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# The periodic law - PowerPoint PPT Presentation

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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### The periodic law

Chapter 5

• To organize the elements

• To show trends

• A repeating pattern

• 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

• 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 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.

• 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.

• 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.

• 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

• An arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column or group.

• 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.

• Are governed by the electron configuration of an atom’s highest energy level

• Determined by the number of electrons than can occupy the sublevels being filled in that period.

• Table 5-1

• Table with f-block in place

1st period

• 1s sublevel being filled

• 1s can hold 2 electrons, so there are 2 elements in the 1st period.

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

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

6th and 7th periods

• Add f-block, which holds 14 electrons

• Fill 6s, 5d, 4f, 6p

• Need 32 electrons

• 32 elements in each period

• Shows blocks

• 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.

• 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)

• Metals in group 1

• Have silvery appearance

• Soft enough to cut with a knife

• React violently with nonmetals

• Melting point decreases as you go down the table

• Group 2

• Harder, denser, and stronger than alkali metals

• Higher melting points than alkalis

• Less reactive

• 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

• 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

• 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

• End in p1 to p6.

• Coefficients are the same as the period

• ns2np1

• Always have a full s-sublevel

• 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

• Lanthanides and actinides

• Endings are f1 to f14

• Coefficients are two less than the period

• 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?

• 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.

• 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.

• 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

• First ionization energy (IE1) – the energy required to remove the most loosely held electron.

• Measured in kJ/mol

• 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

• Energy required to remove other electrons from positive ions.

• IE2, IE3, etc

• Get higher as you remove more electrons

• Less shielding

• 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.

• 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

• 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

• Not as regular

• Usually, electrons add with greater difficulty as we move down

• 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.

• Cation – a positive ion

• Anion – a negative ion

• 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

• Ionic radius increases as you go down the table

• 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

• 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.

• Properties vary less and with less regularity than others

• d-block

• Usual patterns

• f-block (unusual)

• Increase across periods

• Decrease down groups

• Ionization energy

• Increase across periods

• d-block increases down groups (unusual)

• f-block decreases down groups

• 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?

• 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.