The Periodic Table

Periodic Table of Elements• There are 117 elements (January, 2007)

– Your table contains 113• 94 of the elements are naturally occurring,

the rest are man-made• Most of the elements were discovered

between 1735-1843

History/Development• Development of the table has occurred

over 300 years and continues today• Dmitri Mendeleev is commonly credited

with creating the periodic table in 1869

Original Table (Mendeleev)• Classified elements in horizontal rows

based on their atomic mass• When there was a repeat of properties

the elements were placed in the next row– Concluded that similar properties appear at

regular intervals when the elements are listed in order of increasing atomic mass

Original Table (Mendeleev)• Elements with similar properties were

located in the same vertical columns, if no known element had the expected properties to fit the particular space he left that space empty– Assumed that elements not yet discovered

would fit into the empty spaces– Predicted some properties of these unknown

elements (gallium, scandium, and germanium)

Modern Periodic Law - 1914• Henry Moselely bombarded

elements with high speed electrons and they emitted X-rays with a certain wavelength

• He found that each element differed by one proton

• Concluded that the regularity, or periodicity, of the properties is a function of the atomic number (Modern Periodic Law)

Questions

1. The modern periodic table is arranged by increasing _______________________.

2. The original periodic table was arranged by increasing _____________________.

Organization of the Modern Periodic Table

• The periodic table is arranged in order of increasing ATOMIC NUMBER

• Horizontal rows are called PERIODS– There are 7 periods

• Vertical columns are called GROUPS– There are 18 groups

Periods• Horizontal Rows• The number of each period indicates the

principle energy level in which the valence electrons are located

• The number of valence electrons increases as you go from left to right

• The properties of the elements change systematically through a period

Groups (Families)• Vertical Columns• The outermost shell of an atom contains

the same number and arrangement of valence electrons

• Elements that have similar chemical properties are located in the same group

Group/Period Examples1. Are elements in the same period or group

more similar? Explain why. 2. Which elements have the most similar

chemical properties? a. K and Nab. K and Clc. K and Cad. K and S

3. Why do elements of a given group on the periodic table show similar chemical properties?

States of Matter• Solid

– The majority of the elements are solids• Liquid

– The only liquids are Hg and Br, which are found on the right

• Gas– H, O, N, F, Cl and the Noble Gases (Group

18), located on the right

Metals

Metalloids

Nonmetals

Metals/Nonmetals/Metalloids• Where are the metals located?

• Where are the nonmetals located?

• Where are the metalloids located?

– List all of the Metalloids

Properties of Metals and Nonmetals

• Lose electrons to form positive ions

• Solids at STP– Except Hg

• High melting and boiling points

• Good thermal (heat) and electrical conductors

• Luster (shine)• Malleable (bendable)• Ductile – can be made

into wires

• Gain electrons to form negative ions

• Solids, 1 liquid (Br), gases at STP

• Low melting and boiling points

• Poor conductors of heat and electricity

• Dull • Brittle• Not ductile

Metalloids• Properties of both metals and nonmetals• B, Si, As, Te, Ge, and Sb (front + middle

back of staircase)

Metal/Nonmetal Questions1. Atoms of metals tend to

a. lose electrons to form positive ionsb. lose electrons to form negative ionsc. gain electrons to form positive ionsd. gain electrons to form negative ions

2. The majority of elements on the table are classified asa. metalsb. nonmetalsc. metalloids

3. Which property is generally characteristic of metallic elements? a. low electrical conductivityb. high heat conductivityc. existence as brittle solidsd. low melting points

Metal/Nonmetal Examples4. At room temperature, which substance is the best conductor of electricity?

a. nitrogenb. neonc. sulfurd. silver

5. Which element is brittle in the solid phase and a poor conductor of electricity? a. calciumb. strontiumc. sulfurd. copper

6. The majority of elements on the table are in what physical state at STP? a. solidb. liquidc. gas

Atomic Radius• ½ the distance between any two nuclei• Given on Reference Table S in picometers

– 1pm = 1x10-12m

Trend within a Period (Left to Right)

• Atomic Radius decreases• As you move across a period the number of

protons increases, resulting in a stronger nuclear charge therefore, electrons are pulled closer to the nucleus

Trend within a Group (Top to Bottom)

• Atomic Radius increases– Remember period number = PEL

• As you move down a group there are additional rings, therefore the valence electrons are further away from the nucleus, resulting in a larger radius

Atomic Radius Examples

1. Which sequence of elements is arranged in order of decreasing atomic radii?

a. Al, Si, Pb. Li, Na, Kc. Cl, Br, Id. N, C, B

2. What is the radius of Ca? 3. What is the radius of Sr? 4. Explain why Sr has a larger atomic radius than

Ca.

Ionic Radius (IR)• Radius that results from the loss or gain of

electrons

Metals (Left Side)

• Tend to lose 1 or more electrons when forming positive ions

• Radius will decrease• Ionic Radius < Atomic Radius

Ex: Na+ is smaller than Na

Non-metals (Right Side)

• Tend to gain 1 or more electrons when forming negative ions

• Radius will increase• Ionic Radius > Atomic Radius

Ex: Cl- is larger than Cl

Ionic Radius Examples1. Which ion has the largest radius?

a. Na+ b. Mg2+ c. K+ d. Ca2+

2. Which of the following elements has an ionic radius smaller than its atomic radius?

a. Neon b. Nitrogen c. Sodium d. Sulfur

3. The Na+ ion has a smaller radius than the Ne atom, even though they both contain 10 electrons. Explain why this is so.

