अधातुएँ की सम्पूर्ण जानकारी || Non-Metals

बड़ी संख्या में तत्व धातुएँ हैं। केवल 22 अधात्वीय तत्व हैं जिनमें से 11 गैसें हैं, एक द्रव है तथा शेष 10 ठोस हैं।

भौतिक गुण

अधातुएँ सामान्यत: अधिक भंगुर होती हैं तथा इनसे चादरें अथवा तार नहीं बनाए जा सकते हैं। इनमें कोई चमक नहीं होती है तथा इन पर पॉलिश नहीं की जा सकती है।

अधातुएँ सामान्यतः ऊष्मा एवं विद्युत की कुचालक होती हैं। धातुओं की भाँति अधातुओं में स्वतंत्र इलेक्ट्रॉन नहीं होते हैं। कार्बन का एक अपररूप इसका अपवाद है तथा यह विद्युत का अच्छा चालक है।

रासायनिक गुण

धातुओं के विपरीत, अधातुएँ वैद्युत ऋणात्मक होती है। वे इलेक्ट्रानों को आसानी से ग्रहण कर लेती हैं तथा ऋणात्मक आवेशयुक्त आयन बनाती हैं।

ऑक्सीजन के साथ अभिक्रिया

अधातुएँ ऑक्सीजन के साथ सहसंयोजक ऑक्साइड बनाती हैं। इनमें से कुछ ऑक्साइड जल में घुलने के बाद अम्ल बनाते हैं। अन्य ऑक्साइड जैसे CO, N2O आदि उदासीन होते हैं (अर्थात वे लिटमस पेपर के साथ अम्लीय परीक्षण नहीं देते हैं)।

अम्लों के साथ अभिक्रिया

धातुओं की भाँति, अधातुएँ अम्लों में हाइड्रोजन को पुन: स्थापित नहीं करती हैं। इस अभिक्रिया को संपन्न करने के लिए अम्ल के II आयन के लिए इलेक्ट्रॉन उपलब्ध होने चाहिए। क्योंकि अधातु

इलेक्ट्रॉनों को ग्रहण करती हैं अत: वे H आयन को इलेक्ट्रॉन उपलब्ध नहीं कर सकती हैं।

क्लोरीन के साथ अभिक्रिया

क्लोरीन के साथ अधातुएँ सहसंयोजक आबन्ध वाले क्लोराइड बनाती है।

2P2+6Cl 4 PCl3 (फॉस्फोरस ट्राइक्लोराइड)

हाइड्रोजन के साथ अभिक्रिया

अधातुएँ हाइड्रोजन के साथ संयोग कर हाइड्राइड बनती हैं। उदाहरण के लिये H2S.NH, HCI, CH4 आदि। ये हाइड्राइड इलेक्ट्रॉनों की साझेदारी से बनते हैं।

हमने धातुओं एवं अधातुओं दोनों के ही लाक्षणिक गुण पढ़ लिए हैं। धातुओं तथा अधातुओं के गुणों में भिन्नताओं को संक्षेप में जान लेना लाभप्रद होगा।

यद्यपि अधातुएँ संख्या में कम हैं, फिर भी वे वायु, महासागर तथा पृथ्वी की प्रमुख घटक हैं। आपको ज्ञात है कि वायु के मुख्य घटक ऑक्सीजन तथा नाइट्रोजन हैं। क्लोरीन, महासागर में मुख्यतः क्लोराइड के रूप में उपस्थित होती है। भूपर्पटी में बहुलता के अनुसार प्रमुख अधातुएँ ऑक्सीजन, सिलिकान, फॉस्फोरस तथा सल्फर हैं।

सिलिकन

यदि हम कहीं से भी एक मुट्ठी मिट्टी उठाएँ तो हमारे हाथ में सिलिकन का कोई-न-कोई यौगिक अवश्य होगा। ये पृथ्वी पर सर्वाधिक सामान्य तथा वितरित तत्व है। फिर भी, यह कभी भी स्वतंत्र अवस्था में नहीं पाया जाता है। यह सदैव यौगिक के रूप में ही पाया जाता है तथा अधिकांशत: ऑक्सीजन से बन्धित रहता है। ऑक्सीजन तथा सिलिकन का सबसे सरल यौगिक सिलिकन डाइआक्साइड है जिसे सिलिका कहते हैं। यह विभिन्न रूपों में पाया जाता है जैसे रेत, फ्लिंट, स्फटिक तथा दूधिया पत्थर। सिलिकन, यह आवर्ती वर्ग का तत्व है तथा आवर्त सारणी में उसका स्थान कार्बन के नीचे आता है। इसकी परमाणु संख्या 14 है तथा इलेक्ट्रॉन विन्यास (2, 8, 4) हैं। कार्बन के समान ही इसमें अष्टक पूर्ण करने हेतु बाह्य कोष में चार इलेक्ट्रॉन कम हैं। इसलिए कार्बन के समान ही इसकी संयोजकता भी 4 है।

