Editing Alkali metal (section)

== Periodic trends ==
The alkali metals are more similar to each other than the elements in any other [[group (periodic table)|group]] are to each other.<ref name=rsc /> For instance, when moving down the table, all known alkali metals show increasing [[atomic radius]],<ref name=chemguide /> decreasing [[electronegativity]],<ref name=chemguide /> increasing [[reactivity (chemistry)|reactivity]],<ref name=rsc /> and decreasing melting and boiling points<ref name=chemguide /> as well as heats of fusion and vaporisation.<ref name="Greenwood&Earnshaw">{{Greenwood&Earnshaw2nd}}</ref>{{rp|75}} In general, their [[densities]] increase when moving down the table, with the exception that potassium is less dense than sodium.<ref name=chemguide />

=== Atomic and ionic radii ===
[[File:Effective Nuclear Charge.svg|thumb|250px|[[Effective nuclear charge]] on an atomic electron]]
The [[atomic radii]] of the alkali metals increase going down the group.<ref name=chemguide /> Because of the [[shielding effect]], when an atom has more than one [[electron shell]], each electron feels electric repulsion from the other electrons as well as electric attraction from the nucleus.<ref name=shielding>{{cite book |first1=Theodore |last1=L. Brown | first2=H. Eugene Jr. | last2=LeMay |first3=Bruce E. |last3=Bursten |first4=Julia R. |last4=Burdge |year=2003 |title=Chemistry: The Central Science |edition=8th |publisher=Pearson Education |location=US |isbn=978-0-13-061142-0 }}</ref> In the alkali metals, the [[valence electron|outermost electron]] only feels a net charge of +1, as some of the [[nuclear charge]] (which is equal to the [[atomic number]]) is cancelled by the inner electrons; the number of inner electrons of an alkali metal is always one less than the nuclear charge. Therefore, the only factor which affects the atomic radius of the alkali metals is the number of electron shells. Since this number increases down the group, the atomic radius must also increase down the group.<ref name=chemguide />

The [[ionic radii]] of the alkali metals are much smaller than their atomic radii. This is because the outermost electron of the alkali metals is in a different [[electron shell]] than the inner electrons, and thus when it is removed the resulting atom has one fewer electron shell and is smaller. Additionally, the [[effective nuclear charge]] has increased, and thus the electrons are attracted more strongly towards the nucleus and the ionic radius decreases.<ref name=rsc />

=== First ionisation energy ===
[[File:First Ionization Energy blocks.svg|thumb|upright=2.7|Periodic trend for ionisation energy: each period begins at a minimum for the alkali metals, and ends at a maximum for the [[noble gas]]es. Predicted values are used for elements beyond 104.]]
The first [[ionisation energy]] of an [[chemical element|element]] or [[molecule]] is the energy required to move the most loosely held electron from one [[mole (unit)|mole]] of gaseous atoms of the element or molecules to form one mole of gaseous ions with [[electric charge]] +1. The factors affecting the first ionisation energy are the [[nuclear charge]], the amount of [[shielding effect|shielding]] by the inner electrons and the distance from the most loosely held electron from the nucleus, which is always an outer electron in [[main group element]]s. The first two factors change the effective nuclear charge the most loosely held electron feels. Since the outermost electron of alkali metals always feels the same effective nuclear charge (+1), the only factor which affects the first ionisation energy is the distance from the outermost electron to the nucleus. Since this distance increases down the group, the outermost electron feels less attraction from the nucleus and thus the first ionisation energy decreases.<ref name=chemguide /> This trend is broken in francium due to the [[relativistic quantum chemistry|relativistic]] stabilisation and contraction of the 7s orbital, bringing francium's valence electron closer to the nucleus than would be expected from non-relativistic calculations. This makes francium's outermost electron feel more attraction from the nucleus, increasing its first ionisation energy slightly beyond that of caesium.<ref name="Uue" />{{Rp|1729}}<!--Also explain why the alkali metals have the lowest ionization energies in their period.-->

The second ionisation energy of the alkali metals is much higher than the first as the second-most loosely held electron is part of a fully filled [[electron shell]] and is thus difficult to remove.<ref name=rsc />

=== Reactivity ===
The [[Reactivity (chemistry)|reactivities]] of the alkali metals increase going down the group. This is the result of a combination of two factors: the first ionisation energies and [[atomisation energy|atomisation energies]] of the alkali metals. Because the first ionisation energy of the alkali metals decreases down the group, it is easier for the outermost electron to be removed from the atom and participate in [[chemical reaction]]s, thus increasing reactivity down the group. The atomisation energy measures the strength of the [[metallic bond]] of an element, which falls down the group as the atoms increase in [[atomic radius|radius]] and thus the metallic bond must increase in length, making the [[delocalised electrons]] further away from the attraction of the nuclei of the heavier alkali metals. Adding the atomisation and first ionisation energies gives a quantity closely related to (but not equal to) the [[activation energy]] of the reaction of an alkali metal with another substance. This quantity decreases going down the group, and so does the activation energy; thus, chemical reactions can occur faster and the reactivity increases down the group.<ref name="alkaliwater">{{cite web |url=http://www.chemguide.co.uk/inorganic/group1/reacth2o.html#top |title=Reaction of the Group 1 Elements with Water |last=Clark |first=Jim |year=2005 |work=chemguide |access-date=18 June 2012 |archive-date=31 May 2012 |archive-url=https://web.archive.org/web/20120531110524/http://www.chemguide.co.uk/inorganic/group1/reacth2o.html#top |url-status=live }}</ref>

=== Electronegativity ===
[[File:Periodic variation of Pauling electronegativities.svg|thumb|upright=1.25|Periodic variation of Pauling electronegativities as one descends the [[main group element|main groups]] of the periodic table from the [[period 2 element|second]] to the [[period 6 element|sixth period]].]]

