فسفر

از ویکیپدیا، دانشنامه آزاد
رفتن به ناوبری پرش به جستجو

فسفر ،  15 ص
PhosphComby.jpg
سفید مومی (برش زرد) ، قرمز (گرانول در وسط چپ ، قطعه وسط راست) و فسفر بنفش
فسفر
تلفظ/ F ɒ بازدید کنندگان F ər ə بازدید کنندگان / ( FOS -fər-əs )
ظاهرسفید مومی ، زرد ، قرمز ، بنفش ، رنگ مشکی فلزی دارد
وزن اتمی استاندارد A r، std (P) 30.973 761 998 (5) [1]
فراوانی
در  پوسته زمین5.2 (سیلیکون = 100)
فسفر در جدول تناوبی
هیدروژن هلیوم
لیتیوم بریلیوم بور کربن نیتروژن اکسیژن فلورین نئون
سدیم منیزیم آلومینیوم سیلیکون فسفر گوگرد کلر آرگون
پتاسیم کلسیم اسکاندیم تیتانیوم وانادیوم کروم منگنز اهن کبالت نیکل مس فلز روی گالیوم ژرمانیم آرسنیک سلنیوم بروم کریپتون
روبیدیم استرانسیوم ایتریوم زیرکونیوم نیوبیوم مولیبدن تکنسیوم روتنیوم رادیوم پالادیوم نقره اي کادمیوم ایندیوم قلع آنتیموان تلوریم ید زنون
سزیم باریم لانتانیم سریم پرازئودیمیوم نئودیمیم پرومتیم ساماریوم یوروپیوم گادولینیوم تربیوم دیسپروزیم هولمیوم اربیوم تولیوم ایتربیوم لوتیتیم هافنیوم تانتالوم تنگستن رنیوم اسمیوم ایریدیوم پلاتین طلا جیوه (عنصر) تالیوم رهبری بیسموت پولونیوم استاتین رادون
فرانسیوم رادیوم اکتینیوم توریم پروتاکتینیوم اورانیوم نپتونیم پلوتونیوم قاره آمریکا کوریوم برکلیوم کالیفرنیوم اینشتینیم فرمیوم مندلویوم نوبلیوم لورنسیوم رادرفوردیوم دوبنیوم Seaborgium بوهریوم پتاسیم میتنریوم درمستادیم رنتژنیوم کوپرنیکیوم نیهونیوم فلورویوم مسکوویوم لیورموریوم تنسی اوگانسون
N

P

As
سیلیکونفسفرگوگرد
عدد اتمی ( Z )15
گروهگروه 15 (مواد مخدر pnictogens)
عادت زنانهدوره 3
مسدود کردن  بلوک p
ساختار الکترونی[ Ne ] 3s 2 3p 3
الکترون در هر پوسته2 ، 8 ، 5
مشخصات فیزیکی
مرحله در  STPجامد
نقطه ذوبسفید: 317.3 کیلوگرم ( 44.15  درجه سانتی گراد ، 111.5 درجه فارنهایت)
قرمز: -860 کیلوگرم (-590 درجه سانتی گراد ، -1090 درجه فارنهایت) [2]
نقطه جوشسفید: 553.7 کیلوگرم (280.5 درجه سانتی گراد ، 536.9 درجه فارنهایت)
نقطه تصعیدقرمز: -689.2-863 K (-416-590 ° C ، -780.8-1094 ° F)
بنفش: 893 K (620 ° C ، 1148 ° F)
تراکم (نزدیک  rt )سفید: 1.823 گرم/سانتی متر 3
قرمز: .22.2–2.34 گرم/سانتی متر 3
بنفش: 2.36 گرم/سانتی متر 3
سیاه: 2.69 گرم/سانتی متر 3
گرمای همجوشیسفید: 0.66  کیلوژول/مول
گرمای تبخیرسفید: 51.9 کیلوژول بر مول
ظرفیت حرارتی مولارسفید: 23.824 J/(mol · K)
فشار بخار (سفید)
P  (Pa) 1 10 100 1 کیلوگرم 10 کیلوگرم 100 کیلوگرم
در  T  (K) 279 307 342 388 453 549
فشار بخار (قرمز ، فشار خون 431 درجه سانتی گراد)
P  (Pa) 1 10 100 1 کیلوگرم 10 کیلوگرم 100 کیلوگرم
در  T  (K) 455 489 529 576 635 704
خواص اتمی
حالتهای اکسیداسیون−3 ، −2 ، −1 ، 0 ، [3] +1 ، [4] +2 ، +3 ، +4 ، +5 (یکاکسید اسیدی ملایم)
الکترونگاتیویمقیاس پاولینگ: 2.19
انرژی های یونیزاسیون
  • اول: 1011.8 کیلوژول بر مول
  • دوم: 1907 کیلوژول بر مول
  • سوم: 2914.1 کیلوژول بر مول
  • ( بیشتر )
شعاع کوالانسی107 ± 3  بعد از ظهر
شعاع ون در والسساعت 180 بعد از ظهر
Color lines in a spectral range
خطوط طیفی فسفر
سایر خواص
وقوع طبیعیازلی
ساختار کریستالیbody-centered است مکعب (BCC)
Bodycentredcubic crystal structure for phosphorus
رسانایی گرماییسفید: 0.236 W/(m⋅K)
سیاه: 12.1 W/(m⋅K)
ترتیب مغناطیسیسفید ، قرمز ، بنفش ، سیاه: دیامغناطیس [5]
حساسیت مغناطیسی مولار.820.8 × 10 −6  سانتی متر 3 /مول (293 کیلوگرم) [6]
مدول فلهسفید: 5 GPa
قرمز: 11 GPa
شماره CAS7723-14-0 (قرمز)
12185-10-3 (سفید)
تاریخ
کشفبرند هنیگ (1669)
به عنوان عنصر توسط شناخته شده استآنتوان لاووازیه [7] (1777)
ایزوتوپهای اصلی فسفر
ایزوتوپ فراوانی نیمه عمر ( t 1/2 ) حالت پوسیدگی تولید - محصول
31 ص 100٪ پایدار
32 ص پی گیری 14.28 روز β - 32 ثانیه
33 ص پی گیری 25.3 روز β - 33 ثانیه
Category دسته: فسفر
| منابع

فسفر یک عنصر شیمیایی با علامت P و عدد اتمی 15. فسفر عنصری به دو شکل عمده وجود دارد ، فسفر سفید و فسفر قرمز ، اما از آنجا که بسیار واکنش پذیر است ، فسفر هرگز به عنوان یک عنصر آزاد در زمین یافت نمی شود. غلظت آن در پوسته زمین حدود یک گرم بر کیلوگرم است (مس را در حدود 0.06 گرم مقایسه کنید). در مواد معدنی ، فسفر به طور کلی به عنوان فسفات وجود دارد.

