a laboratory. 2. **The reactivity of group 13 to halogens** They react with halogens to form trihalides $MX_3$ at high temperatures. While on the other hand, Ti only produces $TiF_3$and $TiCl_3$. $2M (s) + 3X_2(g) \rightarrow 2MX_3$ 3. **The reactivity of group 13 with Metals** Only boron reacts with metals to form borides, while the other elements in group 13 are generally reluctant to mix with metals. This is because boron is a nonmetal, while the other elements in group 13 are more apprehensive ### Anomalous behaviour of boron Boron exhibits anomalous behavior compared to other elements in Group 13 of the periodic table due to its small size, high electronegativity, and lack of d-orbitals in its valence shell **Size** Boron is the smallest element in Group 13, so its nucleus has a greater influence on its electrons than other elements in the group. **Electronegativity** Boron is more electronegative than other elements in Group 13, so it's more likely to share electrons than lose them. **Ionization enthalpy** Boron has a very high ionization potential, so it doesn't lose electrons easily. **D-orbitals** Boron doesn't have d-orbitals in its valence shell, so its coordination number can only expand to a maximum of four. Other elements in the group can expand their coordination number to a maximum of six. Other anomalous behaviors of boron include: - Boron is a non-metal, while other elements in the group are metals. - Boron has a high melting and boiling point. - Boron forms only covalent and ionic compounds. - Boron oxide is a weak acid, while oxides of other elements in the group are amphoteric or acidic. ## Group 14 Elements-The carbon Family Group 14 Elements: The Carbon Family Group 14 includes carbon (C), silicon (Si), Germanium (Ge), tin (Sn) and lead (Pb). General electronic configuration of carbon family is $ns^2np^2$. | Element | Atomic Number | Electronic Configuration | Group Number | Period Number | |---|---|---|---|---| | Carbon | 6 |[He] $2s^2 2p^2$ | 14 | 2 | | Silicon | 14 | [Ne] $3s^2 3p^2$ | 14 | 3 | | Germanium | 32 | [Ar] $3d^{10} 4s^2 4p^2$ | 14 | 4 | | Tin | 50 | [Kr] $4d^{10} 5s^2 5p^2$ | 14 | 5 | | Lead | 82 | [Xe] $4f^{14} 5d^{10} 6s^2 6p^2$ | 14 | 6 | **Carbon:** Carbon is the seventeenth most abundant element by weight in the earth's crust. (i) It is available as coal, graphite and diamond. In combined state it is present in metal carbonates, hydrocarbons and carbon dioxide gas (0.03%) in air. (ii) Naturally occurring carbon contains two stable iosotopes $^{12}C$ and $^{13}C$ and third isotope $^{14}C$. $^{14}C$ is a radioactive isotope with half-life 5770 years and is used for radiocarbon dating. Silicon is the second (27.7% by mass) most abundant element on the earth's crust and is present in nature in the form of silica and silicates. Silicon is a very important component of ceramics, glass and cement. Germanium exists only in traces. Tin occurs mainly as cassiterite, $SnO_2$ and lead as galena, $PbS$. **Covalent radius:** Covalent radius expected to increase from C to Si. From Si to Pb small increase is found. **Reason:** Due to the addition of a new energy shell in each succeeding element. The increase in covalent radii from Si to Pb is small due to ineffective shielding of the valence electrons by the intervening d- and f orbitals. **Ionization Enthalpy:** The first ionization enthalpies of group 14 elements are higher than those of the corresponding group 13 elements. **Reason:** Because effective nuclear charge increases and size of the atoms becomes smaller. First ionization enthalpy decreases on moving down the group from carbon to tin. The decrease is very sharp from carbon to silicon while there is slight increase in the first ionization enthalpy of lead as compared to that of tin. **Electronegativity:** Group 14 elements