[[Electronegativity]] is a [[chemical property]] that describes the tendency of an [[atom]] or a [[functional group]] to attract [[electron]]s (or [[electron density]]) towards itself.<ref name="definition">{{GoldBookRef|file=E01990|title=Electronegativity}}</ref> If the bond between [[sodium]] and [[chlorine]] in [[sodium chloride]] were [[covalent]], the pair of shared electrons would be attracted to the chlorine because the effective nuclear charge on the outer electrons is +7 in chlorine but is only +1 in sodium. The electron pair is attracted so close to the chlorine atom that they are practically transferred to the chlorine atom (an [[ionic bond]]). However, if the sodium atom was replaced by a lithium atom, the electrons will not be attracted as close to the chlorine atom as before because the lithium atom is smaller, making the electron pair more strongly attracted to the closer effective nuclear charge from lithium. Hence, the larger alkali metal atoms (further down the group) will be less electronegative as the bonding pair is less strongly attracted towards them. As mentioned previously, francium is expected to be an exception.<ref name=chemguide />

Because of the higher electronegativity of lithium, some of its compounds have a more covalent character. For example, [[lithium iodide]] (LiI) will dissolve in [[organic solvent]]s, a property of most covalent compounds.<ref name=chemguide /> [[Lithium fluoride]] (LiF) is the only [[alkali halide]] that is not soluble in water,<ref name=rsc /> and [[lithium hydroxide]] (LiOH) is the only [[alkali metal hydroxide]] that is not [[deliquescent]].<ref name=rsc />

=== Melting and boiling points ===
The [[melting point]] of a substance is the point where it changes [[state of matter|state]] from solid to liquid while the [[boiling point]] of a substance (in liquid state) is the point where the [[vapour pressure]] of the liquid equals the environmental pressure surrounding the liquid<ref>{{cite book |last=Goldberg|first=David E. |title=3,000 Solved Problems in Chemistry|edition=1st|publisher=McGraw-Hill|year=1988|isbn=978-0-07-023684-4}} Section 17.43, page 321</ref><ref>{{cite book |editor1=Theodore, Louis |editor2=Dupont, R. Ryan |editor3=Ganesan, Kumar |title=Pollution Prevention: The Waste Management Approach to the 21st Century|publisher=CRC Press|year=1999|isbn=978-1-56670-495-3|page=15 Section 27}}</ref> and all the liquid changes state to gas. As a metal is heated to its melting point, the [[metallic bond]]s keeping the atoms in place weaken so that the atoms can move around, and the metallic bonds eventually break completely at the metal's boiling point.<ref name=chemguide /><ref name="metallic-bonding">{{cite web |url=http://www.chemguide.co.uk/atoms/bonding/metallic.html |title=Metallic Bonding |last=Clark |first=Jim |year=2000 |work=chemguide |access-date=23 March 2012 |archive-date=25 July 2017 |archive-url=https://web.archive.org/web/20170725104821/http://www.chemguide.co.uk/atoms/bonding/metallic.html |url-status=live }}</ref> Therefore, the falling melting and boiling points of the alkali metals indicate that the strength of the metallic bonds of the alkali metals decreases down the group.<ref name=chemguide /> This is because metal atoms are held together by the electromagnetic attraction from the positive ions to the delocalised electrons.<ref name=chemguide /><ref name="metallic-bonding" /> As the atoms increase in size going down the group (because their atomic radius increases), the nuclei of the ions move further away from the delocalised electrons and hence the metallic bond becomes weaker so that the metal can more easily melt and boil, thus lowering the melting and boiling points.<ref name=chemguide /> The increased nuclear charge is not a relevant factor due to the shielding effect.<ref name=chemguide />

=== Density ===
The alkali metals all have the same [[crystal structure]] ([[body-centred cubic]])<ref name="Greenwood&Earnshaw" /> and thus the only relevant factors are the number of atoms that can fit into a certain volume and the mass of one of the atoms, since density is defined as mass per unit volume. The first factor depends on the volume of the atom and thus the atomic radius, which increases going down the group; thus, the volume of an alkali metal atom increases going down the group. The mass of an alkali metal atom also increases going down the group. Thus, the trend for the densities of the alkali metals depends on their atomic weights and atomic radii; if figures for these two factors are known, the ratios between the densities of the alkali metals can then be calculated. The resultant trend is that the densities of the alkali metals increase down the table, with an exception at potassium. Due to having the lowest atomic weight and the largest atomic radius of all the elements in their periods, the alkali metals are the least dense metals in the periodic table.<ref name=chemguide /> Lithium, sodium, and potassium are the only three metals in the periodic table that are less dense than water:<ref name=rsc /> in fact, lithium is the least dense known solid at [[room temperature]].<ref name="Greenwood&Earnshaw" />{{rp|75}}