Organizing the Periodic Table

How did chemists begin to organize the

known elements?

Chemists used the properties of the

elements to sort them into groups.

The Organizers

JW Dobereiner

grouped the elements into triads.

Problem: Not all of the elements could

be put into triads.

The Organizers

Dmitri Mendeleev

arranged the elements by

increasing atomic mass.

Problem: the chemical properties of

the elements did not line up correctly.

Pg. 156

Figure

6.3

The Organizers

Lother Meyer

also created a periodic table similar to

Mendeleev’s, however Mendeleev

published his table first and could

explain his table better.

Development of Atomic Theory

The structure of the atom was

determined through years of

experimentation. Once the structure

of the atom was determined by Dalton,

Thomson, Rutherford, Chadwick, Bohr,

and others, Henry Moseley

determined the atomic number for

the elements.

Organization of

Modern Periodic Table

The elements are arranged

in order of

increasing atomic number.

Periodic Law

When elements are arranged in

order of increasing atomic

number, there is a periodic

repetition of their physical and

chemical properties.

Three classes of elements:

Metals

Nonmetals

Metalloids

Across a period, the properties of

elements become less metallic and

more nonmetallic.

Properties of Metals

•Solids (except for Hg) at room temperature

•Make up 80% of the periodic table

•Good conductors of heat and electricity

•Shiny

•Ductile (can be made into thin wire)

•Malleable (can be hammered into sheets)

Properties of Nonmetals

Properties vary among the nonmetals.

At room temperature, some are gases,

some are solids, and 1 is a liquid.

•Poor conductors

•Brittle

•Not ductile

•Not malleable

Metalloids (Semi-metals)

B, Si, Ge, As, Sb, Te, At

Behave like metals under certain

conditions.

Behave like nonmetals under

different conditions.

Which of these sets of elements have

similar physical and chemical properties?

a. oxygen, nitrogen, carbon, boron

b. strontium, magnesium, calcium, beryllium

c. nitrogen, neon, nickel, niobium

Identify each element as a metal,

metalloid, or nonmetal.

a. gold

b. silicon

c. sulfur

d. barium

Name two elements that have

properties similar to those of the

element sodium.

What pattern is revealed when the

elements are arranged in a periodic table

in order of increasing atomic number?

Identify each property below as more

characteristic of a metal or a nonmetal.

a. a gas at room temperature

b. brittle

c. malleable

d. poor conductor of electric current

e. shiny

In general, how are metalloids different

from metals and nonmetals?

Representative Elements

Elements in columns

1, 2, 13, 14, 15, 16, 17, 18

OR

Elements in columns

1A, 2A, 3A, 4A, 5A, 6A, 7A, 8A

OR

Elements in blocks s and p

Representative Elements

Called “REPRESENTATIVE”

elements because they display a

wide range of chemical and

physical properties.

Which of the following are symbols for

representative elements:

Na, Mg, Fe, Ni, Cl ?

Transition Metals

Elements in columns

3, 4, 5, 6, 7, 8, 9, 10, 11, 12

OR

Elements in columns

3B, 4B, 5B, 6B, 7B, 8B, 1B, 2B

OR

Elements in the d block

Transition Metals

Characterized by electrons

filling the d orbitals

When transition metal ions form compounds,

the compounds often exhibit various colors.

Iron oxide Red

Cadmium sulfate Orange

Copper(II) chloride Blue-green

Inner Transition Metals (Rare earth elements)

Elements in the f block

Consists of the

Lanthanide (La) series

and the

Actinide (Ac) series.

Inner Transition Metals (Rare earth elements)

Contains some elements that occur

naturally

and some that are synthetic.

Some are radioactive, some are not.

Groups in the Periodic Table

Alkali metals – column 1 (or 1A)

Alkaline earth metals – column 2 (or 2A)

Halogens (salt-formers) – column 17 (7A)

Noble gases – column 18 (or 8A)

Why do the elements potassium and

sodium have similar chemical properties?

Classify each element as a

representative element, transition

metal, or noble gas.

a. 1s22s22p63s23p64s23d104p6

b. 1s22s22p63s23p64s23d6

c. 1s22s22p63s23p2

Which of the following elements are

transition metals:

Cu, Sr, Cd, Au, Al, Ge, Co ?

How many electrons are in the

highest occupied energy level of a

Group 5A element?

Trends

in the

Periodic

Table

Shielding is the result of the core

electrons blocking the ‘pull’ of the

nucleus on the outermost electrons.

Terms to Know

Nuclear charge is the charge

in the nucleus.

Terms to Know

Atomic Radius is one half the distance

between the nuclei of two atoms of the

same element when the atoms are joined.

Pg. 170 Figure 6.13

Atomic Radius

From left to right across a period,

the atomic radius generally decreases.

Why? As you go across a row, there is an increase

in the number of protons and the number of electrons.

However, the electrons being added are added to the

same principle energy level, so the shielding remains

the same. Therefore, the increased nuclear charge

pulls the outermost electrons in closer to the nucleus.

The result is a smaller atom.

(See pg. 171, 3rd paragraph.)

Atomic Radius

From top to bottom down a group, the

atomic radius generally increases.

Why? As you go down a column, there is an

increase in the number of protons and the number of

occupied energy levels. The increase in positive

charge pulls the electrons closer to the nucleus, but

the increase in the number of occupied orbitals

shields the nucleus more. The shielding effect is

greater than the increased nuclear charge. The

result is a larger atom.

(See pg. 171, 1st and 2nd paragraph.)