are smaller in size as compared to group 13 elements that's why this group are slightly more electronegative than group 13. From Si to Pb it is almost same. Small increase in ionization enthalpy from Sn to Pb is due to the effect of increased nuclear charge outweighs the shielding effect due to the presence of additional 4f- and 5d-electrons. ### Physical properties: (i) All the elements of group 14 elements are solids. They are less metallic than group 13. (ii) M.P. and boiling points of group 14 elements are generally high. ### Chemical properties: Carbon and silicon mostly show +4 oxidation state. Germanium forms stable compounds in +4 state and only few compounds in +2 state. Tin forms compounds in both oxidation states. Lead forms compounds in +2 state are stable and in +4 state are strong oxidising agents. **• Anomalous Behaviour of Carbon** Carbon, differs from the rest of the member of its family. The main reason for the anomalous behaviour is: (i) exceptionally small atomic and ionic size (ii) higher ionization enthalpy (iii) absence of d-orbitals in the valence shell. (iv) Higher electronagativity. It can be explained as follows: => Since carbon has only s and p-orbitals it can accommodate only four pairs of electrons; other member can expand their covalence due to the presence of d-orbitals. => Carbon can form Ρπ-Ρπ multiple bonds with itself and other atoms having small size and high electronegativity. For example, C=C, C≡C, C=O, C=S and C=N The order of catenation is C>>Si>Ge≈Sn Heavier elements do not form Ρπ-Pr bonds because their atomic orbitals are too large and diffuse to have effective overlapping. => Carbon atoms have the tendency to link with one another through covalent bonds to form chains and rings. This property is called catenation. Down the group property to show catenation decreases. C>Si>Ge>Sn>Pb Lead does not show catenation.- Allotropy: The important allotropes of carbon are diamond, graphite, and fullerenes. The members of carbon family exhibit +4 and +2 oxidation state. The tendency to show +2 oxidation state increases among heavier elements. Lead in +2 state is stable whereas in +4 oxidation state it is a strong oxidising agent. (i) **Reactivity towards oxygen** All members when heated in oxygen form oxides. There are mainly two types of oxides, i.e, monoxide and dioxide of formula MO and $MO_2$ respectively. SiO only exists at high temperature. Oxides in higher oxidation states of elements are generally more acidic than those in lower oxidation states. The dioxides $CO_2$, $SiO_2$ and $GeO_2$ are acidic, whereas $SnO_2$ and $PbO_2$ are amphoteric in nature. Among monoxides, CO is neutral, GeO is distinctly acidic whereas SnO and PbO are amphoteric. (ii) **Reaction with hydrogen** Carbon forms stable bonds with hydrogen and oxygen, which store a large amount of energy. The formation and breakage of these bonds in the carbon cycle are important for life on Earth. For example, plants use energy from the sun to convert carbon dioxide into carbohydrates like glucose in photosynthesis (iii) **Reactivity towards Halogen** These elements can form halides of formula $MX_2$, and $MX_4$, where X = F, Cl, Br, I. All other members except carbon react directly with halogens under suitable conditions to form halides. The stability of tetrahalides decreases as we move from C to Pb. $CCl_4 > SiCl_4 > GeCl_4 > SnCl_4 > PbCl_4$ Most of the $MX_4$ compounds are covalent in nature. ## Group 15 Elements- Nitrogen Family ž The elements of group 15 - Nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb) bismuth (Bi) ### Electronic Configuration | Element | Atomic Number | Electronic Configuration | Group Number | Period Number | |---|---|---|---|---| | Nitrogen | 7 |[He] $2s^2 2p^3$ | 15 | 2 | | Phosphorus | 15 | [Ne] $3s^2 3p^3$ | 15 | 3 | | Arsenic | 33 | [Ar] $ 3d^{10} 4s^2 4p^3$ | 15 | 4 | | Antimony | 51| [Kr] $4d^{10} 5s^2 5p^3$ | 15 | 5 | | Bismuth | 83 | [Xe] $4f^{14} 5d^{10} 6s^2 6p^3$ | 15 | 6 | - The valence shell electronic configuration $ns^2np^3$. - The s orbital is completely filled and p orbitals are half-filled, making their electronic configuration extra stable. +3 and +5 oxidation state **Atomic and Ionic Radii** Covalent and ionic (in a particular state) radii increase in size down the group. There is a considerable increase in covalent radius from N to P. As to Bi only a small increase in covalent radius is observed due to the presence of completely filled d orbitals and/or f orbitals in heavier members. **Ionisation Enthalpy** Ionisation enthalpy: decreases down the group due to gradual increase in atomic size. Ionisation enthalpy of group 15 elements greater than group 14 elements: Because of the extra stable half-filled p orbitals electronic configuration and smaller size of group 15 elements. Increase in magnitude of effective nuclear charge. **Electronegativity** The electronegativity value, in general, decreases down the group with increasing atomic size. ### Physical Properties Polyatomic nature: Dinitrogen diatomic gas while all others are solids (Polyatomic $P_4$). Metallic character: increases down the group. Nitrogen and phosphorus are non-metals, arsenic and antimony metalloids bismuth is a metal. This is due to decrease in ionisation enthalpy and increase in atomic size. The boiling points, increase from top to bottom in the group. The melting point increases upto arsenic and then decreases upto bismuth. ### Periodic properties | Trends | |---|---| | **Electronegativity:** (the atom's ability of attracting electrons) | Decreases down the group| | **Ionization Enthalpy** (the amount of energy required to remove an electron from the atom in it's gaseous phase) | decreases | | **Atomic Radii** (the radius of the atom) | increases | | **Electron Affinity** (ability of the atom to accept an electron) | decreases | | **Melting Point** (amount of energy required to break bonds to change a solid phase substance to a liquid phase) | increases going down the group | | **Boiling Point** (amount of energy required to break bonds to change a liquid phase substance to a gas) | increases going down the group | ### Anomalous properties of nitrogen **Reason** Anamalous property is due to 1) its small size 2) High electronegativity 3) High ionisation enthalpy 4) Non-availability of d orbitals. - Nitrogen form pп-рп multiple bonds. - Bond enthalpy (941.4 kJ mol-1) is very high. - Heavier elements do not form pn-pn bonds as their atomic orbitals are so large and diffuse that they cannot have effective overlapping. - phosphorus (P-P), arsenic (As-As) and antimony (Sb-Sb)form single bonds and bismuthms metallic bonds in elemental state. N-N single bond is weaker than P-P single bond. **Because** - Bond length is short in N-N. - High interelectronic repulsion of the non-bonding electrons. Therefore catenation is weak in nitrogen . - Nitrogen cannot form bond with transition elements : - Absence of d orbitals in its valence shell. - Nitrogen cannot form dn -pn bond as the heavier elements can e.g., $R_3P=O$ or $R_3P=CH_2$ (R = alkyl group). - Phosphorus and arsenic can form dn -dn bond with transition metals when their compounds like $P(C_2H_5)_3$ and $As(C_6H_5)_3$ act as ligands. ### Reactivity towards hydrogen - Hydride formation: Group 15 form $EH_3$ - ( E = N, P, As, Sb or Bi) - The hydrides show regular gradation in their properties. - Stability: decreases from $NH_3$ to $BiH_3$ due to increase in atomic size , decrease in bond dissociation enthalpy. - Reducing character: Increases, due to small bond dissociation enthalpy, covalent character decreases. - Basicity: decreases $NH_3 > PH_3 > AsH_3 > SbH_3 > BiH_3$. - $NH_3$ Is strong base: Small size and high electron density, has lone pair . ### Reactivity towards metals - Form binary compounds (having +3 oxidation) - $Ca_3N_2$ (Calcium Nitride), $Ca_3P_2$ (Calcium Phosphide) $Na_3As_2$ (Sodium Arsenide), $Zn_3Sb_2$(ZincAntimonide) $Mg_3Bi_2$ (Magnesium Bismuthide) ### Hydrides comparison | Property | Down the group | Reason | |---|---|---| | Basic strength of $MH_3$ | Decreases gradition. $NH_3>PH_3>AsH_3>56H_3>BiH_3$ | As the size of central atom increases electron density decreases. | | Thermal stability of $MH_3$ | decreases gradition $NH_3>PH_3>AsH_3>SbH_3>BiH_3$ | As the size of the central atom increases its tendency to form stable M-H bords decreases. | | Reducing character | Increases | As the stability of hydrides decrease as the reducing character increase. | | Melting and boiling point | Increases (except in N) gradation. $PH_3 < AsH_3 < NH_3 <SbH_3 < BiH_3$ | $NH_3$ has high melting point and boiling point than $PH_3$ due to hydrogen bonding. As the molecular size increases vander waals force increases. | ### Reactivity towards oxygen - Form two types of oxides: $E_2O_3$ and $E_2O_5$. - The oxide in higher oxidation state is more acidic. Their acidic character decreases down the group. - Nitrogen atom has small atomic size, strong pull of electron pair between O-H bond,releases the H+ ion. - Effect decreases as atomic size increases. - The oxides of type $E_2O_3$ of nitrogen and phosphorus are purely acidic, arsenic and antimony amphoteric, bismuth predominantly basic. ### Reactivity towards halogens - Halides formation: $MX_3$ and $MX_5$. - Nitrogen does not form pentahalide due to non-availability of the d orbitals in its valence shell, contains only 1s and 3- p orbitals. - Pentahalides are more covalent than trihalides. - All the trihalides (covalent nature) of these elements except those of nitrogen are stable. - In case of nitrogen, only $NF_3$ is known to be stable. Trihalides except $BiF_3$ are predominantly covalent in nature. ## Group 16 Elements - Oxygen, sulphur, selenium, tellurium and polonium(radioactive). - They are also known as (chalcogens - ore forming) - Derived from Greek word for brass and points to the association of sulphur and its congeners with copper. - Copper minerals contain oxygen or sulphur and other members of the group. - Present in earth crust, gypsum, epsum, pyrite, zinc blend, $H_2S$ in volcanoes, protein, garlic, onion, hair ### Periodic properties | Trends| |---|---| | **Atomic Radii** (the radius of the atom) | increases | | **Electronegativity:** (the atom's ability of attracting electrons) | Decreases down the group | | **Ionization Enthalpy** (the amount of energy required to remove an electron from the atom in its gaseous phase) | decreases | | **Electron Affinity** (ability of the atom to accept an electron) | decreases | | **Melting Point** (amount of energy required to break bonds to change a solid phase substance to a liquid phase) | increases going down the group | | **Boiling Point** (amount of energy required to break bonds to change a liquid phase substance to a gas) | increases going down the group | 1. **Electronic Configuration : $ns^2 np^4$** | Element | At. No. | Electronic Configuration | Oxidation State | |---|---|---|---|---| | Oxygen (O) | 8 | [He] $2s^2 2p^4$ | -2, -1, +1, +2 | | Sulphur (S) | 16 | [Ne] $3s^2 3p^4$ | -2, +2, +4, +6 | | Selenium (Se) | 34 | [Ar] $3d^{10} 4s^2 4p^4$ | -2, +2, +4, +6 | | Tellurium (Te) | 52 | [Kr] $4d^{10} 5s^2 5p^4$ | -2, +2, +4, +6 | | Polonium (Po) | 84 | [Xe] $4f^{14} 5d^{10} 6s^2 6p^4$ | +2, +4, +6 | | Livermorium (Lv) | 116 | [Rn] $5f^{14} 6d^{10} 7s^2 7p^4$ | | 2. **Atomic and Ionic Radii:** Increases Due to increase in the number of shells 3. **Ionisation Enthalpy:** Decreases due to increase in size Grop16 has lower I.E than Group15 . Due to the fact that Group 15 elements have extra stable half- filled p orbitals electronic configurations. 