فسفر عنصری برای اولین بار در سال 1669 به عنوان فسفر سفید جدا شد. فسفر سفید هنگام قرار گرفتن در معرض اکسیژن درخشش ضعیفی از خود ساطع می کند -از این رو این نام ، برگرفته از اساطیر یونانی ، Φωσφόρος به معنی "حامل نور" ( لوسیفر لاتین ) ، با اشاره به " ستاره صبح " ، سیاره ناهید . اصطلاح " فسفرسانس " ، به معنای درخشش پس از روشنایی ، از این خاصیت فسفر نشأت می گیرد ، اگرچه این کلمه از آن به بعد برای فرایند فیزیکی متفاوتی استفاده می شود که درخشندگی ایجاد می کند. درخشش فسفر در اثر اکسیداسیون فسفر سفید (اما نه قرمز) ایجاد می شود - فرایندی که امروزه به آن کمیلومینسانس گفته می شود.به همراه با نیتروژن ، آرسنیک ، آنتیموان و بیسموت ، فسفر به عنوان یک ماده پنیکتوژن طبقه بندی می شود .

فسفر عنصری ضروری برای تداوم زندگی تا حد زیادی از طریق فسفات ها است ، ترکیباتی که حاوی یون فسفات ، PO 4 3− هستند . فسفات ها جزء DNA ، RNA ، ATP و فسفولیپیدها هستند ، ترکیبات پیچیده ای که برای سلول ها اساسی هستند . فسفر عنصری ابتدا از ادرار انسان و خاکستر استخوان جدا شدمنبع فسفات اولیه مهم بود. معادن فسفات حاوی فسیل هستند زیرا فسفات در رسوبات فسیل شده بقایای حیوانات و فضولات وجود دارد. سطوح پایین فسفات محدودیت مهمی برای رشد در برخی از سیستمهای آبزی است. اکثریت قریب به اتفاق ترکیبات فسفر استخراج شده به عنوان کود مصرف می شود . فسفات برای جایگزینی فسفر مورد نیاز گیاهان از خاک مورد نیاز است و تقاضای سالانه آن تقریباً دو برابر رشد جمعیت انسانی افزایش می یابد. سایر کاربردها شامل ترکیبات ارگانوفسفر در مواد شوینده ، آفت کش ها و عوامل عصبی است .

مشخصات

آلوتروپ ها

فسفر سفید در معرض هوا در تاریکی می درخشد
ساختار بلوری فسفر قرمز
ساختار بلوری فسفر سیاه

فسفر دارای آلوتروپهای متعددی است که خواص بسیار متنوعی از خود نشان می دهند. [8] دو رایج ترین آلوتروپ عبارتند از فسفر سفید و فسفر قرمز. [9]

ساختار P 4 مولکول ها، تعیین پراش الکترونی گاز . [10]

از دیدگاه کاربردها و ادبیات شیمیایی ، مهمترین شکل فسفر عنصری فسفر سفید است که اغلب به اختصار WP نامیده می شود. این یک جامد نرم و مومی است که از P چهار ضلعی تشکیل شده است
4
مولکول هایی که در آنها هر اتم با یک پیوند واحد به سه اتم دیگر متصل می شود. این پی
4
تتاهدرون همچنین در فسفر مایع و گازی تا دمای 800 درجه سانتیگراد (1470 درجه فارنهایت) وجود دارد که شروع به تجزیه به P می کند.
2
مولکول ها. [11] فسفر سفید در دو شکل کریستالی وجود دارد: α (آلفا) و β (بتا). در دمای اتاق ، شکل α پایدار است ، که بیشتر رایج است و دارای ساختار بلوری مکعبی است و در دمای 195.2 درجه سانتی گراد (-0.078- درجه سانتیگراد) به شکل β تبدیل می شود که دارای ساختار بلوری شش ضلعی است. این اشکال از نظر جهت گیری نسبی تشکیل دهنده P متفاوت 4 تتراهدرا. [12] [13]

فسفر سفید کمترین پایداری ، واکنشی ترین ، فرارترین ، کم چگال ترین و سمی ترین آلوتروپ ها است. فسفر سفید به تدریج به فسفر قرمز تبدیل می شود. این تغییر با نور و گرما تسریع می شود و نمونه های فسفر سفید تقریباً همیشه مقداری فسفر قرمز دارند و بر این اساس زرد به نظر می رسند. به همین دلیل ، فسفر سفید که پیر شده یا ناخالص است (به عنوان مثال ، WP درجه سلاح ، نه درجه آزمایشگاهی) نیز فسفر زرد نامیده می شود. وقتی فسفر سفید در معرض اکسیژن قرار می گیرد ، در تاریکی با رنگ سبز و آبی بسیار کم نور می درخشد. هنگام تماس با هوا بسیار قابل اشتعال و آتش سوز (خود اشتعال) است. از آنجا که فسفر سفید به دلیل دو طرفه بودن آن به عنوان افزودنی در آن استفاده می شودناپالم . بوی احتراق این شکل دارای بوی مشخص سیر است و نمونه ها معمولاً با " پنتوکسید فسفر سفید " ، که شامل P
4
O
10
تتراهدر با اکسیژن بین اتمهای فسفر و در راس آنها وارد شده است. فسفر سفید در آب نامحلول است اما در دی سولفید کربن محلول است. [14]

تجزیه حرارتی از P 4 در 1100 K می دهد دیفسفور ، P 2 . این گونه به صورت جامد یا مایع پایدار نیست. واحد دیمری شامل یک پیوند سه گانه و به N مشابه است 2 . همچنین می تواند به عنوان یک واسطه گذرا در محلول با ترمولیز معرفهای پیش ساز ارگانوفسفر تولید شود. [15] در درجه حرارت هنوز هم بالاتر، P 2 جدا به ص اتمی [14]

فسفر قرمز از نظر ساختار پلیمری است. می توان آن را به عنوان یک مشتق شده از P مشاهده 4 در آن یک پیوند PP شکسته است، و یک باند اضافی با چهار وجهی مجاور و در نتیجه زنجیره ای مانند ساختار تشکیل شده است. فسفر قرمز ممکن است با حرارت دادن فسفر سفید تا دمای 250 درجه سانتی گراد (482 درجه فارنهایت) یا قرار گرفتن فسفر سفید در معرض نور خورشید ایجاد شود. [16] فسفر پس از این درمان بی شکل است . با گرم شدن بیشتر ، این ماده متبلور می شود. از این نظر ، فسفر قرمز یک آلوتروپ نیست ، بلکه یک فاز میانی بین فسفر سفید و بنفش است و بیشتر خواص آن دارای طیف وسیعی از مقادیر است. به عنوان مثال ، فسفر قرمز تازه آماده شده بسیار واکنش پذیر است و در دمای 300 درجه سانتی گراد (572 درجه فارنهایت) مشتعل می شود ، [17]اگرچه پایدارتر از فسفر سفید است که در دمای 30 درجه سانتی گراد (86 درجه فارنهایت) مشتعل می شود. [18] پس از مدت زمان طولانی گرمایش یا ذخیره سازی ، رنگ تیره می شود (تصاویر مربوط به جعبه را ببینید). محصول حاصله پایدارتر است و خود به خود در هوا شعله ور نمی شود. [19]

