The following notes were taken by me while listening to Chem 1A Lectures 11 and 12 by Professor Heino Nitsche



Atomic Number and Mass: Each element in the table has an assigned atomic number. Hydrogen (H) has atomic number 1, Helium (He) has atomic number 2, etc.

The atomic number indicates the number of protons in the nucleus. The total atomic weight of the element is equal to the sum of the weights of the nucleons (protons and neutrons) and electrons. An element is defined by the number of protons, but can have differing numbers of neutrons, resulting in different isotopes. For example, Carbon-12 has 6 protons and 6 neutrons, Carbon-13 has 6 protons and 7 neutrons, and Carbon-14 has 6 protons and 8 neutrons.

As the atomic number increases, the number of protons, neutrons, and electrons in each subsequent element increases, and thus the mass of the element also increases. Groups: The elements in a given group (column) of the periodic table have similar chemical properties because they have the same number of valence electrons.

In general, only the valence electrons take part in bonding with other atoms; thus, elements in a group have similar bonding behavior since they have the same number of valence electrons.

For example, in the first column (group) of the periodic table, the elements H, Li, and Na all have one valence electron in the s orbital, and thus only have one electron available for bonding with other atoms.

Similarly, in the last group (column) of the Periodic Table, all of the elements have their valence electron shells filled, so they are (for the most part) un-reactive with other elements; for this reason, we call this last group the "noble gases".

Elements within a group are called "homologues", because they exhibit similar chemical behavior. Periods: The elements in a given row (called a "period") all have the same quantum energy level (the "n" quantum number").

For example, all the elements in the first row have n=1, all the elements in the second period have n=2, etc. Each subsequent row (period) has a higher energy level for its ground state, since higher "n" quantum numbers indicate higher energy states.

As you traverse a period from left to right, the valence shells are filled, following the Aufbau principle

In the topmost period, the Hydrogen atom (H) has a single 1s electron, and Helium (He) has a filled valence shell with two electrons in the 1s shell.

In the second period, Li (Lithium) has one electron in the 2s shell, Be has two electrons in the 2s shell, B has three valence electrons (two 2s electrons, and one 2p electron).

Completing the 2nd period, Ne, a noble gas, has a completely filled valence electron shell with two 2s electrons and six 2p electrons.

In the 3rd period, you start with one electron in the 3s orbital (Na), and then add a second electron to the 3s orbital (Mg), and then fill the 3p orbital with six electrons as you progress from Al to Ar.

In the 4th period, there is an exception. After filling the 4s orbital with two electrons (Ca), you then fill the 3d orbitals with 10 electrons (Sc -> Zn).

Because the 3d orbitals are higher in energy than the 4s orbitals, but lower in energy than the 4p orbitals, we must first fill the 3d orbitals.

Similarly, for the 5th period, we fill the 4d orbitals after the 5s orbitals, but before filling the 5p orbitals.

In the sixth period, we first fill the 6s orbitals (Cs and Ba), and then add one electron to the 6p orbital (La). Before filling the rest of the 6p orbitals, we must first fill the 4f orbitals. These elements are shown in a separate row called the "Lathanide Series". One could argue that the Lanthanides, beginning with element number 58, should be listed in the sixth row following element number 57, La. After filling the 4f orbitals, we then continue in period 6 by filling the 5f orbitals. Finally, we finish filling the 6p orbitals, ending on Rn.

The seventh period is analogous to the sixth period. After filling the 7s orbital with two electrons (Ra), and adding a single 7p electron (Ac), we then need to fill the 5f orbitals as they are higher in energy than the 6p1 orbital, but lower in energy than the 6p2 (and 5d) orbitals. These elements are shown separately as the "Actinide series". Then, we fill the 5d orbitals, and finally the 6p orbitals.



Group Names: The elements in the first group (column) of the Periodic Table are called Alkali Metals.

The elements in the second group (column) of the Periodic Table are called Alkaline Earth Metals.

Groups (columns) 3 through 12 are called "Transition Elements".

