PERIODIC TABLE



PERIODIC TABLE

Development of the Periodic Table (pp. 348 - 352)

1. In the late 1790s Lavoisier compiled a list of 23 known elements.

2. During the 1800s many more were discovered as scientists used electricity to separate compounds and newly developed instruments, such as the spectrometer, to identify the elements. By 1870, there were approximately 70 known elements.

3. Chemists were overwhelmed by the volume of new information about the elements. A tool for organizing all this information was needed.

4. A big step came in 1860 when chemists agreed upon a method for accurately determining atomic masses of the elements.

5. The first attempt at organizing the elements was made by John Newlands. He noticed that properties of the elements were repeated every eighth element. He called this the _________________________. This law was not widely accepted because it did not work for all the known elements.

6. Even so, Newlands was correct that the properties of the elements do repeat in a periodic way.

7. Dimitri Mendeleev also arranged the elements by increasing atomic mass. He did not limit the length of his rows. He noticed that the elements fell into columns of elements with similar properties.

8. Mendeleev also left empty spaces to account for elements that had not yet been discovered. He was able to correctly predict the undiscovered element’s properties.

9. There were problems with Mendeleev’s table because as new elements were discovered and atomic masses better determined, some of the elements were placed out of order.

10. Henry Moseley, who had discovered that the atoms of each element contain a unique number of protons in their nuclei, arranged the elements by their atomic number. This adjustment fixed the periodic table.

11. The modern periodic table is based on the _________________________, which states that the physical and chemical properties of the elements tend to change with increasing atomic number in a periodic way.

Modern Periodic Table

1. Each element has its own box containing information about the element.

2. The boxes are arranged in order of increasing atomic number into a series of columns called ____________________, or families, and rows called ____________________.

Regions of the Periodic Table

1. There are five major regions of the periodic table:

2. The metals are the largest region of the periodic table. Elements within this region are solids, generally lustrous, ductile and malleable. They are also excellent conductors of heat and electricity.

3. Nonmetals make up the second largest region of the table. These elements have a variety of physical states and properties. In general they are poor conductors of heat and electricity, brittle and not lustrous.

4. Metalloids are elements that have properties of both metals and nonmetals. These elements are sandwiched between the metals and nonmetals along the “stair step”. There are six of them:

5. The noble gases are the gases located in Group 18. These gases do not have a tendency to react with anything.

6. Hydrogen, because of its unique properties, is also its own region.

FAMILY CHARACTERISTICS

Group 18: Noble Gases

1. The noble gases were once called inert gases because they were thought to be unreactive.

2. No stable compounds of helium, neon and argon have ever been formed.

3. The other noble gases – xenon, krypton, and radon – have very low reactivity. They have been forced to form compounds.

4. Noble gases have full orbitals in the highest energy level, called an ____________________.

5. From this low reactivity we can infer that the noble gas electron configuration is very ____________________.

6. Atoms from the other 17 groups either gain or lose electrons to form compounds. In doing so, the atoms achieve an electron configuration like that of the noble gases.

Group 1: Alkali Metals

1. These elements have metallic properties: soft, shiny, highly reactive, can be cut with a knife, react with atmospheric oxygen, good conductors of heat and electricity.

2. Reactivity comes from their characteristic electron configuration – a single electron in the highest energy level.

3. By losing an electron, an alkali metal atom achieves the stable, nonreactive electron configuration of the noble gas.

Group 2: Alkaline Earth Metals

1. Harder, denser, stronger and have higher melting points than Group 1 elements.

2. Group 2 elements are less reactive than Group 1 elements.

3. This is because they must lose 2 electrons to achieve a noble gas electron configuration.

Groups 3 - 12: Transition Elements

1. These are all metals, but not as reactive as the elements in Groups 1 and 2.

2. They are harder, denser and have higher melting points.

3. There are many irregularities in the electron configurations of the transition elements. These variations occur in the way some of the s-orbitals fill. The transition elements are where the d-orbitals begin to fill.

4. The bottom two rows are known as the _________________________. Each also has its own name. The _________________________, elements 58 – 71, and the _________________________,

elements 90 – 103.

5. The lanthanides are shiny, reactive metals.

6. The actinides have an unstable arrangement of protons and neutrons in the nucleus. They are usually radioactive.

Groups 13 – 18: Main Block Elements

1. These are also known as the _________________________ because they represent a wide range of chemical and physical properties.

