Periodic Trends

Representative Groups: These are the “A” groups on the P.T. Group # with the letter A next to it tells # valence electrons.

Valence Electrons Elements in the same group have similar properties because they have the same number of valence electrons. Valence Electrons – Electrons in the outermost s and p sublevels. What elements are these? What group?

Valence Electrons The number of electrons within an energy level increases by one as you move from left to right across the periodic table. Elements in the same column have the same number of valence electrons. (There are exceptions in the transition elements.)

Valence Electrons Group Number e- ending Valence e- 1A s1 1 2A s2 2 3A 4 5A p3 5 6A p4 6 7A p5 7 8A p6 8 Exception: He

Oxidation Number **Essential Vocabulary** Group Oxidation Number 1A 1+ 2A 2+ 3A 3+ 4A 4+/4- 5A 3- 6A 2- 7A 1- The positive (cation) or negative (anion) charge of an ion. Predicted by the group/family (column).

X Electron Dot Diagram Write the symbol. Put one dot for each valence electron Don’t pair electrons up until they have to pair up. X = hypothetical element X

Write the Electron Dot Structure for Nitrogen. Nitrogen has 5 valence electrons. First we write the symbol. Then add 1 electron at a time to each side. N

Practice: Write the electron dot diagram. Na Mg F Ne

Shielding The “interference” of the nuclear attraction from lower level electrons. Reduced the nuclear pull on electrons on the outer most (valence) energy level. Nuclear refers to nucleus not necessarily radioactive elements.

Atomic Radii BECAUSE: This is the size of an atom (atomic radius). Remember: The positive nucleus is attracting the electrons. Size decreases as you go across a period (left  right). BECAUSE: There is an increased positive charge in the nucleus because of more protons. This exerts a stronger pull on electrons (within the same energy level), thus drawing them toward the nucleus.

Atomic Radii Remember: The positive nucleus is attracting the electrons. The Size increases as you move down a group. This is caused by the addition of energy levels as you go down a group. This gives the electrons more “space to expand”. The outermost electrons are not as attracted by the nucleus (they are further away and “shielded” by the inner electrons).

Practice: Atomic Radii Does Na or Mg have a greater atomic radius? Does Na or K have a smaller atomic radius?

Ion Radii: Loss of gain of an electron creates an ion (to achieve a stable octet). Positive Ions (cation) = loss of electrons Negative Ions (anion) = gain of electrons Ion Radii: Size of an ion (an atom that has lost or gained electrons). Losing electrons (cations) causes ions to be smaller because protons outnumber electrons and there is less shielding – protons are pulled in closer to the nucleus. The size of an ion increases as you go down a group. Gaining electrons (anions) causes ions to be larger because more electrons increases repulsion between the electrons and electrons outnumber protons.

Ionization Energy Ionization energy is the energy needed to remove an electron from an atom. Removing one electron makes a +1 ion. The energy required is called the first ionization energy. What does it mean? Tells how difficult it is to remove 1 electron.

Ionization Energy Generally increases as you go across the periodic table due to the pull of the increased nuclear charge. Stronger the pull the harder it is to take an electron. Generally decreases as you go down a group due to distance from the positive nucleus and shielding of the electrons. Must take 3 things into consideration: Energy level # valence electrons in s &/or p sublevel Shielding

He has a greater IE than H same shielding (Same E level) greater nuclear charge (more p+) He H First Ionization energy Atomic number

more shielding (2nd E level). outweighs greater nuclear charge Li has lower IE than H more shielding (2nd E level). outweighs greater nuclear charge He H First Ionization energy Li Atomic number

same shielding (same E level) greater nuclear charge (more p+) Be has higher IE than Li same shielding (same E level) greater nuclear charge (more p+) He First Ionization energy H Be Li Atomic number

greater nuclear charge By removing an electron we make s orbital full. B has lower IE than Be same shielding greater nuclear charge By removing an electron we make s orbital full. He First Ionization energy H Be B Li Atomic number

First Ionization energy He First Ionization energy H C Be B Li Atomic number

First Ionization energy He N First Ionization energy H C Be B Li Atomic number

First Ionization energy Breaks the pattern because removing an electron gets to 1/2 filled p orbital He N H C O First Ionization energy Be B Li Atomic number

First Ionization energy He F N First Ionization energy H C O Be B Li Atomic number

First Ionization energy Ne has a lower IE than He Both are full, Ne has more shielding He Ne F N O First Ionization energy H C Be B Li Atomic number

Na has a lower IE than Li Both are s1 Na has more shielding He Ne F N H C O First Ionization energy Be B Li Na Atomic number

IE for 1st 37 Elements First Ionization energy Atomic number Period 4 transition metals Atomic number

Practice: Ionization Energy Which element has the greatest ionization energy? F or Br? Ca or K? Na or Ne?

Electronegativity This is the tendency of atom to attract electrons in a chemical bond.

Electronegativity Electronegativity increases across a period due to the increased positive charge of the nucleus. Electronegativity decreases down the group due to the distance of outermost electrons and increased shielding.

Reactivity The number of valence electrons, ionization energy and electronegativity of an atom are all indicators of reactivity.

Reactivity Metals are more reactive if they have a low number of valence electrons and low ionization energy. The most reactive metal is Francium. Nonmetals are more reactive in they have larger numbers of valence electrons and high electronegativity. The most reactive non metal is Fluorine.

Element Phases at Room Temperature Liquid at room temperature: Bromine and Mercury.

Gases at Room Temperature. Note the Noble Gases Gases at Room Temperature *Note the Noble Gases* Note that all are non-metals Hydrogen Nitrogen Oxygen Fluorine Chlorine *Helium *Neon *Argon *Krypton *Xenon *Radon

Solids at Room-Temperature The rest of the elements are solid at room temperature. Yellow – Solid Red – Liquid Purple - Gas

*Diatomic Elements* Covalent Bonding can occur between two atoms of the same element. As pure elements, these elements exist as diatomic molecules (bonded to another atom of the same element). H2, N2, O2, F2, Cl2, Br2, I2 (Lucky 7)

The Periodic Table The End