4. **Electron Gain Enthalpy:** Because of the compact nature of oxygen atom (small size) e-e repulsion, it has less negative electron gain enthalpy than sulphur. However, from sulphur onwards the value again becomes less negative upto polonium due to increase in size. 5. **Electronegativity : F>O>N** Electronegativity decreases with an increase in atomic number or size. Metallic character increases from oxygen to polonium. ### Physical Properties - Radioactive - Exhibit allotropy - Melting and boiling point increases due to increase in atomic mass. ### Chemical Properties 1. **Oxidation state:** a) -2,-1,+2,+4,+6 b) +2 $OF_2$ c) Oxygen does not show +4 and +6 O.S due to lack of d-orbitals d) Stability of +6 oxidation state in higher elements due to inert pair effect. ### Anomalous behaviour of oxygen - Due to Small size, high I.E. and high electronegativity. - The absence of d orbitals in oxygen limits its covalency to four, rarely exceeds two. - On the other hand, in case of other elements of the group, the valence shells can be expanded and covalence exceeds four. ### Reactivity with hydrogen - Hydrides of the type $H_2E$ (E = O, S, Se, Te, Po). - Acidic character: increases from $H_2O$ to $H_2Te$. Due to decrease in bond (H-E) dissociation enthalpy. - Thermal stability: decrease bond (H-E) dissociation enthalpy decreases. - Reducing property: character increases from $H_2S$ to $H_2Te$. Bond length increases ### Reactivity with oxygen - Oxides: $EO_2$ and $EO_3$ (E = S, Se, Te, Po) - Ozone ($O_3$) and sulphur dioxide ($SO_2$) and ($SO_3$) are gases while selenium dioxide ($SeO_2$) is solid. - Reducing property of dioxide decreases from $SO_2$ to $TeO_2$ . Besides $EO_2$ type, sulphur, selenium and tellurium also form $EO_3$ type oxides ($SO_3$, $SeO_3$, $ТеO_3$). Both types of oxides are acidic in nature. ### Reactivity towards the halogens - Type: $EX_6$, $EX_4$ and $EX_2$. - Stability: decreases in the order F-> Cl- > Br¯ > I- . - Hexahalides : hexafluorides are only stable halides. gaseous in nature, octahedral structure sp³d². Eg.$SF_6$. - Tetrafluorides : $SF_4$ gas, $SeF_4$-liquid and $TeF_4$ - solid. Sp³d hybridisation ,have trigonal bipyramidal, having lone pair of electrons at equitorial position. - All elements except selenium form dichlorides and dibromides. sp³ hybridisation tetrahedral structure. ### Reactivity towards the Metals Metals react with O, S, Se to form oxides, sulphides and selenides respectively. - $4Al + 3O_2 \rightarrow 2Al_2O_3$ - Aluminium oxide - $Cu+S \rightarrow CuS$ - Copper sulphide - $Mg + Se \rightarrow MgSe$ - Magnesium selenide ## Group 17 Elements - Fluorine, chlorine, bromine, iodine and astatine (radioactive). - halogens (salt forming or salt producers). - Highly reactive, non-metallic elements. **Occurance:** - Fluorine: Fluorspar $CaF_2$, Cryolite $Na_3AlF_6$ - Cl, Br, I:Sea water as salt of Na, K, Mg, Ca, ### Electronic Configuration: ($ns^2 np^5$) | Element | Atomic Number | Electronic Configuration | |---|---|---| | Fluorine | 9 | [He] $2s^2 2p^5$ | | Chlorine | 17 | [Ne] $3s^2 3p^5$ | | Bromine | 35 | [Ar] $3d^{10} 4s^2 4p^5$ | | Iodine | 53 | [Kr] $4d^{10} 5s^2 5p^5$ | | Astatine | 85 | [Xe] $4f^{14} 5d^{10} 6s^2 6p^5$ | - Atomic and Ionic Radii: smallest atomic radii due to maximum effective nuclear charge. Atomic and ionic radii increase due to increasing number of quantum shells. - Ionisation Enthalpy: Little tendency to lose electron due to very high ionisation enthalpy. Due to increase in atomic size, ionisation enthalpy decreases down the group. - Electron Gain Enthalpy: Maximum :only one electron less than stable noble gas configurations. - Negative electron