2. Within each group the properties may vary systematically.

Group 17: Halogens

1. The halogens combine easily with metals, especially the alkali metals, to form compounds known as ____________________.

2. The halogens are the most reactive nonmetal elements. Their electron configuration is one electron short of the noble gas configuration.

Hydrogen

1. Most common element in the universe.

2. It has only one electron and reacts very rapidly with most other elements.

***** Identify the region from which each of the following elements comes:

Ba I

Sb Rn

***** For each of the given elements, list two other elements with similar chemical properties:

Fe Br

Se Rb

CLASSIFICATION OF THE ELEMENTS

Organizing by Electron Configuration

1. Elements within the same group have the same electron configuration for their outermost energy level.

2. These electrons are the valence electrons.

3. Atoms in the same group have similar chemical properties because they have the same number of valence electrons.

4. The energy level of an element’s valence electrons indicates the period in the periodic table in which it is found.

The s-, p-, d- and f- Block Elements

1. The s-block consists of groups 1 and 2 and hydrogen and helium. In this block, the valence electrons occupy only s-orbitals.

2. The p-block elements consist of groups 13 – 18. As one progresses from left to right, one more electron is added to the p-orbital until it is filled with six electrons. The noble gases exhibit great stability because both their s- and p-orbitals are filled.

3. The d-block contains the transition elements and it is the largest of the blocks. These elements are characterized by a filled outermost s-orbital and a filled, or partially filled, d-orbital.

4. The f-block elements contain the inner transition elements from the lanthanide and actinide series.

***** Strontium has an abbreviated electron configuration of [Kr]5s2. Without using the periodic table, determine the group, period and block in which strontium is located.

***** Write the abbreviated electron configuration of the element in Group 12, period 4.

PERIODIC TRENDS

1. The _________________________ states that the physical and chemical properties of the elements are periodic functions of their atomic numbers.

2. Because the elements are arranged side by side in order of increasing atomic number, one can view certain important vertical and horizontal trends.

Atomic Radius

1. Recall from Rutherford’s experiments that the nucleus was found to occupy a small fraction of the atom’s entire volume.

2. It is the electron cloud surrounding the nucleus that determines the boundaries of the atom.

3. Because the electrons travel in this “cloud” region, it is difficult to measure the size of an atom.

4. The atomic radius is determined in two primary ways. One is to measure the distance between centers of like atoms joined together in a diatomic molecule. The other is to measure the bond lengths of atoms in compounds.

5. Looking at the periodic table one can notice things that affect the size of an atom.

6. As one moves down through a group, a new principal energy level is added. Each new level is farther from the nucleus.

7. As each energy level is added, the energy levels and electrons closer to the nucleus ____________________ it from the effects of the positive charge. This _________________________ allows the outer electrons to be farther away.

8. When crossing a period from left to right each atom gains one more proton and one more electron.

9. No principal energy level is added so the electrons enter the same energy level. The additional protons in the nucleus provide more pull on the electrons bringing them closer to the nucleus.

10. Trend:

Ionic Radius

1. Atoms can gain or lose electrons to form charged particles.

2. Because electrons are negatively charged, atoms that gain or lose electrons acquire a net charge. These particles are called _______________.

3. When atoms lose an electron and form positively charged ions, they become smaller.

4. When atoms gain an electron and form negatively charged ions, they become larger.

Ionization Energy

1. To form a positive ion, an electron must be removed from a neutral atom.

2. This requires energy to overcome the attraction between the positive charge in the nucleus and the negative charge of the electron.

3. This energy is known as the _________________________.

4. Ionization energy is defined as the energy required to remove an electron from a gaseous atom.

5. The energy required to remove the first electron from an atom is called the first ionization energy.

6. The process is:

7. As one moves down a column it becomes easier to remove electrons because they are both shielded and farther away from the nucleus.

8. As one moves left to right across a period it becomes harder to remove electrons because there is no increase, or change, in the shielding effect and the electrons are closer to the nucleus.

9. Trend:

Electronegativity

1. The ____________________ of an element indicates the relative ability of its atoms to attract electrons in a chemical bond.

2. These values are calculated based on a number of factors and have values of 3.98 or less.

3. The noble gases are ignored since they are basically inert.

4. Fluorine is the most electronegative element with a value of 3.98 while cesium and francium are the least electronegative.

