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In July 2021, Microsoft announced it will start selling subscriptions to virtualized Windows desktops as part of a new "Windows 365" service in the following month. It is not a standalone version of Windows, but a web service that provides access to Windows 10 and Windows 11 built on top of Azure Virtual Desktop. The new service will allow for cross-platform usage, aiming to make the operating system available for both Apple and Android users. The subscription service will be accessible through any operating system with a web browser. The new service is an attempt at capitalizing on the growing trend, fostered during the COVID-19 pandemic, for businesses to adopt a hybrid remote work environment, in which "employees split their time between the office and home". As the service will be accessible through web browsers, Microsoft will be able to bypass the need to publish the service through Google Play or the Apple App Store.
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Interface languages for the operating system are free for download, but some languages are limited to certain editions of Windows. Language Interface Packs (LIPs) are redistributable and may be downloaded from Microsoft's Download Center and installed for any edition of Windows (XP or later) they translate most, but not all, of the Windows interface, and require a certain base language (the language which Windows originally shipped with). This is used for most languages in emerging markets. Full Language Packs, which translates the complete operating system, are only available for specific editions of Windows (Ultimate and Enterprise editions of Windows Vista and 7, and all editions of Windows 8, 8.1 and RT except Single Language). They do not require a specific base language, and are commonly used for more popular languages such as French or Chinese. These languages cannot be downloaded through the Download Center, but available as optional updates through the Windows Update service (except Windows 8).
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Windows 8 and Windows Server 2012 introduces a new Language Control Panel where both the interface and input languages can be simultaneously changed, and language packs, regardless of type, can be downloaded from a central location. The PC Settings app in Windows 8.1 and Windows Server 2012 R2 also includes a counterpart settings page for this. Changing the interface language also changes the language of preinstalled Windows Store apps (such as Mail, Maps and News) and certain other Microsoft-developed apps (such as Remote Desktop). The above limitations for language packs are however still in effect, except that full language packs can be installed for any edition except Single Language, which caters to emerging markets.
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Including personal computers of all kinds (e.g., desktops, laptops, mobile devices, and game consoles), Windows OSes accounted for 32.67% of usage share in August 2021, compared to Android (highest, at 46.03%), iOS's 13.76%, iPadOS's 2.81%, and macOS's 2.51%, according to Net Applications and 30.73% of usage share in August 2021, compared to Android (highest, at 42.56%), iOS/iPadOS's 16.53%, and macOS's 6.51%, according to StatCounter.
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While the Windows 9x series offered the option of having profiles for multiple users, it had no concept of access privileges, allowing any user to edit others' files. In addition, while it ran separate 32-bit applications in separate address spaces, protecting an application's code and data from being read or written by another application, it did not protect the first megabyte of memory from userland applications for compatibility reasons. This area of memory contains code critical to the functioning of the operating system, and by writing into this area of memory an application can crash or freeze the operating system. This was a source of instability as faulty applications could accidentally write into this region, potentially corrupting important operating system memory, which usually resulted in some form of system error and halt.
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Windows NT was far more secure, implementing access privileges and full memory protection, and, while 32-bit programs meeting the DoD's C2 security rating, yet these advantages were nullified by the fact that, prior to Windows Vista, the default user account created during the setup process was an administrator account; the user, and any program the user launched, had full access to the machine. Though Windows XP did offer an option of turning administrator accounts into limited accounts, the majority of home users did not do so, partially due to the number of programs which required administrator rights to function properly. As a result, most home users still ran as administrator all the time. These architectural flaws, combined with Windows's very high popularity, made Windows a frequent target of computer worm and virus writers.
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Windows Vista introduced a privilege elevation system called User Account Control. When logging in as a standard user, a logon session is created and a token containing only the most basic privileges is assigned. In this way, the new logon session is incapable of making changes that would affect the entire system. When logging in as a user in the Administrators group, two separate tokens are assigned. The first token contains all privileges typically awarded to an administrator, and the second is a restricted token similar to what a standard user would receive. User applications, including the Windows shell, are then started with the restricted token, resulting in a reduced privilege environment even under an Administrator account. When an application requests higher privileges or "Run as administrator" is clicked, UAC will prompt for confirmation and, if consent is given (including administrator credentials if the account requesting the elevation is not a member of the administrators group), start the process using the unrestricted token.
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In August 2019, computer experts reported that the BlueKeep security vulnerability, , that potentially affects older unpatched Windows versions via the program's Remote Desktop Protocol, allowing for the possibility of remote code execution, may now include related flaws, collectively named "DejaBlue", affecting newer Windows versions (i.e., Windows 7 and all recent versions) as well. In addition, experts reported a Microsoft security vulnerability, , based on legacy code involving Microsoft CTF and ctfmon (ctfmon.exe), that affects all Windowsversions from Windows XP to the then most recent Windows 10 versions; a patch to correct the flaw is currently available.
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All Windows versions from Windows NT 3 have been based on a file system permission system referred to as AGDLP(Accounts, Global, Domain Local, Permissions) in which file permissions are applied to the file/folder in the form of a 'local group' which then has other 'global groups' as members. These global groups then hold other groups or users depending on different Windows versions used. This system varies from other vendor products such as Linux and NetWare due to the 'static' allocation of permission being applied directly to the file or folder. However using this process of AGLP/AGDLP/AGUDLP allows a small number of static permissions to be applied and allows for easy changes to the account groups without reapplying the file permissions on the files and folders.
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Wikis are enabled by wiki software, otherwise known as wiki engines. A wiki engine, being a form of a content management system, differs from other web-based systems such as blog software, in that the content is created without any defined owner or leader, and wikis have little inherent structure, allowing structure to emerge according to the needs of the users. Wiki engines usually allow content to be written using a simplified markup language and sometimes edited with the help of a rich-text editor. There are dozens of different wiki engines in use, both standalone and part of other software, such as bug tracking systems. Some wiki engines are open-source, whereas others are proprietary. Some permit control over different functions (levels of access); for example, editing rights may permit changing, adding, or removing material. Others may permit access without enforcing access control. Other rules may be imposed to organize content.
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A wiki enables communities of editors and contributors to write documents collaboratively. All that people require to contribute is a computer, Internet access, a web browser, and a basic understanding of a simple markup language (e.g. MediaWiki markup language). A single page in a wiki website is referred to as a "wiki page", while the entire collection of pages, which are usually well-interconnected by hyperlinks, is "the wiki". A wiki is essentially a database for creating, browsing, and searching through information. A wiki allows non-linear, evolving, complex, and networked text, while also allowing for editor argument, debate, and interaction regarding the content and formatting. A defining characteristic of wiki technology is the ease with which pages can be created and updated. Generally, there is no review by a moderator or gatekeeper before modifications are accepted and thus lead to changes on the website. Many wikis are open to alteration by the general public without requiring registration of user accounts. Many edits can be made in real-time and appear almost instantly online, but this feature facilitates abuse of the system. Private wiki servers require user authentication to edit pages, and sometimes even to read them. Maged N. Kamel Boulos, Cito Maramba, and Steve Wheeler write that the open wikis produce a process of Social Darwinism. "... because of the openness and rapidity that wiki pages can be edited, the pages undergo an evolutionary selection process, not unlike that which nature subjects to living organisms. 'Unfit' sentences and sections are ruthlessly culled, edited and replaced if they are not considered 'fit', which hopefully results in the evolution of a higher quality and more relevant page."
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Some wikis have an edit button or link directly on the page being viewed if the user has permission to edit the page. This can lead to a text-based editing page where participants can structure and format wiki pages with a simplified markup language, sometimes known as wikitext, wiki markup or wikicode (it can also lead to a WYSIWYG editing page; see the paragraph after the table below). For example, starting lines of text with asterisks could create a bulleted list. The style and syntax of wikitexts can vary greatly among wiki implementations, some of which also allow HTML tags.
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Wikis can also make WYSIWYG editing available to users, usually through a JavaScript control that translates graphically entered formatting instructions into the corresponding HTML tags or wikitext. In those implementations, the markup of a newly edited, marked-up version of the page is generated and submitted to the server transparently, shielding the user from this technical detail. An example of this is the VisualEditor on Wikipedia. WYSIWYG controls do not, however, always provide all the features available in wikitext, and some users prefer not to use a WYSIWYG editor. Hence, many of these sites offer some means to edit the wikitext directly.
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Many wiki implementations, such as MediaWiki, the software that powers Wikipedia, allow users to supply an "edit summary" when they edit a page. This is a short piece of text summarizing the changes they have made (e.g. "Corrected grammar" or "Fixed formatting in table"). It is not inserted into the article's main text but is stored along with that revision of the page, allowing users to explain what has been done and why. This is similar to a log message when making changes in a revision-control system. This enables other users to see which changes have been made by whom and why, often in a list of summaries, dates and other short, relevant content, a list which is called a "log" or "history".
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Within the text of most pages, there are usually many hypertext links to other pages within the wiki. This form of non-linear navigation is more "native" to a wiki than structured/formalized navigation schemes. Users can also create any number of index or table-of-contents pages, with hierarchical categorization or whatever form of organization they like. These may be challenging to maintain "by hand", as multiple authors and users may create and delete pages in an ad hoc, unorganized manner. Wikis can provide one or more ways to categorize or tag pages to support the maintenance of such index pages. Some wikis, including the original, have a backlink feature, which displays all pages that link to a given page. It is also typically possible in a wiki to create links to pages that do not yet exist, as a way to invite others to share what they know about a subject new to the wiki. Wiki users can typically "tag" pages with categories or keywords, to make it easier for other users to find the article. For example, a user creating a new article on cold-weather biking might "tag" this page under the categories of commuting, winter sports and bicycling. This would make it easier for other users to find the article.
