PERIODICITY

SC4b. Compare and contrast trends in the chemical and physical properties of elements

and their placement on the Periodic Table

The Periodic Table

• Started by Dmitri Mendeleev – arranged by atomic mass – Wrote out info known about 63 elements on cards– Arranged by similar properties– Predicted existence of 3 unknown elements– w/I 4 years, 2 had been discovered

• Since Mendeleev, evidence has supported that the arrangement of elements should be based on the elements atomic number (# of protons)

Periodic Law

• Periodic Law – the physical and chemical properties of the elements are periodic functions of their atomic numbers.

• Periodic Law allows properties of elements to be predicted based on their position in the periodic table.

The Periodic Table

• Period – horizontal row across the PT– Elements in a row have the same number of major

energy levels– i.e. A Row 4 element has 4 major energy levels

• Group – vertical column in the PT– Elements in the same group have the same # of

valence electrons– Valence e-’s are the outermost, bonding electrons

(always in the S and P orbitals)

By the NUMBERS

Based on the Noble Gases He, Ne, Ar, Kr, Xe, and Rn, is there a repeatable

pattern of atomic numbers?

A: 2, 8, 8, 18, 18, 32

The Periodic Table

• Chemical and Physical properties– Vary across a period– Similar down a group

• Elements are either solids or gases (only 2 liquid elements at room temperature, Hg & Br)

METALS

• Metals – majority of the PT (78 %)– Properties include high conductivity, high density,

solid at room temp., malleable, ductile, and lusterous.

• Primarily belong in sublevels S, D, and F.

METALLOIDS

• Metalloids – have properties b/w metals and non-metals; small portion (7 %)

NON-METALS

• Non-metals – make up right side of PT (15%)• Properties include being non-conductive, low

density, mostly gases at room temp., brittle, non-lusterous

• 7 non-metals form DIATOMIC (2 atoms) molecules when pure elements

Quantum Numbers

• The Principal Quantum Number (n)– Main energy level

• Sublevels (s, p, d, f)• The principal quantum number is followed by

the sublevel and the number of electrons within the sublevel

Sublevels

SUBLEVEL # ORBITALS # ELECTRONS

S 1 2P 3 6D 5 10F 7 14

Electron Configuration

What period are the follow elements in based on their electron configurations?

1) Nitrogen2) Aluminum3) Calcium

Valence Electrons

Valence Electrons • electrons in the highest

energy level of atoms of an element

• Each element in a group has same # of valence e-’s

• Electrons available to be lost, gained, or shared in bonding

Group # of Valence e-’s

1 1

2 2

3-12 Varies b/w 1,2,3

13 3

14 4

15 5

16 6

17 7

18 8

Group Names

• Alkali metals – group 1 of the PT• Alkaline Earth metals – group 2 of the PT• Transition Metals – groups 3-12 of the PT• Halogens – group 17 of the PT• Noble gases – group 18 of the PT• Lanthanides – elements 58 – 71• Actinides – elements 90 - 103

Groups 1 & 2

Alkali Metals

• Group 1• Very reactive• 1 valence electron• ns1

• Soft• Silvery

Alkali Earth Metals

• Group 2• Reactive• 2 valence electrons• ns2

• Soft • Silvery

Halogens

• Group 17• Most reactive group of elements• Why? They have 7 valence electrons and want

1 to have a full octet• Based on electronegativity values• Have an electron configuration of ns2 np5

• Fluorine has the greatest Electronegativity

Noble Gases

• Group 18• 8 valence electrons• ns2 np6

• Not reactive• All gases

PERIODICITY

SC4a. Use the Periodic Table to predict periodic trends including atomic radii,

ionic radii, ionization energy, and electronegativity of various elements.

Periodic Trends

• Atomic Radius• Electron Affinity• Ionic Radius

• Ionization Energy• Electronegativity

Atomic Radius

• Atomic radius ½ the distance b/w the nuclei of identical atoms that are bonded together

• Radii decrease across a period• Radii increase down a group

Electron Affinity

• Electron affinity – the amount of energy absorbed when an electron is added to an isolated gaseous atom to form an ion with a 1- charge.

Cl(g) + e- Cl- (g) EA = -349 kJ/mol

Electron Affinity

Ionic Radius

• Cation – a positively charged ion (i.e. Na1+)

• Anion – a negatively charged ion (i.e. F1-)

• Ionic radius increases from right to left across a period and increases from top to bottom down a group.

Ionization Energy

• Ionization Energy - energy needed to remove 1 e- from a neutral atom

• Increases from bottom to top in a group.

• Increases from left to right across a period.

Electronegativity

• Electronegativity is a measure of the relative tendency of an atom to attract electrons to itself when it is chemically combined with another atom.

• Elements w/ HI Electronegativity’s (nonmetals) gain electrons to form anions.

• Elements w/ LOW Electronegativity’s (metals) often lose electrons to form cations.