5. In a chemical bond, the atom with the greater electronegativity more strongly attracts the bond’s electrons.

6. Trend:

***** Which has the largest atomic radius? The smallest?

Mg Si S Na

***** Which has he largest atomic radius? The smallest?

He Xe Ar

***** For each of the following properties, indicate whether fluorine or bromine has a larger value:

a. electronegativity b. ionic radius

c. atomic radius d. ionization energy

IONIC COMPOUNDS

FORMATION OF IONS (pp 71 - 74)

1. Many of the properties of the elements are due to the valence electrons.

2. These same electrons are involved in the formation of chemical bonds between two atoms.

3. Electron dot structures help keep track of the valence electrons.

4. Remember that ionization energy refers to how easily an atom loses an electron.

5. Another property, ______________________________, refers to how much attraction an atom has for electrons.

6. Elements gain, lose, or share valence electrons so their valence shell mimics that of a noble gas.

7. Noble gases have a full outermost energy level ____________________, which gives them a stable ____________________.

Positive Ions

1. A positive ion forms when an atom loses one or more valence electrons.

2. A positive ion is called a ____________________.

3. Losing electrons gives the resulting cation an electron configuration similar to that of a noble gas, but it does not change the element into a noble gas.

4. The protons that establish the character of the element remain in the nucleus.

5. The reactivity of metals is based on the ease with which they lose valence electrons.

6. Transition elements lose electrons from both the _____ and _____ orbitals, thus giving them a wide variety of cations.

Negative Ions

1. A negative ion forms when an atom gains one or more electrons.

2. Nonmetals have a great attraction for electrons and form a stable octet by gaining electron.

3. A negative ion is called an _________________________.

IONIC BONDS (pp 75 - 78)

1. Oppositely charged ions attract one another.

2. The electrostatic forces that hold oppositely charged particles together in a compound are called an ______________________________.

3. Any compound formed in this manner is called an ______________________________.

4. Ionic compounds have no resultant charge.

5. The process for sodium chloride looks like:

Properties of Ionic Compounds

1. The chemical bonds that occur between the atoms in a compound determine many of the physical properties of the compound.

2. In an ionic compound, the positive and negative ions are packed into a regular repeating pattern that balances the forces of attraction and repulsion.

3. This particle packing forms an ionic lattice.

4. The three-dimensional crystal lattice gives the crystal its shape and properties.

5. Ionic compounds are: very strong, rigid, have high melting and boiling pints, and require a large amount of energy to be broken.

NAMES AND FORMULAS FOR IONIC COMPOUNDS (pp 85 - 94)

1. Since the generic term “salt” can mean any one of thousands of chemical compounds, a better system is needed to describe an ionic compound.

2. The system is based on the chemical symbols on the periodic table.

Monatomic Ions

1. A monatomic ion is either a cation or an anion formed from a single atom.

2. To write the chemical symbol for a monatomic ion you must indicate both the symbol for the element and its charge.

***** Write the symbol for the ions formed from:

Cs F

Rb O

Be N

Al S

3. The charge of a monatomic ion is its ______________________________. This is found in the upper right corner of the element’s box.

4. Groups 1, 2, 13-18 generally have one oxidation number while the transition elements typically have more than one.

5. Some elements have both positive and negative oxidation numbers. The oxidation state used will depend on the other elements bonding with it.

6. The primary oxidation number is listed first.

7. The oxidation number, or oxidation state, of an element in an ionic compound equals the number of electrons transferred from one atom and accepted by the other atom.

8. The nomenclature for monatomic ions is fairly simple.

9. A monatomic cation is named using the element’s name followed by the word __________.

10. A monatomic anion is named by dropping the ending of the element’s name and adding the suffix __________.

***** Name the following ions:

H+ H-

O2- S2-

Ca2+ F-

Li+ Ba2+

I- P3-

Nomenclature for Multiple Charge Ions

1. The transition elements form many different cations depending upon the number of electrons given up.

2. To distinguish between the various charges, Roman numerals are used to represent the number of electrons given up.

*****Name the following ions

Fe2+ Cr6+

Fe3+ Pt4+

Cr2+ Cr3+

Zn2+

3. Anions, even though they may have multiple negative charge states, _______________ use Roman numerals.

Writing Formulas for Ionic Compounds

1. An ionic compound is _______________ even though it is composed of charged ions.

2. Steps for writing a chemical formula:

a. List the symbol for the cation and its charge.

b. List the symbol for the anion and its charge.

c. Write the symbols for the ions, side-by-side, with the cation first.

d. If the charges don’t balance, criss-cross them.

e. The charge number – no sign – becomes the subscript for the cation or anion.