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Links are created using a specific syntax, the so-called "link pattern". Originally, most wikis used CamelCase to name pages and create links. These are produced by capitalizing words in a phrase and removing the spaces between them (the word "CamelCase" is itself an example). While CamelCase makes linking easy, it also leads to links in a form that deviates from the standard spelling. To link to a page with a single-word title, one must abnormally capitalize one of the letters in the word (e.g. "WiKi" instead of "Wiki"). CamelCase-based wikis are instantly recognizable because they have many links with names such as "TableOfContents" and "BeginnerQuestions". A wiki can render the visible anchor of such links "pretty" by reinserting spaces, and possibly also reverting to lower case. This reprocessing of the link to improve the readability of the anchor is, however, limited by the loss of capitalization information caused by CamelCase reversal. For example, "RichardWagner" should be rendered as "Richard Wagner", whereas "PopularMusic" should be rendered as "popular music". There is no easy way to determine which capital letters should remain capitalized. As a result, many wikis now have "free linking" using brackets, and some disable CamelCase by default.
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Wikipedia became the most famous wiki site, launched in January 2001 and entering the top ten most popular websites in 2007. In the early 2000s (decade), wikis were increasingly adopted in enterprise as collaborative software. Common uses included project communication, intranets, and documentation, initially for technical users. Some companies use wikis as their only collaborative software and as a replacement for static intranets, and some schools and universities use wikis to enhance group learning. There may be greater use of wikis behind firewalls than on the public Internet. On March 15, 2007, the word "wiki" was listed in the online "Oxford English Dictionary".
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In the late 1990s and early 2000s, the word "wiki" was used to refer to both user-editable websites and the software that powers them; the latter definition is still occasionally in use. Wiki inventor Ward Cunningham wrote in 2014 that the word "wiki" should not be used to refer to a single website, but rather to a mass of user-editable pages or sites so that a single website is not "a wiki" but "an instance of wiki". He wrote that the concept of wiki federation, in which the same content can be hosted and edited in more than one location in a manner similar to distributed version control, meant that the concept of a single discrete "wiki" no longer made sense.
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Wikis are generally designed with the philosophy of making it easy to correct mistakes, rather than making it difficult to make them. Thus, while wikis are very open, they provide a means to verify the validity of recent additions to the body of pages. The most prominent, on almost every wiki, is the "Recent Changes" page—a specific list showing recent edits, or a list of edits made within a given time frame. Some wikis can filter the list to remove minor edits and edits made by automatic importing scripts ("bots"). From the change log, other functions are accessible in most wikis: the revision history shows previous page versions and the diff feature highlights the changes between two revisions. Using the revision history, an editor can view and restore a previous version of the article. This gives great power to the author to eliminate edits. The diff feature can be used to decide whether or not this is necessary. A regular wiki user can view the diff of an edit listed on the "Recent Changes" page and, if it is an unacceptable edit, consult the history, restoring a previous revision; this process is more or less streamlined, depending on the wiki software used.
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In case unacceptable edits are missed on the "recent changes" page, some wiki engines provide additional content control. It can be monitored to ensure that a page, or a set of pages, keeps its quality. A person willing to maintain pages will be warned of modifications to the pages, allowing them to verify the validity of new editions quickly. This can be seen as a very pro-author and anti-editor feature. A watchlist is a common implementation of this. Some wikis also implement "patrolled revisions", in which editors with the requisite credentials can mark some edits as not vandalism. A "flagged revisions" system can prevent edits from going live until they have been reviewed.
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Critics of publicly editable wiki systems argue that these systems could be easily tampered with by malicious individuals ("vandals") or even by well-meaning but unskilled users who introduce errors into the content, while proponents maintain that the community of users can catch such malicious or erroneous content and correct it. Lars Aronsson, a data systems specialist, summarizes the controversy as follows: "Most people when they first learn about the wiki concept, assume that a Web site that can be edited by anybody would soon be rendered useless by destructive input. It sounds like offering free spray cans next to a grey concrete wall. The only likely outcome would be ugly graffiti and simple tagging and many artistic efforts would not be long lived. Still, it seems to work very well." High editorial standards in medicine and health sciences articles, in which users typically use peer-reviewed journals or university textbooks as sources, have led to the idea of expert-moderated wikis. Some wikis allow one to link to specific versions of articles, which has been useful to the scientific community, in that expert peer reviewers could analyse articles, improve them and provide links to the trusted version of that article. Noveck points out that "participants are accredited by members of the wiki community, who have a vested interest in preserving the quality of the work product, on the basis of their ongoing participation." On controversial topics that have been subject to disruptive editing, a wiki author may restrict editing to registered users.
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Wikis tend to take a "soft-security" approach to the problem of vandalism, making damage easy to undo rather than attempting to prevent damage. Larger wikis often employ sophisticated methods, such as bots that automatically identify and revert vandalism and JavaScript enhancements that show characters that have been added in each edit. In this way, vandalism can be limited to just "minor vandalism" or "sneaky vandalism", where the characters added/eliminated are so few that bots do not identify them and users do not pay much attention to them. An example of a bot that reverts vandalism on Wikipedia is ClueBot NG. ClueBot NG can revert edits, often within minutes, if not seconds. The bot uses machine learning in lieu of heuristics.
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Edit wars can also occur as users repetitively revert a page to the version they favor. In some cases, editors with opposing views of which content should appear or what formatting style should be used will change and re-change each other's edits. This results in the page being "unstable" from a general user's perspective, because each time a general user comes to the page, it may look different. Some wiki software allows an administrator to stop such edit wars by locking a page from further editing until a decision has been made on what version of the page would be most appropriate. Some wikis are in a better position than others to control behavior due to governance structures existing outside the wiki. For instance, a college teacher can create incentives for students to behave themselves on a class wiki they administer by limiting editing to logged-in users and pointing out that all contributions can be traced back to the contributors. Bad behavior can then be dealt with under university policies.
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The English Wikipedia has the largest user base among wikis on the World Wide Web and ranks in the top 10 among all Web sites in terms of traffic. Other large wikis include the WikiWikiWeb, Memory Alpha, Wikivoyage, and Susning.nu, a Swedish-language knowledge base. Medical and health-related wiki examples include Ganfyd, an online collaborative medical reference that is edited by medical professionals and invited non-medical experts. Many wiki communities are private, particularly within enterprises. They are often used as internal documentation for in-house systems and applications. Some companies use wikis to allow customers to help produce software documentation. A study of corporate wiki users found that they could be divided into "synthesizers" and "adders" of content. Synthesizers' frequency of contribution was affected more by their impact on other wiki users, while adders' contribution frequency was affected more by being able to accomplish their immediate work. From a study of thousands of wiki deployments, Jonathan Grudin concluded careful stakeholder analysis and education are crucial to successful wiki deployment.
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In 2005, the Gartner Group, noting the increasing popularity of wikis, estimated that they would become mainstream collaboration tools in at least 50% of companies by 2009. Wikis can be used for project management. Wikis have also been used in the academic community for sharing and dissemination of information across institutional and international boundaries. In those settings, they have been found useful for collaboration on grant writing, strategic planning, departmental documentation, and committee work. In the mid-2000s, the increasing trend among industries toward collaboration placed a heavier impetus upon educators to make students proficient in collaborative work, inspiring even greater interest in wikis being used in the classroom.
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Wikis have found some use within the legal profession and within the government. Examples include the Central Intelligence Agency's Intellipedia, designed to share and collect intelligence, DKospedia, which was used by the American Civil Liberties Union to assist with review of documents about the internment of detainees in Guantánamo Bay; and the wiki of the United States Court of Appeals for the Seventh Circuit, used to post court rules and allow practitioners to comment and ask questions. The United States Patent and Trademark Office operates Peer-to-Patent, a wiki to allow the public to collaborate on finding prior art relevant to the examination of pending patent applications. Queens, New York has used a wiki to allow citizens to collaborate on the design and planning of a local park. Cornell Law School founded a wiki-based legal dictionary called Wex, whose growth has been hampered by restrictions on who can edit.
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A city wiki (or local wiki) is a wiki used as a knowledge base and social network for a specific geographical locale. The term 'city wiki' or its foreign language equivalent (e.g. German 'Stadtwiki') is sometimes also used for wikis that cover not just a city, but a small town or an entire region. A city wiki contains information about specific instances of things, ideas, people and places. Much of this information might not be appropriate for encyclopedias such as Wikipedia (e.g. articles on every retail outlet in a town), but might be appropriate for a wiki with more localized content and viewers. A city wiki could also contain information about the following subjects, that may or may not be appropriate for a general knowledge wiki, such as:
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Joint authorship of articles, in which different users participate in correcting, editing, and compiling the finished product, can also cause editors to become tenants in common of the copyright, making it impossible to republish without permission of all co-owners, some of whose identities may be unknown due to pseudonymous or anonymous editing. Where persons contribute to a collective work such as an encyclopedia, there is, however, no joint ownership if the contributions are separate and distinguishable. Despite most wikis' tracking of individual contributions, the action of contributing to a wiki page is still arguably one of jointly correcting, editing, or compiling, which would give rise to joint ownership. Some copyright issues can be alleviated through the use of an open content license. Version 2 of the GNU Free Documentation License includes a specific provision for wiki relicensing; Creative Commons licenses are also popular. When no license is specified, an implied license to read and add content to a wiki may be deemed to exist on the grounds of business necessity and the inherent nature of a wiki, although the legal basis for such an implied license may not exist in all circumstances.
