Chemical Periodicity

Chemical Periodicity

Discover it!

Elements / Atomic
Number / Melting
Point / Boiling
Point / Density
g/cm3
Li
Na /
K /
Rb
Cs
Elements / Atomic
Number / Melting
Point / Boiling
Point / Density
g/cm3
F
Cl
Br
I

Chemical Periodicity

Discover it!

Elements / Atomic
Number / Melting
Point / Boiling
Point / Density
g/cm3
Li / 3 / 180.54 / 1342 / 0.534
Na / 11 / 97.81 / 882.9 / 0.971
K / 19 / 63.25 / 760 / 0.862
Rb / 37 / 38.89 / 686 / 1.532
Cs / 55 / 28.40 / 669.3 / 1.873
Elements / Atomic
Number / Melting
Point / Boiling
Point / Density
g/cm3
F / 9 / -219.62 / -188.54 / 0.001696
Cl / 17 / -100.98 / -34.6 / 0.003214
Br / 35 / -7.2 / 58.78 / 3.12
I / 53 / 113.5 / 184.35 / 4.93

TP: Ch. 4 and 9CP: Ch. 6

Periodicity- the tendency to recur at regular intervals

Periodic Law – The statement that the physical and chemical properties of the elements repeat in a regular pattern when they are arranged in order of increasing atomic number.

Dimitri Mendeleevwas the Russian chemist who proposed the basis for the first periodic table.

Group (or family) is found in the vertical columns.

List the elements are in the following groups:

Periods are the horizontal rows on the periodic table.

List the elements in the following periods:

Three broad classes of representative elements: Metals, non-metals, and metalloids:

1. Metals -solid at room temperature, high luster when clean, high heat and electrical conductivity, loosely held valence electrons, solids are easily deformed. Only one is a liquid metal – Hg.

2. Non-metals – generally non-lustrous, various colors, generally poor conductors of electricity and heat, tightly held valence electrons, solids may be hard or soft, usually brittle.

3. Metalloids–have physical and chemical properties of both metals and non-metals, lie along the dark, stair-step border, and valence electrons are tightly held. Some are semiconductors.

Semiconductors are elements that does not conduct electricity as well as metals, but does conduct better than non-metals. Examples are Si, Ge, As

Semiconductors:

n-typed (negatively charged, extra electrons)

p-typed (“act like positive”, shortage of electrons)

- used in computers, tvs, handheld games, “miniature electric circuits”

What are three ways the elements on the Periodic Table are categorized?

1.) electronic configuration

2.) physical state

3.) metals, non-metals, and metalloids

Four categories of elements according to their electronic configurations:

I. Noble gases

II. Representative elements

III. Transition metals

IV. Inner Transition metals

I. Noble Gases:

Outermost(valence) s and p sublevels are filled and are in Group 0, 18 or 8A.

s2p6Neon 1s22s22p6

inert gases/unreactive

II. Representative Elements:

outermost s and p sublevel is only partially filled are called Group A elements.
7 group A’s (not including 8A)

1. Group IA (1A):

1 electron in outermost energy level:s1

Alkali metals – Li, Na, K, Rb, Cs, Fr

Na – 1s22s22p63s1

Reacts vigorously with water, solid at room temperature, high luster when clean, high electrical conductivity.

2. Group IIA(2A):

2 electrons in outermost energy level:s2

Alkaline Earth metals – Be, Mg to Ra(radium)

Be - 1s22s2

- solid at room temperature, high luster when clean, high electrical conductivity.

3. Group IIIA(3A, 13):

3 electrons in outermost energy level: s2p1

B, Al, Ga, In, Tl

Al - 1s22s22p63s23p1

4. Group IVA (4A,14):

4 electrons in outermost energy level: s2p2

C, Si, to Pb

Si - 1s22s22p63s23p2

5. Group VA (5A, 15):

5 electrons in outermost energy level:s2p3

N, P, to Bi

P - 1s22s22p63s23p3

6. Group VIA (6A, 16):

6 electrons in outermost energy level:s2p4

O, S, Se, to Po

S - 1s22s22p63s23p4

7. Group VII (7A, 17):

7 electrons in outermost energy level:s2p5

Halogens – F, Cl, to At

Cl - 1s22s22p63s23p5

III. Transition Metals:

Outermost s sublevel and nearby d sublevelcontain electrons and are the Group B elements. These are characterized by adding electrons to d sublevel.

Mn (manganese) - 1s22s22p63s23p64s23d5

IV. Inner Transition Metals:

Outermost s sublevel and nearby f sublevel contain electrons. These are characterized by adding electrons to f sublevel.
Lanthanide series – Ce to Lu
Actinide series – Th to Lr

Review: Give examples of the following:

Representative Elements

Nonmetals MetalloidsMetals

Noble Gases Alkali Metals

Alkaline Earth

Transition Metals

Inner Transition Metals

SolidsLiquidsGases

Ions: An atom or group of combined atoms that has a charge because of the loss or gain of electrons.

Cation – positive ion (because it loses electrons)

Anion – negative ion (because it gains electrons)

ProtonElectronsTotal charge

5+5-0 (neutral)

5+4-1+ (cation)

5+6-1- (anion)

Looking at electrons:

Electron configuration – Total electron
Lewis Dot Diagram (EDD or LDD) – Valence electrons ( coordinates with Group # - A only)

Octet Rule – Most atoms would like to have 8 electrons in their outer shell.

Group 1A – 1 electron in outer shell: gives it away.

Trends in Elements: (CP:Ch 6-- TP: Ch4, 9)

Standard 2.2 and 2.3 Periodic Trends

I. Atomic radius: one-half the distance between the nuclei in a molecule consisting of identical atoms.

Atomic (radius) size:

Across the periods (left to right), it decreases because the principal energy level remains the same. As one more proton and electron are added, the attractions from the nucleus grow stronger and it pulls the electrons closer in.

It increases down a group, because electrons are added to successively higher energy levels. More electrons, more energy levels  larger atomic radius.

II. Ionization Energy: the energy required to overcome the attraction of the nuclear charge and remove an electron from a gaseous atom.

Ionization Energy:

Across the period, the IE generally increases because of the greater attraction of the nucleus for the electrons.

Down a group, the outermost electrons are further from the nucleus, the attraction between the nucleus and electrons are weaker, thus the IE generally decreases. This means the electrons are easier to remove.

III. Ionic Size: ions are atoms or a group of atoms that gain or lose electrons.

Cations – positive ions (atoms lose electrons)

Anions – negative ions (atoms gain electrons)

Cations are always smaller than the neutral atoms from which they form.

Example: Na+ion is 95 pm (lost 1 e-)

Na atom is 186 pm (no lost e-)

The nuclear attraction to the remaining e-, increases.

If you let the air out of a balloon, the balloon gets smaller.

Anions are always larger than neutral atom from which they form.

Example: Cl- ion is 181 pm (gain e-)

Cl atom is 99 pm

The effective nuclear attraction (charge) is less or decreases. If you add more air to the balloon, the balloon is going to get bigger.

IV. Electronegativity

The tendency for the atoms of an element to attract electrons when they are chemically combined with atoms of other elements.

Electronegativity increases across the periods. It decreases as it goes down a group or family.

Elements on the left side would rather give electrons away. Elements on the top right side want to take electrons.

Cs (cesium’s value is 0.7) loses a “tug-of-war” to F (fluorine’s value is 4.0)

Electron Affinity is the energy change associated with the addition of an electron to a gaseous atom.