فسفر بنفش شکلی از فسفر است که می تواند با بازپخت یک روزه فسفر قرمز در دمای 550 درجه سانتی گراد تولید شود. در سال 1865 ، هیتورف کشف کرد که وقتی فسفر از سرب مذاب تبلور مجدد پیدا می کند ، شکل قرمز/بنفش به دست می آید. بنابراین ، این شکل گاهی اوقات به عنوان "فسفر هیتورف" (یا بنفش یا فسفر فلزی α) شناخته می شود. [20]

فسفر سیاه کمترین آلوتروپ واکنشی است و از نظر ترمودینامیکی پایدار زیر 550 درجه سانتی گراد (1022 درجه فارنهایت) است. همچنین به عنوان فسفر β- فلزی شناخته می شود و ساختاری شبیه به گرافیت دارد . [21] [22] با گرم کردن فسفر سفید تحت فشارهای زیاد (حدود 12000 اتمسفر استاندارد یا 1.2 گیگاپاسکال) به دست می آید. همچنین می تواند در شرایط محیطی با استفاده از نمک های فلزی مانند جیوه به عنوان کاتالیزور تولید شود. [23] از نظر ظاهر ، خواص و ساختار ، شبیه گرافیت است ، سیاه و پوسته پوسته ، رسانای الکتریسیته ، و دارای صفحاتی از اتم های پیوند خورده است. [24]

شکل دیگر ، فسفر قرمز مایل به قرمز ، با اجازه تبخیر محلول فسفر سفید در دی سولفید کربن در زیر نور خورشید به دست می آید . [20]

خواص برخی از آلوتروپهای فسفر [8] [20]
فرم سفید (α) سفید (β) بنفش سیاه
تقارن مکعب بدن محور تریکلینیک مونوکلینیک اورترهومبیک
نماد پیرسون aP24 mP84 oS8
گروه فضایی من 4 3 متر P 1 شماره 2 P2/c شماره 13 Cmca شماره 64
تراکم (g/cm 3 ) 1.828 1.88 2.36 2.69
فاصله باند (eV) 2.1 1.5 0.34
ضریب شکست 1.8244 2.6 2.4

شیمی نوری

هنگامی که برای اولین بار جدا شد ، مشاهده شد که درخشش سبز ناشی از فسفر سفید برای مدتی در یک شیشه درپوش باقی می ماند ، اما سپس متوقف می شود. رابرت بویل در دهه 1680 آن را به "تضعیف" هوا نسبت داد. در واقع ، اکسیژن در حال مصرف است. در قرن 18 ، مشخص شد که در اکسیژن خالص ، فسفر به هیچ وجه نمی درخشد. [25] تنها طیف وسیعی از فشارهای جزئی وجود دارد که به آن فشار وارد می کند. گرما را می توان برای هدایت واکنش در فشارهای بیشتر اعمال کرد. [26]

در سال 1974 ، درخشش توسط RJ van Zee و AU Khan توضیح داده شد. [27] [28] واکنش با اکسیژن در سطح فسفر جامد (یا مایع) صورت می گیرد و مولکولهای کوتاه مدت HPO و P را تشکیل می دهد.
2
O
2
که هر دو نور مرئی ساطع می کنند. واکنش آهسته است و فقط مقدار بسیار کمی از واسطه ها برای ایجاد لومینسانس مورد نیاز است ، بنابراین زمان طولانی درخشش در یک شیشه درپوش دار ادامه می یابد.

از زمان کشف آن ، فسفر و فسفرسانس برای توصیف موادی که در تاریکی بدون سوختن می درخشند ، به طور شل مورد استفاده قرار گرفت. اگرچه اصطلاح فسفرسانس از فسفر مشتق شده است ، اما واکنشی که به فسفر درخشندگی می بخشد به درستی کمیلومنسانس (درخشش ناشی از واکنش شیمیایی سرد) نامیده می شود ، نه فسفرسانس (بازتاب مجدد نوری که قبلاً روی یک ماده سقوط کرده و آن را برانگیخته است). [29]

ایزوتوپ ها

23 ایزوتوپ فسفر شناخته شده است [30] که از آنها متفاوت است25
P
به47
پ
. [31] فقط31
P
پایدار است و بنابراین در فراوانی 100٪ وجود دارد. چرخش هسته ای نیم عدد صحیح و فراوانی زیاد 31 P طیف سنجی NMR فسفر-31 را به یک ابزار تحلیلی بسیار مفید در مطالعات نمونه های حاوی فسفر تبدیل می کند.

دو ایزوتوپ رادیواکتیو فسفر دارای نیمه عمر مناسب برای آزمایشات علمی بیولوژیکی هستند. اینها هستند:

  • 32
    P
    ، یک بتا -emitter (1.71 مگا الکترون ولت) با نیمه عمر 14.3 روز است که به طور معمول در آزمایشگاه زندگی علم استفاده می شود، در درجه اول به تولید نشاندار DNA و RNA پروب ، به عنوان مثال برای استفاده در نابودی شمالی یا نابودی جنوبی .
  • 33
    P
    ، یک انتشار دهنده بتا (0.25 MeV) با نیمه عمر 25.4 روز. از آن در آزمایشگاههای علوم زندگی در مواردی استفاده می شود که در آن انتشار بتا با انرژی کمتر مانند تعیین توالی DNA مفید است .

ذرات بتا با انرژی بالا از 32
P
به پوست و قرنیه و هر گونه دیگری نفوذ می کند32
P
خورده ، استنشاق یا جذب می شود و به آسانی در استخوان و اسیدهای نوکلئیک گنجانده می شود . به همین دلایل ، اداره ایمنی و بهداشت شغلی در ایالات متحده و م institutionsسسات مشابه دیگر کشورهای توسعه یافته به پرسنلی نیاز دارند که با آنها کار کنند32
از
کت آزمایشگاهی ، دستکش یکبار مصرف و عینک یا عینک ایمنی برای محافظت از چشم استفاده کنید و از کار مستقیم روی ظروف باز خودداری کنید. نظارت بر آلودگی شخصی ، لباس و سطح نیز لازم است. حفاظت نیاز به توجه ویژه دارد. انرژی زیاد ذرات بتا باعث انتشار ثانویه اشعه ایکس از طریق Bremsstrahlung (تابش ترمز) در مواد محافظ متراکم مانند سرب می شود. بنابراین ، تشعشع باید با مواد چگالی کم مانند اکریلیک یا پلاستیک دیگر ، آب یا (در صورت عدم نیاز به شفافیت) ، حتی چوب محافظت شود. [32]

وقوع

کائنات

در سال 2013 ، ستاره شناسان فسفر را در Cassiopeia A کشف کردند ، که تأیید کرد که این عنصر در ابرنواخترها به عنوان محصول جانبی هسته نوترکیب ابرنواختر تولید می شود . نسبت فسفر به آهن در مواد بقایای ابرنواختر می تواند تا 100 برابر بیشتر از راه شیری به طور کلی باشد. [33]

