Chapter 6

The Periodic Table and

Periodic Trends

The Periodic Table

1. Mendeleev - arranged elements in order ofincreasing atomic mass.

2. Moseley - used x-rays to determine the atomic number of the elements.

- arranged elements with most similar properties in columns

-elements arranged by atomic number show a clear “periodicity” of properties.

Periodic Law

• There is a periodic repetition of the element’s chemical and physical properties when they are arranged in order of increasing atomic number. This repetition is known as the Periodic Law.

Arrangement of the Periodic Table

1. Period - horizontal row of elements - period # = # of energy levels an

element has2. Group (or family) - vertical column of

elements- Number (AorB) above a column is its

family designation.- elements in a group have much more similar

properties than elements in a period.

Group A Elements (Representative Elements)

contain metals, non-metals and metalloids

• s and p blocks.• Group # = # of electrons in the outermost

energy level.• All shells beneath the outermost have stable octet.

– Except: Helium (He), Lithium (Li) & Beryllium (Be). First energy level is

full with 2 electrons.

Group B Elements (Transition Metals)

• d block.

• May have more than 1 ionic charge because electrons in the shell beneath the outermost get involved in bonding.– Shells beneath the outermost may have more

than 8 electrons.

Classification of the Elements

• There are 3 classes of elements; metals, nonmetals and metalloids.

• Metals - make up about 80% of the table– Good conductors of heat and electricity– Have a high luster or sheen. (shiny)– Solids at room temp. except for Mercury (Hg)– Ductile; drawn into wires– Malleable; hammered into thin sheets.

Nonmetals • Found in the upper right corner of the table• Properties of nonmetals are not as similar as

they are for metals.• Most are gases at room temperature.• Some are solids like sulfur and phosphorus• Bromine is a dark red liquid.• Generally have properties opposite of the

metals.• Not good conductors (carbon exception),

dull and brittle if they are solids.

Metalloids

• Along the staircase that separates the metals and nonmetals.

• Have properties that are similar to metals and nonmetals under certain conditions.

• Aluminum (Al) is a metal.

Classification of Elements Continued

1. Metals – lose electrons to form positive ions (cations) when they bond.

S block:a) Group 1A (ns1) - Alkali or active metals

- form +1 ions

- very reactive, stored in kerosene

S block:a) Group 2A (ns2) - Alkaline earth metals

- form +2 ions

Note – Elements with the same outer shell electronsconfiguration have the same physical and chemical properties. (same number of valence electrons.)

Metals (continued)

• D block – these metals are all similar because they all have the same s2 electron configuration in their outermost shell.

– Transition Metals – Group B metals

– May have more than 1 ionic charge (designated by Roman #)

Metals (continued)

• Hund’s Rule – 1/2 filled and completely filled sublevels have the lowest energy are the most stable, and therefore the preferred electron arrangement.

Write the electron configuration for Fe

• # 26 1s2 2s2 2p6 3s2 3p6 4s2 3d6

Forms Fe+2 when it loses 2 electrons

Forms Fe+3 when this additional electron is lost the ½ filled is d is more stable.

Write the electron configuration for Cu

• # 29 1s2 2s2 2p6 3s2 3p6 4s2 3d9

Forms Cu+2 when it loses 2 electronsForms Cu+2 when it loses 2 electrons

Hund’s Rule Shift: 1 e- shifts from 4s to 3d. The 3d is now completely filled and Cu will only lose 1 electron to form Cu+1.

Metals (continued)• F block – Inner Transition Metals

– 2 rows:

• 1st row is Lanthanide series – “rare earth” elements (natural).

• 2nd row is Actinide series – mostly synthetic (manmade).

–Parentheses around mass # means radioactive.

same

Write configurations for #58 and #95. Then write the condensed

configurations.• Ce # 58

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 5d1 4f 1

• Am # 95

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f 14

5d10 6p6 7s2 6d1 5f6

2

8

18

19

9

2

2

8

18

32

24

9

2

Inner transition metals are most similar b/c their last 2 energy level config. are identical.

2. Semi-metals

• Located on either side of the “stairs” in the p block separating the metals from the non-metals.

• Have properties of both metals and nonmetals.

– Aluminum (Al) is a metal

3. Non-metals• in upper right corner of p block

(except last column).

• gain electrons to achieve stable octet

• form negative ions - anions

• group 5, 6, and 7 gain 3, 2, or 1 electrons

Halogens - Group 7A (ns2p5)Gains 1 e- and has a –1 charge

Question: Which are more similar?

F, Cl, and Br or N, O, and F

Answer: F, Cl, and Br

Because in the same group or column,they have the same outer shell electronconfiguration, therefore react similarly.

4. Noble Gases (ns2p6)

• Last column on the right in p block.

• All but He already have their stable 8 in the outer shell. Therefore they are called:

– Inert – means non-reactive

• Isoelectric- same electron configuration as a noble gas.– Atoms react to become isoelectric with a

noble gas.

i.e. Sulfide ion (S-2) ion

is isoelectric to Argon (Ar) atom.

Periodic Trends

• Trends or patterns on the Periodic Table are all basically the result of atomic size and whether an atom of an element is a metal or a non-metal.

1. Atomic size (atomic radii) – size of the atom

• Nuclear charge – pull of protons in nucleus on electrons

– increase number of protons = increase the nuclear charge.

Atomic Size

•Atomic Radius = half the distance between two nuclei of a diatomic molecule.