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Wikis and their users can be held liable for certain activities that occur on the wiki. If a wiki owner displays indifference and forgoes controls (such as banning copyright infringers) that he could have exercised to stop copyright infringement, he may be deemed to have authorized infringement, especially if the wiki is primarily used to infringe copyrights or obtains a direct financial benefit, such as advertising revenue, from infringing activities. In the United States, wikis may benefit from Section 230 of the Communications Decency Act, which protects sites that engage in "Good Samaritan" policing of harmful material, with no requirement on the quality or quantity of such self-policing. It has also been argued, however, that a wiki's enforcement of certain rules, such as anti-bias, verifiability, reliable sourcing, and no-original-research policies, could pose legal risks. When defamation occurs on a wiki, theoretically, all users of the wiki can be held liable, because any of them had the ability to remove or amend the defamatory material from the "publication." It remains to be seen whether wikis will be regarded as more akin to an internet service provider, which is generally not held liable due to its lack of control over publications' contents, than a publisher. It has been recommended that trademark owners monitor what information is presented about their trademarks on wikis, since courts may use such content as evidence pertaining to public perceptions. Joshua Jarvis notes, "Once misinformation is identified, the trademark owner can simply edit the entry."
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Sexual reproduction is a process exclusive to eukaryotes in which organisms produce offspring that possess a selection of the genetic traits of each parent. Genetic traits are contained within the deoxyribonucleic acid (DNA) of chromosomes. The cells of eukaryotes have a set of paired homologous chromosomes, one from each parent, and this double-chromosome stage is called "diploid". During sexual reproduction, diploid organisms produce specialized haploid sex cells called gametes via meiosis, each of which has a single set of chromosomes. Meiosis involves a stage of genetic recombination via chromosomal crossover, in which regions of DNA are exchanged between matched pairs of chromosomes, to form new chromosomes each with a new and unique combination of the genes of the parents. Then the chromosomes are separated into single sets in the gametes. Each gamete in the offspring thus has half of the genetic material of the mother and half of the father. The combination of chromosomal crossover and fertilization, bringing the two single sets of chromosomes together to make a new diploid zygote, results in new organisms that contain different sets of the genetic traits of each parent.
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In animals, the haploid stage only occurs in the gametes, the haploid cells that are specialized to fuse to form a zygote that develops into a new diploid organism. In plants, the diploid organism produces haploid spores by meiosis that are capable of undergoing repeated cell division to produce multicellular haploid organisms. In either case, gametes may be externally similar (isogamy) as in the green alga "Ulva" or may be different in size and other aspects (anisogamy). The size difference is greatest in oogamy, a type of anisogamy in which a small, motile gamete combines with a much larger, non-motile gamete.
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In most birds, both excretion and reproduction is done through a single posterior opening, called the cloaca—male and female birds touch cloaca to transfer sperm, a process called "cloacal kissing". In many other terrestrial animals, males use specialized sex organs to assist the transport of sperm—these male sex organs are called intromittent organs. In humans and other mammals this male organ is the penis, which enters the female reproductive tract (called the vagina) to achieve insemination—a process called sexual intercourse. The penis contains a tube through which semen (a fluid containing sperm) travels. In female mammals the vagina connects with the uterus, an organ which directly supports the development of a fertilized embryo within (a process called gestation).
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The flowers of flowering plants contain their sexual organs. Flowers are usually hermaphroditic, containing both male and female sexual organs. The female parts, in the center of a flower, are the pistils, each unit consisting of a carpel, a style and a stigma. Two or more of these reproductive units may be merged to form a single compound pistil, the fused carpels forming an ovary. Within the carpels are ovules which develop into seeds after fertilization. The male parts of the flower are the stamens: these consist of long filaments arranged between the pistil and the petals that produce pollen in anthers at their tips. When a pollen grain lands upon the stigma on top of a carpel's style, it germinates to produce a pollen tube that grows down through the tissues of the style into the carpel, where it delivers male gamete nuclei to fertilize an ovule that eventually develops into a seed.
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The biological cause of an organism developing into one sex or the other is called "sex determination". The cause may be genetic, environmental, haplodiploidy, or multiple factors. Within animals and other organisms that have genetic sex-determination systems, the determining factor may be the presence of a sex chromosome. In plants that are sexually dimorphic, such as "Ginkgo biloba", the liverwort "Marchantia polymorpha" or the dioecious species in the flowering plant genus "Silene", sex may also be determined by sex chromosomes. Non-genetic systems may use environmental cues, such as the temperature during early development in crocodiles, to determine the sex of the offspring.
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Humans and most other mammals have an XY sex-determination system: the Y chromosome carries factors responsible for triggering male development, making XY sex determination mostly based on the presence or absence of the Y chromosome. It is the male gamete that determines the sex of the offspring. In this system XX mammals typically are female and XY typically are male. However, individuals with XXY or XYY are males, while individuals with X and XXX are females. Unusually, the platypus, a monotreme mammal, has ten sex chromosomes; females have ten X chromosomes, and males have five X chromosomes and five Y chromosomes. Platypus egg cells all have five X chromosomes, whereas sperm cells can either have five X chromosomes or five Y chromosomes.
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In many animals and some plants, individuals of male and female sex differ in size and appearance, a phenomenon called sexual dimorphism. Sexual dimorphism in animals is often associated with sexual selection—the mating competition between individuals of one sex vis-à-vis the opposite sex. In many cases, the male of a species is larger than the female. Mammal species with extreme sexual size dimorphism tend to have highly polygynous mating systems—presumably due to selection for success in competition with other males—such as the elephant seals. Other examples demonstrate that it is the preference of females that drives sexual dimorphism, such as in the case of the stalk-eyed fly.
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The table is divided into four roughly rectangular areas called blocks. The rows of the table are called periods, and the columns are called groups. Elements from the same group of the periodic table show similar chemical characteristics. Trends run through the periodic table, with nonmetallic character (keeping their own electrons) increasing from left to right across a period, and from down to up across a group, and metallic character (surrendering electrons to other atoms) increasing in the opposite direction. The underlying reason for these trends is electron configurations of atoms. The periodic table exclusively lists electrically neutral atoms that have an equal number of positively charged protons and negatively charged electrons and puts isotopes (atoms with the same number of protons but different numbers of neutrons) at the same place. Other atoms, like nuclides and isotopes, are graphically collected in other tables like the tables of nuclides (often called Segrè charts).
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The first periodic table to become generally accepted was that of the Russian chemist Dmitri Mendeleev in 1869: he formulated the periodic law as a dependence of chemical properties on atomic mass. Because not all elements were then known, there were gaps in his periodic table, and Mendeleev successfully used the periodic law to predict properties of some of the missing elements. The periodic law was recognized as a fundamental discovery in the late 19th century, and it was explained with the discovery of the atomic number and pioneering work in quantum mechanics of the early 20th century that illuminated the internal structure of the atom. With Glenn T. Seaborg's 1945 discovery that the actinides were in fact f-block rather than d-block elements, a recognisably modern form of the table was reached. The periodic table and law are now a central and indispensable part of modern chemistry.
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The periodic table continues to evolve with the progress of science. In nature, only elements up to atomic number 94 exist; to go further, it was necessary to synthesise new elements in the laboratory. Today, all the first 118 elements are known, completing the first seven rows of the table, but chemical characterisation is still needed for the heaviest elements to confirm that their properties match their positions. It is not yet known how far the table will stretch beyond these seven rows and whether the patterns of the known part of the table will continue into this unknown region. Some scientific discussion also continues regarding whether some elements are correctly positioned in today's table. Many alternative representations of the periodic law exist, and there is some discussion as to whether there is an optimal form of the periodic table.
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Atoms can be subdivided into different types based on the number of protons (and thus also electrons) they have. This is called the atomic number, often symbolised "Z" (for "Zahl" — German for "number"). Each distinct atomic number therefore corresponds to a class of atom: these classes are called the chemical elements. The chemical elements are what the periodic table classifies and organises. Hydrogen is the element with atomic number 1; helium, atomic number 2; lithium, atomic number 3; and so on. Each of these names can be further abbreviated by a one- or two-letter chemical symbol; those for hydrogen, helium, and lithium are respectively H, He, and Li. Neutrons do not affect the atom's chemical identity, but do affect its weight. Atoms with the same number of protons but different numbers of neutrons are called isotopes of the same chemical element. Naturally occurring elements usually occur as mixes of different isotopes; since each isotope usually occurs with a characteristic abundance, naturally occurring elements have well-defined atomic weights, defined as the average mass of a naturally occurring atom of that element.
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Today, 118 elements are known, the first 94 of which are known to occur naturally on Earth at present. Of the 94 natural elements, eighty have a stable isotope and one more (bismuth) has an almost-stable isotope (with a half-life over a billion times the age of the universe). Two more, thorium and uranium, have isotopes undergoing radioactive decay with a half-life comparable to the age of the Earth. The stable elements plus bismuth, thorium, and uranium make up the 83 primordial elements that survived from the Earth's formation. The remaining eleven natural elements decay quickly enough that their continued trace occurrence rests primarily on being constantly regenerated as intermediate products of the decay of thorium and uranium. All 24 known artificial elements are radioactive.
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An electron can be thought of as inhabiting an atomic orbital, which characterises the probability it can be found in any particular region of the atom. Their energies are quantised, which is to say that they can only take discrete values. Furthermore, electrons obey the Pauli exclusion principle: different electrons must always be in different states. This allows classification of the possible states an electron can take in various energy levels known as shells, divided into individual subshells, which each contain one or more orbitals. Each orbital can contain up to two electrons: they are distinguished by a quantity known as spin, conventionally labeled "up" or "down". In a cold atom (one in its ground state), electrons arrange themselves in such a way that the total energy they have is minimised by occupying the lowest-energy orbitals available. Only the outermost electrons (so-called valence electrons) have enough energy to break free of the nucleus and participate in chemical reactions with other atoms. The others are called core electrons.
