What characterizes the electron configurations of transition metals?
Transition metals are characterized by the filling of an inner electron orbital as atomic number increases.
How do the electron configurations of the transition metals differ from those of the metals in groups 1 and 2?
How do the electron configurations of the lanthanide and actinide elements differ from the electron configurations of the other transition metals?
transition elements are not as reactive as alkali metals or alkaline earth metals.. this is so because of the valence electronic configurations. alkali metals have ns1 configurations which have easy chances of loosing electrons so as to gain the stable state. Attaining a stable state is the law of nature. hence to attain it they have to redily donate electron. this is not so easily possible for the transition elements..
The physical properties of transition metals are determined by their electron configurations. Most transition metals are hard solids with relatively high melting and boiling points. Differences in properties among transition metals are based on the ability of unpaired d electrons to move into the valence level. The more unpaired electrons in the d sublevel, the greater the hardness and the higher the melting and boiling points.
Early transition metals are does starting at the beginning of the transition metals (i.e. Sc) and going through about d5 which would be Mn. These metals are less electron rich as compared to the so-called "late" transition metals and the chemistry of each is somewhat different and definitely unique. Hardness and softness of the each of these groups changes (see Hard Soft Acid Base Theory) as does the stable oxidation states and coordination numbers.
Because they have different amounts of outer electrons and electron configurations. The electron configurations define the structure and bonding of the elements which defines the properties. However, elements in the same groups have the same amount of outer electrons and therefore the same structure and bonding ( except metals and non- metals which have different types of bonding) and therefore have similar properties.
IUPAC define it as; "an element whose atom has an incomplete d sub-shell, or which can give rise to cations with an incomplete d sub-shell." Transition metals are all; paramagnetic when they have one or more unpaired d shell electron Transition metals are also able to exist in a number of oxidation states. Metallic bonding resulting from the delocalised d electrons means that transition metals have high boiling and melting points.
How do the stable ions of the transition metals differ from those formed by the metals in group one and two?
In the Group 1 and Group 2 elements, these metals want to loan out electrons to achieve what is called inert gas electron configuration, which is a full outer electron shell or valence shell. Because of their electron configurations as elemental metals, they are in a big hurry to do this (they are very reactive). In the ion, the metal has already loaned an electron (in the case of the Group 1 metals) or two…
All the transition elements have 2 electrons in their outer shell (4th energy level). If one element has one more electron than another, it is because it has one more electron in its inner shell (third energy level). So instead of adding an electron to the outer shell, transition metals with more electrons add it to their 3rd energy level. Therefore the number of electrons in the outer shell stays the same (2) and since…
They do! Bonding with transition metals can definitely be drawn using Lewis dot structures. They are generally not taught until you reach high level chemistry courses however because they do not follow the 8-electron or octet rule as most other elements as taught in high school. Transition metals have valence bonding electrons in 3 different orbitals: s, p, and d (as opposed to say, carbon, oxygen, nitrogen, etc. that only have s and p valence…
Metals tend to be on the left of the transition elements and therefore they have less valence electrons in their outer shell. Nonmetals are on the the right of the transition metals and those nonmetals almost have in their electron outer shell full and the number they have in the outer shell represents the valence electrons.
The alkali and alkali-earth metals want to lose electrons. Alkali metals need to only lost one electron to have a full outer shell (8). Alkali-earth metals need to lose 2 electrons. Transition metals are different, but remember that the outermost shell in a transition metal is NOT a d orbital. Take for example the electron configuration of the transition metal Cobalt. Co[Ar] 4s23d7 The outer most shell is the 4s orbital and not the 3d…
How is the number assigned to a family related to the electron configuration of each element in that family?
The number assigned to a family is related to the electron configuration of each element in that family, in that each family contains elements with the same valence electron configuration. The most common element families used are alkali metals, alkaline earth metals, transition metals, halogens, and noble gases.
Some early transition metals are scandium, titanium, yttrium, and zirconium. Also on the list of early transition metals are hafnium and rutherfordium. Some early transition metals are scandium, titanium, yttrium, and zirconium. Also on the list of early transition metals are hafnium and rutherfordium.