در سال 2020 ، اخترشناسان با تجزیه و تحلیل داده های ALMA و ROSINA از منطقه عظیم ستاره ساز AFGL 5142 ، مولکول های فسفر زا و نحوه حمل آنها را در دنباله دارها به زمین اولیه شناسایی کردند. [34] [35]

پوسته و منابع آلی

غلظت فسفر در پوسته زمین حدود یک گرم بر کیلوگرم است (مس را در حدود 0.06 گرم مقایسه کنید). در طبیعت رایگان یافت نمی شود ، اما به طور گسترده در بسیاری از مواد معدنی ، معمولاً به عنوان فسفات ها توزیع می شود. [9] سنگ فسفات معدنی ، که تا حدی از آپاتیت (گروهی از مواد معدنی ، به طور کلی ، فلوراید پنتاکالسیوم تریورتوفسفات (هیدروکسید)) ساخته شده است ، امروزه منبع اصلی تجاری این عنصر است. طبق گزارش زمین شناسی ایالات متحده (USGS) ، حدود 50 درصد از ذخایر فسفر جهانی در کشورهای عربی است. [36] 85٪ از ذخایر شناخته شده زمین در مراکش و ذخایر کوچکتر در چین است ،روسیه ، [37] فلوریدا ، آیداهو ، تنسی ، یوتا و جاهای دیگر. [38] آلبرایت و ویلسون در انگلستان و گیاه نیاگارا ، به عنوان مثال ، از سنگ فسفات در دهه 1890 و 1900 در تنسی فلوریدا و ایل دو کانتابل ( منابع فسفات جزیره گوانو ) استفاده می کردند. تا سال 1950 ، آنها از سنگ فسفات عمدتا در تنسی و شمال آفریقا استفاده می کردند. [39]

منابع ارگانیک ، یعنی ادرار ، خاکستر استخوان و (در قرن نوزدهم اخیر) گوانو ، از نظر تاریخی دارای اهمیت بودند اما موفقیت تجاری محدودی داشتند. [40] از آنجا که ادرار حاوی فسفر است ، دارای ویژگی های باروری است که امروزه هنوز در برخی از کشورها ، از جمله سوئد ، با استفاده از روش هایی برای استفاده مجدد از دفع ، مورد استفاده قرار می گیرد . برای این منظور می توان از ادرار به عنوان کود به شکل خالص یا بخشی از مخلوط شدن آن با آب به صورت فاضلاب یا لجن فاضلاب استفاده کرد .

ترکیبات

فسفر (V)

ساختار چهارضلعی P 4 O 10 و P 4 S 10 .

شایع ترین ترکیبات فسفر مشتقات فسفات (PO 4 3− ) ، آنیون چهار ضلعی است. [41] فسفات پایه مزدوج اسید فسفریک است که در مقیاس وسیع برای استفاده در کودها تولید می شود. از آنجا که اسید فسفریک سه گانه است ، مرحله به مرحله به سه پایه مزدوج تبدیل می شود:

H 3 PO 4 + H 2 O ⇌ H 3 O + + H 2 PO 4 -       K a1  = 7.25 × 10 −3
H 2 PO 4 - + H 2 O ⇌ H 3 O + + HPO 4 2 −       K a2  = 6.31 × 10 − 8
HPO 4 2− + H 2 O ⇌ H 3 O + + PO 4 3−        K a3  = 3.98 × 10 − 13

فسفات تمایل به تشکیل زنجیره و حلقه های حاوی پیوندهای POP را نشان می دهد. بسیاری از پلی فسفات ها شناخته شده اند ، از جمله ATP . پلی فسفاتها با کم شدن آب از فسفاتهای هیدروژن مانند HPO 4 2− و H 2 PO 4 - بوجود می آیند . به عنوان مثال ، پنتازدیم تری فسفات مهم صنعتی (همچنین به عنوان تری پلی فسفات سدیم ، STPP شناخته می شود) توسط این واکنش تراکم به صورت صنعتی توسط مگاتون تولید می شود :

2 Na 2 [(HO) PO 3 ] + Na [(HO) 2 PO 2 ] → Na 5 [O 3 P-OP (O) 2 -O-PO 3 ] + 2 H 2 O

پنتوکسید فسفر (P 4 O 10 ) انیدرید اسید اسید فسفریک است ، اما چندین واسطه بین این دو مشخص است. این جامد سفید مومی به شدت با آب واکنش نشان می دهد.

با کاتیون های فلزی ، فسفات انواع نمک ها را تشکیل می دهد. این جامدات پلیمری هستند و دارای اتصالات POM هستند. هنگامی که کاتیون فلزی دارای بار 2+ یا 3+ است ، نمک ها به طور کلی نامحلول هستند ، بنابراین به عنوان مواد معدنی معمولی وجود دارند. بسیاری از نمک های فسفات از هیدروژن فسفات (HPO 4 2− ) مشتق می شوند .

PCl 5 و PF 5 ترکیبات متداول هستند. PF 5 یک گاز بی رنگ است و مولکولها دارای هندسه دو ضلعی سه ضلعی هستند. PCl 5 یک ماده جامد بی رنگ است که دارای فرمول یونی PCl 4 + PCl 6 است - اما هنگام ذوب شدن یا در مرحله بخار ، هندسه دو ضلعی سه ضلعی را اتخاذ می کند . [14] PBr 5 یک ماده جامد ناپایدار است که به صورت PBr 4 + Br فرموله شده است - و PI 5 شناخته شده نیست. [14] پنتاکلراید و پنتافلوراید اسیدهای لوئیس هستند. With fluoride, PF5 forms PF6, an anion that is isoelectronic with SF6. The most important oxyhalide is phosphorus oxychloride, (POCl3), which is approximately tetrahedral.

Before extensive computer calculations were feasible, it was thought that bonding in phosphorus(V) compounds involved d orbitals. Computer modeling of molecular orbital theory indicates that this bonding involves only s- and p-orbitals.[42]

Phosphorus(III)

All four symmetrical trihalides are well known: gaseous PF3, the yellowish liquids PCl3 and PBr3, and the solid PI3. These materials are moisture sensitive, hydrolysing to give phosphorous acid. The trichloride, a common reagent, is produced by chlorination of white phosphorus:

P4 + 6 Cl2 → 4 PCl3

The trifluoride is produced from the trichloride by halide exchange. PF3 is toxic because it binds to haemoglobin.

Phosphorus(III) oxide, P4O6 (also called tetraphosphorus hexoxide) is the anhydride of P(OH)3, the minor tautomer of phosphorous acid. The structure of P4O6 is like that of P4O10 without the terminal oxide groups.

Phosphorus(I) and phosphorus(II)

A stable diphosphene, a derivative of phosphorus(I).

These compounds generally feature P–P bonds.[14] Examples include catenated derivatives of phosphine and organophosphines. Compounds containing P=P double bonds have also been observed, although they are rare.