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Elements are known with up to the first seven shells occupied. The first shell contains only one orbital, a spherical s orbital. As it is in the first shell, this is called the 1s orbital. This can hold up to two electrons. The second shell similarly contains a 2s orbital, and it also contains three dumbbell-shaped 2p orbitals, and can thus fill up to eight electrons (2×1 + 2×3 = 8). The third shell contains one 3s orbital, three 3p orbitals, and five 3d orbitals, and thus has a capacity of 2×1 + 2×3 + 2×5 = 18. The fourth shell contains one 4s orbital, three 4p orbitals, five 4d orbitals, and seven 4f orbitals, thus leading to a capacity of 2×1 + 2×3 + 2×5 + 2×7 = 32. Higher shells contain more types of orbitals that continue the pattern, but such types of orbitals are not filled in the ground states of known elements. The subshell types are characterised by the quantum numbers. Four numbers describe an orbital in an atom completely: the principal quantum number "n", the azimuthal quantum number ℓ (the orbital type), the magnetic quantum number "m", and the spin quantum number "s".
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Here the sign ≪ means "much less than" as opposed to < meaning just "less than". Phrased differently, electrons enter orbitals in order of increasing "n" + ℓ, and if two orbitals are available with the same value of "n" + ℓ, the one with lower "n" is occupied first. In general, orbitals with the same value of "n" + ℓ are similar in energy, but in the case of the s-orbitals (with ℓ = 0), quantum effects raise their energy to approach that of the next "n" + ℓ group. Hence the periodic table is usually drawn to begin each row (often called a period) with the filling of a new s-orbital, which corresponds to the beginning of a new shell. Thus, with the exception of the first row, each period length appears twice:
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Starting from the third element, lithium, the first shell is full, so its third electron occupies a 2s orbital, giving a 1s 2s configuration. The 2s electron is lithium's only valence electron, as the 1s subshell is now too tightly bound to the nucleus to participate in chemical bonding to other atoms. Thus the filled first shell is called a "core shell" for this and all heavier elements. The 2s subshell is completed by the next element beryllium (1s 2s). The following elements then proceed to fill the 2p subshell. Boron (1s 2s 2p) puts its new electron in a 2p orbital; carbon (1s 2s 2p) fills a second 2p orbital; and with nitrogen (1s 2s 2p) all three 2p orbitals become singly occupied. This is consistent with Hund's rule, which states that atoms usually prefer to singly occupy each orbital of the same type before filling them with the second electron. Oxygen (1s 2s 2p), fluorine (1s 2s 2p), and neon (1s 2s 2p) then complete the already singly filled 2p orbitals; the last of these fills the second shell completely.
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Starting from element 11, sodium, the second shell is full, making the second shell a core shell for this and all heavier elements. The eleventh electron begins the filling of the third shell by occupying a 3s orbital, giving a configuration of 1s 2s 2p 3s for sodium. This configuration is abbreviated [Ne] 3s, where [Ne] represents neon's configuration. Magnesium ([Ne] 3s) finishes this 3s orbital, and the following the six elements aluminium, silicon, phosphorus, sulfur, chlorine, and argon fill the three 3p orbitals ([Ne] 3s 3p through [Ne] 3s 3p). This creates an analogous series in which the outer shell structures of sodium through argon are analogous to those of lithium through neon, and is the basis for the periodicity of chemical properties that the periodic table illustrates: at regular but changing intervals of atomic numbers, the properties of the chemical elements approximately repeat.
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The first eighteen elements can thus be arranged as the start of a periodic table. Elements in the same column have the same number of valence electrons and have analogous valence electron configurations: these columns are called groups. The single exception is helium, which has two valence electrons like beryllium and magnesium, but is typically placed in the column of neon and argon to emphasise that its outer shell is full. There are eight columns in this periodic table fragment, corresponding to at most eight outer-shell electrons. A period begins when a new shell starts filling. Finally, the colouring illustrates the blocks: the elements in the s-block (coloured red) are filling s-orbitals, while those in the p-block (coloured yellow) are filling p-orbitals.
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Starting the next row, for potassium and calcium the 4s subshell is the lowest in energy, and therefore they fill it. Potassium adds one electron to the 4s shell ([Ar] 4s), and calcium then completes it ([Ar] 4s). However, starting from scandium ([Ar] 3d 4s) the 3d subshell becomes the next highest in energy. The 4s and 3d subshells have approximately the same energy and they compete for filling the electrons, and so the occupation is not quite consistently filling the 3d orbitals one at a time. The precise energy ordering of 3d and 4s changes along the row, and also changes depending on how many electrons are removed from the atom. For example, due to the repulsion between the 3d electrons and the 4s ones, at chromium the 4s energy level becomes slightly higher than 3d, and so it becomes more profitable to have a [Ar] 3d 4s configuration than an [Ar] 3d 4s one. A similar anomaly occurs at copper. These are violations of the Madelung rule. Such anomalies however do not have any chemical significance, as the various configurations are so close in energy to each other that the presence of a nearby atom can shift the balance. The periodic table therefore ignores these and considers only idealised configurations.
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At zinc ([Ar] 3d 4s), the 3d orbitals are completely filled with a total of ten electrons. Next come the 4p orbitals, completing the row, which are filled progressively by gallium ([Ar] 3d 4s 4p) through krypton ([Ar] 3d 4s 4p), in a manner analogous to the previous p-block elements. From gallium onwards, the 3d orbitals form part of the electronic core, and no longer participate in chemistry. The s- and p-block elements, which fill their outer shells, are called main-group elements; the d-block elements (coloured blue below), which fill an inner shell, are called transition elements (or transition metals, since they are all metals).
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The sixth row of the table likewise starts with two s-block elements: caesium and barium. After this, the first f-block elements (coloured green below) begin to appear, starting with lanthanum. These are sometimes termed inner transition elements. As there are now not only 4f but also 5d and 6s subshells at similar energies, competition occurs once again with many irregular configurations; this has resulted in some dispute about where exactly the f-block is supposed to begin, but most who study the matter agree that it starts at lanthanum in accordance with the Aufbau principle. Even though lanthanum does not itself fill the 4f subshell as a single atom, because of repulsion between electrons, its 4f orbitals are low enough in energy to participate in chemistry. At ytterbium, the seven 4f orbitals are completely filled with fourteen electrons; thereafter, a series of ten transition elements (lutetium through mercury) follows, and finally six main-group elements (thallium through radon) complete the period. From lutetium onwards the 4f orbitals are in the core, and from thallium onwards so are the 5d orbitals.
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The seventh row is analogous to the sixth row: 7s fills (francium and radium), then 5f (actinium to nobelium), then 6d (lawrencium to copernicium), and finally 7p (nihonium to oganesson). Starting from lawrencium the 5f orbitals are in the core, and probably the 6d orbitals join the core starting from nihonium. Again there are a few anomalies along the way: for example, as single atoms neither actinium nor thorium actually fills the 5f subshell, and lawrencium does not fill the 6d shell, but all these subshells can still become filled in chemical environments. For a very long time, the seventh row was incomplete as most of its elements do not occur in nature. The missing elements beyond uranium started to be synthesised in the laboratory in 1940, when neptunium was made. The row was completed with the synthesis of tennessine in 2010 (the last element oganesson had already been made in 2002), and the last elements in this seventh row were given names in 2016.
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Under an international naming convention, the groups are numbered numerically from 1 to 18 from the leftmost column (the alkali metals) to the rightmost column (the noble gases). The f-block groups are ignored in this numbering. Groups can also be named by their first element, e.g. the "scandium group" for group 3. Previously, groups were known by Roman numerals. In America, the Roman numerals were followed by either an "A" if the group was in the s- or p-block, or a "B" if the group was in the d-block. The Roman numerals used correspond to the last digit of today's naming convention (e.g. the group 4 elements were group IVB, and the group 14 elements were group IVA). In Europe, the lettering was similar, except that "A" was used if the group was before group 10, and "B" was used for groups including and after group 10. In addition, groups 8, 9 and 10 used to be treated as one triple-sized group, known collectively in both notations as group VIII. In 1988, the new IUPAC (International Union of Pure and Applied Chemistry) naming system (1–18) was put into use, and the old group names (I–VIII) were deprecated.
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Atomic radii (the size of atoms) are dependent on the sizes of their outermost orbitals. They generally decrease going left to right along the main-group elements, because the nuclear charge increases but the outer electrons are still in the same shell. However, going down a column, the radii generally increase, because the outermost electrons are in higher shells that are thus further away from the nucleus. The first row of each block is abnormally small, due to an effect called kainosymmetry or primogenic repulsion: the 1s, 2p, 3d, and 4f subshells have no inner analogues that they would have to be orthogonal to. Higher s-, p-, d-, and f-subshells experience strong repulsion from their inner analogues, which have approximately the same angular distribution of charge, and must expand to avoid this. This makes significant differences arise between the small 2p elements, which prefer multiple bonding, and the larger 3p and higher p-elements, which do not. Similar anomalies arise for the 1s, 2p, 3d, 4f, and the hypothetical elements: the degree of this first-row anomaly is highest for the s-block, is moderate for the p-block, and is less pronounced for the d- and f-blocks.