gain enthalpy of fluorine is less than that of chlorine due to small size of fluorine atom, strong interelectronic repulsions in 2p orbitals of fluorine, experience less attraction. - Electronegativity: Very high due to increase nuclear charge. Decreases down the group due to increase atomic radia. ### Physical Properties - F, Cl- gases, Br - liquid, I solid. - Melting and boiling points increase with atomic number. - Coloured: Due to absorption of radiations in visible region which results in the excitation of outer electrons to higher energy level. By absorbing different quanta of radiation, they display different colours. Eg. $F_2$ - yellow, $Cl_2$ greenish yellow, $Br_2$ red and $I_2$-violet colour. - Bond dissociation enthalpy: $F_2 <Cl_2>Br_2 > I_2$ - 1. $F_2$ has smaller bond dissociation enthalpy than $Cl_2$ Due to large electron-electron repulsion among the lone pairs in $F_2$ - 2. Much closer to each other than $Cl_2$. ### Chemical Properties - Oxidation states, all the halogens exhibit -1 oxidation state. However, chlorine, bromine and iodine exhibit + 1, +3, + 5 and + 7 oxidation states ### Anomalous behaviour of fluorine 1. Small size 2. ionisation enthalpy, electronegativity, and electrode potentials are higher. 3. Non availability of d orbitals in valence shell. 4. ionic and covalent radii, m.p. and b.p.,low F-F bond dissociation enthalpy and electron gain enthalpy are quite lower than expected. ### Reactivity towards hydrogen 1. Affinity for hydrogen decreases from fluorine to iodine. 2. Acidic strength: HF <HCl < HBr < HI. due to decrease in bond (H-X) dissociation enthalpy 3. Reducing character:HF < HCI < HBr < HI. due to decrease in bond (H-X) dissociation enthalpy 4. Stability: due to decrease in bond (H-X) dissociation enthalpy. H-F > H-Cl > H-Br > H-I. ### Reactivity towards oxygen 1. Fluorine: $OF_2$ and $O_2F_2$. 2. Chlorine, bromine and iodine form oxides in +1 to +7 3. Chlorine oxides, $Cl_2O$, $ClO_2$, $Cl_2O_6$ and $Cl_2O_7$. $ClO_2$used as bleaching agent for paper pulp and textiles and water treatment. 4. Bromine oxides, $Br_2O$, $BrO_2$, $BrO_3$. 5. Iodine oxides, $I_2O_4$, $I_2O_5$, $I_2O_7$. $I_2O_5$ is very good oxidising agent and used in estimation of carbon monoxide. ### Reactivity towards metals - Metal halides: - $Mg(s) + Br_2 (1) → MgBr_2(s)$ - The ionic character decreases: MF >MCI >MBr>MI (M is a monovalent metal)+1 - Halides in higher oxidation state - covalent. - Eg. $SnCl_4$, $PbCl_4$, $SbC1_5$ and $UF_6$ are more covalent than $SnCl_2$, $PbCl_2$, $SbCl_3$ and $UF_4$ respt. ## Group 18 Elements - Elements of 18 group are helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe) and radon (Rn) (radioactive). - All of these are gases and chemically unreactive. Down the group their ionisation enthalpy decreases, therefore Xe reacts with oxygen and fluorine and forms different compounds. Xenon and radon are the rarest elements of the group. First prepared noble gas compound by Neil Bartlett in 1962 is $XePtF_6$. ### Physical properties - Noble gases are monatomic gases with weak van der Waals forces between their atoms. As you move down the group, the atomic radius increases, which results in an increase in melting point, boiling point, and enthalpy of vaporization. | He | Ne | Ar | Kr | Xe | Rn | |---|---|---|---|---|---| | increase in density | | | | | | | increase in atomic radius | | | | | | | increase in number of filled electron shells | | | | | | | increase in melting points and boiling points | | | | | | - **Electronic configuration** - General Electronic configuration ($ns^2 np^6$) | Element | Symbol | Electronic configuration | |---|---|---| | Helium | 2He | $1s^2$ | | Neon | 10Ne | [He] $2s^22p^6$ | | Argon | 18At | [Ne] $3s^23p^6$ | | Krypton | 36Kr | [Ar] $3d^{10} 4s^2 4p^6$ | | Xenon | 54 Xe | [Kr] $4d^{1