Phosphides and phosphines

Phosphides arise by reaction of metals with red phosphorus. The alkali metals (group 1) and alkaline earth metals can form ionic compounds containing the phosphide ion, P3−. These compounds react with water to form phosphine. Other phosphides, for example Na3P7, are known for these reactive metals. With the transition metals as well as the monophosphides there are metal-rich phosphides, which are generally hard refractory compounds with a metallic lustre, and phosphorus-rich phosphides which are less stable and include semiconductors.[14] Schreibersite is a naturally occurring metal-rich phosphide found in meteorites. The structures of the metal-rich and phosphorus-rich phosphides can be complex.

Phosphine (PH3) and its organic derivatives (PR3) are structural analogues of ammonia (NH3), but the bond angles at phosphorus are closer to 90° for phosphine and its organic derivatives. It is an ill-smelling, toxic compound. Phosphorus has an oxidation number of −3 in phosphine. Phosphine is produced by hydrolysis of calcium phosphide, Ca3P2. Unlike ammonia, phosphine is oxidised by air. Phosphine is also far less basic than ammonia. Other phosphines are known which contain chains of up to nine phosphorus atoms and have the formula PnHn+2.[14] The highly flammable gas diphosphine (P2H4) is an analogue of hydrazine.

Oxoacids

Phosphorous oxoacids are extensive, often commercially important, and sometimes structurally complicated. They all have acidic protons bound to oxygen atoms, some have nonacidic protons that are bonded directly to phosphorus and some contain phosphorus - phosphorus bonds.[14] Although many oxoacids of phosphorus are formed, only nine are commercially important, and three of them, hypophosphorous acid, phosphorous acid, and phosphoric acid, are particularly important.

Oxidation state Formula Name Acidic protons Compounds
+1 HH2PO2 hypophosphorous acid 1 acid, salts
+3 H2HPO3 phosphorous acid 2 acid, salts
+3 HPO2 metaphosphorous acid 1 salts
+3 H3PO3 (ortho)phosphorous acid 3 acid, salts
+4 H4P2O6 hypophosphoric acid 4 acid, salts
+5 (HPO3)n metaphosphoric acids n salts (n = 3,4,6)
+5 H(HPO3)nOH polyphosphoric acids n+2 acids, salts (n = 1-6)
+5 H5P3O10 tripolyphosphoric acid 3 salts
+5 H4P2O7 pyrophosphoric acid 4 acid, salts
+5 H3PO4 (ortho)phosphoric acid 3 acid, salts

Nitrides

The PN molecule is considered unstable, but is a product of crystalline phosphorus nitride decomposition at 1100 K. Similarly, H2PN is considered unstable, and phosphorus nitride halogens like F2PN, Cl2PN, Br2PN, and I2PN oligomerise into cyclic Polyphosphazenes. For example, compounds of the formula (PNCl2)n exist mainly as rings such as the trimer hexachlorophosphazene. The phosphazenes arise by treatment of phosphorus pentachloride with ammonium chloride:

PCl5 + NH4Cl → 1/n (NPCl2)n + 4 HCl

When the chloride groups are replaced by alkoxide (RO), a family of polymers is produced with potentially useful properties.[43]

Sulfides

Phosphorus forms a wide range of sulfides, where the phosphorus can be in P(V), P(III) or other oxidation states. The three-fold symmetric P4S3 is used in strike-anywhere matches. P4S10 and P4O10 have analogous structures.[44] Mixed oxyhalides and oxyhydrides of phosphorus(III) are almost unknown.

Organophosphorus compounds

Compounds with P-C and P-O-C bonds are often classified as organophosphorus compounds. They are widely used commercially. The PCl3 serves as a source of P3+ in routes to organophosphorus(III) compounds. For example, it is the precursor to triphenylphosphine:

PCl3 + 6 Na + 3 C6H5Cl → P(C6H5)3 + 6 NaCl

Treatment of phosphorus trihalides with alcohols and phenols gives phosphites, e.g. triphenylphosphite:

PCl3 + 3 C6H5OH → P(OC6H5)3 + 3 HCl

Similar reactions occur for phosphorus oxychloride, affording triphenylphosphate:

OPCl3 + 3 C6H5OH → OP(OC6H5)3 + 3 HCl

History

Etymology

The name Phosphorus in Ancient Greece was the name for the planet Venus and is derived from the Greek words (φῶς = light, φέρω = carry), which roughly translates as light-bringer or light carrier.[16] (In Greek mythology and tradition, Augerinus (Αυγερινός = morning star, still in use today), Hesperus or Hesperinus (΄Εσπερος or Εσπερινός or Αποσπερίτης = evening star, still in use today) and Eosphorus (Εωσφόρος = dawnbearer, not in use for the planet after Christianity) are close homologues, and also associated with Phosphorus-the-morning-star).

According to the Oxford English Dictionary, the correct spelling of the element is phosphorus. The word phosphorous is the adjectival form of the P3+ valence: so, just as sulfur forms sulfurous and sulfuric compounds, phosphorus forms phosphorous compounds (e.g., phosphorous acid) and P5+ valence phosphoric compounds (e.g., phosphoric acids and phosphates).

Discovery

The discovery of phosphorus, the first element to be discovered that was not known since ancient times,[45] is credited to the German alchemist Hennig Brand in 1669, although other chemists might have discovered phosphorus around the same time.[46] Brand experimented with urine, which contains considerable quantities of dissolved phosphates from normal metabolism.[16] Working in Hamburg, Brand attempted to create the fabled philosopher's stone through the distillation of some salts by evaporating urine, and in the process produced a white material that glowed in the dark and burned brilliantly. It was named phosphorus mirabilis ("miraculous bearer of light").[47]

Brand's process originally involved letting urine stand for days until it gave off a terrible smell. Then he boiled it down to a paste, heated this paste to a high temperature, and led the vapours through water, where he hoped they would condense to gold. Instead, he obtained a white, waxy substance that glowed in the dark. Brand had discovered phosphorus. We now know that Brand produced ammonium sodium hydrogen phosphate, (NH
4
)NaHPO
4
. While the quantities were essentially correct (it took about 1,100 litres [290 US gal] of urine to make about 60 g of phosphorus), it was unnecessary to allow the urine to rot first. Later scientists discovered that fresh urine yielded the same amount of phosphorus.[29]

Brand at first tried to keep the method secret,[48] but later sold the recipe for 200 thalers to D. Krafft from Dresden,[16] who could now make it as well, and toured much of Europe with it, including England, where he met with Robert Boyle. The secret that it was made from urine leaked out and first Johann Kunckel (1630–1703) in Sweden (1678) and later Boyle in London (1680) also managed to make phosphorus, possibly with the aid of his assistant, Ambrose Godfrey-Hanckwitz, who later made a business of the manufacture of phosphorus.

Boyle states that Krafft gave him no information as to the preparation of phosphorus other than that it was derived from "somewhat that belonged to the body of man". This gave Boyle a valuable clue, so that he, too, managed to make phosphorus, and published the method of its manufacture.[16] Later he improved Brand's process by using sand in the reaction (still using urine as base material),

4 NaPO
3
+ 2 SiO
2
+ 10 C → 2 Na
2
SiO
3
+ 10 CO + P
4

Robert Boyle was the first to use phosphorus to ignite sulfur-tipped wooden splints, forerunners of our modern matches, in 1680.[49]

Phosphorus was the 13th element to be discovered. Because of its tendency to spontaneously combust when left alone in air, it is sometimes referred to as "the Devil's element".[50]

Bone ash and guano

Guano mining in the Central Chincha Islands, ca. 1860.

