UNIT 3 NOTES Fall 2013

 Elements in the same column had the same:  physical and chemical properties  Valence number  Elements in the same row had the same:  Number of electron shells Metalloids – elements along zigzag line in periodic table  Have properties of metals and nonmetals

Alkali Metals Group 1A s orbital contains last electron form cations, always 1+ soft metals most reactive

Alkaline Earth Metals Group 2A s orbital is last orbital filled form cations, always 2+ metals reactive

Transition Metals Groups 1B – 8B Last electrons are in d orbitals Form cations Many transition metals can form several cations with different charges. Most are metals with high density and melting points.

Boron Group Group 3A Contain one p orbital electron. Semimetals and metals All form cations 3+ charge

Carbon Group Group 4A Contain 2 p orbital electrons. C and Si generally form covalent bonds. Ge forms a cation with 2+ charge. Sn and Pb form 2+ and 4+ cations. Nonmetals, metalloids, and metals.

Nitrogen Group Group 5A Contain one electron in each p orbital. All except for Bi can be a 3- anion or a 3+ cation. Many can form 5+ cations. Nonmetals, metalloids, and metals.

Oxygen Group Group 6A Contain one set of paired and two sets of unpaired p orbital electrons. Many form 2- anions, all but O can form 2+ and 4+ cations. Nonmetals, metalloids, and metals.

Halogens Group 7A Halogen means “salt former.” Contain two sets of paired p orbital electrons and one unpaired electron. Commonly form diatomic molecules. F 2 Cl 2 Br 2 I 2 Nonmetals that form anions with 1- charge.

Noble Gases Group 8A All s and p orbitals contain paired electrons. All are generally unreactive gases. Noble gases do not commonly form ions.

Ionization Energy Ionization energy is the energy required to remove a negative electron and leave an atom with a positive charge – as an ion.

Valence Electron  An electron in the outermost energy level for an atom.  The electrons that interact when atoms form bonds. HLiNa CNFCNF

The Octet Rule  Atoms tend to gain, lose or share electrons in order to have a full set of valence electrons (usually 8)

Valence Electrons  So, every atom will either gain or lose electrons to get an electron configuration like the closest noble gas  What does Fluorine want to do?  (gain 1 electron)…  What do you suppose would happen if you brought a sodium and fluorine atom together in a bond?  What would the electrons do to make each atom “happy” like a noble gas configuration?  (sodium would lose one, fluorine would gain it)

Electron Orbital Shapes  s-spherical  p-dumbbell  d-clover leaf  f-multi-lobed

Filling Orbitals

Atomic Model of Matter Bohr’s Model  Proposed that electrons move in definite orbits around the nucleus

Atomic Model of Matter Quantum Mechanical Model  An atom has a small positively charged nucleus surrounded by a large region (electron cloud) containing enough electrons to make the atom neutral

Periodic Trends  Trends in atomic radius, ionization energy, & electronegativity are determined by:  The number of energy levels present.  The attraction between the positive nucleus and the outer shell electrons.  Interfering “shielding” by electrons on inner shells.  How close an atom is to completing the stable octet of outer “valence” electrons.

Electronegativity (1 of 3)  Electronegativity is the ability of an atom to attract electrons that are shared in a covalent bond.

Electronegativity (2 of 3)  What are the trends in electronegativity?

Electronegativity (3 of 3)  Electronegativity increases up & to the right.  This trend corresponds to stronger attractions to the nucleus.  Less shielding effect strengthens attractions to the nucleus in upper rows.

Atomic Radius (1 of 3)  Alkali metals are the largest atoms.  Noble gases are the smallest atoms.

Atomic Radius (2 of 3) Atomic radius trends: 1) Atomic radius increases down a group or column. 2) Atomic radius decreases across a period or row.

Atomic Radius (3 of 3) How do we explain the trends? 1. Atomic radius increases down a group: Each row adds an energy level. Interior electrons interfere with attraction of valence electrons toward the nucleus “shielding effect” 2. Atomic radius decreases across a row even while the atomic number increases: While in the same energy level, the nucleus becomes more positive & attractive.

Ionization Energy (1 of 4) Ionization energy is the energy required to remove a negative electron and leave an atom with a positive charge – as an ion. Occurs in solar cells, geiger counters & smoke detectors with Amerecium 241

Ionization Energy (2 of 4)  Alkali metals lose their electrons most easily.  Noble gases hold their electrons most tightly.

Ionization Energy (3 of 4) Removing an electron becomes more difficult across a row. Removing electrons becomes easier down a column.

Ionization Energy (4 of 4)  Removing electrons is more difficult across a row as the nuclear attractions become stronger.  Removing electrons is easier down a column as each additional energy level increases the distance from the nucleus and weakens the nuclear attraction.  Repulsive shielding by interior electrons also decreases the attraction for each added level.