PERIODIC PATTERNS

Unit 3 – Periodic Table

What patterns exist on the periodic table?

Lesson Essential Question:

METALLIC TREND

INCREASE

S

INC

REA

SE

S

ATOMIC RADIUS

Radius is the distance from the center of the nucleus to the “edge” of the electron cloud.

Atomic radii are usually measured in picometers (pm) or angstroms (Å). An angstrom is

1 x 10-10 m.

ATOMIC RADIUS

BROMINE = Br2

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.

2.86 Å1.43 Å 1.43 Å

ATOMIC RADII TRENDS

As you go down a family the number of energy levels increases making the radius larger.

DOWN A FAMILY OR GROUP WHY?

INC

REA

SES

ATOMIC RADII TRENDS

As you go across a period the number of protons increases, (nuclear charge) pulling the electrons in tighter making the radius smaller.

ACROSS A PERIOD WHY?DECREASES

IONS - remember

Metals Lose electrons

becoming positive.

Calcium (Ca)Loses 2 electrons

becomingCa+2 and [Ar] Noble

gas Configuration. (Octet Rule)

Nonmetals Gain electrons

becoming negative.

Chlorine (Cl)Gains one e-

becoming Cl-1 and [Ar] Noble gas configuration.

(Octet Rule)

IONS – How can I remember?

Metals Nonmetals

This is Ann ion - ANIONShe is unhappy and

negative.

This is Cat-ion - CATIONHe is a “plussy”

cat!

IONIC RADII TRENDS

As you go down a family the number of electron shells increases making the radius larger.

DOWN A FAMILY OR GROUP WHY?

INC

REA

SES

IONIC RADII TRENDS

For the metals the nuclear charge is greater than then number of electrons pulling them in tighter making the radius smaller.

At the nonmetals the radius gets larger because the ion has gained electrons.

ACROSS A PERIOD WHY?DECREASES then INCREASE

METALLIC ATOM AND ION COMPARISON

NONMETALLIC ATOM AND ION COMPARISON

Why do the Noble Gases not have an ionic Radius?

ATOM AND ION COMPARISON

Why does Hydrogen not have an ionic Radius?

As more electrons are added to atoms, the inner layers of electrons shield the outer electrons from the nucleus.The effective nuclear charge on those outer electrons is less, and so the outer electrons are less tightly held

Shielding Effect

Example of Shielding Effect

The energy required to remove an electron from an atom. (measured in kilojoules, kJ)

Ionization Energy

IONIZATION TREND

INC

REA

SE

S

Why?• Closer to

nucleus (more +) • Electrons

less likely to be removed• Requires

more energy to form ion• Less

shielding

IONIZATION ENERGY

The larger the atom is, the easier its electrons are to remove. (Why?)

Ionization energy and atomic radius are inversely proportional.

Ionization energy is always endothermic, that is energy is added to the atom to remove the electron.

IONIZATION TREND

INCREASES

NC

REA

SE

SWhy?• Elements in

alkali metals have 1 valence electron so what to remove that electron, they therefore take the least amount of energy to remove an electron

is a measure of the tendency of an atom to attract a bonding pair of electrons.

Electronegativity

Electronegativity

Why?•Closer to nucleus (more +) so electrons are more attracted

INC

REA

SE

S

http://www.thecatalyst.org/electabl.html

Electronegativity

Why?• Elements in

halogens only need 1 more electron to have a full valence shell so are MOST likely to attract electrons. As you move to left elements are more likely to LOSE electrons.

INCREASES

http://www.thecatalyst.org/electabl.html

In Summary….Electronegativity

Ele

ctro

negati

vit

y