***** Determine the formulas for the following compounds:

aluminum nitride aluminum oxide

strontium phosphide iron(II) oxide

calcium chloride tin(IV) fluoride

rubidium iodide copper(II) arsenide

3. Chemical formulas must reflect the actual composition of the compound.

Naming Ionic Compounds

1. Ionic compounds consisting of two elements are known as ____________________.

2. To name a binary compound, you just combine the names of the ions; cation first, then the anion. If the cation is a transition element that has multiple oxidation numbers, then a Roman numeral must be used.

***** Name the following ionic compounds:

MgBr2 MnCl4

Al2S3 VI3

Ni2O3 CrO3

Re2O7 NpF5

Polyatomic Ions

1. A ______________________________ is an ion made of two or more atoms bonded together that function as a single ion.

2. Polyatomic ions have special names.

3. Polyatomic ions can be either cations or anions.

4. Compounds formed by polyatomic ions must be neutral.

5. Polyatomic ions have charges just like the monatomic ions. The charge sign is written to the right of the ion’s formula. Remember that this charge applies to the _______________ ion, not to any individual element.

6. Parentheses are used to group polyatomic ions. The subscript outside the parentheses refers to everything within the parentheses.

Compounds With Polyatomic Ions

1. Ionic compounds containing polyatomic ions are named in the same manner as binary ionic compounds. The cation is listed first, followed by the anion.

2. In some cases, two different polyatomic ions are formed by the same two elements. One of the two elements in every case is _______________.

3. When naming these compounds, the ions with the larger number of oxygens are named with the ________ ending and the ions with the smaller number of oxygens are named with the _______ ending.

4. The ending does not tell you how many oxygen atoms are present in the polyatomic ion.

5. The rules for writing formulas for ionic compounds involving polyatomic ions are similar to those for binary compounds.

6. When a polyatomic ion has a subscript, use parentheses to make it clear that the subscript refers to the entire polyatomic ion.

7. Parentheses are not needed if only a single polyatomic ion is used.

8. Parentheses can be used with polyatomic cations and anions.

Rules For Writing Formulas For Polyatomic Ions

1. Identify the cation and anion names in the compound.

2. Write the symbols for the cation and the anion side by side.

3. If the cation name is followed by Roman numerals, then assign that amount of charge to the cation and determine the anion charge.

4. If the cation name is not followed by Roman numerals, then determine the ions’ charges.

5. Then determine the least common multiple of the ion’s charges. (Criss-cross and simplify)

6. If polyatomic ions are not present, use subscripts to indicate how many of each ion would be necessary to have the amount of charge designated by the least common multiple.

7. If polyatomic ions are present, use subscripts and parentheses to indicate how many of each ion would be necessary to have the amount of charge designated by the least common multiple.

***** Write the formula for the following compounds:

aluminum sulfate hydrogen peroxide

magnesium hydroxide iron (III) sulfide

copper (II) acetate lead (II) phosphate

iron (II) chlorate sodium bicarbonate

dimercury(I) acetate ammonium hydroxide

copper (II) phosphate aluminum oxalate

Rules For Naming Polyatomic Ions

1. Identify the cation and the anion.

2. If the cation is a metal that can have more than one charge, then determine the cation charge.

3. Write the cation name using Roman numerals for the charge.

4. Write the anion name.

5. If the cation is either a polyatomic ion or a metal that can have only one charge, then write the cation name first and the anion name last.

***** Name the following compounds:

NaNO3 Ca(NO2)2

Fe(OH)3 AlAsO4

(NH4)2SO4

METALLIC BONDS (p 459)

1. Metals do not bond ionic ally, but they often form lattices in the solid state.

2. Metal atoms do not share their valence electrons with neighboring atoms nor do they lose electrons to form ions.

3. Instead, the outer energy levels overlap.

4. The ______________________________ proposes that all the metal atoms in a metallic solid contribute their electrons to form a “sea” of electrons.

5. The electrons are free to move from one atom to another. They are often referred to as ______________________________.

6. A ______________________________ is the attraction of a metallic cation for delocalized electrons.

Properties of Metals

1. Delocalized electrons moving around positive metallic ions explain why metals are such good conductors.

2. The delocalized electrons interact with light, absorbing and releasing photons, thereby creating the property of luster.

3. As the number of delocalized electrons increases, so do the properties of hardness and strength.

Metal Alloys

1. Due to the nature of the metallic bond, it is easy to introduce other elements into a metallic crystal.

2. An ____________________ is a mixture of elements that has metallic properties.

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