Atomic radius has two general trends which you should remember:

1. Atomic radius will decrease when moving left to right along a period.

2. Atomic radius will increase when moving down a group.

Since potassium (K) and calcium (Ca) are farther down the group than argon (Ar) and chlorine (Cl), we conclude that they are the largest atoms in the set. Because Ca is to the right of K, it is slightly smaller than K. As a result, K has the largest atomic radius in the set followed by Ca. Since Ar is to the right of Cl, Cl has a larger atomic radius than Ar.

The decreasing order is K, Ca, Cl, Ar.

First, you should see that all four of these ions have the same amount of electrons, resulting in a stable \[Kr\] electron shell. These ions differ, however, by the number of protons in their nuclei. Since the heavier ions have more protons pulling on the same amount of electrons, the atomic radius will be smaller, as the negative electrons are drawn inward toward the positive nucleus.

Sr2+ has the most protons in its nucleus out of this set, so it will have the smallest atomic radius. The atomic radius will increase with decreasing atomic number.

The correct order from largest to smallest is Se2-, Br-, Rb+, and Sr2+.

Elements within a group have the same number of valence electrons, but in increasing energy levels. Elements toward the bottom of a group have valence electrons with higher energies in larger orbitals. This results in a larger radius and a weaker attractive force between the nucleus and outer electrons. The ionization energy decreases as the electrons are more removed from the attraction of the nucleus.

When moving down a group, atomic radius increases and ionization energy decreases.

Energy level increases moving down a group of the periodic table. As energy level increases, the outer valence shell becomes more distant from the nucleus, causing atomic radius to increase.

Energy level remains constant across a period, but electrons are added within the same orbitals. When new electrons are added within the same orbital, additional protons are also added to the nucleus. This increases the effective nuclear charge, pulling the electrons closer to the nucleus. The trend for atomic radius is to decrease as we move right along a row.

This means that the general trend for atomic radius is to increase as one moves to the left and downward on the periodic table.

Nitrogen, phosphorous, antimony, and bismuth are all in the same group (column) of the periodic table.

The atomic radius increases from the top of a group to the bottom, due to increased principle shell number (n). As one travels down a group, another s shell is added, meaning that electrons are added in another orbit farther from the nucleus. This serves to increase the atomic radius of the atom.

Atomic radius increases with increasing effective nuclear charge (Z). Elements toward the right and toward the top of the periodic table have the highest Z values. Protons and electrons are added in pairs as we traverse the periodic table from left to right. A attractive force is established between the positively-charged nucleus and the negatively-charged electron cloud, which increases as the number of particles grows.

When electrons are added or taken away without the same happening to a proton, an imbalance of charge accumulates. When more electrons are present than normal, the electron cloud sags farther away from the nucleus. When fewer electrons are present than normal, the electron cloud is drawn in more tightly toward the nucleus. Atoms with extra electrons (a negative charge) will have larger nuclei than their neutral counterparts. A chloride ion will thus has a larger atomic radius than argon, a potassium ion, or a calcium ion.

Ionization energy describes the energy required to remove a single electron from an atom. The periodic table trend for ionization energy is to increase from left to right or a period, and decrease from top to bottom of a group. Typically, elements that have seven valence electrons, one electron away from having a complete valence shell, have very high ionization energies. These elements are found adjacent to noble gases on the periodic table: the halogens. Chlorine would thus have a very high ionization energy.

Alkali metals have the lowest ionization energies, as removing a single electron will give them a full octet configuration.

Ionization energy refers to the amount of energy necessary to remove a single electron from an atom or ion in the gas state, resulting in an increase in charge.

To find our answer, we need to know that ionization energy increases as we move left to right across a period and decreases as we travel down a group. Alkali metals, on the left side of the table, gain stability with the removal of a single electron, while halogens lose stability. Atoms or ions with a electron configurations similar to those of the noble gases are very stable. Due to this stability, larger amounts of energy are required to alter the electron configuration.

When a halogen gains an electron, it achieves a stable valence octet; thus, the fluoride and bromide ions are incredibly stable and will have very high first ionization energies. Bromine is lower on the periodic table than fluorine, so its electrons are separated from its positively-charged nucleus by more energy shells. This effect, called shielding, leads to a smaller amount of energy required to remove one electron from bromine. , therefore, is our solution.

The correct answer is chlorine. The most electronegative elements are those in the upper right of the periodic table, with the exception of the noble gases. Electronegativity describes how easily an element will gain an electron. The halogens (second to last group) "want" an extra electron to complete their valence shell. Iodine and chlorine are both halogens. Chlorine, however, has a smaller atomic radius, and therefore a smaller distance between the protons and outer electrons. Chlorine thus has a stronger attraction for an additional electron due to the greater effective nuclear attraction.

As you ionize atoms, you generate charged ionic species. These charges will resist further ionization (cations will more strongly attract the electrons you are trying to pull away). Noble gas configurations are particularly stable, so you would expect a large increase in needed energy to ionize away from this state.