Atoms and Elements - AP Chemistry

Cadmium normally has electrons, but only has electrons.

The normal electron configuration for is as follows:

Since the cadmium is losing electrons, it must lose them from the highest energy shell. In this case, this means the cadmium will be losing its electrons in .

Thus, the electron configuration for is .

The normal electron configuration for is as follows:

Since it is losing electrons, the element must lose the electrons from the highest energy shell first. Thus, the element loses electrons from and electron for .

The electron configuration for is then

Start by finding the noble gas core. For iron, this will be argon as this is the noble gas that is closest to it.

Next, recall that since the orbitals are higher in energy that the orbitals, electrons will be lost from the orbital first.

The normal electron configuration for is as follows:

has lost electrons. It will lose the first two electrons from the shell, then it will lose electron from the shell, giving it the following electron configuration:

Start by finding the noble gas core. For tungsten, this will be xenon as this is the noble gas that is closest to it.

The normal electron configuration for is as follows:

Recall that electrons are lost in the highest energy level subshell first.

has lost electrons. It will lose the first two electrons from the shell, then it will lose electron from the shell, giving it the following electron configuration:

Electrons of an atom are located within electronic orbitals around a nucleus. The electrons of each atoms have their own specific energy level called principal energy level. When electrons are excited by absorbing energy the electrons can jump to a high energy level. Then when an electron drops back to a lower energy level the electron emits the energy. Therefore, when an atom moves from a lower energy state to a higher energy state. the electrons absorb energy.

Each element has a unique electron configuration that represents the arrangement of electrons in orbital shells and sub shells. There are four different orbitals, s, p, d, and f that each contain two electrons. The p, d, and f orbitals contain subshells that allow them to hold more electrons. The orbitals for an element can be determined using the periodic table. The s-block consists of group 1 and 2 (the alkali metals) and helium. The p-block consists of groups 3-18. The d-block consists of groups 3-12 (transition metals), and the f-block contains the lanthanides and actinides series. Using this information we can determine the full electron configuration of sodium.

To do this, start at hydrogen located at the top left of the periodic table. Hydrogen and helium are in the first s orbital and account for . Next, we move to the second s-orbital that contains lithium (Li) and beryllium (Be), which accounts for . Then we move to boron, carbon, nitrogen, oxygen, fluorine, and neon, which are all in the p-block and account for . There is no 1p orbital. Finally, we are at sodium, which is in the s-block and accounts for . Therefore the full electron configuration of sodium is .

Each element has a unique electron configuration that represents the arrangement of electrons in orbital shells and subshells. There are four different orbitals, s, p, d, and f that each contain two electrons. The p, d, and f orbitals contain subshells that allow them to hold more electrons. The orbitals for an element can be determined using the periodic table. The s-block consists of group 1 and 2 (the alkali metals) and helium. The p-block consists of groups 3-18. The d-block consists of groups 3-12 (transition metals), and the f-block contains the lanthanides and actinides series. Using this information we can determine the electron configuration of iodine in nobel gas configuration.

The nobel gas configuration is a short hand to writing out the full electron configuration. To do this, start at the nobel gas that come before the element of interest. In the case of iodine, the nobel gas is krypton. Therefore, the electron configuration will begin with , and this will be the new starting place for the electron configuration.

After krypton comes the s-block, which contains elements with the atomic numbers 37 and 38 that account for . Then comes the d-block containing elements 39-48 that account for . Finally comes the p-block containing elements 49-53 that account for . Therefore, the electron configuration of iodine in nobel gas configuration is .

Atomic radius increases down the columns and to the left. The furthest left elements are Na and Cs. Cs is furthest down so it is the biggest.

Atomic radius expands down the columns and to the LEFT.

Atomic radius INCREASES DOWN and TO THE LEFT in the periodic table

Electronegativity increases across a period (going right) and decreases down a group \[not including noble gases\]. The element closest to the top right is Phosphorous

If we look at elements of the same period the prinicipal quantum number for each one of these elements is the same. Thus the outermost energy level for the electrons of each atom is the same. If we consider moving through the periodic table in numerical order, the left side of the periodic table features atoms that have just begun to add electrons to new energy levels. On the right side of the periodic table the elements are moving closer to filling the energy level. The differences in atomic radii are the result of differing amounts of protons between atoms whose electrons are in the same energy level. K has one electron in the fouth energy level and Kr has eight electrons in the fouth energy level. K has 19 protons with which to generate pull on the electrons. Kr has 36 protons with which to generate pull on the electrons. Thus Kr has a smaller atomic radius becasue of its ability to have a tighter grasp on its electrons becasue of the stronger charge generated by its nucleus.

Electronegativity measures the ability an atom to attract shared electrons
in bond. It follows a trend that it increases moving from left to right and down to up across
the periodic table.

The trend for radius is that it increases down and to the left

Atomic radius decreases as one moves across the periodic table from left to right, since effective nuclear charge increases and the electrons are held more tightly to the nucleus.

The alkali metals are found in the first group (column) of the periodic table, on the leftmost side. They have only 1 loosely bound electron in their outermost shells, and their effective nuclear charge values are low, giving them the largest atomic radii of all the elements in their periods.

The transition metals are capable of losing various numbers of electrons from the s and d orbitals of the valence shell. Metals such as Cu, Fe, and Mn have various oxidation states and can form many different ionic compounds.

Electronegativity increases as one moves across a period (row) from left to right, or up a group (column) from bottom to top. Following these trends, fluorine is the most electronegative element.

Effective nuclear charge increases as one moves across a period (row) from left to right, so Cl will have a higher effective nuclear charge. Anions of an element have a lower effective nuclear charge than their parent atom, since they have more electrons than protons and feel less of a pull from the protons in the nucleus. Thus, Cl will have the highest effective nuclear charge.

Ionization energy is the energy required to completely remove an electron from an atom. The electron that would be removed is the atoms most loosely held electron. Some atoms are more likely to give up an electron to get to a more stable electron configuration. This is the reason that Group I elements have the lowest first ionization energy because after loosing one electron, these elements have now acheived an octet configuration. The second ionization energy is the energy to remove a second electron from an atom. This will be lowest for the Group II elements because these elements acheive an octet configuration after loosing two electrons.