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The periodic law. Chapter 5. Why do we need a table?. To organize the elements To show trends. Periodic. A repeating pattern. Mendeleev’s table. 1869 – Dmitri Mendeleev – Russian Arranged the elements in order of increasing mass and noticed that chemical properties were periodic

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why do we need a table
Why do we need a table?
  • To organize the elements
  • To show trends
  • A repeating pattern
mendeleev s table
Mendeleev’s table
  • 1869 – Dmitri Mendeleev – Russian
  • Arranged the elements in order of increasing mass and noticed that chemical properties were periodic
  • Put the elements into groups according to properties
mendeleev vs meyer
Mendeleev vs. Meyer
  • 1860s Mendeleev and German Lothar Meyer each made an eight column table.
  • Mendeleev left some blanks in his table in order for all the columns to have similar properties – he predicted elements that hadn’t been discovered yet.
why similar properties
Why similar properties?
  • Why did they group according to properties and mass and not atomic number or number of outer level electrons?
  • Mendeleev’s blank spots and his ability to predict future elements helped his table win acceptance.
mendeleev s table1
Mendeleev’s table
  • Elements arranged in order of increasing mass.
  • Properties are repeated in an orderly, periodic, fashion.
  • Mendeleev’s periodic law – the properties of the elements are a periodic function of their masses.
mass mistakes
Mass mistakes?
  • In order for Mendeleev to arrange his elements by properties, he had to put tellurium and iodine in the wrong order.
  • He explained this by assuming that their masses hadn’t been measured very accurately.
more mass mistakes
More mass mistakes?
  • Nickel and cobalt
  • Argon and potassium
  • Better mass measurements just confirmed the discrepancy
  • 1913 – Henry Moseley
  • X-ray experiments revealed the atomic number was the number of protons
  • Modern periodic law – the properties of the elements are a periodic function of their atomic numbers
modern periodic table
Modern periodic table
  • An arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column or group.
noble gases
Noble gases
  • Not discovered on Earth until 1894 - 1900.
  • Group 18 was added to the table
  • Hard to separate
  • All have similar properties
  • Added to the table in the early 1900s
  • Discovered later
  • Also all have similar properties
  • Elements in the same group (column) have similar properties.
chemical properties of an element
Chemical properties of an element
  • Are governed by the electron configuration of an atom’s highest energy level
period length
Period length
  • Determined by the number of electrons than can occupy the sublevels being filled in that period.
  • Table 5-1
full periodic table
Full periodic table
  • Table with f-block in place
1 st period
1st period
  • 1s sublevel being filled
  • 1s can hold 2 electrons, so there are 2 elements in the 1st period.
2 nd and 3 rd periods
2nd and 3rd periods
  • 2s and 2p or 3s and 3p being filled
  • s and p sublevels can hold 8 total, so there are eight elements in these periods
4 th and 5 th periods
4th and 5th periods
  • Add d sublevels, which can hold 10 electrons
  • Need to fill 4s, 3d, and 4p – 18 electrons
  • 18 elements in each period
6 th and 7 th periods
6th and 7th periods
  • Add f-block, which holds 14 electrons
  • Fill 6s, 5d, 4f, 6p
  • Need 32 electrons
  • 32 elements in each period
figure 5 5
Figure 5-5
  • Shows blocks
electron configurations
Electron configurations
  • Elements in columns 1, 2, and 13-18 have their last electron added in an s or p orbital.
  • Elements in columns 3-12 have their last electron added in a d level.
the s block elements groups 1 and 2
The s-block elements: Groups 1 and 2
  • Chemically reactive metals
  • Group 1
    • Have 1 electron in outer s orbital
      • Coefficient represents period
        • Row 2: 2s1, Row 3: 3s1, etc. (ns1)
  • Group 2
    • Have 2 electrons in outer s orbital
      • Coefficient represents period
        • Row 2: 2s2, Row 3: 3s2, etc. (ns2)
alkali metals
Alkali metals
  • Metals in group 1
  • Have silvery appearance
  • Soft enough to cut with a knife
  • Not found alone in nature
  • React violently with nonmetals
  • Melting point decreases as you go down the table
alkaline earth metals
Alkaline-earth metals
  • Group 2
  • Harder, denser, and stronger than alkali metals
  • Higher melting points than alkalis
  • Less reactive
  • Not found alone in nature
hydrogen and helium
Hydrogen and helium
  • Hydrogen
    • Located above group 1 because of its electron configuration
    • Not really in group 1, because its properties don’t match
  • Helium
    • Has an electron configuration like group 2 elements
    • In group 18 because it is unreactive
  • Page 133
  • Sample problem 5-1 and practice problems
  • Without looking at the periodic table, give the group, period, and block in which the element with the electron configuration [Rn] 7s1 is located.
    • Group 1, 7th period, s block
  • Without looking at the periodic table, give the group, period, and block in which the element with the electron configuration [He]2s2 is located.
    • Group 2, second period, s block
d block elements groups 3 12
d-block elements: Groups 3-12
  • End in d1 to d10.
    • Coefficients are one less than the period
      • Example: Fe is in the 6th column of transition elements in the 4th period, ends in 3d6
transition elements
Transition elements
  • Groups 3-12
  • Typical metallic properties
    • Good conductors
    • High luster
  • Less reactive than alkalis and alkaline-earths
  • Some are unreactive enough to appear in nature
p block elements groups 13 18
p-block elements: groups 13-18
  • End in p1 to p6.
