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Chapter 6. The Periodic Table and Periodic Law. 6.1 Development of the Modern Periodic Table. 1. Antoine Lavoisier (1743-1794) Compiled a list of all known elements 23 know elements at that time. 6.1 Development of the Modern Periodic Table. 2. John Newlands (1837 – 1898)

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Chapter 6

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## Chapter 6

The Periodic Table and

Periodic Law

### 6.1 Development of the Modern Periodic Table

• 1. Antoine Lavoisier (1743-1794)

• Compiled a list of all known elements

• 23 know elements at that time

### 6.1 Development of the Modern Periodic Table

• 2. John Newlands (1837 – 1898)

• Law of Octaves = every 8th element repeats a common set of properties

• Not widely accepted due to missing elements and his use of musical terminology

### 6.1 Development of the Modern Periodic Table

• 3. Dimitri Mendeleev (1834 – 1907)

• 1869 he published 1st periodic table by atomic mass and chemical properties

• Predicted the properties of missing elements: scandium, gallium, and germanium

### 6.1 Development of the Modern Periodic Table

• 4. Henry Moseley (1887 – 1915)

• 1913 equated number of protons with atomic number

• Reordered the P.T. by atomic # fising some of the elements that didn’t fit their spots based on properties

### 6.1 Development of the Modern Periodic Table

• 5. periodic law:

• The repeating pattern of chemical and physical properties when elements are arranged by atomic number

### 6.1 Development of the Modern Periodic Table

• 6. organization

• Groups/families:

• Columns of/on the P.T.

### 6.1 Development of the Modern Periodic Table

• Periods:

• Rows of/on the P.T.

### 6.1 Development of the Modern Periodic Table

• Main group/representative elements:

• Elements in groups where the number is followed with an “A”

• Have a wide range of chemical and physical properties

### TABLE:

• Metals:

• Location = to the left of the staircase except H

• Properties = shiny, mostly solids, good conductors of heat/electricity, and malleable/ductile

• Examples = copper (Cu), gold (Au), iron (Fe)

### TABLE:

• Nonmetals:

• Location = to the right of the staircase plus H

• Properties = brittle (when solid), mostly gases, poor conductors

• Examples =helium (He), oxygen (O), Iodine (I)

### TABLE:

• Semimetals/metalloids:

• Location = along the staircase except Al

• Properties = properties of both metals and nonmetals

• Examples = B, Si, Ge, As, Sb, Te, Po, and At

### TABLE:

• Alkali metals:

• Location = 1A (except H)

• Valence e- and charge = 1 ve- and +1

• Properties = highly reactive; soft, gray solids

### TABLE:

• Alkaline Earth metals:

• Location = 2A

• Valence e- and charge = 2 ve- and +2

• Properties = very reactive; soft, gray solids

### TABLE:

• Transition metals:

• Location = “B” groups

• Valence e- and charge = 2 ve- and +1,+2,+3 or +4

• Properties = shiny, good conductors, can be polyvalent (means can have more than 1 possible charge)

### TABLE:

• Halogens:

• Location = 7A

• Valence e- and charge = 7 ve- and -1

• Properties = highly reactive; can be solids, liquids or gases

### TABLE:

• Noble Gases:

• Location = 8A

• Valence e- and charge = 8 ve- and no ion formation

• Properties = extremely unreactive, gases, full outer energy level

### TABLE:

• Rare Earth metals:

• Location = Bottom double rows

• Valence e- and charge = 2 ve-

• Properties = often used as phosphors (elements that emit light when struck by electrons)

• Also know as the “Lanthanides” and “Actinides”

### 6.2 Classification of the Elements

• 1. Valence Electrons = electrons in the highest principle energy level

• Within a period: elements have the same # of energy levels as the period # where they are found

• Within a group: elements have the same # of ve-s as their group (representative elements) and all transition and rare earth elements have 2 ve-s

### 6.2 Classification of the Elements

• 2. s, p, d, f blocks

• s block: group 1A, 2A, hydrogen and helium

• p block: groups 3A-8A but not He

• d block: transition elements; all have 2 ve-s because d’s are 1 energy level behind

• f block: rare earth elements; all have 2 ve-s because f’s are 2 energy levels behind

### 6.3 Periodic Trends

• Periodicity:

• The repeating nature of the properties of the elements creating common groups (periodic law)

### 6.3 Periodic Trends

• DEFINITION: relative size; distance from the center of the atom to the edge of the e- shell

• PERIOD TREND:

• GROUP TREND:

### 6.3 Periodic Trends

• DEFINITION: relative size; distance from the center of the ion to the edge of the e- shell

• PERIOD TREND:

• GROUP TREND:

### 6.3 Periodic Trends

• Ionization Energy

• DEFINITION: the energy required to remove an electron from a gaseous atom

• PERIOD TREND:

• GROUP TREND:

### 6.3 Periodic Trends

• Electronegativity

• DEFINITION: the ability of an atom to attract electrons while in a chemical bond

• PERIOD TREND:

• GROUP TREND:

### 6.3 Periodic Trends

• Lower Left Large (atomic/ionic radius)

• Lower Left Low (ionization E./electroneg.)