Chapter 5 – The Periodic Law

I. History of the Periodic Table• A. 60 elements were known by 1860

• 1. Chemists had to memorize the properties of all these

• 2. No method existed for determining mass or number of atoms of an element in a compound at this time

• 3. Different chemists used

• different masses and got

• different compositions

• for the same compound

• B. In September of 1860, the First International Congress of Chemists was held in Germany

• 1. It settled the issue of atomic mass

• 2. Cannizzaro presented a method for accurately measuring the relative mass of atoms

• 3. A search was begun to

• find relationships between

• atomic mass and other

• properties of elementsCannizzaroreaction

• C. Mendeleev and Chemical Periodicity• 1. Demitri Mendeleev put the name or each

element and all the information he knew about it on a card

• a. Atomic mass• b. Physical properties• c. Chemical properties• 2. He arranged the cards according to the

properties and looked for trends or patterns• a. When arranged by atomic mass,

properties repeated at regular intervals• b. This kind of pattern is called periodic• 1) days of the week 2) months of the

year

• 3. Mendeleev crated a table to show these patterns, and it became the 1st Periodic Table of Elements

• a. When arranged by repeating properties, occasionally the atomic mass did not follow the correct increasing sequence

• b. Also, he left blank spaces, where he predicted new elements would

• be discovered, to fill in missing• properties or gaps in the table• 1) scandium• 2) gallium• 3) germanium

• D. Moseley and the Periodic Law

• 1. Henry Moseley examined the spectra of 38 different metals

• a. He discovered a new pattern

• b. The elements fit into more exact patterns when arranged by their atomic numbers (number of protons)

• 2. Developed the definition for atomic #

• a.Followed Mendeleev’s principle for repeating properties in stating the periodic law

• b.Periodic Law – the physical and chemical properties of the elements are periodic functions of their atomic numbers

• E. The Modern Periodic Table

• 1. The Periodic Table is an arrangement of the elements in order of their atomic numbers, so that elements with similar properties fall in the same column or group

• 2. Noble Gases

• a. discovered later than most other elements because of their inactivity

• b. placed in a new group after Group VII (17)

Helium Neon Argon Krypton Xenon Radon

• 3. The Lanthanides (lanthanoids)• a. 14 elements from atomic number 58 to

atomic number 71• b. Chemical and physical properties are so

similar that it is difficult to separate them into individual elements

• 4. The Actinides (actinoids)• a. 14 elements from atomic number 90 to

atomic number 103• b. Are pulled to the bottom of the periodic

table, along with the lanthanides, in order to save space

• c. Lanthanides and actinides are also known as the rare earth elements

• 5. Periodicity

• a. All elements in any particular column of group have similar chemical and physical properties

• b. The similarities are due to the number of electrons in the outer level only of the atom

• c. Outer level electrons determine how atoms behave chemically

II. Electron Configuration & the Periodic Table

• A. The Noble Gases rarely undergo a chemical reaction

• 1. This is the result of their electron configuration

• 2. Their highest energy levels contain stable octets (except for Helium which has only one energy level that is completely filled with two electrons

• 3. An atom’s highest energy level determines its chemical properties

• B. Periods and Blocks of the Periodic Table

• 1. Elements are arranged vertically on the periodic table in groups

• 2. Elements are arranged horizontally on the periodic table in periods

• 3. There are 8

• groups and 7

• periods of

• elements on

• the periodic

• table

• 4. The group number tells how many electrons are in the highest (or outer) energy level

• 5. The period number tells how many energy levels the atom has

• a. The 1st period has only one sublevel (s), which can hold only 2 electrons; therefore, this period can have only 2 elements

• b. The 2nd period has two sublevels (s & p, and p has 3 orientations); therefore, this period can hold 8 total electrons and 8 elements

• c. Period 3 is the same as Period 2, with 8 elements and 8 electrons

• d. Periods 4 & 5 have 2 electrons in the s orbital, 10 electrons in the d orbitals, and 6 electrons in the p orbitals, for a total of 18 elements in these periods

• e. Periods 6 & 7 have 2 electrons in the s orbital, 14 electrons in the f orbitals, 10 electrons in the d orbitals, and 6 electrons in the p orbitals for a total of 32 elements in these periods

• C. The s block elements (Groups I & II)

• 1. The elements in Group I are called alkali metals

• 2. They all have one electron in their outer level, as well as other similar properties

• a. silvery

• b. soft

• c. extremely reactive (never free in nature)

• d. must be stored in kerosene or oil

lithium Sodium or potassiumrubidium cesium

francium

• 3. The elements in Group II are called the alkaline earth metals

• 4. They all have 2 electrons in their outer level and have similar properties

• a. harder, stronger, & denser than alkali metals

• b. have higher melting points than alkali metals

• c. less reactive than Group I, but still never found free in nature

beryllium

Ca strontiumbarium

strontium

radium

• D. Hydrogen & Helium

• 1. Hydrogen has 1 electron and Helium has 2 electrons; therefore, each of them has only one orbital, so they are placed in the s block

