Unit VI: The Periodic Table



Unit VI: The Periodic Table

I. Naming Chemical Compounds

A. Monatomic Ions – ions consisting of only one atom.

- ionic charges can often be determined by

using the periodic table.

Group 1 – Form cations (positive ion), have a +1 charge

Group 2 – Form cations, have a +2 charge

Group 6 – Form anions (negative ion), have –2 charge

Group 7 - Form anions (negative ion), have –1 charge

* When assigning charges, # first and charge second. Fe2+ , not

Fe+2

D-Block or Transition Elements – more than one common ionic

Charge

Ex) Fe2+ Fe3+

If an element has more than one common ion, use a roman

numeral to indicate the charge. (Stock naming System)

Ex) Fe2+ = Iron(II) Fe3+ = Iron(III)

* OLD METHOD: If there are two possible charges the

endings –ous or –ic may be used. –ous is the lower number, -ic

is the higher. Need to know the latin name.

Fe2+ - ferrous ion Fe3+ - Ferric ion

Cu+ - cuprous ion Cu2+ - cupric ion

B. Polyatomic Ions

1. Tightly bound groups of atoms that behave as a unit

and carry a charge.

Ex) SO42- = Sulfate ion

NO31- = Nitrate ion

2. Most polyatomic ions end in –ite or –ate

-ite ending indicates one less Oxygen atom than

the –ate ending.

Ex) SO32- = sulfite SO42- = sulfate

3. Three important exceptions

NH4+ = ammonium cation

CN- = cyanide

OH- = hydroxide

4. Front Page of Reference Table has complete list of

Polyatomic Ions.

II. Writing/Naming Chemical Formulas

A. Writing Binary Ionic Compounds.

- Binary compounds- compounds that are made up of only 2

different elements.

1. The positive charge of the cation must balance the negative

charge of the anion.

*Net ionic charge of the formula must be 0

Ex) KCl = K+ and Cl-

2. Cation is always written first, the anion ends in –ide

Ex) KCl = Potassium chloride

3. Crisscross method

ex)what is the formula for rust (iron(III) oxide)?

Fe3+ O2-

Fe2O3

Ex) Write the formula for a compound such as calcium

sulfide (which contains Ca2+ and S2-)

Ca2+ S2-

Ca2S2 which reduces to CaS

B. Naming Binary Ionic Compounds

1. CuO – how do you name it?

Copper oxide? NO

2. You must know the charge of each of the ions.

3. Cu can form Cu+ and Cu2+

4. Oxygen always has –2 charge, so Cu must be Cu2+

5. So the name is Copper(II) Oxide

C. Ternary Ionic Compounds

1. Compound that contains atoms of three different elements.

2. Usually contain a polyatomic ion.

Ex) CaCO3

3. Write the formula for Calcium nitrate (Ca2+ and NO31-)

Ca2+ NO31-

Ca(NO3)2

*HOMEWORK – Naming ionic compounds

*Quiz – writing chemical formulas

D. Binary Molecular Compounds

1. Composed of two nonmetallic elements.

2. Ionic charges are NOT used in naming

3. The name of the element farthest to the left in the periodic

table is usually written first. EXCEPTION: Oxygen is

always last unless with F.

4. If both elements are in the same group, lower one is first.

5. Prefixes help distinguish between different compounds.

Mono- 1 hexa- 6

di- 2 hepta- 7

tri- 3 octa- 8

tetra- 4 nona- 9

penta- 5 deca- 10

6. When prefix ends in –a or –o and second element begins

with a vowel, the –a or –o is dropped.

ex. CO has one oxygen or (mono) thus Carbon monoxide

CO2 has two oxygen or (di) thus Carbon dioxide

7. All binary molecular compounds end in –ide (like ionic)

*HOMEWORK: Naming molecular compounds

E. Naming common acids

1. acids are compounds that produce hydrogen ions when

dissolved in water.

2. MEMORIZE

A. HCl – hydrochloric acid

B. H2SO4 – sulfuric acid

C. HNO3 – Nitric acid

D. HC2H3O2 – acetic acid

E. H3PO4 – phosphoric acid

F. H2CO3 – Carbonic Acid

*Quiz: Naming chemical compounds #1

*Quiz: Naming/Writing chemical compounds #1

III. Development of the Periodic Table

A. The Russian scientist Dmitri Mendeleev first organized the

elements into a table based upon properties of the elements that

were recurrent (periodic) in 1869. Elements in the same column

(Group) generally have similar properties.

*1864 – John Newlands – early table likened periodicity to the

musical octave – “law of octaves” – properties repeated with every eighth element.

