Periodic Table Development



Periodic Table Development

Late 1700’s – 30 elements discovered

1820 – Dobereiner’s Triads

• Grouped elements into sets of 3 (triads)

(Li, Na, K) ( Ca, Sr, Ba ) ( Cl, Br, I )

Properties of the middle element are averages of the 1st and 3rd

1865 – Newland’s Octaves

• Organized elements into repeating groups of 8

Law of Octaves – elements arranged by increasing atomic mass;

properties of the 8th element are similar to the 1st

1869 – Mendeleev

• Arranged elements by increasing atomic mass

• Observed periodic (repeating) element properties

• Produced 1st Periodic Table

• Elements in the same column had similar properties

• Predicted properties of undiscovered elements

1913 – Moseley

• Arranged elements by atomic number

Periodic Law –

Elements arranged in order of increasing atomic number show a periodic pattern in their physical and chemical properties

Reading the Periodic Table

Information in each square:

• Element name

• Element symbol

• Atomic # = ( # protons or # electrons)

• Atomic Weight = (weighted average of isotope masses)

• Electron configuration

Group = elements in vertical column with similar properties. (Families)

Period = Horizontal row of elements

18 labeled groups and 7 periods

Labeling and Naming Groups

Group number Family Name

Group 1A = Alkali metals

Group 2A = Alkaline Earth metals

Group 7A = Halogens

Group 8A = Noble gases

Metals, Nonmetals, Semimetals

Metal Properties

• Metallic Luster (shine)

• Good conductor of heat and electricity

• Solids at Room temperature ( except Mercury)

• Malleable ( hammered into thin sheets without shattering)

• Ductile (drawn into fine wires)

Examples: Cu, Ag, Al, Au, Zn

Most elements are metals

Nonmetals

• No metallic luster (dull)

• Poor conductors of electricity and heat

• Neither malleable nor ductile (brittle)

• Gases and solids at room temperature

Bromine is a liquid at room temperature

Located on far right side of the Periodic Table

C, N, O ,P, S, F, Cl, Br, I. He, Ne, Ar, Kr, Xe

Semimetals or Metalloids

Have properties of both metals and nonmetals

Boron, Silicon, Germanium, Arsenic, Antimony

Why do elements in a group have similar properties?

Similar electron arrangement

Valence electrons

Electrons in the outermost energy level of an atom are

responsible for an atom’s chemical behavior

Elements in the same group have valence electrons in similar electron configurations

Group 1A (alkali metals)

all have one valence electron in the S orbital

Abbreviated electron configurations

• Focus on valence electrons

Inner core electrons represented by the symbol of the nearest noble gas with a lower atomic #

Group 1A elements:

H = 1S1

Li = [He] 2S1 [He] = 1S2

Na = [Ne] 3S1

K = [Ar] 4S1

Rb = [Kr] 5S1

Cs = [Xe] 6S1

• Each Group 1A element has a single valence electron in the s orbital

• Principal quantum # of the s orbital = Element’s row or period

S, P, d, f – block elements

Periodic Table is divided into 4 blocks

s – block

• Group 1A and 2 A

• Alkali metals & Alkaline Earth metals

• Valence electrons in S orbitals only

• Group 1A, each ecn ends in S1

• Group 2A each ecn ends in S2

p – block

• Group 3A to Group 8A of any period

• Valence electrons in P1 to P6 orbitals

• P sublevel can hold 6 electrons

d- block

• Middle of Periodic Table

• d sublevel can hold 10 electrons

• 3d orbitals start with Sc (atomic # 21)

• d-block elements are called transition metals

f – block

• electrons start to be located in f orbitals

• f orbitals can hold 14 electrons

• start filling 4f orbitals on 6 period with La (atomic # 57)

f – block elements called inner transition metals

Periodic Table shape is due to the way electrons fill s,p,d, f orbitals of different energy levels

Periodic Trends

Systematic changes of element’s properties throughout the periodic table

Properties determined by an atom’s electron configuration

Periodic Trends include:

• Atomic radius

Ion Size

• Ionization Energy

• Electron Affinity

• Electronegativity

Atomic radius

Distance from center of an atom’s nucleus to its valance electrons

Atomic radii increase moving down a group

Why?

There are more electrons down a group,

Energy levels holding those electrons are farther away from the nucleus

With increasing distance, there is less attractive force exerted by the nucleus on the electrons.

Therefore, atomic radius increases

Atomic radii decrease moving across a row

from left to right

Why?

Across a period, the increasing numbers of protons exert a stronger pull of the electrons.

Valence electrons are attracted to the nucleus

This attraction shrinks the electron orbital to reduce the atomic radius

Ion Size

As an atom loses electron(s) to form a positive ion, it becomes smaller

Li atom’s radius is 0.152 nm

Li +1 ion’s radius is 0.060 nm

Loss of electrons vacates the largest orbital

Atom gains an electron to form a negative ion, it becomes larger

Fluorine atom’s radius = 0.064 nm

Fluorine ion’s radius = 0.136 nm

Increasing the # electrons increases repulsive force to spread electron cloud

Periodic trend, elements in the same group form ions of the same charge

Ionization Energy ( I.E.)

Energy needed to remove an electron from an atom

Li (vapor) + Energy ---------- Li +1 (vapor) + electron

I.E. represents how strongly an atom holds onto its valence electrons

Periodic Trend:

I.E. decreases down a group

I.E. Increases across a period ( Left to Right)

Both atomic radii and I.E. depends on how strongly an atom’s electrons are attracted to the nucleus

Electron Affinity ( E.A.)

Energy released when an atom gains an extra electron

Ne (gas) + electron --------- Ne –1 (gas) + Energy

E.A. represents the atom’s attraction for an extra electron

An atom’s E.A. is related to the # of electrons it needs to fill its outer energy level

Nonmetals E.A. > Metal E.A.

Electronegativity

Ability of an atom to attract an electron in a chemical bond

Periodic Trend:

Electronegativity decreases down a group

Electronegativity increases across a row (left to right)

Fluorine is the most electronegative atom

Cesium atom has a low electronegativity

................
................

In order to avoid copyright disputes, this page is only a partial summary.

Google Online Preview   Download