An Introduction to the Periodic Table


  • Tell us how many valence electrons are in the outermost s and p shells
  • Elements with similar properties are found in the same group


  • Tell us which shell the valence electrons are found; each row has valence electrons in same energy lvl

Groups of Periodic Table Group 1: Alkali Metals

  • Low density, soft, silver, very reactive, all salts, form strong bonds with water

Group 2: Alkali Earth Metals

  • Stronger, denser, less reactive, one more valence electron

Group 3-12: Transition metals

  • D-block elements Diagonal row: Metalloids
  • Poor conductors, both properties Next: Nonmetals
  • Nonconductors, dull, brittle

Group 7: Halogens

  • Most reactive cuz have 7 ve

Group 8: Noble Gasses

  • Non reactive cuz have 8 valence electrons and are stable

              ○ 2 in the s and 6 in the p

  • Most gasses at room temperature
  • When solid take crystal form
  • Can run electricity thru noble gasses but won’t react

Polyelectronic Atoms

  • Polyelectronic atoms: atoms with more than one electron
  • Three energy contributions in the description of an atom

○ The kinetic energy of the electrons as they move around the nucleus

○ Effective Nuclear Charge

○ The potential energy of repulsion between electrons.

Effective Nuclear Charge

Effective nuclear charge (Zeff): attraction between the nucleus and the valence electrons

      ○ Note: nuclear charge is the total (+) charge of the protons and attraction to all the 

      ○ 2nd Note: Nuclear charge increases both down a group and across a period; Zeff weakens down a group and increases across a period

              ■ Zeff depends more on distance than # of protons

 ● When justifying trends talk about Zeff and why

Weakens moving down a group

Attraction between valence electrons and nucleus weakens bcuz of increasing number of filled energy lvls and thus distance between them

  • Valence electrons are shielded more from the nucleus
  • Coulomb’s law: attraction between 2 charged particles is inversely related to the distance

Strengthens from left to right across a period

Electrons are being added to the same principal energy lvl & are more strongly attracted to the nucleus due to additional protons

  • Valence electrons are less shielded from the nucleus (due to added e- pair repulsions)
  • Coulomb’s law: attraction between 2 charged particles in directly related to the magnitude of their charges

Atomic Radius

Size of an atom, distance from the nucleus to the outermost electrons (half the distance between two nuclei of a molecule)

Is determined by how much the electrons are attracted to the (+) nucleus

  • Number of protons and electrons determines the size of atoms and ions

Atomic Radius Trends

Group Trend: increases as you go down

  • Bcuz Zeff decreases (less attraction = larger size)

Periodic Trend: decreases as you go across (From left to right)

  • Bcuz Zeff increases → (more (+) nuclear charge = more attraction= smaller size)

Ionic Radius

Size of an atom when it is an ion

Ionic Radius Trend

Metals lose electrons → highest occupied energy shell of atom is further away from highest occupied energy shell in ion (Stronger Zeff so…)

  • Ionic radius<atomic radius

Nonmetals gain → more e- that p+ (weaker Zeff bcuz of stronger electron-electron repulsions so…)

  • Ionic radius>atomic radius

Group Trend: increases as you go down

Period Trend: decreases as you go across

Explain Trends

  • Big jump between Carbon and nitrogen ions (10 e-) bcuz they have different number of energy lvls
  • Decrease nitrogen to fluorine bcuz increasing Zeff

Isoelectronic Ions

Ions containing the same number of total electrons:

  • All have 10 electrons and the electron configuration

Ionization Energy

  • Energy needed to remove a valence electron and form an ion
  • First ionization energy (I1): the energy required to remove the highest-energy (outermost) electron of an atom

             ○ The amount of energy required to remove a second or third electron from an atom is always more than the amount of energy required to remove the first electron

                        ■ I1 < I2 < I3

             ○ Why? Bcuz the Zeff increases as you remove electrons bcuz of reduced electron-pair repulsions & e- are closer to nucleus→ attraction that remaining electrons feel is greater

  • Very important when explaining ionization energy trends

            ○ It takes more energy to remove core electrons than valence e- (in outermost s and p)

            ○ Takes more energy to remove an electron from a half or fully filled electron shell

  • Exceptions in Ionization Energy trend

            ○ Slight decrease in IE from beryllium to boron (group 2 to group 13) bcuz more energy is required to remove an electron from a filled 3s sublvl than from a partially filled 3p sublvl

            ○ Slight decrease in IE from from nitrogen to oxygen (group 15 to 16) bcuz it takes more energy to remove an electron from a half filled 2p sublvl than from a partially filled 2p sublvl (due to the added electron repulsions)

  • Ex: Explain this trend for Mg

           ○ Have to answer…

                 ■ a) basic reason for trend (see above)

                 ■ b) why such a big jump between I2 and I3

  • Bcuz after removing I2 you are no longer removing valence electrons but core electrons from a filled energy lvl → takes much more energy
  • Note: AP exam often focuses on this trend

Ionization Energy Trend

  • Group Trend: decreases down a group

                    ○ Because Zeff is weakening

  • Periodic Trend: increases across a period

                    ○ Because Zeff strengthens


  • Measure of how much an atom in a molecule attracts shared electrons to itself; affected by atomic

radius and atomic number

                ○ Associated with the production of a negative ion Electronegativity Trend

  • Group Trend: decreases as you go down

                 ○ Zeff decreases

  • Period Trend: increases across

               ○ Zeff increases; atom less likely to give up electron

Electron Affinity

  • The energy change associated with adding an electron to a gaseous atom

        ○ The more negative the energy, the more energy is released (exothermic)


  • Metals tend to have positive electron affinities while nonmetals tend to have negative electron



  • “Ability of an atom or compound to undergo a chemical reaction with another atom”
  • Metals lose e- when react so reactivity based on low ionization energy

                    ○ Low I.E = high reactivity

  • Nonmetals gain e- when react so reactivity based on high electronegativity

                  ○ High electronegativity = high reactivity

Properties of Nonmetals

  • The ability to gain one or more electrons to form an anion when reacting with a metal. Thus nonmetals are elements with large ionization energies and the most negative electron affinities.
  • Non-conductors (no flow of electrons)
  • Dull, brittle
  • Most gasses at room temperature and form crystals when solid Metalloids
  • exhibit both metallic and nonmetallic properties under certain circumstances

The Properties of a Group: The Alkali Metals

Alkali Metals

  • Most reactive of metals, low density, soft, silver, all salts, form strong bonds with water
  • Note: even though hydrogen is in Group 1A of the periodic table, it behaves as a nonmetal, bcuz of its very small size → electron in the small 1s orbital is bound tightly to the nucleus