Ionization Energy• The energy required to remove the most

loosely bound electron from an atom• Low IE = greater tendency to lose

electrons and form positive ions• High IE = greater tendency to gain

electrons and form negative ions• Given on Reference Table S in kilojoules

per mole (kJ/mol)

IE within a Period (Left to Right)

• IE increases• Number of protons increases, resulting in

a stronger nuclear charge• The nucleus has a better hold on the

electrons, therefore more energy is required to remove an electron

IE within a Group (Top to Bottom)

• IE decreases• The principle energy levels increase, so

the valence electrons are further away• Protons cannot hold onto the valence

electrons as well, therefore less energy is required to remove an electron

Electronegativity (e-neg)

• Scale that measures the ability of an atom to attract electrons from another atom

• Reference Table S• Scale ranges from 0.0-4.0• Fluorine is the highest = 4.0• Difference between electronegativity

between two atoms can be used to determine the type of bond

Trend within a Period (Left to Right)

• Electronegativity Increases• More protons, resulting in a stronger

nucleus, therefore the nucleus is better able to attract electrons

Trend within a Group (Top to Bottom)

• Electronegativity decreases• The atom is larger, so the nucleus is

further away from the valence shell, therefore the nucleus is less able to attract electrons

IE/e-neg Examples

1. Which element will lose electrons the easiest? a. Na b. Cl c. K d. F

2. Which element would be most likely to gain electrons? a. Na b.Cl c. K d. F

ReactivityMetals • Most reactive = loses electrons the easiest

(low ionization energy) – Lower left corner of the table (Fr)

Nonmetals• Most reactive = gains electrons the easiest

(high electronegativity)– Upper right corner of the table (F, Cl, O) – not

group 18

Periodic Trends Questions

1. What is the ionization energy of K? 2. What is the ionization energy of Ca? 3. Explain why K has a lower ionization

energy than Ca. 4. According to the reference table, which

of the following elements has the smallest radius? a. Ni b. Co c. Ca d. K

Periodic Trends Questions5. An element with high ionization energy would most likely

be? a. A nonmetal with low electronegativityb. A nonmetal with high electronegativityc. A metal with low electronegativityd. A metal with high electronegativity

6. What happens to S when it becomes S2-?a. It loses two electrons and the radius increasesb. It loses two electrons and the radius decreasesc. It gains two electrons and the radius increasesd. It gains two electrons and the radius decreases

Periodic Trends Questions

7. Which of the following would have the largest radius? a. Na b. Na+1 c. Cl Cl-1

8. Which of the following has the greater ionization energy, Na or Na+? Explain your answer.

Groups 1 and 2• Properties: typical metallic characteristics• High reactivity (valence electrons are easily lost)

– Only occur in nature as compounds• Reactivity increases as you move down the

group• Group 1 elements are more reactive than group

2 elements• For metals – Low IE = high reactivity (electrons

are easily lost) – Fr is the most reactive• Na and Water• K and Water

Examples:

1. Which atom is the most reactive? a. Na b. Mg c. K d. Ca

2. Which group 15 element has the least metallic character?

a. N b. P c. As d. Sb

3. Explain why reactivity increases as you move down group 1.

Group 17 - Halogens• Typical Nonmetals• High electronegativity – F is the highest• High ionization energy• Are so reactive that they cannot exist as in the

monoatomic form• Exist in nature as diatomics (HOFBrINCl)

– F2 and Cl2 are gases– Br2 is a liquid– I2 and At2 are solids

• F is the most reactive nonmetal• For nonmetals – high electronegativity = high

reactivity

Group 18 – Noble Gases

• Exist as gases at STP• Exist as monatomic molecules (not

combined with anythingExample: He, Ne, Ar

• The outermost ring is complete, therefore they are VERY STABLE (unreactive)

Groups 3-12 (Transition Elements)

• Hard solids• High melting points (except mercury)• Multiple Positive Oxidation States

– They can react with electrons from both s and d sublevels

– Different numbers of electrons can be lost• Colored Ions

– Easily excited (since d and s sublevels are close)

Group Examples1. The presence of which ion usually produces a colored

solution? a. K+ b. F- c. Fe2+ d. S2-

2. Which solution would be colored?a. KNO3 b. Ca(NO3)2 c. Cu(NO3)2 d. Al(NO3)3

3. Which element at STP exists as monatomic molecules?

a. Ne b. N c. Cl d. O4. Which is a solid at STP?

a. F2 b. Cl2 c. Br2 d. I2

5. Which element in Period 4 is classified as an active nonmetal?

a. Ga b. Ge c. Br Kr

Group Examples6. Which noble gas would most likely form a

compound with fluorine? a. He b. Ne c. Ar d. Kr

7. Which element in Period 3 is the most reactive metal?

a. Na b. Mg c. N d. Cl

8. Which element in Group 15 has the most metallic character? a. Bi b. As c. P d. N

Group Examples

9. Why is hydrogen not considered to be a member of Group 1?

10.Why is hydrogen considered to be a member of Group 1?

11.Why is it unlikely for sodium to form the Na2+ ion?

Allotropes• Some nonmetals can exist in 2 or more forms

in the same phase• Allotropes have different physical and chemical

properties because their atoms are arranged differentlyExamples: – oxygen and ozone (O2 and O3) – Graphite and Diamond (carbon)

Diamond

Graphite