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PERIODIC TRENDS Elemental Properties and Patterns.

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Presentation on theme: "PERIODIC TRENDS Elemental Properties and Patterns."— Presentation transcript:


2 PERIODIC TRENDS Elemental Properties and Patterns

3 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.

4 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?

5 Ionization Energy (first)  Definition:  energy needed to remove the most loosely held e- Valence electron high E electron Outermost electron

6 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”

7 Ion ization energy  Hint: energy needed to make an ion by losing electrons

8 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

9 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?

10 Ionization E trend?  As move across a period…IE increases. WHY? 3 p+ 9 p+

11  As you go across a period  a large proton and a tiny electron are being added.  more p+ hold the e- tighter

12  As we move DOWN a group … IE decreases. WHY? Period 1 Period 2 Period 3 Period 4 Let’s figure it out...

13 Lithium 3 p+ +

14 Sodium 11 p+ +

15 Potassium 19 p+ +

16 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!

17 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.

18 Electronegativity  Ability of an atom to attract electrons of other atoms.  In other words...  Atoms with high electronegativity are bullies that steal electrons!

19 Electronegativity  Determine the periodic trend INCREASES EXCEPT for Noble gases. WHY? F

20 Atoms with high electronegativities also have high ionization E

21 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.


23 Atomic Radius  Radius  distance from the center of a circle (Nucleus) to the outermost edge (valence shell) R

24 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.

25 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


27 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

28 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.

29 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.

30 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.

31  What happens to atomic radius as you create a + ion? radius decreases What happens to atomic radius as you create a - ion? radius increases

32 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 +

33 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.

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