धातुओं तथा अधातुओं के गुणों की तुलना
धातुअधातु
भौतिक गुण
1. ऊष्मा एवं विद्युत सुचालक हैसामान्यत:ऊष्मा एवं विद्युत के कुचालक। ग्रेफाइट अपवाद तथा विद्युत सुचालक है।
2. आघातवर्ध्य तथा तन्यसामान्यत: भंगुर, न ही आघातवर्ध्य न ही तन्य।
3. चमकदार तथा पॉलिश की जा सकती है।सामान्यत: चमकदार नहीं।
4. ठोस रासायनिक गुणठोस, द्रव, गैस।
रासायनिक गुण
1. धातुएँ क्षारीय ऑक्साइड बनाती हैं जिनमें से कुछ क्षार बनाते हैंअधातुएँ, अम्लीय अथवा उदासीन, ऑक्साइड बनाती हैं।
2. धातुएँ अम्लों में से हाइड्रोजन पुन:स्थापित करती हैं तथा लवण बनाती हैं।अधातुएँ, अम्लों में से हाइड्रोजन को पुन: स्थापित नहीं करती हैं।
3. धातुएँ क्लोरीन से संयोग कर क्लोराइड बनाती हैं, जो वैद्युत संयोजक होते हैं।अधातुएँ, क्लोरीन से संयोग कर क्लोराइड बनाती हैं, जो सहसंयोजक हैं।
4. कुछ धातुएँ हाइड्रोजन के साथ संयोग कर हाइड्राइड बनाती हैं, जो विद्युत संयोजक होते हैं।अधातुएँ, हाइड्रोजन के साथ अनेक स्थाई हाइड्राइड बनाती हैं, जो सहसंयोजक होते हैं।

 

सिलिकन को इसके सामान्य रूप से उपलब्ध यौगिक सिलिकन डाइऑक्साइड से प्राप्त किया जाता है। सिलिकन को ऑक्सीजन से पृथक करने की प्रक्रिया में कोई अपचायक उपयोग करना होता है। सबसे कम खर्चीला अपचायक कोक है। सिलिकन डाइऑक्साइड को कोक के साथ विद्युत भट्टी में गर्म करके सिलिकन प्राप्त किया जा सकता है।

SiO2 + C  Si + CO2

सिलिकन हाइड्रॉक्साइड + कोक  सिलिकन

सिलिकन के गुण

तत्व सिलिकन कक्ष ताप पर ठोस होता है। यह कड़ा, सिलेटी रंग का तथा चमकदार पदार्थ है। यह बहुत उच्च ताप 1410°C पर पिघलता है। यह क्रिस्टलीय ठोस पदार्थ है जिसमें परमाणु व्यवस्था हीरे के समान होती है। प्रत्येक सिलिकन परमाणु चार Si परमाणुओं से घिरा रहता है। समस्त Si परमाणु परस्पर सहसंयोजक बन्धों के जाल द्वारा जुड़े रहते हैं।

कक्ष ताप पर सिलिकन लगभग निष्क्रिय होती है। किन्तु उच्च ताप पर क्रियाशील हो जाती है।

सिलिकन गर्म करने पर जलती है तथा ऑक्सीजन के साथ संयोग कर श्वेत रंग की ठोस सिलिका बनाती है।

Si + O2  SiO2

अक्रिस्टलीय सिलिकन 450°C ताप पर क्लोरीन से संयोग कर सिलिकन टेट्राक्लोराइड बनाती है।

Si + 2Cl2  SiCl2

लाल गर्म होने पर यह जल वाष्प का अपघटन कर सिलिकन डाइऑक्साइड तथा हाइड्रोजन उत्पन्न करती है।

Si + 2H2 SiO2 +2H2

सिलिकन हाइड्रोक्लोरिक अम्ल के साथ अभिक्रिया कर के सिलिकन टेट्राक्लोराइड तथा हाइड्रोजन बनाती है।

Si + 4HCl  SiCl4 + 2H2

सिलिकन गर्म सोडियम हाइड्रोक्साइड घोल में घुल कर हाइड्रोजन मुक्त करती है।

2NaOH + Si + H2 Na2SiO3 +2H2

सिलिकन के उपयोग

लोहा, एलुमीनियम, तांबा तथा मैंगनीज की मिश्र धातुओं को बनाने में सिलिकान बनाने में उपयोग होता है। इसका महत्वपूर्ण उपयोग अर्धचालक युक्तियों में होता है।

