Presentation on theme: "PERIODIC TRENDS Elemental Properties and Patterns."— Presentation transcript:
PERIODIC TRENDS Elemental Properties and Patterns
The Periodic Law Dimitri Mendeleev was the first scientist to publish an organized periodic table of the known elements. Mendeleev understood the ‘Periodic Law’ which states: When arranged by increasing atomic number, the chemical elements display a regular and repeating pattern of chemical and physical properties.
The Periodic Law Atoms with similar properties appear in groups or families (vertical columns) on the periodic table. They are similar because they all have the same number of valence (outer shell) electrons, which governs their chemical behavior. Do you remember how to tell the number of valence electrons for elements in the s- and p- blocks?
Ionization Energy (first) Definition: energy needed to remove the most loosely held e- Valence electron high E electron Outermost electron
What’s holding the e - in place? Nuclear charge There is a force of attraction between protons (+) in the nucleus and electrons (-) in the orbitals “Opposites attract”
Ion ization energy Hint: energy needed to make an ion by losing electrons
MOST elements want 8 e- (octet rule) Elements with only a few valence electrons will tend to have lower ionization E In other words… It doesn’t take a lot of E to remove their e- Metals
MOST elements want 8 e- (octet rule) Elements with almost 8 valence electrons won’t give them up so easy It takes a lot of E to remove their e- Non-metals have high Ionization E Who has the highest IEs?
Ionization E trend? As move across a period…IE increases. WHY? 3 p+ 9 p+
As you go across a period a large proton and a tiny electron are being added. more p+ hold the e- tighter
As we move DOWN a group … IE decreases. WHY? Period 1 Period 2 Period 3 Period 4 Let’s figure it out...
Lithium 3 p+ +
Sodium 11 p+ +
Potassium 19 p+ +
SHIELDING Reason why IE decreases as you go down a group the distance increases between the p+ and valence e- AND e- in outer shells repel each other WHICH IS WHY low IE are so reactive!
IE - Summary Going across a period left to right, IE increases. Going down a group, IE decreases. Metals have lower IE than nonmetals. Metals are more likely to form positive ions. Fr (francium) has the lowest ionization energy on the periodic table. It is the most reactive metal.
Electronegativity Ability of an atom to attract electrons of other atoms. In other words... Atoms with high electronegativity are bullies that steal electrons!
Electronegativity Determine the periodic trend INCREASES EXCEPT for Noble gases. WHY? F
Atoms with high electronegativities also have high ionization E
Electronegativity - Summary Going across a period left to right, e-neg increases. Going down a group, e-neg decreases. Metals have lower electronegativity than nonmetals. The noble gases do not have values for e-neg. Nonmetals tend to form negative ions due to their high e- neg. F (fluorine) has the highest electronegativity on the periodic table. It is the most reactive nonmetal.
Atomic Radius Radius distance from the center of a circle (Nucleus) to the outermost edge (valence shell) R
Atomic Radius Since a cloud’s edge is difficult to define, scientists use define covalent radius, or half the distance between the nuclei of 2 bonded atoms. Atomic radii are usually measured in picometers (pm) or angstroms (Å). An angstrom is 1 x m.
Covalent Radius Two Br atoms bonded together are 2.86 angstroms apart. So, the radius of each atom is 1.43 Å Å 1.43 Å
Why????? Why does atomic radius INCREASE as you move down a group? Why does atomic radius DECREASE as you move across a period? Gaining layers of e- Increasing the # of p+ holds the e- in tighter Increasing NUCLEAR CHARGE
Summary – Atomic Radius Going across a period left to right, atomic radius decreases. Going down a group, atomic radius increases. The element helium (He) has the smallest atomic radius. The element francium (Fr) has the largest atomic radius.
Ions When an atom gains an electron, it becomes negatively charged (more electrons than protons ) and is called an anion. Metal atoms can lose electrons. They become positively charged cations.
Ionic Radius Cations are always smaller than the original atom because the entire outer principal energy level is removed during ionization. Conversely, anions are always larger than the original atom because electrons are added to the outer principal energy level.
What happens to atomic radius as you create a + ion? radius decreases What happens to atomic radius as you create a - ion? radius increases
Cation Formation 11p+ Na atom 1 valence electron Valence e- lost in ion formation Effective nuclear charge on remaining electrons increases. Remaining e- are pulled in closer to the nucleus. Ionic size decreases. Result: a smaller sodium cation, Na +
Anion Formation 17p+ Chlorine atom with 7 valence e- One e- is added to the outer shell. Effective nuclear charge is reduced and the e- cloud expands. A chloride ion is produced. It is larger than the original atom.