B. The modern periodic table basically arranges the elements

according to an increasing number of protons (atomic number).

C. Where did the elements come from?

1. Big Bang – created all the matter in the universe, most of the

Hydrogen and Helium.

- at one time the universe was smaller than an

atom.

- A few minutes after the big bang 95% of the

atoms were H2, 5% Helium, trace amounts of Li.

2. Small Stars – Fuse H2 into He, He into Carbon, Nitrogen,

Oxygen.

- A stars gravity pushes in, while the fusion in the

center pushes out. Once the fuel is used up,

fusion stops and the stars falls in on itself.

-This causes the H2 in the outer layers to fuse and

the star “puffs up” – this is Red Giant.

- The star still continues to collapse which fuses

the Helium into Carbon.

- When there is no material left to fuse, the star

collapses again, creating heat in the center,

which blows off the stars outer layers.

- The remaining mass continues to collapse. The

star is now a white dwarf. The density is so

high, that one tablespoon would weight 1 ton.

3. Large Stars – Fuse heavy elements as well as light elements.

ex) the Ca in your bones, Fe in your blood,

S in your hair.

4. SuperNova – In rare cases, a white dwarf can detonate in a

massive explosion called a supernova.

ex) Magnesium in chlorophyll, Gold, Titanium

5. Cosmic Rays – High-energy particles traveling throughout

space close to the speed of light. Lithium comes primarily from these rays, used in watch batteries, cell phone batteries.

IV. Trends of the Elements in the Periodic Table

A. Arrangement

1. The horizontal rows of the periodic table are called periods

(or rows)

a. The number of the period is the same as the outermost

principle energy level.

2. The vertical column are called groups or families.

a. The elements in a group exhibit similar properties.

B. Atomic Radius

1. ½ the distance of closest approach of 2 adjacent nuclei.

2. Along a period, the atomic radius generally decreases with

increasing atomic number (due primarily to an increasing

nuclear charge which pulls the electrons closer to the

nucleus)

3. In a group, the atomic radius generally increases with

increasing atomic number (due primarily to an increasing

number of principle energy levels and the shielding effect of

the non-valence electrons.)

4. SNOWMAN EXAMPLE

EX) Arrange in order of increasing atomic radius: Na, Be, Mg

ans. Be < Mg < Na

5. Different types of atomic radii

a. Covalent radius

1. The radius of an atom when it has formed a

covalent bond.

2. ½ the distance between nuclei of 2 atoms

covalently bonded together in the crystalline

(solid) phase

b. Van der Waals radius

1. ½ the distance of closest approach between the

nuclei of 2 atoms that are not bonded together.

c. Atomic radius in metals

1. ½ the distance between the nuclei of atomsi n a

crystalline (solid) metal

C. Ionic Radius

1. The loss of electrons causes the radius of the ion to be

smaller than that of the corresponding atom (due to the fact

that there are now more protons than electrons thus causing

the electrons to be pulled in closer to the nucleus.

2. The gain of electrons causes the radius of the ion to be larger

than that of the corresponding atom (due to the fact there

are now more electrons than protons and the force of

attraction for each electron is less causing the electrons to

be farther from the nucleus.

D. Ionization Energy

1. The energy required to remove an electron from a gaseous

atom is called the ionization energy.

Na(g) ( Na+(g) + e-

2. The energy required to remove the outermost electron is

called the first ionization energy. To remove the outermost

electron from a 1+ ion is called the second ionization

energy.

3. As you move down a group, ionizaton energy DECREASES

a. size of atom is increasing, so outmost electron is

farther from the nucleus.

4. As you move across a period, ionization energy

INCREASES

a. Nuclear charge is increasing, so there is a greater

attraction for the electrons to the nucleus.

5. Snowman is opposite for I.E.

ex) Using the periodic table arrange in order of increasing first

IE: Ne, Na, P, Ar, K

ans. K < Na < P < Ar < Ne

E. Electronegativity

1. Tendency for atoms to attract electrons to itself.

2. Electronegativity DECREASES as you go down a group.

3. Electronegativty INCREASES as you go across a period.

F. Metals

1. Left hand side of the periodic table

2. Metallic characteristics increase down a column and are thus

greatest in the lower left hand corner of the periodic table.

3. More than 2/3 of all the elements are metals

4. Properties:

a. tend to lose electrons when forming compounds (low

ion. Energy and low electro.)

b. good thermal and electrical conductors

c. exhibit a metallic luster when polished

d. malleable (can hammer into thin sheets)

e. ductile (can be drawn into wire)

f. Hg is the only metal that is liquid at room temperature.

