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.

2- Ex) SO4 = Sulfate ion 1- NO3 = Nitrate ion

2. Most polyatomic ions end in –ite or –ate -ite ending indicates one less Oxygen atom than the –ate ending.

1 2- 2- Ex) SO3 = sulfite SO4 = 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

2 C. Ternary Ionic Compounds 1. Compound that contains atoms of three different elements. 2. Usually contain a polyatomic ion.

Ex) CaCO3 2+ 1- 3. Write the formula for Calcium nitrate (Ca and NO3 ) 2+ 1- Ca NO3 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

3 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

4 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

5 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-

6 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

7 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

8 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

9 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

10 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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