In 1769, Johan Gottlieb Gahn and Carl Wilhelm Scheele showed that calcium phosphate (Ca
3
(PO
4
)
2
) is found in bones, and they obtained elemental phosphorus from bone ash. Antoine Lavoisier recognised phosphorus as an element in 1777.[51] Bone ash was the major source of phosphorus until the 1840s. The method started by roasting bones, then employed the use of clay retorts encased in a very hot brick furnace to distill out the highly toxic elemental phosphorus product.[52] Alternately, precipitated phosphates could be made from ground-up bones that had been de-greased and treated with strong acids. White phosphorus could then be made by heating the precipitated phosphates, mixed with ground coal or charcoal in an iron pot, and distilling off phosphorus vapour in a retort.[53] Carbon monoxide and other flammable gases produced during the reduction process were burnt off in a flare stack.

In the 1840s, world phosphate production turned to the mining of tropical island deposits formed from bird and bat guano (see also Guano Islands Act). These became an important source of phosphates for fertiliser in the latter half of the 19th century.[54]

Phosphate rock

Phosphate rock, which usually contains calcium phosphate, was first used in 1850 to make phosphorus, and following the introduction of the electric arc furnace by James Burgess Readman in 1888[55] (patented 1889),[56] elemental phosphorus production switched from the bone-ash heating, to electric arc production from phosphate rock. After the depletion of world guano sources about the same time, mineral phosphates became the major source of phosphate fertiliser production. Phosphate rock production greatly increased after World War II, and remains the primary global source of phosphorus and phosphorus chemicals today. See the article on peak phosphorus for more information on the history and present state of phosphate mining. Phosphate rock remains a feedstock in the fertiliser industry, where it is treated with sulfuric acid to produce various "superphosphate" fertiliser products.

Incendiaries

White phosphorus was first made commercially in the 19th century for the match industry. This used bone ash for a phosphate source, as described above. The bone-ash process became obsolete when the submerged-arc furnace for phosphorus production was introduced to reduce phosphate rock.[57][58] The electric furnace method allowed production to increase to the point where phosphorus could be used in weapons of war.[27][59] In World War I, it was used in incendiaries, smoke screens and tracer bullets.[59] A special incendiary bullet was developed to shoot at hydrogen-filled Zeppelins over Britain (hydrogen being highly flammable).[59] During World War II, Molotov cocktails made of phosphorus dissolved in petrol were distributed in Britain to specially selected civilians within the British resistance operation, for defence; and phosphorus incendiary bombs were used in war on a large scale. Burning phosphorus is difficult to extinguish and if it splashes onto human skin it has horrific effects.[14]

Early matches used white phosphorus in their composition, which was dangerous due to its toxicity. Murders, suicides and accidental poisonings resulted from its use. (An apocryphal tale tells of a woman attempting to murder her husband with white phosphorus in his food, which was detected by the stew's giving off luminous steam).[27] In addition, exposure to the vapours gave match workers a severe necrosis of the bones of the jaw, known as "phossy jaw". When a safe process for manufacturing red phosphorus was discovered, with its far lower flammability and toxicity, laws were enacted, under the Berne Convention (1906), requiring its adoption as a safer alternative for match manufacture.[60] The toxicity of white phosphorus led to discontinuation of its use in matches.[61] The Allies used phosphorus incendiary bombs in World War II to destroy Hamburg, the place where the "miraculous bearer of light" was first discovered.[47]

Production

Mining of phosphate rock in Nauru

Most production of phosphorus-bearing material is for agriculture fertilisers. For this purpose, phosphate minerals are converted to phosphoric acid. It follows two distinct chemical routes, the main one being treatment of phosphate minerals with sulfuric acid. The other process utilises white phosphorus, which may be produced by reaction and distillation from very low grade phosphate sources. The white phosphorus is then oxidised to phosphoric acid and subsequently neutralised with base to give phosphate salts. Phosphoric acid produced from white phosphorus is relatively pure and is the main route for the production of phosphates for all purposes, including detergent production.

In the early 1990s, Albright and Wilson's purified wet phosphoric acid business was being adversely affected by phosphate rock sales by China and the entry of their long-standing Moroccan phosphate suppliers into the purified wet phosphoric acid business.[62]

Peak phosphorus

In 2017, the USGS estimated 68 billion tons of world reserves, where reserve figures refer to the amount assumed recoverable at current market prices; 0.261 billion tons were mined in 2016.[63] Critical to contemporary agriculture, its annual demand is rising nearly twice as fast as the growth of the human population.[37]

The production of phosphorus may have peaked already (as per 2011), leading to the possibility of global shortages by 2040.[64] In 2007, at the rate of consumption, the supply of phosphorus was estimated to run out in 345 years.[65] However, some scientists now believe that a "peak phosphorus" will occur in 30 years and that "At current rates, reserves will be depleted in the next 50 to 100 years."[66] Cofounder of Boston-based investment firm and environmental foundation Jeremy Grantham wrote in Nature in November 2012 that consumption of the element "must be drastically reduced in the next 20-40 years or we will begin to starve."[37][67] According to N.N. Greenwood and A. Earnshaw, authors of the textbook, Chemistry of the Elements, however, phosphorus comprises about 0.1% by mass of the average rock, and consequently the Earth's supply is vast, although dilute.[14]

Elemental phosphorus

Presently, about 1,000,000 short tons (910,000 t) of elemental phosphorus is produced annually. Calcium phosphate (phosphate rock), mostly mined in Florida and North Africa, can be heated to 1,200–1,500 °C with sand, which is mostly SiO
2
, and coke (refined coal) to produce vaporised P
4
. The product is subsequently condensed into a white powder under water to prevent oxidation by air. Even under water, white phosphorus is slowly converted to the more stable red phosphorus allotrope. The chemical equation for this process when starting with fluoroapatite, a common phosphate mineral, is:

4 Ca5(PO4)3F + 18 SiO2 + 30 C → 3 P4 + 30 CO + 18 CaSiO3 + 2 CaF2

Side products from this process include ferrophosphorus, a crude form of Fe2P, resulting from iron impurities in the mineral precursors. The silicate slag is a useful construction material. The fluoride is sometimes recovered for use in water fluoridation. More problematic is a "mud" containing significant amounts of white phosphorus. Production of white phosphorus is conducted in large facilities in part because it is energy intensive. The white phosphorus is transported in molten form. Some major accidents have occurred during transportation; train derailments at Brownston, Nebraska and Miamisburg, Ohio led to large fires. The worst incident in recent times was an environmental contamination in 1968 when the sea was polluted from spillage and/or inadequately treated sewage from a white phosphorus plant at Placentia Bay, Newfoundland.[68]

Another process by which elemental phosphorus is extracted includes calcining tricalcium phosphate at high temperatures (1500 °C):[69]