Which element in each pair has atoms

with a larger atomic radius?

a. sodium, lithium

b. strontium, magnesium

c. carbon, germanium

d. selenium, oxygen

Explain why fluorine has a smaller atomic

radius than both oxygen and chlorine.

Terms to Know

Ion is a neutral atom that has either lost

or gained an electron.

Atoms that have gained electrons are

called

Atoms that have lost electron(s) are called

Positive Ions or Cations

Negative Ions or Anions

Cation is an ion that has lost one or

more electrons resulting in a net

positive charge.

Anion is an ion that has gained one

or more electrons resulting in a net

negative charge.

Ion Size

Cations are always

SMALLER than the atoms

from which they form.

Why? When a metal atom loses an electron,

the attraction between the remaining electrons

and the nucleus is increased. The electrons are

drawn closer to the nucleus.

(See pg. 176, 2nd paragraph)

Ion Size

Anions are always

LARGER than the atoms

from which they form.

Why? As the number of electrons

increases, the attraction of the nucleus

for any one electron decreases.

(See pg. 176, 3rd paragraph)

Li F O N C B Be

Li1+ F1- O2- N3- C B3+ Be2+

Ion Trends

From left to right across a period, the

cation size decreases.

From left to right across a period, the

anion size decreases.

(See figure 6.20 on pg. 176)

Pg. 176 Figure 6.20

From top to bottom down a group,

the ionic size increases.

(Pg. 172)

Ion Trends

Which particle has the larger

radius in each atom/ion pair?

a. Na, Na+

b. S, S2-

c. I, I-

d. Al, Al3+

In each pair, which ion is larger?

a. Ca2+, Mg2+

b. Cl - , P3-

c. Cu+, Cu2+

The atomic radius _____________ from

left to right across the periodic table.

The atomic radius ____________

from top to bottom on the period table.

Aside from the noble gases, what is the

smallest atom on the periodic table?

Aside from the noble gases, what is the

largest atom on the periodic table?

The atom is _________ than its cation.

The atom is _________ than its anion.

Cations are _________ than the

anions that are in the same row.

The cation trend from left to right is

that the ionic radius ____________.

The anion trend from left to right is

that the ionic radius _____________.

List the following from smallest to

largest according atomic radius.

a. Mg, S, Na, Si

b. Be, Ba, Ca, Ra

c. F, As, Br, Ga, Cl

List the following from smallest to

largest according ionic radius.

a. Mg2+, S2-, Na+, Cl-

b. Be2+, Ba2+, Ca2+, Ra2+

c. F-, As3-, Br-, Ga3+, Cl-

Terms to Know

Ionization

Energy

is the energy

required to

remove an

electron

from an

atom. (Pg. 173-175)

Explain the difference between the terms

“first ionization energy” and “second

ionization energy” of an element.

The first ionization energy is the energy

needed to remove a first electron from an

atom.

The second ionization energy is the energy

needed to remove a second electron.

Why is there a large increase between the

first and second ionization energies of the

alkali metals?

It is relatively easy to remove the first

electron from an alkali metal atom because

they have only one valence electron;

it is much more difficult to remove the

second because it is being removed from a

full energy level.

There is a large jump between the

second and third ionization energies

of magnesium. There is a large jump

between the third and fourth

ionization energies of aluminum.

Explain these observations.

It is relatively easy to remove two electrons

from magnesium because it has 2 valence

electrons; it is much more difficult to remove a

third electron because it comes from a full

energy level.

It is relatively easy to remove three electrons

from aluminum because it has 3 valence

electrons; it is much more difficult to remove a

fourth electron because it comes from a full

energy level.

Ionization Energy

From left to right across a period,

the ionization energy generally

increases.

Why? As you go across a row, there is an increase

in the nuclear charge, but the shielding effect remains

the same. There is an increase in the attraction of the

nucleus for an electron. Thus it takes more energy to

remove an electron from the atom.

(See pg. 174, 2nd paragraph.)

From top to bottom down a group,

the ionization energy generally

decreases.

Why? As the size of the atom increases and the

amount of shielding electrons increases, nuclear

charge has a smaller effect on the electrons in the

outermost energy level. Therefore, less energy is

required to remove an electron from the outermost

energy level.

(See pg. 174, 1st paragraph.)

Ionization Energy

Which element in each pair has a

greater first ionization energy?

a. lithium, boron

b. magnesium, strontium

c. cesium, aluminum

Would you expect metals or nonmetals in

the same period to have higher ionization

energies? Give a reason for your answer.

Explain each of the following comparisons.

Calcium has a smaller second ionization

energy than does potassium.

Explain each of the following comparison.

Lithium has a larger first ionization

energy than does cesium.

Explain the following comparison.

Magnesium has a larger third

ionization energy than does aluminum.

Electronegativity

Electronegativity

is the ability of an atom to attract

electrons when the atom is in a

compound.

(See pg. 177, 1st paragraph)

For the representative elements:

From top to bottom down a group,

the electronegativity values decrease.

(See pg. 177, 3rd paragraph)

From left to right across a period,

the electronegativity values increase.

Electronegativity

Electronegativity

Fluorine is the most electronegative

element on the periodic table.

Why are the noble gases NOT included

in electronegativity table?

Because they are UNREACTIVE and

do NOT want to bond with any atoms.

Which element in each pair has a

higher electronegativity value?

a. Cl, F

b. C, N

c. Mg, Ne

d. As, Ca

When the elements in each pair

are chemically combined, which

element in each pair has a

greater attraction for electrons?

a. Ca or O

b. O or F

c. H or O

d. K or S

Pg. 171

Figure 6.14

Pg. 174

Figure 6.17