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In the transition elements, an inner shell is filling, but the size of the atom is still determined by the outer electrons. The increasing nuclear charge across the series and the increased number of inner electrons for shielding somewhat compensate each other, so the decrease in radius is smaller. The 4p and 5d atoms, coming immediately after new types of transition series are first introduced, are smaller than would have been expected, because the added core 3d and 4f subshells provide only incomplete shielding of the nuclear charge for the outer electrons. Hence for example gallium atoms are slightly smaller than aluminium atoms. Together with kainosymmetry, this results in an even-odd difference between the periods (except in the s-block) that is sometimes known as secondary periodicity: elements in even periods have smaller atomic radii and prefer to lose fewer electrons, while elements in odd periods (except the first) differ in the opposite direction. Thus for example many properties in the p-block show a zigzag rather than a smooth trend along the group. For example, phosphorus and antimony in odd periods of group 15 readily reach the +5 oxidation state, whereas nitrogen, arsenic, and bismuth in even periods prefer to stay at +3.
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Thallium and lead atoms are about the same size as indium and tin atoms respectively, but from bismuth to radon the 6p atoms are larger than the analogous 5p atoms. This happens because when atomic nuclei become highly charged, special relativity becomes needed to gauge the effect of the nucleus on the electron cloud. These relativistic effects result in heavy elements increasingly having differing properties compared to their lighter homologues in the periodic table. Spin–orbit interaction splits the p-subshell: one p-orbital is relativistically stabilised and shrunken (it fills in thallium and lead), but the other two (filling in bismuth through radon) are relativistically destabilised and expanded. Relativistic effects also explain why gold is golden and mercury is a liquid at room temperature. They are expected to become very strong in the late seventh period, potentially leading to a collapse of periodicity. Electron configurations are only clearly known until element 108 (hassium), and experimental chemistry beyond 108 has only been done for 112 (copernicium), 113 (nihonium), and 114 (flerovium), so the chemical characterisation of the heaviest elements remains a topic of current research.
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The first ionisation energy of an atom is the energy required to remove an electron from it. This varies with the atomic radius: ionisation energy increases left to right and down to up, because electrons that are closer to the nucleus are held more tightly and are more difficult to remove. Ionisation energy thus is minimised at the first element of each period – hydrogen and the alkali metals – and then generally rises until it reaches the noble gas at the right edge of the period. There are some exceptions to this trend, such as oxygen, where the electron being removed is paired and thus interelectronic repulsion makes it easier to remove than expected.
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The valence of an element can be defined either as the number of hydrogen atoms that can combine with it to form a simple binary hydride, or as twice the number of oxygen atoms that can combine with it to form a simple binary oxide (that is, not a peroxide or a superoxide). The valences of the main-group elements are directly related to the group number: the hydrides in the main groups 1–2 and 13–17 follow the formulae MH, MH, MH, MH, MH, MH, and finally MH. The highest oxides instead increase in valence, following the formulae MO, MO, MO, MO, MO, MO, MO. Today the notion of valence has been extended by that of the oxidation state, which is the formal charge left on an element when all other elements in a compound have been removed as their ions.
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The electron configuration suggests a ready explanation from the number of electrons available for bonding, although a full explanation requires considering the energy that would be released in forming compounds with different valences rather than simply considering electron configurations alone. For example, magnesium forms Mg rather than Mg cations when dissolved in water, because the latter would spontaneously disproportionate into Mg and Mg cations. This is because the enthalpy of hydration (surrounding the cation with water molecules) increases in magnitude with the charge and radius of the ion. In Mg, the outermost orbital (which determines ionic radius) is still 3s, so the hydration enthalpy is small and insufficient to compensate the energy required to remove the electron; but ionising again to Mg uncovers the core 2p subshell, making the hydration enthalpy large enough to allow magnesium(II) compounds to form. For similar reasons, the common oxidation states of the heavier p-block elements (where the ns electrons become lower in energy than the np) tend to vary by steps of 2, because that is necessary to uncover an inner subshell and decrease the ionic radius (e.g. Tl uncovers 6s, and Tl uncovers 5d, so once thallium loses two electrons it tends to lose the third one as well). Analogous arguments based on orbital hybridisation can be used for the less electronegative p-block elements. For example, GaCl requires gallium's 4s orbital to mix with only one 4p orbital, whereas GaCl and GaCl would both require it to mix with two. Hence no extra energy is needed between these last two steps to involve more orbitals, so gallium(II) is unstable while gallium(III) is stable.
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For transition metals, common oxidation states are nearly always at least +2 for similar reasons (uncovering the next subshell); this holds even for the metals with anomalous ds or ds configurations (except for silver), because repulsion between d-electrons means that the movement of the second electron from the s- to the d-subshell does not appreciably change its ionisation energy. Because ionising the transition metals further does not uncover any new inner subshells, their oxidation states tend to vary by steps of 1 instead. The lanthanides and late actinides generally show a stable +3 oxidation state, removing the outer s-electrons and then (usually) one electron from the (n−2)f-orbitals, that are similar in energy to ns. The common and maximum oxidation states of the d- and f-block elements tend to depend on the ionisation energies. As the energy difference between the (n−1)d and ns orbitals rises along each transition series, it becomes less energetically favourable to ionise further electrons. Thus, the early transition metal groups tend to prefer higher oxidation states, but the +2 oxidation state becomes more stable for the late transition metal groups. The highest formal oxidation state thus increases from +3 at the beginning of each d-block row, to +7 or +8 in the middle (e.g. OsO), and then to +2 at the end. The lanthanides and late actinides usually have high fourth ionisation energies and hence rarely surpass the +3 oxidation states, whereas early actinides have low fourth ionisation energies and so for example neptunium and plutonium can reach +7.
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As elements in the same group share the same valence configurations, they usually exhibit similar chemical behaviour. For example, the alkali metals in the first group all have one valence electron, and form a very homogeneous class of elements: they are all soft and reactive metals. However, there are many factors involved, and groups can often be rather hetereogeneous. For instance, hydrogen also has one valence electron and is in the same group as the alkali metals, but its chemical behaviour is quite different. The stable elements of group 14 comprise a nonmetal (carbon), two semiconductors (silicon and germanium), and two metals (tin and lead); they are nonetheless united by having four valence electrons. This often leads to similarities in maximum and minimum oxidation states (e.g. sulfur and selenium in group 16 both have maximum oxidation state +6, as in SO and SeO, and minimum oxidation state −2, as in sulfides and selenides); but not always (e.g. oxygen is not known to form oxidation state +6, despite being in the same group as sulfur and selenium).
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Another important property of elements is their electronegativity. Atoms can form covalent bonds to each other by sharing electrons in pairs, creating an overlap of valence orbitals. The degree to which each atom attracts the shared electron pair depends on the atom's electronegativity – the tendency of an atom towards gaining or losing electrons. The more electronegative atom will tend to attract the electron pair more, and the less electronegative (or more electropositive) one will attract it less. In extreme cases, the electron can be thought of as having been passed completely from the more electropositive atom to the more electronegative one, though this is a simplification. The bond then binds two ions, one positive (having given up the electron) and one negative (having accepted it), and is termed an ionic bond.
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A simple substance is a substance formed from atoms of one chemical element. The simple substances of the more electronegative atoms tend to share electrons (form covalent bonds) with each other. They form either small molecules (like hydrogen or oxygen, whose atoms bond in pairs) or giant structures stretching indefinitely (like carbon or silicon). The noble gases simply stay as single atoms, as they already have a full shell. Substances composed of discrete molecules or single atoms are held together by weaker attractive forces between the molecules, such as the London dispersion force: as electrons move within the molecules, they create momentary imbalances of electrical charge, which induce similar imbalances on nearby molecules and create synchronised movements of electrons across many neighbouring molecules.
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The more electropositive atoms, however, tend to instead lose electrons, creating a "sea" of electrons engulfing cations. The outer orbitals of one atom overlap to share electrons with all its neighbours, creating a giant structure of molecular orbitals extending over all the atoms. This negatively charged "sea" pulls on all the ions and keeps them together in a metallic bond. Elements forming such bonds are often called metals; those which do not are often called nonmetals. Some elements can form multiple simple substances with different structures: these are called allotropes. For example, diamond and graphite are two allotropes of carbon.
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The metallicity of an element can be predicted from electronic properties. When atomic orbitals overlap during metallic or covalent bonding, they create both bonding and antibonding molecular orbitals of equal capacity, with the antibonding orbitals of higher energy. Net bonding character occurs when there are more electrons in the bonding orbitals than there are in the antibonding orbitals. Metallic bonding is thus possible when the number of electrons delocalised by each atom is less than twice the number of orbitals contributing to the overlap. This is the situation for elements in groups 1 through 13; they also have too few valence electrons to form giant covalent structures where all atoms take equivalent positions, and so almost all of them metallise. The exceptions are hydrogen and boron, which have too high an ionisation energy. Hydrogen thus forms a covalent H molecule, and boron forms a giant covalent structure based on icosahedral B clusters. In a metal, the bonding and antibonding orbitals have overlapping energies, creating a single band that electrons can freely flow through, allowing for electrical conduction.
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In group 14, both metallic and covalent bonding become possible. In a diamond crystal, covalent bonds between carbon atoms are strong, because they have a small atomic radius and thus the nucleus has more of a hold on the electrons. Therefore, the bonding orbitals that result are much lower in energy than the antibonding orbitals, and there is no overlap, so electrical conduction becomes impossible: carbon is a nonmetal. However, covalent bonding becomes weaker for larger atoms and the energy gap between the bonding and antibonding orbitals decreases. Therefore, silicon and germanium have smaller band gaps and are semiconductors: electrons can cross the gap when thermally excited. The band gap disappears in tin, so that tin and lead become metals.