सिलिकन कार्बाइड, हीरे की तरह कठोर पदार्थ है, जिसका उपयोग कटिंग तथा ग्राइडिंग के औजार बनाने में किया जाता है। सिलिकन डाइऑक्साइड का उपयोग सीमेन्ट तथा काँच बनाने के लिये भी किया जाता है।

फॉस्फोरस

फॉस्फोरस पृथ्वी पर पर्याप्त मात्रा में उपलब्ध है। वह समस्त जैव जीवों-पौधे तथा जन्तु दोनों में ही पाया जाता है। हमारे शरीर में फॉस्फोरस, हड्डियों, मांसपेशियों तथा स्नायु ऊतकों में होता है। फॉस्फोरस वर्ग V का तत्व है तथा आवर्त सारणी में इसका स्थान नाइट्रोजन के नीचे आता है। नाइट्रोजन के समान इसके बाह्य कोश में पाँच इलेक्ट्रॉन होते हैं। इसके अष्टक को पूर्ण करने हेतु तीन और इलेक्ट्रॉनों की आवश्यकता है। इसकी संयोजकता तीन अथवा पांच होती है अर्थात् यह दो संयोजकता प्रदर्शित करता है।

फॉस्फोरस अत्यधिक क्रियाशील तत्व है, अतः स्वतंत्र अवस्था में नहीं पाया जाता है। फॉस्फोरस तत्व को अस्थियों अथवा खनिज फॉस्फोरस से प्राप्त किया जाता है। रॉक फॉस्फेट को रेत तथा कोक के साथ गर्म करके भी फॉस्फोरस प्राप्त किया जाता है।

फॉस्फोरस के अपररूप

फॉस्फोरस दो प्रमुख अपररूपों में पाया जाता है। श्वेत या पीला फॉस्फोरस अस्थायी होता है तथा इसे जल में भण्डारित करना आवश्यक है। खुली हवा में रख देने पर यह ऑक्सीकरण के कारण सुलगने लगता है और इस कारण यह अंधेरे में प्रकाशमान हो जाता है। फॉस्फोरस के इस गुण को बहुधा स्फुरदीप्ति कहते हैं। इसके विपरीत लाल फॉस्फोरस स्थायी होता है। यह सामान्यतः माचिस उद्योग में उपयोग होता है।

लाल तथा पीले फॉस्फोरस के गुणों में भिन्नता केवल इस तथ्य के कारण है कि लाल फॉस्फोरस में फॉस्फोरस परमाणुओं के मध्य बन्ध सशक्त है, जबकि पीले फॉस्फोरस में यह अपेक्षाकृत दुर्बल है।

सुपर फास्फोरस उर्वरक

हरे पौधों को स्वस्थ रहने के लिए फॉस्फोरस की आवश्यकता पड़ती है। पौधे, फास्फोरस की अनुपस्थिति में प्रकाश संश्लेषण की क्रिया सम्पन्न नहीं कर सकते हैं। इसकी कमी से पहले पतियाँ रंगहीन हो जाती हैं। यदि फॉस्फोरस प्राप्त न हो पाए तो अंततः पौधे मर जाते हैं। फसल की पैदावार भी मृदा में फास्फोरस की उपलब्धता पर निर्भर करती है। पौधे केवल फॉस्फोरस के घुलनशील यौगिकों को उपयोग कर सकते हैं। अत: उर्वरक उद्योग उन उर्वरकों को बनाने में लगे हुए हैं, जिनमें फॉस्फेट आयन हों।

फॉस्फेट उर्वरक, फॉस्फेट चट्टान से बनते हैं, जिसमें कैल्सियम फॉस्फेट होता है। कैल्सियम फॉस्फेट की सल्फ्यूरिक अम्ल की उचित मात्रा के साथ अभिक्रिया से सुपर फॉस्फेट बनाते हैं। इनका उर्वरक के रूप में उपयोग होता है।

ऑक्सीजन से अभिक्रिया

पीला फॉस्फोरस, ऑक्सीजन के साथ शीघ्रता से संयोग करता है। हल्का-सा गर्म करते ही अभिक्रिया प्रारम्भ करने हेतु पर्याप्त है। यदि इसे हवा में खुला रख दिया जाए तब भी यह जलना प्रारंभ कर देता है। लाल फॉस्फोरस ऑक्सीजन के साथ अभिक्रिया करने के लिए गर्म करना पड़ता है। दोनों ही प्रकार के फॉस्फोरस ऑक्सीजन से संयोग कर ट्राइआक्साइड या पेन्टाआक्साइड बनाते हैं।