2 Ca3(PO4)2 + 6 SiO2 + 10 C → 6 CaSiO3 + 10 CO + P4

Historically, before the development of mineral-based extractions, white phosphorus was isolated on an industrial scale from bone ash.[70] In this process, the tricalcium phosphate in bone ash is converted to monocalcium phosphate with sulfuric acid:

Ca3(PO4)2 + 2 H2SO4 → Ca(H2PO4)2 + 2 CaSO4

Monocalcium phosphate is then dehydrated to the corresponding metaphosphate:

Ca(H2PO4)2 → Ca(PO3)2 + 2 H2O

When ignited to a white heat (~1300C) with charcoal, calcium metaphosphate yields two-thirds of its weight of white phosphorus while one-third of the phosphorus remains in the residue as calcium orthophosphate:

3 Ca(PO3)2 + 10 C → Ca3(PO4)2 + 10 CO + P4

Applications

Fertiliser

Phosphorus is an essential plant nutrient (the most often limiting nutrient, after nitrogen),[71] and the bulk of all phosphorus production is in concentrated phosphoric acids for agriculture fertilisers, containing as much as 70% to 75% P2O5. That led to large increase in phosphate (PO43−) production in the second half of the 20th century.[37] Artificial phosphate fertilisation is necessary because phosphorus is essential to all living organisms; it is involved in energy transfers, strength of root and stems, photosynthesis, the expansion of plant roots, formation of seeds and flowers, and other important factors effecting overall plant health and genetics.[71]

Natural phosphorus-bearing compounds are mostly inaccessible to plants because of the low solubility and mobility in soil.[72] Most phosphorus is very stable in the soil minerals or organic matter of the soil. Even when phosphorus is added in manure or fertilizer it can become fixed in the soil. Therefore, the natural cycle of phosphorus is very slow. Some of the fixed phosphorus is released again over time, sustaining wild plant growth, however, more is needed to sustain intensive cultivation of crops.[73] Fertiliser is often in the form of superphosphate of lime, a mixture of calcium dihydrogen phosphate (Ca(H2PO4)2), and calcium sulfate dihydrate (CaSO4·2H2O) produced reacting sulfuric acid and water with calcium phosphate.

Processing phosphate minerals with sulfuric acid for obtaining fertiliser is so important to the global economy that this is the primary industrial market for sulfuric acid and the greatest industrial use of elemental sulfur.[74]

Widely used compounds Use
Ca(H2PO4)2·H2O Baking powder and fertilisers
CaHPO4·2H2O Animal food additive, toothpowder
H3PO4 Manufacture of phosphate fertilisers
PCl3 Manufacture of POCl3 and pesticides
POCl3 Manufacture of plasticiser
P4S10 Manufacturing of additives and pesticides
Na5P3O10 Detergents

Organophosphorus

White phosphorus is widely used to make organophosphorus compounds through intermediate phosphorus chlorides and two phosphorus sulfides, phosphorus pentasulfide and phosphorus sesquisulfide.[75] Organophosphorus compounds have many applications, including in plasticisers, flame retardants, pesticides, extraction agents, nerve agents and water treatment.[14][76]

Metallurgical aspects

Phosphorus is also an important component in steel production, in the making of phosphor bronze, and in many other related products.[77][78] Phosphorus is added to metallic copper during its smelting process to react with oxygen present as an impurity in copper and to produce phosphorus-containing copper (CuOFP) alloys with a higher hydrogen embrittlement resistance than normal copper.[79]

Matches

Match striking surface made of a mixture of red phosphorus, glue and ground glass. The glass powder is used to increase the friction.

The first striking match with a phosphorus head was invented by Charles Sauria in 1830. These matches (and subsequent modifications) were made with heads of white phosphorus, an oxygen-releasing compound (potassium chlorate, lead dioxide, or sometimes nitrate), and a binder. They were poisonous to the workers in manufacture,[80] sensitive to storage conditions, toxic if ingested, and hazardous when accidentally ignited on a rough surface.[81][82] Production in several countries was banned between 1872 and 1925.[83] The international Berne Convention, ratified in 1906, prohibited the use of white phosphorus in matches.

In consequence, phosphorous matches were gradually replaced by safer alternatives. Around 1900 french chemists Henri Sévène and Emile David Cahen invented the modern strike-anywhere match, wherein the white phosphorus was replaced by phosphorus sesquisulfide (P4S3), a non-toxic and non-pyrophoric compound that ignites under friction. For a time these safer strike-anywhere matches were quite popular but in the long run they were superseded by the modern safety match.

Safety matches are very difficult to ignite on any surface other than a special striker strip. The strip contains non-toxic red phosphorus and the match head potassium chlorate, an oxygen-releasing compound. When struck, small amounts of abrasion from match head and striker strip are mixed intimately to make a small quantity of Armstrong's mixture, a very touch sensitive composition. The fine powder ignites immediately and provides the initial spark to set off the match head. Safety matches separate the two components of the ignition mixture until the match is struck. This is the key safety advantage as it prevents accidental ignition. Nonetheless, safety matches, invented in 1844 by Gustaf Erik Pasch and market ready by the 1860s, didn't gain consumer acceptance until the prohibition of white phosphorus. Using a dedicated striker strip was considered clumsy.[17][75][84]

Water softening

Sodium tripolyphosphate made from phosphoric acid is used in laundry detergents in some countries, but banned for this use in others.[19] This compound softens the water to enhance the performance of the detergents and to prevent pipe/boiler tube corrosion.[85]

Miscellaneous

Biological role

Inorganic phosphorus in the form of the phosphate PO3−
4
is required for all known forms of life.[89] Phosphorus plays a major role in the structural framework of DNA and RNA. Living cells use phosphate to transport cellular energy with adenosine triphosphate (ATP), necessary for every cellular process that uses energy. ATP is also important for phosphorylation, a key regulatory event in cells. Phospholipids are the main structural components of all cellular membranes. Calcium phosphate salts assist in stiffening bones.[14] Biochemists commonly use the abbreviation "Pi" to refer to inorganic phosphate.[90]

Every living cell is encased in a membrane that separates it from its surroundings. Cellular membranes are composed of a phospholipid matrix and proteins, typically in the form of a bilayer. Phospholipids are derived from glycerol with two of the glycerol hydroxyl (OH) protons replaced by fatty acids as an ester, and the third hydroxyl proton has been replaced with phosphate bonded to another alcohol.[91]

An average adult human contains about 0.7 kg of phosphorus, about 85–90% in bones and teeth in the form of apatite, and the remainder in soft tissues and extracellular fluids (~1%). The phosphorus content increases from about 0.5 weight% in infancy to 0.65-1.1 weight% in adults. Average phosphorus concentration in the blood is about 0.4 g/L, about 70% of that is organic and 30% inorganic phosphates.[92] An adult with healthy diet consumes and excretes about 1-3 grams of phosphorus per day, with consumption in the form of inorganic phosphate and phosphorus-containing biomolecules such as nucleic acids and phospholipids; and excretion almost exclusively in the form of phosphate ions such as H
2
PO
4
and HPO2−
4
. Only about 0.1% of body phosphate circulates in the blood, paralleling the amount of phosphate available to soft tissue cells.