}Radius

1. Atomic size (atomic radii) – size of the atom

• Across a period – decreases due to increasing nuclear charge which pulls the electrons on the same energy level in closer to the nucleus.

• Down a group – increases due to the addition of energy levels which are farther from the nucleus.

1. Atomic size

Down a group:More energy levels

Across a period:Greater nuclear charge – same energy level

2. Metallic vs. Non-metallic Character

Across a period –decreases because the smaller the atom, the stronger the pull from the nucleus so the harder it is to lose electrons.– Also the # of e- to lose is increasing as you

approach the non-metals.

a) Metallic Character – tendency to lose electrons- larger atom – easier to lose electrons.

2. Metallic vs. Non-metallic Character

• Down a group – increases because larger atoms have less pull from nucleus on outer electrons.– More energy levels easier to lose electrons.

a) Metallic Character – tendency to lose electrons- larger atom – easier to lose electrons.

2. Metallic vs. Non-metallic Character

Across a period –increases because as atomic size decreases its easier to gain electrons due to increased pull from nucleus.

- also needs to gain fewer electrons.

b) Non-metallic Character – tendency to gain e-- smaller atom – easier to gain electrons

2. Metallic vs. Non-metallic Character

Down a group – decreases because atoms are larger (more energy levels).

- less pull from nucleus- harder to gain electrons.

b) Non-metallic Character – tendency to gain e-- smaller atom – easier to gain electrons

3. Shielding Effect

Each additional energy level acts as another layer of electrons that shields the outermost electrons from the pull of the nucleus.

3. Shielding Effect

• Across the period – unchanged; on the same energy level

– With each addition electron, there is an additional proton.

• Down a group – increases due to additional energy levels between the nucleus and outer electrons.

4. Ionization Energy

• Amount of energy required to remove an electron from an atom.

– Metals have low ionization energies because they lose electrons easily.

– Non-metals have high ionization energies because they do not lose electrons easily.

4. Ionization Energy

• For group A metals: as electrons are removed 1 by 1 from an atom, there is a steady rise in the ionization energy until the structure of a noble gas is attained.

• Then there is a dramatic rise when a stable octet is disturbed.

4. Ionization Energy• 1st ionization energy – amount of energy

needed to remove the 1st electron from an atom.

• 2nd ionization energy – amount of energy needed to remove the 2nd electron from an atom.

• 3rd ionization energy – amount of energy needed to remove the 3rd electron from an atom.

4. Ionization Energy

1st ion. 2nd ion. 3rd ion. 4th ion.

Na 119 1090 1652 2281

Mg 176 347 1848 2519

Al 138 434 656 2767

283

1s2 2s2 2p6 3s1

1s2 2s2 2p6 3s2

1s2 2s2 2p6 3s2 3p1

281

282

4. Ionization Energy• Across a period – 1st ionization energy

increases (across group A) because nuclear charge is increasing, holding electrons in tighter and closer to nucleus.– For transition metals (pds 4, 5, & 6) in D

block there is a sharp drop in ionization energy between the last transition element and where the next energy level’s p sublevel begins because the p sublevel is farther from the nucleus – becomes easier to lose electrons.

4. Ionization Energy

• Down a group – 1st ionization energy decreases because larger atoms have less pull on outer electrons due to increased shielding effect (see # 3).

5. Electronegativity or Electron Affinity (attraction)

• Energy change due to the addition of an electron to an atom.

• Metals have low electron affinities.

• Non-metals have high electron affinities because by gaining electrons, they achieve the stable configuration of a noble gas.

Halogens in group 7A have the highest electron affinity because they only need to gain 1 electron.

5. Electronegativity or Electron Affinity (attraction)

• Across a period – increases because smaller atoms have a greater attraction for electrons and also approaching the non-metals that want to gain.

• Down a group – decrease because larger atoms have less attraction to gain electrons (outer shell farther from nucleus).

6. Ionic Size• Ion – atom with unequal # of protons and

electrons and has charge.

• Metals lose electrons to form positive ions that are smaller than their respective atoms because when outer electrons are lost, the protons in the nucleus have a great pull on the remaining electrons.

6. Ionic Size• Ion – atom with unequal # of protons and

electrons and has charge.

• Non-metals gain electrons to form negative ions that are larger than their respective atoms because when electrons are gained, there are less protons than electrons, so the electrons are not held as tightly.

6. Ionic Size• Across a period – gradual decrease in the

size of positive ions as they lose more electrons and then at group 5 begin the large negative ions that gradually decrease in size as you continue across.

LOSE electrons GAIN electrons 1 2 3 4 3 2 1

7. Ionic Size

• Down a group – increases due to additional energy levels

– Holds true for both positive and negative ions.

Summation of Periodic Trends

Going across a period.

Going down a group.

Atomic size

Metallic character

Non-metallic character

decreases increases

decreases increases

increases decreases

Summation of Periodic Trends 2

Going across a period.

Going down a group.

Nuclear charge

Shielding effect

Speed of reaction – metals

increases increases

no change increases

decreases increases

Summation of Periodic Trends 3

Going across a period.

Going down a group.

Speed of reaction – non-metals

Ionization energy

Electron affinity

increases decreases

increases decreases

increases decreases

Summation of Periodic Trends 4Going across a period.

Going down a group.

Ionic size of

+ ions(cation)

Ionic size of

- ions(anions)

Decreases thru

group 4A

increases

At group 5,

large increases

then a gradual

decrease

increases