chapter 5 periodic law. section 1 history of the periodic table
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Chapter 5
Periodic Law
SECTION 1 History of the Periodic Table
Mendeleev and Chemical Periodicity
When the Russian chemist Dmitri Mendeleev heard about the new atomic masses he decided to include the new values in a chemistry textbook that he was writing
Mendeleev hoped to organize the elements according to their properties
He placed the name of each known element on a card, together with the atomic mass of the element and a list of its observed physical and chemical properties
He then arranged the cards according to various properties and looked for trends or patterns
Mendeleev noticed that when the elements were arranged in order of increasing atomic mass, certain similarities in their chemical properties appeared at regular intervals
Such a repeating pattern is referred to as periodic
Mendeleev’s procedure left several empty spaces in his periodic table
In 1871, he predicted the existence and properties of the elements that would fill three of the spaces
By 1886, all three elements had been discovered
scandium, Sc, gallium, Ga, and germanium, Ge Their properties are very similar to those
predicted by Mendeleev
Success of Mendeleev’s predictions persuaded most chemists to accept his periodic table and earned him credit as the discoverer of the periodic law
Two questions remained (1) Why could most of the elements be arranged
in the order of increasing atomic mass but a few could not?
(2) What was the reason for chemical periodicity?
Moseley and the Periodic Law
In 1911 English scientist Henry Moseley examined the spectra of 38 different metals
Moseley discovered a previously unknown pattern
The elements in the periodic table fit into patterns better when they were arranged in increasing order according to nuclear charge, or the number of protons in the nucleus
Moseley’s work led to both the modern definition of atomic number and the recognition that atomic number, not atomic mass, is the basis for the organization of the periodic table
SECTION 2
Electron Configuration and the Periodic Table
The Modern Periodic Table
“Periodic” - Repeating patterns
Listed in order of increasing number of protons (atomic #)
Properties of elements repeat Periodic Law-Periodic Law- “the physical and chemical properties of
the elements are periodic functions of their atomic numbers.”
Periods and Blocks of the Periodic Table
Elements are arranged vertically in the periodic table in groups that share similar chemical properties
They are also organized horizontally in rows, or periods
The length of each period is determined by the number of electrons that can occupy the sublevels being filled in that period
main group elements
Metals Most solids (Hg is liquid) Luster – shiny. Ductile – drawn into thin wires. Malleable – hammered into
sheets. Conductors of heat and
electricity. Include transition metals –
“bridge” between elements on left & right of table
Non-Metals
Properties are generally opposite of metals
Poor conductors of heat and electricity
Low boiling points Many are gases at room
temperature Solid, non-metals are brittle (break
easily) Chemical properties vary
Metalloids
stair-step pattern Have properties similar to
metals and non-metals Ability to conduct heat
and electricity varies with temp Better than non-metals but
not metals semiconductors
Group 1 – The Alkali Metals
The elements of Group 1 of the periodic table (lithium, sodium, potassium, rubidium, cesium, and francium) are known as the alkali metals
In their pure state, all of the alkali metals have a silvery appearance and are soft enough to cut with a knife
Because they are so reactive, alkali metals are not found in nature as free elements
They combine strongly with most nonmetals
And they react strongly with water to produce hydrogen gas and aqueous solutions of substances known as alkalis
Because of their extreme reactivity with air or moisture, alkali metals are usually stored in kerosene or oil
Group 2 – The Alkaline-Earth Metals The elements of
Group 2 of the periodic table are called the alkaline-earth metals
Atoms of alkaline-earth metals contain a pair of electrons in their outermost s sublevel
Group 2 metals are harder, denser, and stronger than the alkali metals
They also have higher melting points Less reactive than the alkali metals, but also
too reactive to be found in nature as free elements
Transition Elements
Good conductors of electricity and have a high luster
They are typically less reactive than the alkali metals and the alkaline-earth metals
Some are so unreactive that they do not easily form compounds, existing in nature as free elements
TungstenMercury
Vanadium
Rare Earth ElementsLanthanide series (period 6)Actinide Series (period 7)
Some radioactive Separated from table to make easy to
read/print silver, silvery-white, or gray metals. Conduct electricity
uranium
Halogen Family (“salt-former”)
-7 Valence Electrons-most active nonmetals-never found pure in nature-react with alkali metals
easily (forms salts)-F most active halogen
Halogens cont…
F compounds in toothpasteCl kills bacteriaI keeps thyroid gland
working properly
bromine
The Noble Gases (Inert Gases)
- non-reactive- outermost e-
shell is full (8 VE)
- In “neon” lights-in earth’s
atmosphere (less than 1%)
Neon
Section 5.3Electron Configuration and
Periodic Properties
Periodic Trends
In periodic table, there is a DECREASE in atomic radii across the periods from left to right
Caused by increasing positive charge of nucleus (more protons = more positive charge)
Group Trends
Radii of elements decrease as you go UP a group
Electrons occupy sublevels in consecutively higher main energy levels (located further away from nucleus)
In general, the atomic radii of the main-groups elements decrease from the bottom to the top of a group
2. Ionization Energy
Electrons can be removed from an atom if enough energy is supplied
Using A as a symbol for an atom of ANY element, the process can be expressed as follows:
A + energy A+ + e-
A+ represents an ion of element A with single positive charge (a 1+ ion)
Ion an atom or group of bonded atoms that have a positive or negative charge
Ionization any process that results in the formation of an ion
Period Trends
In general, ionization energies of the main-group elements INCREASE across each period
Caused by increasing nuclear charge
Higher charge more strongly attracts electrons in same energy level
Group Trends
Ionization energies generally INCREASE going UP a group
Electrons going down in group are in higher energy levels, so further away from the nucleus
Removed more easily Also more electrons between outermost
electrons and the nucleus (shields them from attraction to positive nucleus)
What are Valence electrons?
outermost e-’s Responsible for chem props Elements in same group… same # of VE ALL atoms want full outer energy level (usually 8 VE) To get full outer energy level, some elements:
lose e- (metals) gain e- (non-metals) share electrons (some non-metals & metalloids)
Main-group elements – valence e- are in outermost s and p sublevels
Inner e- held too tightly by nucleus to be involved in compound formation
6. Electronegativity
Valence e- hold atoms together in compounds
In many compounds, negative charge centered around one atom more than another
Uneven distribution of charge has effect on compound’s properties
Useful to have measurement of how strongly one atom attracts e- of another atom
Electronegativity measure of the ability of an atom in a chemical compound to attract electrons
Most e-neg element (fluorine) – randomly assigned value of 4
Other values calculated in relation to F
Period Trends
e-negs tend to INCREASE across each period
There are exceptions (of course)
Alkali and alkaline-earth metals are least e-neg
In compounds, their atoms have low attraction for e-
Group Trends
Electronegativities tend to INCREASE going UP a group or stay the same
At higher energy levels electrons being added are further away from the nucleus
Therefore, less attraction to the nucleus Also more electrons between outermost
electrons and the nucleus (shields them from attraction to positive nucleus)
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