Periodic Table. Antoine Lavoisier (1700’s) Listed all known elements (33) at the time 4 groups:...
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Transcript of Periodic Table. Antoine Lavoisier (1700’s) Listed all known elements (33) at the time 4 groups:...
Periodic Table
Antoine Lavoisier (1700’s)
• Listed all known elements (33) at the time
• 4 groups: gases, metals, nonmetals, and earths
Dobereiner (early 1800’s)
• Dobereiner arranged the elements into triads (groups of three elements) based on similarities in properties.
John Newlands (1864)
• Arranged elements by increasing atomic mass (70)
• Noticed a repeating pattern of properties
• Created the law of octaves (repeating patterns at every eighth element)
Newlands
Lothar Meyer (1869)
• Identified and proved that there was a connection between atomic mass and the property of the element
• Arranged the elements by increasing atomic mass (added the new ones)
Dmitri Mendeleev (1869)• Proved a connection
between atomic mass and element properties
• Arranged elements by increasing atomic mass
• Predicted the existence and properties of elements yet to be discovered
Henry Moseley (1913)
• Discovered atomic number
• Arranged elements by increasing atomic number
• By doing this a pattern of properties was discovered and fixed previous problems
Periodic Law
• When elements are arranged by increasing atomic number, there is a periodic repetition of physical and chemical properties.
Modern Periodic Table
• Periods (rows)- contain a variety of elements ranging from metals to nonmetals to Noble gases. There are 7.
• Groups or Family (columns)- contain elements that share similar properties. There are 18.
Periods
Representative (Main) Elements
• Marked by “A” on most groups.
• Elements in the ‘s’ and ‘p’ block
• Wide range of characteristics
• This is Newlands octaves
Transition Elements (B)
• Consists of only metals.
• Found in the center of the period table.
• Elements of the “d” block
Metals, Nonmetals, and Metalloids
Metals
• Make up most of the periodic table
• Solid at room temperature (except Mercury)
• Good conductors of heat and electricity
• Ductile and malleable• Have Luster (shiny)
Nonmetals
• Gases or solids at room temperature (except Br, it is a liquid)
• Poor conductors of heat or electricity
• Brittle• Dull
Metalloids
• Combination of characteristics of both metals and nonmetals
• Silicon and Germanium both used in computer chips
Answer the following questions about iron:
1. In what period is iron found?
2. In what group is iron found?
3.How many protons does an atom of iron have?
4.Write the electron configuration for iron.
5.Write the dot notation for iron.
6. Is iron a representative or transition element?
7. Is iron a metal or nonmetal?
8.List three properties of iron.
9.Describe one way in which classification was used in biology.
Elements ColorHydrogen Aqua
Group 1: Alkali Metals Red
Group 2: Alkali Earth Metals Yellow
Groups 3 -12: Transition Elements Blue
Group 13: Boron Group Grey
Group 14: Carbon Group Peach
Group 15: Nitrogen Group Brown
Group 16: Oxygen Group Purple
Group 17: Halogens Orange
Group 18: Noble Gases Green
Rare Earth Elements: Lanthanide Series
Pink
Rare Earth Elements: Actinide Series Magenta
Color Code the Periodic Table
S- Block• Alkali Metals
– 1 valence e-. This makes them highly reactive
– Exist only as compounds
– Silvery white in color– Often bond with
halogens– Used in salts and
batteries– Forms ions with a 1+
charge.
• Alkaline Earth Metals– 2 valence e-. Makes
them highly reactive– Ca and Mg are
important components of living cells
– Silvery in color– Used to make laptop
casings– Forms ions with a 2+
charge.
S- Block
P- Block (Families 13-18)
• Boron Family (13)– 3 valence e-– Tends to give its
valence e- away– Most are metals– Not as reactive as
group one and two– Forms ions with a 3+
charge
• Carbon Family (14)– 4 valence e-– Can either give away
its valence electrons or take additional electrons
– Sn and Pb will form ions with 4+charges
P- Block
• Nitrogen Family (15)– 5 valence e-, but will
form 3- (it prefers to gain 3 e- rather than give away 5)
– N and P are reactive and found in many molecular compounds
• Oxygen Family (16)– 6 valence e-– Forms 2- ions. It
prefers to gain 2 e- rather than give away 6)
– O and S are reactive and found in many compounds
P- Block
• Halogens (17)– 7 valence e-– Form 1- ions (gains 1
e-)– Highly reactive
nonmetals– Will often bond with
metals to make salts
• Noble Gases (18)– 8 valence e-, full p
sublevel– Does not form ions– Inert gases
(unreactive)– They do not bond with
other elements because they do not need anymore e-
D- Block (3-12)
• All transition metals
• Most are hard metals
• All can exist as free elements in nature
• Will form a variety of charged ions due to the fact that the s and d sublevels are close in energy amounts
D- Block (3-12)
F- Block (Period 6 and 7)
• Lanthanide Series– Shiny metals– Highly reactive– Fits in period 6
• Actinide Series – Radioactive– The first 4 are
naturally occurring the rest are lab created
– Fits in period 7
F- Block
Identify each of the following elements described below:
1. Nonmetal of the second period and group 4A.
2. The noble gas in period 3.
3. This element has two more protons than phosphorus.
4. The only nonmetal in group 1A.
5. Metal in period 7 with two valence electrons.
6. The element whose electron configuration ends with 3p1.
7. The nonreactive element consisting of 4 energy levels.
8. The metalloid with three valence electrons.
9. The only noble gas that does not have 8 valence electrons.
10. The element in group 2A that has fewer energy levels than magnesium.
Periodic Trends
• Patterns in the periodic table that can be determined by comparing a period or a group
• Atomic Radius, Ionic Radius, Ionization energy, Electronegativity
Atomic Radius
• Size of the atom• Half the distance
between two identical nuclei
increases
incr
ease
s
Atomic Radius Increases:
• Top to Bottom within a group because as you move down a column, the number of energy levels is increasing.
• Right to Left because as you move left across a period the atomic number decreases which results in a larger atomic radius
Ionization Energy
• Energy needed to remove the outermost electron.
• Follows this trend because small atoms have a high ionization energy
increases
incr
ease
s
Formation of Ions
• A positive ion (called a cation) results when an atom loses electrons.
• Metals have low I.E. and therefore, form positive ions.
• Positive ions are smaller than the parent atom.• A negative ion (called an anion) results when an
atom gains electrons.• Nonmetals have high I.E. and therefore, gain
electrons.• Negative ions are larger than the parent atom.
Octet Rule
• Every Element wants 8 valence e-
Ionic Radius• Increases • The
formation of an ion
• When atoms lose e- they become smaller
• When atoms gain e- they become larger
incr
ease
s
Ionization Energy
• Energy required to remove an e- from a gaseous atom. (Pulling an e- off)
Electronegativity
• The ability of an atom to attract an electron while in a chemical bond