The elements in the second to last column (group) of the Periodic Table are called Halogens.

The elements in the last group (column) of the Periodic Table are called Noble Gases. With few exceptions, these elements are not very reactive, as their valence shells are completely filled. Atomic Radius Atomic radius decreases within a period, and increases within a group.

Within a period, as more protons are added, the protons bind more tightly to the electrons, reducing the radius.

Within a group, the effective charge of the protons is reduced by the shielding of the filled electron shells, and the outermost electrons are bound less tightly resulting in a larger atomic radius. Ionization Energy: The ionization energy is the amount of energy needed to extract a single electron from an atom, forming a positive ion (cation).

To remove an electron from an atom, one must add energy to that atom (endothermic).

Looking at a Group (column), one expects the ionization energy to decrease as one goes down the column. For example, Na (Sodium) has a higher ionization energy than Rb, because Rb has more filled electron shells shielding the protons, and thus has a lower effective positive charge with which it can bind the outermost electron.

By choosing the "Ionization Energy" radio button above, and then moving the scrollbar, you can see that the ionization energy decreases within a column (group) .

. Within a period, the ionization energy generally increases as the number of protons increases. There are some exceptions as the angular momentum quantum number changes (i.e. as we switch from s orbitals to p orbitals), and also as the first electron with anti-parallel spin is added while following Hund's Rule.

By choosing the "Ionization Energy" radio button above, and then moving the scrollbar, you can observe that the ionization energy increases within a row (period) .

. Note that the 1st ionization energy (energy required to remove the 1st electron) is less than the 2nd ionization energy (energy required to remove the second electron) because after removing the 1st electron, the remaining electrons are more tightly bound by the protons, causing the ionization energy to increase for each subsequent electron removal. Electron Affinity: The electron affinity indicates how likely an atom is to acquire another electron, forming a negative ion (anion).

When an atom acquires another electron, energy is released (exothermic).

Within a group, one expects the electron affinity to decrease as one goes down the column. For example, Cl has a higher electron affinity (is more likely to pull an electron from another atom) than Br, because Br has an additional filled electron shell shielding the protons, and thus has a lower effective positive charge with which to attract and bind to electrons.

By choosing the "Electron Affinity" radio button above, and then moving the scrollbar, you can observe how the electron affinity decreases within a column (group) .

. Similarly, you can observe how electron affinity increases within a period (row) as the number of protons increase. Electronegativity Electronegativity is the ability of an atom within a molecule to pull electrons away from a binding partner (another atom in the molecule).

Electronegativity increases within a Period as one travels from left to right in the Periodic Table.

Electronegativity decreases as one traverses a group (column) from top to bottom.

Francium (Fr) is the least electronegative element (we often say it is “electropositive”).

Flourine (F) is the most electronegative element; in a molecule, it will try to pull the binding electrons from other atoms in the molecule towards itself.

Looking at the difference between the electronegativity of two bonded atoms If the difference is > 2.0, the bond is ionic. For example, NaCl is an ionic bond, as Sodium (Na) has “given” its electron to Chlorine (Cl). If the difference is between .4 and 2.0, the bond is polarized covalent (partly ionic). For example, in HCl, the bond is covalent, but the shared electron spends most of its time closer to the Cl atom than the H atom, forming a dipole If the difference is less than .4, the bond is covalent, and the electron is shared equally. For example, Cl-Cl is a perfectly covalent bond since both partners have the same electronegativity

If two atoms have a large difference in electronegativites, a large amount of energy will be released upon forming the bond. If two atoms have a small difference in electronegativities, the energy released on formation of the bond will be small. Observations Halides, having high ionization energies and high Electron Affinity, are more likely to acquire electrons from other atoms.

Metals, having low ionization energies, and low Electron Affinity, are more likely to give electrons to other atoms.

Non-metals often form anions (high electron affinity)

In general, elements on the right side of the periodic table are more likely to form anions (high electron affinity combined with high ionization energy), and elements on the left side of the period table are more likely to form cations (lower ionization energy and lower electron affinity).