P4+3O2 → 2P2O2

P+5O 2P2O3

क्लोरीन से अभिक्रिया

फॉस्फोरस के दोनों अपररूप क्लोरीन गैस के साथ अभिक्रिया करके ट्राइक्लोराइड तथा पेन्टाक्लोराइड बनाते हैं। पीला फॉस्फोरस कक्ष ताप पर ही शीघ्रता से अभिक्रिया करता है जबकि लाल फॉस्फोरस गर्म करने पर ही अभिक्रिया करता है।

P4+6C 4PC3

P4+ 10Cl 4PCI5

सोडियम अथवा पोटेशियम हाइड्रोक्साइड से अभिक्रिया

सोडियम या पोटेशियम हाइड्रोक्साइड के गर्म सांद्र घोल के साथ लाल फॉस्फोरस कोई अभिक्रिया नहीं करता है। इस अभिक्रिया में फास्फीन गैस उत्पन्न होती है। यह अत्यधिक दम घोटने वाली गैस है।

Р4 + 3NaOH+3H2 PH3 + 3NaH2РO2

व्यवहारिक जीवन में रसायन

  1. दियासलाई बनाने में लाल फॉस्फोरस अथवा फॉस्फोरस ट्राई सल्फाइड का तथा चूहानाशक विष एवं कई प्रकार के विस्फोटों में श्वेत फॉस्फोरस का प्रयोग किया जाता है।
  2. फॉस्फोरस, फॉस्फोरस ब्रान्ज नामक मिश्र धातु (जो फॉस्फोरस, कॉपर और टीन का मिश्रित कप है) बनाने में भी प्रयोग किया जाता है।
  3. फॉस्फोरस के यौगिकों में प्रमुख हैं- फॉस्फीन (PH3)। इसे प्रयोगशाला में कास्टिक सोडे के सान्द्र विलयन का फॉस्फोरस से संयोग कराकर प्राप्त किया जाता है।

4Р +3NaOH+3H2 3NаH2 +РО2+РН3

सल्फर

प्रकृति में स्वतंत्र एवं संयुक्त, दोनों अवस्थाओं में उपलब्ध है। सिलिकन तथा फॉस्फोरस के विपरीत, संसार के कुछ भागों में प्राकृतिक सल्फर के बड़े, भण्डार पाए जाते हैं। यौगिक रूप में यह अनेक महत्वपूर्ण धात्वीय खनिजों, जैसे- सिनाबार, जिंक ब्लेंड, कॉपर पायराइट आदि का घटक हैं। सल्फर VI वर्ग के तत्वों के अन्तर्गत आती है तथा आवर्त सारणी में इसका स्थान ऑक्सीजन के नीचे आता है। इसके बाह्य कोष में 6 इलेक्ट्रॉन होते हैं तथा परिवर्ती संयोजकता दर्शाती है जो 2, 4, तथा 6 हो सकती है। सामान्य ताप पर यह पीला क्रिस्टलीय ठोस है, जिसका गलनांक 159°C होता है।

भूमिगत सल्फर एक विशेष प्रक्रिया द्वारा बाहर निकलती जाती है जिसको फ्रास प्रक्रिया कहते हैं। यह प्रक्रिया इस तथ्य पर आधारित है कि सल्फर का गलनांक अपेक्षाकृत कम होता है। फ्रास प्रक्रिया में पृथ्वी की सतह से सल्फर तल तक भेदन करके छिद्र बनाया जाता है। इस छिद्र के माध्यम से तीन संक्रेदित पाइपों को अन्दर डाला जाता है। सबसे बाहरी पाइप से उच्च दाब अत्यधिक गर्म पानी नीचे की ओर भेजा जाता है। इस अति गर्म पानी के कारण तल में सल्फर पिघल जाती है। सबसे अन्दर के पाइप के द्वारा अत्यधिक उच्च दाब पर वायु भेजी जाती है, जो पिघली हुई सल्फर को बीच की पाइप में धकेल देती है, जिससे होती हुई वह सतह तक पहुँच जाता है।

अम्लों द्वारा ऑक्सीकरण

सल्फ्यूरिक अम्ल तथा नाइट्रिक अम्ल जैसे आक्सीकरण अम्लों द्वारा सल्फर, आक्सीकृत हो जाती है। यह गर्म एवं सांद्र सल्फ्यूरिक अम्ल के साथ सल्फर डाईआक्साइड बनाती है।

S+2H2SO 2H2O+3SO2

सांद्र नाइट्रिक अम्ल के साथ यह सल्फ्यूरिक अम्ल बनाती है।

S+6HNO H2SO4+6NO2+2H2O

सल्फर का प्रमुख उपयोग सल्फ्यूरिक अम्ल बनाने के लिये होता है। इसका उपयोग त्वचा क्रीम, एण्टीसेप्टिक, तथा बारूद बनाने और रबड़ के वल्कनीकरण में होता है।

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