Bone and teeth enamel

The main component of bone is hydroxyapatite as well as amorphous forms of calcium phosphate, possibly including carbonate. Hydroxyapatite is the main component of tooth enamel. Water fluoridation enhances the resistance of teeth to decay by the partial conversion of this mineral to the still harder material called fluoroapatite:[14]

Ca
5
(PO
4
)
3
OH
+ F
Ca
5
(PO
4
)
3
F
+ OH

Phosphorus deficiency

In medicine, phosphate deficiency syndrome may be caused by malnutrition, by failure to absorb phosphate, and by metabolic syndromes that draw phosphate from the blood (such as in refeeding syndrome after malnutrition[93]) or passing too much of it into the urine. All are characterised by hypophosphatemia, which is a condition of low levels of soluble phosphate levels in the blood serum and inside the cells. Symptoms of hypophosphatemia include neurological dysfunction and disruption of muscle and blood cells due to lack of ATP. Too much phosphate can lead to diarrhoea and calcification (hardening) of organs and soft tissue, and can interfere with the body's ability to use iron, calcium, magnesium, and zinc.[94]

Phosphorus is an essential macromineral for plants, which is studied extensively in edaphology to understand plant uptake from soil systems. Phosphorus is a limiting factor in many ecosystems; that is, the scarcity of phosphorus limits the rate of organism growth. An excess of phosphorus can also be problematic, especially in aquatic systems where eutrophication sometimes leads to algal blooms.[37]

Nutrition

Dietary recommendations

The U.S. Institute of Medicine (IOM) updated Estimated Average Requirements (EARs) and Recommended Dietary Allowances (RDAs) for phosphorus in 1997. If there is not sufficient information to establish EARs and RDAs, an estimate designated Adequate Intake (AI) is used instead. The current EAR for phosphorus for people ages 19 and up is 580 mg/day. The RDA is 700 mg/day. RDAs are higher than EARs so as to identify amounts that will cover people with higher than average requirements. RDA for pregnancy and lactation are also 700 mg/day. For children ages 1–18 years the RDA increases with age from 460 to 1250 mg/day. As for safety, the IOM sets Tolerable upper intake levels (ULs) for vitamins and minerals when evidence is sufficient. In the case of phosphorus the UL is 4000 mg/day. Collectively the EARs, RDAs, AIs and ULs are referred to as Dietary Reference Intakes (DRIs).[95]

The European Food Safety Authority (EFSA) refers to the collective set of information as Dietary Reference Values, with Population Reference Intake (PRI) instead of RDA, and Average Requirement instead of EAR. AI and UL defined the same as in United States. For people ages 15 and older, including pregnancy and lactation, the AI is set at 550 mg/day. For children ages 4–10 years the AI is 440 mg/day, for ages 11–17 640 mg/day. These AIs are lower than the U.S RDAs. In both systems, teenagers need more than adults.[96] The European Food Safety Authority reviewed the same safety question and decided that there was not sufficient information to set a UL.[97]

For U.S. food and dietary supplement labeling purposes the amount in a serving is expressed as a percent of Daily Value (%DV). For phosphorus labeling purposes 100% of the Daily Value was 1000 mg, but as of May 27, 2016 it was revised to 1250 mg to bring it into agreement with the RDA.[98][99] Compliance with the updated labeling regulations was required by 1 January 2020 for manufacturers with US$10 million or more in annual food sales, and by 1 January 2021 for manufacturers with lower volume food sales.[100][101] A table of the old and new adult daily values is provided at Reference Daily Intake.

Food sources

The main food sources for phosphorus are the same as those containing protein, although proteins do not contain phosphorus. For example, milk, meat, and soya typically also have phosphorus. As a rule, if a diet has sufficient protein and calcium, the amount of phosphorus is probably sufficient.[102]

Precautions

Phosphorus explosion

Organic compounds of phosphorus form a wide class of materials; many are required for life, but some are extremely toxic. Fluorophosphate esters are among the most potent neurotoxins known. A wide range of organophosphorus compounds are used for their toxicity as pesticides (herbicides, insecticides, fungicides, etc.) and weaponised as nerve agents against enemy humans. Most inorganic phosphates are relatively nontoxic and essential nutrients.[14]

The white phosphorus allotrope presents a significant hazard because it ignites in air and produces phosphoric acid residue. Chronic white phosphorus poisoning leads to necrosis of the jaw called "phossy jaw". White phosphorus is toxic, causing severe liver damage on ingestion and may cause a condition known as "Smoking Stool Syndrome".[103]

In the past, external exposure to elemental phosphorus was treated by washing the affected area with 2% copper sulfate solution to form harmless compounds that are then washed away. According to the recent US Navy's Treatment of Chemical Agent Casualties and Conventional Military Chemical Injuries: FM8-285: Part 2 Conventional Military Chemical Injuries, "Cupric (copper(II)) sulfate has been used by U.S. personnel in the past and is still being used by some nations. However, copper sulfate is toxic and its use will be discontinued. Copper sulfate may produce kidney and cerebral toxicity as well as intravascular hemolysis."[104]

The manual suggests instead "a bicarbonate solution to neutralise phosphoric acid, which will then allow removal of visible white phosphorus. Particles often can be located by their emission of smoke when air strikes them, or by their phosphorescence in the dark. In dark surroundings, fragments are seen as luminescent spots. Promptly debride the burn if the patient's condition will permit removal of bits of WP (white phosphorus) that might be absorbed later and possibly produce systemic poisoning. DO NOT apply oily-based ointments until it is certain that all WP has been removed. Following complete removal of the particles, treat the lesions as thermal burns."[note 1][citation needed] As white phosphorus readily mixes with oils, any oily substances or ointments are not recommended until the area is thoroughly cleaned and all white phosphorus removed.

People can be exposed to phosphorus in the workplace by inhalation, ingestion, skin contact, and eye contact. The Occupational Safety and Health Administration (OSHA) has set the phosphorus exposure limit (Permissible exposure limit) in the workplace at 0.1 mg/m3 over an 8-hour workday. The National Institute for Occupational Safety and Health (NIOSH) has set a Recommended exposure limit (REL) of 0.1 mg/m3 over an 8-hour workday. At levels of 5 mg/m3, phosphorus is immediately dangerous to life and health.[105]

US DEA List I status

Phosphorus can reduce elemental iodine to hydroiodic acid, which is a reagent effective for reducing ephedrine or pseudoephedrine to methamphetamine.[106] For this reason, red and white phosphorus were designated by the United States Drug Enforcement Administration as List I precursor chemicals under 21 CFR 1310.02 effective on November 17, 2001.[107] In the United States, handlers of red or white phosphorus are subject to stringent regulatory controls.[107][108][109]

See also

Notes

  1. ^ WP, (white phosphorus), exhibits chemoluminescence upon exposure to air and if there is any WP in the wound, covered by tissue or fluids such as blood serum, it will not glow until it is exposed to air, which requires a very dark room and dark-adapted eyes to see clearly

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