Chapter 5: The Periodic Table - Bartlett High School
Chapter 5 Periodic Law History of the Periodic Table SECTION 1 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 Mendeleevs 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 Mendeleevs 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 Moseleys 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 Electron Configuration and the Periodic Table
SECTION 2 The Modern Periodic Table Periodic - Repeating patterns
Listed in order of increasing number of protons (atomic #) Properties of elements repeat 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 alkalineearth metals Atoms of alkalineearth 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 Mercury Tungsten
Vanadium uranium Rare Earth Elements Lanthanide 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 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 toothpaste Cl kills bacteria I keeps thyroid gland working properly
bromine The Noble Gases (Inert Gases) - non-reactive - outermost eshell is full (8 VE) - In neon lights -in earths atmosphere (less than 1%)
Neon Section 5.3 Electron 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 maingroups 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 compounds 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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