** Alloys – mixtures composed of two or more elements, at least one of

which is a metal

ex) brass is an alloy of copper and zinc

sterling silver – 92.5% silver/7.5% copper

steel – Fe, C, B, Cr, Mn, Mo, Ni, W, V

G. Non-metals

1. Right hand side of the periodic table

2. Non-metallic characteristics decrease down a column and

are thus greatest in the upper right hand corner of the

periodic table.

3. Non-metals tend to be gases, molecular, or network solids.

4. Properties:

a. tend to gain electrons when forming compounds (high

ionization energy and high electronegativities.)

b. poor thermal and electrical conductors (good

insulators)

c. lack metallic luster

d. brittle

e. Br is only non-metallic element that is liquid at room

temp.

H. Metalloids

1. Elements that have about ½ metallic and ½ non-metallic

characteristics.

2. Metalloids are those elements which are between the metals

and the non-metals on the periodic table (along the solid

jagged line)

3. Best examples: B, Si, As, Te

V. Chemistry of the Groups (Families)

A. General Properties

1. Groups are labeled 1-18

2. Elements in a group have similar properties because they

have the same number of valence electrons.

3. Elements in a group react similarly with elements from other

groups.

As atomic number increases (going down a group)

4. the radius of the atom increases

5. ionization energy decreases

6. electronegativity decreases.

7. metallic characteristics increase

B. Groups 1 & 2

1. Group 1 – Alkali metals (H is not an alkali metal)

*comes from Arabic word meaning “ashes”

*7-Up originally contained lithium, 1927 removed in 1950

2. Group 2 - Alkaline earth metals

3. Most reactive metals (occur in nature only in compounds)

4. In the same period, the group 1 metal is more reactive than

the group 2 metal.

5. Generally, reactivity increases as atomic number increases.

6. Low ionization energies and electronegativities and

therefore often form ionic compounds.

7. Elements are usually produced through the electrolysis of

their fused (liquefied/molten) compounds.

8. 2M(s) + 2H2O(l) ( 2MOH(aq) + H2(g)

C. Group 15

1. Progress from non-metals to metals

a. N & P – non-metals

b. As – metalloid

c. Bi – metal

2. Nitrogen

a. relatively inactive gas

b. diatomic molecule containing a triple bond

c. N compounds are essential for all living matter.

d. N compounds can be unstable and are thus useful as

explosives (TNT)

3. Phosphorus

a. tetratomic molecule

b. also very essential for all living matter

D. Group 16

1. Progress from non-metals to metals

a. O & S – non-metals

b. Te – metalloid

c. Po – metal

2. Oxygen

a. very reactive (forms compounds with most elements)

b. diatomic (double bond)

** Diatomic elements – two of the same element bonded together. Mr.

BrINClHOF

c. continually produced by plants during photosynthesis.

d. second highest electronegativity (always shows a

negative oxidation state except when combined with

Fluorine)

** Allotropes: Different structural forms of the same element

- Cdiamond & Cgraphite

- O2 & O3 (ozone)

3. Sulfur

a. has several allotropic forms (S8)

4. Se & Te

a. rare elements

5. Polonium

a. radioactive

E. Group 17

1. Halogens

*greek words “halos” and “gennao” ( salt formers

2. All are typical non-metals

3. High electronegativities as a group

a. Fluorine

1. highest electro. (always shows a negative

oxidation number in compounds.)

4. F & Cl are gases at room temperature

5. Br is a liquid

6. I is a solid

7. At is radioactive

8. all are diatomic

F. Group 18

1. Noble gases (inert gases)

2. Completely filled outer (valence) shells

3. Very unreactive

4. Monoatomic – single atoms

Trivia: Helium – Named for greek god of the sun, Helios

most inert of ALL substances

At absolute zero He is still a liquid, while all

other substances are solid.

-Below 2.2K it conducts heat perfectly, no

part of the liquid is warmer than another

part.

He(l) has practically no viscosity, it flows

more easily than a gas. Can go through

openings that would stop a gas.

Argon – derived from greek word meaning “lazy” for its

refusal to react.

G. Groups 3-12 (the “d” block elements)

1. Transition elements (groups 3-11)

2. Partially filled outer “d” orbitals

3. Multiple positive oxidation states

4. Transition element ions are usually colored in solutions and

in solid compounds.

VI. Chemistry of a Period

A. Trends (properties)

As atomic number increases (from left to right)

1. atomic radius decreases

2. ionization energy increases

3. electronegativity increases

4. transition from positive to negative oxidation states

5. metallic characteristics decrease

B. Lanthanoid (Lanthanide) & Actinoid (Actinide) Series

1. similar chemical properties

2. “f” block elements

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