• 2. Neither of them, however, are metals like the other elements in Groups I and II

• E. d block elements (between Groups II – III)

• 1. These are also known as the transition elements

• 2. They are less reactive than other metals

• F. p block elements (Groups III-VIII)

• 1. These are called the main group elements along with the s block

• 2. These are mostly nonmetals & metalloids

• 3. The metalloids are boron, silicon, germanium, arsenic, antimony, tellurium, polonium, and astatine

• 4. The elements in Group VII are called the halogens

fluorine

chlorine

bromineiodine

astatine

• G. f block elements (lanthanides & actinides)

• 1. Also known as the rare earth elements

• 2. Lanthanides should fit in period 6 between 56 & 72

• 3. Actinides should fit in period 7 between 88 & 104

lanthanum

promethium erbium

uranium

actinium

plutonium einsteinium

nobelium

III. Electron Configuration & Periodic Properties

• A. Atomic Radii

• 1. Ideally, the size of an atom should be from its nucleus to the outer edge of its highest orbital, but the outer edge of the orbital is not exact due to the electron’s fluctuating levels of energy

• 2. Atomic radius – one half

• the distance between the

• nuclei of identical atoms that

• are bonded together

• B. Period Trends

• 1. There is a gradual decrease in atomic radii across the 2nd period from lithium to neon

• 2. The trend toward smaller atoms across a period (from left to right) is caused by the increasing positive charge of the nucleus

• 3. Electrons (-) are pulled closer to the nucleus(+) as the nucleus becomes more positive

• C. Group Trends

• 1. The radii of the atoms of main group elements increase as you go from the top to the bottom of a group

• 2. This is due to the addition of one energy level per atom as you progress down the group

• 3. This does not always

• hold true, however, for

• those atoms preceded

• by the d block elements

• (Ex.- Al & Ga)

• D. Ionization Energy

• 1. An electron can be removed from an atom if enough energy is supplied

• atom A + energy → A+ (positive ion A) + e-

• 2. An ion is an atom or group of bonded atoms that have a positive or negative charge

• 3. The process that results in the formation of an ion or ions is called ionization

• 4. The energy required

• to remove one electron

• from a neutral atom is

• called ionization energy+ -

• E. Period Trends

• 1. Ionization energy for the main group elements increases as you move from left to right across the periodic table

• 2. This is due to a greater number of protons having a greater attraction for the electrons

• F. Group Trends

• 1. For main group elements, ionization energy usually decreases as you move down the group

• 2. This is due to the outer electrons being further and further from the attraction of the nucleus as each new energy level is added

• G. Removing Electrons from Positive Ions• 1. The amount of energy required to remove

the first electron from an atom is referred to as the first ionization energy (IE1)

• 2. IE1 will be less than IE2 or IE3, etc. because as each electron is removed, the attractive force of the nucleus becomes greater and greater in relationship to the electrons

• 3. Therefore, IE2 is • greater than IE1 and• IE3 is greater than • IE2 and so on

• H. Electron Affinity• 1. The energy change that occurs when an

electron is acquired by a neutral atom is called electron affinity

• 2. Most atoms release energy when they acquire an electron (A + e- → A- + energy)

• 3. Some atoms must be “forced” to take on another electron, and energy

• must be added in order to • make this happen • (A + e- + energy → A-) • This type of ion will be • unstable

• I. Trends

• 1. In general electrons are gained more easily as you go from left to right on the table (across a period)

• 2. An exception is the noble gases, which require huge amounts of energy to add electrons, since they already have a full octet

• 3. Usually it is more difficult to add electrons as you move down a group, because as you move down the group the electrons are getting further from the attractive force of the nucleus

• 4. It is also more difficult to add a second electron to ion that is already negative

• J. Ionic Radii

• 1. A positive ion is a cation, and its radius will be smaller than its atom’s radius

• 2. A negative ion is an anion, and its radius will be larger than its atom’s radius

• 3. Atomic radii tend

• to decrease across

• a period, and

• increase down a

• group

• K. The electrons available to be gained, lost, or shared are called valence electrons

• L. A measure of the ability of an atom to attract electrons is called electronegativity

• 1. Electronegativity tends to increase across a period

• 2. Electronegativity tends to decrease down a group

LAB – Flame Test• 1. How can a flame test be used as a confirmatory test for an • unknown metal?• 2. What is one purpose for which these metal ions are used?• 3. What kind of ions color the flame when glass is melted?• 4. Why do ions produce colors when placed in a flame?• 5. What condition is the atom in before it is placed in the • flame?• 6. After placing the atom in the flame, what condition is it then • in?• 7. What condition is the atom in when a color is produced in• the flame?• 8. When atoms produce colors in a flame, what series in the • electromagnetic spectrum is this called?• 9. How was the color of the flame affected when viewed • through the cobalt glass?• 10. Imagine and describe a scenario in forensics in which a • flame test could be used for identification of something.