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Elements in groups 15 through 17 have too many electrons to form giant covalent molecules that stretch in all three dimensions. For the lighter elements, the bonds in small diatomic molecules are so strong that a condensed phase is disfavoured: thus nitrogen (N), oxygen (O), white phosphorus (P), sulfur (S), and the stable halogens (F, Cl, Br, and I) readily form covalent molecules with few atoms. The heavier ones tend to form long chains (e.g. red phosphorus, grey selenium, tellurium) or layered structures (e.g. carbon as graphite, black phosphorus, grey arsenic, grey antimony, bismuth) that only extend in one or two rather than three dimensions. As these structures do not use all their orbitals for bonding, they end up with bonding, nonbonding, and antibonding bands in order of increasing energy. Similarly to group 14, the band gaps shrink for the heavier elements and free movement of electrons between the chains or layers becomes possible. Thus for example black phosphorus, black arsenic, grey selenium, tellurium, and iodine are semiconductors; grey arsenic, grey antimony, and bismuth are semimetals (exhibiting quasi-metallic conduction, with a very small band overlap); and polonium and probably astatine are true metals. Finally, the natural group 18 elements all stay as individual atoms.
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The dividing line between metals and nonmetals is roughly diagonal from top left to bottom right, with the transition series appearing to the left of this diagonal (as they have many available orbitals for overlap). This is expected, as metallicity tends to be correlated with electropositivity and the willingness to lose electrons, which increases right to left and up to down. Thus the metals greatly outnumber the nonmetals. Elements near the borderline are difficult to classify: they tend to have properties that are intermediate between those of metals and nonmetals, and may have some properties characteristic of both. They are often termed semimetals or metalloids. The term "semimetal" used in this sense should not be confused with its strict physical sense having to do with band structure: bismuth is physically a semimetal, but is generally considered a metal by chemists.
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The following table considers the most stable allotropes at standard conditions. The elements coloured yellow form simple substances that are well-characterised by metallic bonding. Elements coloured light blue form giant covalent structures, whereas those coloured dark blue form small covalently bonded molecules that are held together by weaker van der Waals forces. The noble gases are coloured in violet: their molecules are single atoms and no covalent bonding occurs. Greyed-out cells are for elements which have not been prepared in sufficient quantities for their most stable allotropes to have been characterised in this way. Theoretical considerations and current experimental evidence suggest that all of those elements would metallise if they could form condensed phases, except perhaps for oganesson.
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Generally, metals are shiny and dense. They usually have high melting and boiling points due to the strength of the metallic bond, and are often malleable and ductile (easily stretched and shaped) because the atoms can move relative to each other without breaking the metallic bond. They conduct electricity because their electrons are free to move in all three dimensions. Similarly, they conduct heat, which is transferred by the electrons as extra kinetic energy: they move faster. These properties persist in the liquid state, as although the crystal structure is destroyed on melting, the atoms still touch and the metallic bond persists, though it is weakened. Metals tend to be reactive towards nonmetals. Some exceptions can be found to these generalisations: for example, manganese, arsenic, antimony, and bismuth are brittle; chromium is extremely hard; gallium, rubidium, caesium, and mercury are liquid at or close to room temperature; and noble metals such as gold are chemically very inert.
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Nonmetals exhibit different properties. Those forming giant covalent crystals exhibit high melting and boiling points, as it takes considerable energy to overcome the strong covalent bonds. Those forming discrete molecules are held together mostly by dispersion forces, which are more easily overcome; thus they tend to have lower melting and boiling points, and many are liquids or gases at room temperature. Nonmetals are often dull-looking. They tend to be reactive towards metals, except for the noble gases, which are inert towards most substances. They are brittle when solid as their atoms are held tightly in place. They are less dense and conduct electricity poorly, because there are no mobile electrons. Near the borderline, band gaps are small and thus many elements in that region are semiconductors, such as silicon, germanium, selenium, and tellurium. Again there are exceptions; for example, diamond has the highest thermal conductivity of all known materials, greater than any metal.
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There are some other relationships throughout the periodic table between elements that are not in the same group, such as the diagonal relationships between elements that are diagonally adjacent (e.g. lithium and magnesium). Some similarities can also be found between the main groups and the transition metal groups, or between the early actinides and early transition metals, when the elements have the same number of valence electrons. Thus uranium somewhat resembles chromium and tungsten in group 6, as all three have six valence electrons. Relationships between elements with the same number of valence electrons but different types of valence orbital have been called secondary or isodonor relationships: they usually have the same maximum oxidation states, but not the same minimum oxidation states. For example, chlorine and manganese both have +7 as their maximum oxidation state (e.g. ClO and MnO), but their respective minimum oxidation states are −1 (e.g. HCl) and −3 (K[Mn(CO)]). Elements with the same number of valence vacancies but different numbers of valence electrons are related by a tertiary or isoacceptor relationship: they have similar minimum but not maximum oxidation states. For example, hydrogen and chlorine both have −1 as their minimum oxidation state (in hydrides and chlorides), but hydrogen's maximum oxidation state is +1 (e.g. HO) while chlorine's is +7.
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Many other physical properties of the elements exhibit periodic variation in accordance with the periodic law, such as melting points, boiling points, heats of fusion, heats of vaporisation, atomisation energy, and so on. Similar periodic variations appear for the compounds of the elements, which can be observed by comparing hydrides, oxides, sulfides, halides, and so on. Chemical properties are more difficult to describe quantitatively, but likewise exhibit their own periodicities. Examples include the variation in the acidic and basic properties of the elements and their compounds, the stabilities of compounds, and methods of isolating the elements. Periodicity is and has been used very widely to predict the properties of unknown new elements and new compounds, and is central to modern chemistry.
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Many terms have been used in the literature to describe sets of elements that behave similarly. The group names "alkali metal", "alkaline earth metal", "pnictogen", "chalcogen", "halogen", and "noble gas" are acknowledged by IUPAC; the other groups can be referred to by their number, or by their first element (e.g., group 6 is the chromium group). Some divide the p-block elements from groups 13 to 16 by metallicity, although there is neither an IUPAC definition nor a precise consensus on exactly which elements should be considered metals, nonmetals, or semi-metals (sometimes called metalloids). Neither is there a consensus on what the metals succeeding the transition metals ought to be called, with "post-transition metal" and "poor metal" being among the possibilities having been used. Some advanced monographs exclude the elements of group 12 from the transition metals on the grounds of their sometimes quite different chemical properties, but this is not a universal practice.
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The "lanthanides" are considered to be the elements La–Lu, which are all very similar to each other: historically they included only Ce–Lu, but lanthanum became included by common usage. The "rare earth elements" (or rare earth metals) add scandium and yttrium to the lanthanides. Analogously, the "actinides" are considered to be the elements Ac–Lr (historically Th–Lr), although variation of properties in this set is much greater than within the lanthanides. IUPAC recommends the names "lanthanoids" and "actinoids" to avoid ambiguity, as the -ide suffix typically denotes a negative ion; however "lanthanides" and "actinides" remain common. Some authors who consider lutetium and lawrencium to be group 3 elements prefer to define the lanthanides as La–Yb and the actinides as Ac–No, matching the f-block.
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The scope of terms varies significantly between authors. For example, according to IUPAC, the noble gases extend to include the whole group, including the very radioactive superheavy element oganesson. However, among those who specialise in the superheavy elements, this is not often done: in this case "noble gas" is typically taken to imply the unreactive behaviour of the lighter elements of the group. Since calculations generally predict that oganesson should not be particularly inert due to relativistic effects, and may not even be a gas at room temperature if it could be produced in bulk, its status as a noble gas is often questioned in this context. Furthermore, national variations are sometimes encountered: in Japan, alkaline earth metals often do not include beryllium and magnesium as their behaviour is different from the heavier group 2 metals.
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In 1817, German physicist Johann Wolfgang Döbereiner began to formulate one of the earliest attempts to classify the elements. In 1829, he found that he could form some of the elements into groups of three, with the members of each group having related properties. He termed these groups triads. Chlorine, bromine, and iodine formed a triad; as did calcium, strontium, and barium; lithium, sodium, and potassium; and sulfur, selenium, and tellurium. Today, all these triads form part of modern-day groups. Various chemists continued his work and were able to identify more and more relationships between small groups of elements. However, they could not build one scheme that encompassed them all.
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German chemist Lothar Meyer noted the sequences of similar chemical and physical properties repeated at periodic intervals. According to him, if the atomic weights were plotted as ordinates (i.e. vertically) and the atomic volumes as abscissas (i.e. horizontally)—the curve obtained a series of maximums and minimums—the most electropositive elements would appear at the peaks of the curve in the order of their atomic weights. In 1864, a book of his was published; it contained an early version of the periodic table containing 28 elements, and classified elements into six families by their valence—for the first time, elements had been grouped according to their valence. Works on organizing the elements by atomic weight had until then been stymied by inaccurate measurements of the atomic weights. In 1868, he revised his table, but this revision was published as a draft only after his death.
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The definitive breakthrough came from the Russian chemist Dmitri Mendeleev. Although other chemists (including Meyer) had found some other versions of the periodic system at about the same time, Mendeleev was the most dedicated to developing and defending his system, and it was his system that most impacted the scientific community. On 17 February 1869 (1 March 1869 in the Gregorian calendar), Mendeleev began arranging the elements and comparing them by their atomic weights. He began with a few elements, and over the course of the day his system grew until it encompassed most of the known elements. After finding a consistent arrangement, his printed table appeared in May 1869 in the journal of the Russian Chemical Society. When elements did not appear to fit in the system, he boldly predicted that either valencies or atomic weights had been measured incorrectly, or that there was a missing element yet to be discovered. In 1871, Mendeleev published a long article, including an updated form of his table, that made his predictions for unknown elements explicit. Mendeleev predicted the properties of three of these unknown elements in detail: as they would be missing heavier homologues of boron, aluminium, and silicon, he named them eka-boron, eka-aluminium, and eka-silicon ("eka" being Sanskrit for "one").
