Periodic Trends

  • Chemical Reactions 

    • either electron removal, electron addition, or electron sharing 


  • Valence Electrons 

    • are the electrons in the highest occupied principal energy level of an atom

    • are directly involved in the element’s reactivity or in formation of compounds


  • Inner-shell (kernel) electrons

    • are not involved directly in element’s reactivity or in formation of compounds

Electron Dot Diagrams

  • aka Lewis structures, Lewis dot formulas, etc. 

  • group 1: 1 valence electron

  • group 2: 2 valence electrons

  • p - block: group number minus 10

  • group 13: 3 valence electrons

  • group 14: 4 valence electrons

  • …

  • group 18: 8 valence electrons

    • noble gases

  • must keep track of electrons during chemical bonding

  • electron dot diagrams show valence e- only around element symbol

Octet Rule

  • noble gases are unreactive

  • has full outer shell and cannot incorporate any more electrons into the valence shell

  • Gilbert Lewis (1875-1946): observed types of ions and molecules that are formed by other elements

  • states that atoms tend to form compounds in ways that give them eight valence electrons — same elec. config as noble gases

  • exceptions: He, H, Li

2 Methods to Satisfy Octet Rule

  1. By sharing valence electrons with other atoms


  1. By transferring valence electrons from one atom to another

    • metals tend to lose all valence electrons

    • nonmetals tend to gain electrons to fill octet

    • cation and anion predicted by octet rule

Metals

  • an element that is a good conductor of heat and electricity 

  • malleable, ductile, shiny/high luster

  • solid at room temp. (excluding Hg)

  • wide variance in melting point

Nonmetals

  • elements that don’t have metal properties

  • generally poor heat and electricity conductor

  • wider variation in properties than among the metals

  • exist in all three states of matter

    • majority gases (N, O)

    • liquid (Br)

    • solid (C, S)

      • brittle, not lustrous

  • lower melting points than metals

Metalloids

  • has properties that are intermediate between metals and nonmetals 

  • aka semimetals

  • stair-step line

  • excludes aluminum

  • silicon, antimony, arsenic

Periods and Blocks

  • horizontal rows

  • length determined by the number of electrons that are capable of occupying the sublevels that are filled within that period

  • periodic table can be divided into blocks denoting sublevel that is being filled

Atomic Radius

  • size of atom is important to explaining behavior of atoms or compounds 

  • defined by edge of its orbital from the nucleus 

    • but this varies under different conditions

  • units: picometers

  • generally decreases from left to right across a period

    • exceptions: i.e. Bi and Po

  • within period, protons added to nucleus as electrons are being added to the same principal energy level

  • electrons pulled closer to the nucleus because of increases positive charge of the nucleus

Group Trend

  • generally increases from top to bottom within a group

  • as atomic number increases down a group, increase in positive nuclear charge and in number of occupied principal energy levels

  • higher principal energy levels have orbitals that are larger in size (i.e. 2s bigger than 1s)

  • effect of increase in nuclear charge outweighed by effect of greater amount of principal energy levels


Ionic Radius

  • is the effective distance from the nucleus of an ion to this outer level of electrons 

  • when electrons are lost during ion formation the resulting ion is positive (cation) and smaller than its atom

  • why a cation is smaller than atom when valence electron(s) removed:

    • one fewer occupied principal energy level 

    • protons outnumber electrons

  • metal atom radius > metal ion radius

    • Li0 > Li+1 

  • when electrons are gained during ion formation the resulting ion is negative (anion) and larger than its atom

  • why a anion is larger than atom when valence electron(s) gained:

    • electrons outnumber protons — attractive force of protons decreased 

    • electron cloud spreads — more electrons means more repulsion between electrons

  • nonmetal atom radius < nonmetal ion radius

    • S0 < S-2 

Ionization Energy

  • is the amount of energy required to remove the most loosely held electron from the valence level of an atom in the gas phase to form a positive ion

  • unit: kJ/mol

  • valence electrons have lower ionization energy

  • Why?

    • electron shielding: as more electrons are added, the outer electrons are shielded from the pull of the nucleus by the inner shell electrons

  • increases left to right across period

  • consider nuclear charge of the atom

  • more protons in the nucleus, stronger attraction of the nucleus to electrons

  • stronger attraction makes it more difficult to remove electrons   

  • exceptions: Group 13 lower IE than Group 2 — electron shielding; Group 16 lower IE than Group 15— Hund’s rule

  • down a group, IE decreases as the size of the atom increases

  • first electron removed is farther from the nucleus as number of protons (at. number) increases

  • being farther from the positive attraction makes it easier for that electron to be pulled off


Electron Sheilding

  • increase in size, number of protons and electrons — these changes influence how nucleus attracts electrons

  • outer electrons are partially shielded from the attractive force of the protons in the nucleus by inner electrons

  • between sublevels within same principal energy level

    • s sublevel capable of shielding electrons in p sublevel because s orbital is spherical

Electronegativity

  • is the measure of the electron attracting ability of an element

  • the greater the number of valence electrons the greater the ability to attract electrons — the closer to the nucleus the stronger the attracting power

  • important in chemical bonds when forming compounds

  • measured on a relative scale not in energy units

    • elements are compared to one another

  • Fluorine attracts electrons better than any other ion — EN = 3.98 (~4.0)

  • Cesium the lowest — EN = 0.79

  • generally increase across a period due to increase in nuclear charge

    • Alkali metals (Group 1) have lowest and halogens (Group 17) highest

  • Noble gases left out because most do not form compounds — do not have electronegativities  

  • little variation among transition metals

  • generally decrease down a group due to larger atomic size

  • can use these values to predict what happens when certain elements combine

  • if difference between EN is greater than ~1.7

    • complete exchange of electrons occur

    • typically between metal and nonmetal

  • sodium and chlorine

    • form compound and each ion becomes isoelectronic to nearest noble gas

    • |0.93 - 3.2| = 2.27 — complete exchange