    • Coefficients are the same as the period
      • ns2np1
      • Always have a full s-sublevel
p block elements
p-block elements
  • Properties vary greatly
  • Includes all nonmetals except hydrogen and helium
    • Solids, liquids and gases
  • Includes all the metalloids
    • Between metals and nonmetals
    • Brittle solids
    • Semiconductors – can conduct under certain conditions
  • Includes some metals
    • Less reactive than alkalis and alkaline-earths
  • Group 17
  • Most reactive nonmetals
  • Form compounds called salts
f block elements
f-block elements
  • Lanthanides and actinides
    • Endings are f1 to f14
    • Coefficients are two less than the period
  • All actinides are radioactive
  • Those after neptunium are synthetic
  • Sample problems and practice problems on pages 136, 138, and 139
  • With your group first, then join with another group.
  • Do you have any questions?
atomic radius
Atomic radius
  • Ideally, the distance from the center of the atom to the edge of it’s orbital.
    • But, atoms are “fuzzy”, not clearly defined.
  • Defined as one-half the distance between the nuclei of identical atoms that are bonded together.
period trends see figure 5 13
Period trends – see figure 5-13
  • As we move from left to right across the table, we gain protons.
  • There is a greater positive charge on the nucleus.
  • This greater charge pulls harder on the outer electrons, pulling them in closer.
  • The atom gets smaller.
group trends
Group trends
  • As we move down the table, the principle quantum number increases.
  • When the principle quantum number increases, the electron cloud gets bigger.
  • The size of the atoms gets bigger.
  • Which of the elements Li, Rb, K, and Na has the smallest atomic radius? Why?
    • Li, it is highest on the table
  • Which of the elements Zr, Rb, Mo, and Ru has the largest atomic radius? Why?
    • Rb, it is farthest to the left on the table
  • An atom or group of bonded atoms that has a positive or negative charge
  • Any process that makes ions
ionization energy ie
Ionization energy (IE)
  • First ionization energy (IE1) – the energy required to remove the most loosely held electron.
  • Measured in kJ/mol
ionization energy see figure 5 15
Ionization energy – see figure 5-15
  • Experimentally determined.
    • From isolated atoms in the gas phase
  • Tends to increase as you move across a row from left to right
    • Why group 1 is most reactive
    • Caused by higher charge
  • Tends to decrease as you move down a column
    • Electrons are farther from nucleus
    • Shielding from inner electrons
other ionization energies see table 5 3
Other Ionization Energies – see Table 5-3
  • Energy required to remove other electrons from positive ions.
  • IE2, IE3, etc
  • Get higher as you remove more electrons
    • Less shielding
noble gases1
Noble Gases
  • Have High ionization energies
  • When a positive ion of another element reaches a noble gas configuration, its ionization energy goes up.
    • Example: When K loses one electron, it has Ar’s electron configuration
    • This makes it stable
    • Its IE2 is much higher than its IE1
  • State in words the general trends in ionization energies down a group and across a period of the periodic table.
electron affinity
Electron affinity
  • The energy change that occurs when an electron is gained by a neutral atom
    • Most atoms release energy
      • Represented by a negative number
    • Some atoms gain energy
      • Represented by a positive number
      • These ions will be unstable
  • KJ/mol
period trends see figure 5 17
Period trends – see figure 5-17
  • Group 17 has most negative electron affinity.
  • Tends to get more negative (release more energy) as we move to the right
  • Exceptions:
    • groups with full or half-full sublevels are more stable
group trends1
Group trends
  • Not as regular
  • Usually, electrons add with greater difficulty as we move down
adding additional electrons
Adding additional electrons
  • Second electron affinities are all positive because it is more difficult to add electrons to a negative ion.
  • If a noble gas configuration has been reached, it is even more difficult.
  • State in words the general trends in electron affinities down a group and across a period of the periodic table.
ionic radii
Ionic Radii
  • Cation – a positive ion
    • Ionic radius smaller than atomic radius
  • Anion – a negative ion
    • Ionic radius is larger
period trends see figure 5 19
Period Trends – see figure 5-19
  • Metals form cations by losing electrons
    • Ions are smaller
    • Radius decreases as we move across
  • Nonmetals form anions by gaining electrons
    • Ions are larger
    • Radius decreases as we move across
group trends2
Group trends
  • Ionic radius increases as you go down the table
valence electrons
Valence electrons
  • Available to be lost, gained or shared in the formation of chemical compounds
  • In highest energy levels
  • For s-block, the group number is the same as the number of valence electrons
  • For the p-block, the group number is 10 more than the number of valence electrons
  • The measure of the ability of an atom in a compound to attract electrons
    • The atom with higher electronegativity pulls the electrons closer to itself
electronegativity trends figure 5 20
Electronegativity trends (figure 5-20)
  • Increases left to right across the rows
  • Decreases down the columns
  • Explain why elements with high (more negative) electron affinities are also the most electronegative.
d and f block elements
d- and f-block elements
  • Properties vary less and with less regularity than others
  • Atomic radii
    • d-block
      • Usual patterns
    • f-block (unusual)
      • Increase across periods
      • Decrease down groups
d and f block elements1
d- and f-block elements
  • Ionization energy
    • Increase across periods
    • d-block increases down groups (unusual)
    • f-block decreases down groups
  • Ionic radii
    • Cations have smaller radii
  • Electronegativity
    • d-block follows normal rules
    • f-block all have similar electronegativities
  • Among the main-group elements, what is the relationship between group number and the number of valence electrons?
  • In general, how do the periodic properties of the d-block elements compare with those of the main-group elements?
prelab notes
Prelab notes
  • Precipitate – solid that falls out of a solution
  • The formation of a precipitate indicates there has been a chemical change.
  • This means that there were ions present that were free to react.