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In 1875, the French chemist Paul-Émile Lecoq de Boisbaudran, working without knowledge of Mendeleev's prediction, discovered a new element in a sample of the mineral sphalerite, and named it gallium. He isolated the element and began determining its properties. Mendeleev, reading de Boisbaudran's publication, sent a letter claiming that gallium was his predicted eka-aluminium. Although Lecoq de Boisbaudran was initially sceptical, and suspected that Mendeleev was trying to take credit for his discovery, he later admitted that Mendeleev was correct. In 1879, the Swedish chemist Lars Fredrik Nilson discovered a new element, which he named scandium: it turned out to be eka-boron. Eka-silicon was found in 1886 by German chemist Clemens Winkler, who named it germanium. The properties of gallium, scandium, and germanium matched what Mendeleev had predicted. In 1889, Mendeleev noted at the Faraday Lecture to the Royal Institution in London that he had not expected to live long enough "to mention their discovery to the Chemical Society of Great Britain as a confirmation of the exactitude and generality of the periodic law". Even the discovery of the noble gases at the close of the 19th century, which Mendeleev had not predicted, fitted neatly into his scheme as an eighth main group.
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The same year, English physicist Henry Moseley using X-ray spectroscopy confirmed van den Broek's proposal experimentally. Moseley determined the value of the nuclear charge of each element from aluminium to gold and showed that Mendeleev's ordering actually places the elements in sequential order by nuclear charge. Nuclear charge is identical to proton count and determines the value of the atomic number ("Z") of each element. Using atomic number gives a definitive, integer-based sequence for the elements. Moseley's research immediately resolved discrepancies between atomic weight and chemical properties; these were cases such as tellurium and iodine, where atomic number increases but atomic weight decreases. Although Moseley was soon killed in World War I, the Swedish physicist Manne Siegbahn continued his work up to uranium, and established that it was the element with the highest atomic number then known (92). Based on Moseley and Siegbahn's research, it was also known which atomic numbers corresponded to missing elements yet to be found.
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Bohr called his electron shells "rings" in 1913: atomic orbitals within shells did not exist at the time of his planetary model. Bohr explains in Part 3 of his famous 1913 paper that the maximum electrons in a shell is eight, writing, "We see, further, that a ring of n electrons cannot rotate in a single ring round a nucleus of charge ne unless n < 8." For smaller atoms, the electron shells would be filled as follows: "rings of electrons will only join if they contain equal numbers of electrons; and that accordingly the numbers of electrons on inner rings will only be 2, 4, 8." However, in larger atoms the innermost shell would contain eight electrons: "on the other hand, the periodic system of the elements strongly suggests that already in neon N = 10 an inner ring of eight electrons will occur." His proposed electron configurations for the light atoms (shown to the right) do not always accord with those now known.
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The first one to systematically expand and correct the chemical potentials of Bohr's atomic theory was Walther Kossel in 1914 and in 1916. Kossel explained that in the periodic table new elements would be created as electrons were added to the outer shell. In Kossel's paper, he writes: "This leads to the conclusion that the electrons, which are added further, should be put into concentric rings or shells, on each of which ... only a certain number of electrons—namely, eight in our case—should be arranged. As soon as one ring or shell is completed, a new one has to be started for the next element; the number of electrons, which are most easily accessible, and lie at the outermost periphery, increases again from element to element and, therefore, in the formation of each new shell the chemical periodicity is repeated."
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Prompted by Bohr, Wolfgang Pauli took up the problem of electron configurations in 1923. Pauli extended Bohr's scheme to use four quantum numbers, and formulated his exclusion principle which stated that no two electrons could have the same four quantum numbers. This explained the lengths of the periods in the periodic table (2, 8, 18, and 32), which corresponded to the number of electrons that each shell could occupy. In 1925, Friedrich Hund arrived at configurations close to the modern ones. As a result of these advances, periodicity became based on the number of chemically active or valence electrons rather than by the valences of the elements. The Aufbau principle that describes the electron configurations of the elements was first empirically observed by Erwin Madelung in 1926, though the first to publish it was Vladimir Karapetoff in 1930. In 1961, Vsevolod Klechkovsky derived the first part of the Madelung rule (that orbitals fill in order of increasing "n" + ℓ) from the Thomas–Fermi model; the complete rule was derived from a similar potential in 1971 by Yury N. Demkov and Valentin N. Ostrovsky.
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The quantum theory clarified the transition metals and lanthanides as forming their own separate groups, transitional between the main groups, although some chemists had already proposed tables showing them this way before then: the English chemist Henry Bassett did so in 1892, the Danish chemist Julius Thomsen in 1895, and the Swiss chemist Alfred Werner in 1905. Bohr used Thomsen's form in his 1922 Nobel Lecture; Werner's form is very similar to the modern 32-column form. In particular, this supplanted Brauner's asteroidal hypothesis. The exact position of the lanthanides, and thus the composition of group 3, remained under dispute for decades longer because their electron configurations were initially measured incorrectly. In 2021 IUPAC released a provisional report suggesting that group 3 should contain scandium, yttrium, lutetium, and lawrencium. This matches the classification Bassett and Werner adopted over a century earlier as well as the Madelung rule.
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By 1936, the pool of missing elements from hydrogen to uranium had shrunk to four: elements 43, 61, 85, and 87 remained missing. Element 43 eventually became the first element to be synthesised artificially via nuclear reactions rather than discovered in nature. It was discovered in 1937 by Italian chemists Emilio Segrè and Carlo Perrier, who named their discovery technetium, after the Greek word for "artificial". Elements 61 (promethium) and 85 (astatine) were likewise produced artificially; element 87 (francium) became the last element to be discovered in nature, by French chemist Marguerite Perey. The elements beyond uranium were likewise discovered artificially, starting with Edwin McMillan and Philip Abelson's 1940 discovery of neptunium (via bombardment of uranium with neutrons). Glenn T. Seaborg and his team at the Lawrence Berkeley National Laboratory (LBNL) continued discovering transuranium elements, starting with plutonium, and discovered that contrary to previous thinking, the elements from actinium onwards were congeners of the lanthanides rather than transition metals. Bassett (1892), Werner (1905), and the French engineer Charles Janet (1928) had previously suggested this, but their ideas did not then receive general acceptance. Seaborg thus called them the actinides. Elements up to 101 (named mendelevium in honour of Mendeleev) were synthesised either through neutron or alpha-particle irradiation, or in nuclear explosions in the cases of 99 (einsteinium) and 100 (fermium).
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A significant controversy arose with elements 102 through 106 in the 1960s and 1970s, as competition arose between the LBNL team (now led by Albert Ghiorso) and a team of Soviet scientists at the Joint Institute for Nuclear Research (JINR) led by Georgy Flyorov. Each team claimed discovery, and in some cases each proposed their own name for the element, creating an element naming controversy that lasted decades. These elements were made by bombardment of actinides with light ions. IUPAC at first adopted a hands-off approach, preferring to wait and see if a consensus would be forthcoming. Unfortunately, it was also the height of the Cold War, and it became clear after some time that this would not happen. As such, IUPAC and the International Union of Pure and Applied Physics (IUPAP) created a Transfermium Working Group (TWG, fermium being element 100) in 1985 to set out criteria for discovery, which were published in 1991. After some further controversy, these elements received their final names in 1997, including seaborgium (106) in honour of Seaborg.
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The TWG's criteria were used to arbitrate later element discovery claims from LBNL and JINR, as well as from research institutes in Germany (GSI) and Japan (Riken). Currently, consideration of discovery claims is performed by a IUPAC/IUPAP Joint Working Party. After priority was assigned, the elements were officially added to the periodic table, and the discoverers were invited to propose their names. By 2016, this had occurred for all elements up to 118, therefore completing the periodic table's first seven rows. The discoveries of elements beyond 106 were made possible by techniques devised by Yuri Oganessian at the JINR: cold fusion (bombardment of lead and bismuth by heavy ions) made possible the 1981–2004 discoveries of elements 107 through 112 at GSI and 113 at Riken, and he led the JINR team (in collaboration with American scientists) to discover elements 114 through 118 using hot fusion (bombardment of actinides by calcium ions) in 1998–2010. The heaviest known element, oganesson (118), is named in Oganessian's honour. Element 114 is named flerovium in honour of his predecessor and mentor Flyorov.
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In celebration of the periodic table's 150th anniversary, the United Nations declared the year 2019 as the International Year of the Periodic Table, celebrating "one of the most significant achievements in science". The discovery criteria set down by the TWG were updated in 2020 in response to experimental and theoretical progress that had not been foreseen in 1991. Today, the periodic table is among the most recognisable icons of chemistry. IUPAC is involved today with many processes relating to the periodic table: the recognition and naming of new elements, recommending group numbers and collective names, determining which elements belong to group 3, and the updating of atomic weights.
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Like the group 1 metals, hydrogen has one electron in its outermost shell and typically loses its only electron in chemical reactions. It has some metal-like chemical properties, being able to displace some metals from their salts. But hydrogen forms a diatomic nonmetallic gas at standard conditions, unlike the alkali metals which are reactive solid metals. This and hydrogen's formation of hydrides, in which it gains an electron, brings it close to the properties of the halogens which do the same (though it is rarer for hydrogen to form H than H). Moreover, the lightest two halogens (fluorine and chlorine) are gaseous like hydrogen at standard conditions. Some properties of hydrogen are not a good fit for either group: hydrogen is neither highly oxidising nor highly reducing and is not reactive with water. Hydrogen thus has properties corresponding to both those of the alkali metals and the halogens, but matches neither group perfectly, and is thus difficult to place by its chemistry. Therefore, while the electronic placement of hydrogen in group 1 predominates, some rarer arrangements show either hydrogen in group 17, duplicate hydrogen in both groups 1 and 17, or float it separately from all groups.
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Helium is an unreactive noble gas at standard conditions, and has a full outer shell: these properties are like the noble gases in group 18, but not at all like the reactive alkaline earth metals of group 2. Therefore, helium is nearly universally placed in group 18 which its properties best match. However, helium only has two electrons in its outermost shell, whereas the other noble gases have eight; and it is an s-block element, whereas all other noble gases are p-block elements. Also, solid helium crystallises in a hexagonal close-packed structure, which matches beryllium and magnesium in group 2, but not the other noble gases in group 18. In these ways helium better matches the alkaline earth metals and it is occasionally placed to head group 2. Some chemists have advocated placing helium in group 2 on the grounds of the first-row anomaly trend, as helium as the first s element before the alkaline earth metals stands out as anomalous in a way that helium as the first noble gas does not. Thus for example a large difference in atomic radii between the first and second members of each main group is seen in groups 1 and 13–17: it exists between neon and argon, and between helium and beryllium, but not between helium and neon. Moving helium to group 2 makes this trend consistent in groups 2 and 18 as well. Tables that float both hydrogen and helium outside all groups may also rarely be encountered.
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Published periodic tables show variation regarding the heavier members of group 3, which begins with scandium and yttrium. They are commonly lanthanum and actinium, but many physicists and chemists have argued that they should be lutetium and lawrencium. The spaces below yttrium are sometimes left blank as a third option, but there is confusion in the literature on whether this format implies that group 3 contains only scandium and yttrium, or if it also contains all the lanthanides and actinides. The dispute arose because the electron configurations of the lanthanides were first measured incorrectly. The incorrect configurations first measured suggested that Sc-Y-La-Ac was better, but the correct ones now known suggest instead Sc-Y-Lu-Lr. While it has been argued that atoms of lanthanum and actinium lack f-electrons and hence that they cannot be f-block elements, the same is true of thorium which is never disputed as an f-block element. Additionally, lanthanum and actinium have valence f-orbitals, whereas lutetium and lawrencium do not. While there is also a relationship between yttrium and lanthanum, it can thus be thought of as a secondary relationship between elements with the same number of valence electrons but different kinds of valence orbitals, such as that between chromium and uranium. The Sc-Y-Lu-Lr table is in accordance with the Madelung rule, and as no other exceptions to the Madelung rule have ever been used to argue for changes to the periodic table, it has been argued that lanthanum and actinium should be treated likewise.
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In 2015, a IUPAC project chaired by Eric Scerri was set up to decide the question, giving only Sc-Y-La-Ac and Sc-Y-Lu-Lr as possible resolutions. In 2021, it decided on Sc-Y-Lu-Lr on the basis of three "desiderata": displaying all elements in order of increasing atomic number, avoiding a split of the d-block into "two highly uneven portions", and having the blocks follow the widths quantum mechanics demands of them (2, 6, 10, and 14). The Sc-Y-La-Ac form forces a split in the d-block between lanthanum and hafnium (and between actinium and rutherfordium), and the form with blank spaces under yttrium makes the f-block 15 elements wide even though quantum mechanics requires it to be 14 elements wide. While it was noted that 15-element-wide f-blocks are supported by some practitioners of a specialised branch of relativistic quantum mechanics focusing on the properties of superheavy elements, the report's opinion was that such interest-dependent concerns should not have any bearing on how the periodic table is presented to "the general chemical and scientific community".
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The most recently named elements – nihonium (113), moscovium (115), tennessine (117), and oganesson (118) – completed the seventh row of the periodic table. Future elements would have to begin an eighth row. These elements may be referred to either by their atomic numbers (e.g. "element 119"), or by the IUPAC systematic element names which directly relate to the atomic numbers (e.g. "ununennium" for element 119, derived from Latin "unus" "one", Greek "ennea" "nine", and the traditional "-ium" suffix for metallic elements). All attempts to synthesise such elements have failed so far. An attempt to make element 119 has been ongoing since 2018 at the Riken research institute in Japan. The Joint Institute for Nuclear Research in Russia also plans to make its own attempts at synthesising the first few period 8 elements.
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Currently, discussion continues regarding whether this future eighth period should follow the pattern set by the earlier periods or not, as calculations predict that by this point relativistic effects should result in significant deviations from the Madelung rule. Various different models have been suggested. All agree that the eighth period should begin like the previous ones with two 8s elements, and that there should then follow a new series of g-block elements filling up the orbitals, but the precise configurations calculated for these elements vary widely between sources. Beyond this series, calculations do not agree on what exactly should follow. Filling of the , 6f, 7d, and 8p shells is expected to occur in approximately that order, but they are likely to be intermingled with each other and with the 9s and 9p subshells, so that it is not clear which elements should go in which groups anymore. Scerri has raised the question of whether an extended periodic table should take into account the failure of the Madelung rule in this region, or if such exceptions should be ignored. The shell structure may also be fairly formal at this point: already the electron distribution in an oganesson atom is expected to be rather uniform, with no discernible shell structure.
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Nuclear stability will likely prove a decisive factor constraining the number of possible elements. It depends on the balance between the electric repulsion between protons and the strong force binding protons and neutrons together. Protons and neutrons are arranged in shells, just like electrons, and so a closed shell can significantly increase stability: the known superheavy nuclei exist because of such a shell closure. They are probably close to a predicted island of stability, where superheavy nuclides should have significantly longer half-lives: predictions range from minutes or days, to millions or billions of years. However, as the number of protons increases beyond about 126, this stabilising effect should vanish as a closed shell is passed. It is not clear if any further-out shell closures exist, due to an expected smearing out of distinct nuclear shells (as is already expected for the electron shells at oganesson). Furthermore, even if later shell closures exist, it is not clear if they would allow such heavy elements to exist. As such, it may be that the periodic table practically ends around element 120, as elements become too short-lived to observe; the era of discovering new elements would thus be close to its end.
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The periodic law may be represented in multiple ways, of which the standard periodic table is only one. Within 100 years of the appearance of Mendeleev's table in 1869, Edward G. Mazurs had collected an estimated 700 different published versions of the periodic table. Many forms retain the rectangular structure, including Charles Janet's left-step periodic table (pictured below), and the modernised form of Mendeleev's original 8-column layout that is still common in Russia. Other periodic table formats have been shaped much more exotically, such as spirals (Otto Theodor Benfey's pictured to the right), circles, triangles, and even elephants.
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The many different forms of the periodic table have prompted the questions of whether there is an optimal or definitive form of the periodic table, and if so, what it might be. There are no current consensus answers to either question. Janet's left-step table is being increasingly discussed as a candidate for being the optimal or most fundamental form; Scerri has written in support of it, as it clarifies helium's nature as an s-block element, increases regularity by having all period lengths repeated, faithfully follows Madelung's rule by making each period correspond to one value of + , and regularises atomic number triads and the first-row anomaly trend. He notes that its placement of helium atop the alkaline earth metals can be seen a disadvantage from a chemical perspective, but counters this by appealing to the first-row anomaly, pointing out that the periodic table "fundamentally reduces to quantum mechanics", and that it is concerned with "abstract elements" and hence atomic properties rather than macroscopic properties.
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The various sub-fields of AI research are centered around particular goals and the use of particular tools. The traditional goals of AI research include reasoning, knowledge representation, planning, learning, natural language processing, perception, and the ability to move and manipulate objects. General intelligence (the ability to solve an arbitrary problem) is among the field's long-term goals. To solve these problems, AI researchers have adapted and integrated a wide range of problem-solving techniques – including search and mathematical optimization, formal logic, artificial neural networks, and methods based on statistics, probability and economics. AI also draws upon computer science, psychology, linguistics, philosophy, and many other fields.
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By the 1950s, two visions for how to achieve machine intelligence emerged. One vision, known as Symbolic AI or GOFAI, was to use computers to create a symbolic representation of the world and systems that could reason about the world. Proponents included Allen Newell, Herbert A. Simon, and Marvin Minsky. Closely associated with this approach was the "heuristic search" approach, which likened intelligence to a problem of exploring a space of possibilities for answers. The second vision, known as the connectionist approach, sought to achieve intelligence through learning. Proponents of this approach, most prominently Frank Rosenblatt, sought to connect Perceptron in ways inspired by connections of neurons. James Manyika and others have compared the two approaches to the mind (Symbolic AI) and the brain (connectionist). Manyika argues that symbolic approaches dominated the push for artificial intelligence in this period, due in part to its connection to intellectual traditions of Descarte, Boole, Gottlob Frege, Bertrand Russell, and others. Connectionist approaches based on cybernetics or artificial neural networks were pushed to the background but have gained new prominence in recent decades.
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Unsupervised learning finds patterns in a stream of input. Supervised learning requires a human to label the input data first, and comes in two main varieties: classification and numerical regression. Classification is used to determine what category something belongs in – the program sees a number of examples of things from several categories and will learn to classify new inputs. Regression is the attempt to produce a function that describes the relationship between inputs and outputs and predicts how the outputs should change as the inputs change. Both classifiers and regression learners can be viewed as "function approximators" trying to learn an unknown (possibly implicit) function; for example, a spam classifier can be viewed as learning a function that maps from the text of an email to one of two categories, "spam" or "not spam".
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