Periodic properties of the elements
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Periodic Properties of the Elements. Chapter 7. 7.1 Development of the Periodic Table. 1 st developed by Dmitri Mendeleev (Russia) & Lothar Meyer (Germany) on the basis of the similarity in chemical and physical properties Mendeleev … started by organizing elements by increasing mass.

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Periodic properties of the elements

Periodic Properties of the Elements

Chapter 7


7 1 development of the periodic table

7.1 Development of the Periodic Table

  • 1st developed by Dmitri Mendeleev (Russia) & Lothar Meyer (Germany) on the basis of the similarity in chemical and physical properties

  • Mendeleev …

    • started by organizing elements by increasing mass.

    • Recognized a repetition of pattern.

    • Placed elements by same column  same properties

    • Predicted correctly about the existence of new elements

  • Henry Moseley

    • established that each element has a unique atomic number, which added more order to the periodic table

    • Identified the atomic number with the # of protons in the nucleus of the atom & the # of electrons in the atom.


7 2 electron shells and the sizes of atoms

7.2 Electron Shells and the Sizes of Atoms

  • Atoms aren’t hard spheres with well-defined shells of electrons

  • The edges of atoms are a bit “fuzzy”

  • The quantum mechanical model of the atom supports the notion of electron shells: certain distances from the nucleus at which there is a higher likelihood of finding an electron


Atomic sizes

Atomic Sizes

  • The size of an atom can be gauged by its bonding atomic radius, based on measurements of the distances separating atoms in their chemical combinations with other atoms

  • Measure the atomic radius from the center of the nucleus to the outermost electron.

  • Atom size increases going down a group.

  • Atomic size decreases going left to right across the period.


7 3 ionization energy

7.3 Ionization Energy

  • Ionization energy – the minimum energy required to remove an electron from the ground state of the isolated gaseous atom or ion

  • 1st ionization energy (I1) – The energy needed to remove the first electron from a neutral atom, forming a cation

  • 2nd ionization energy (I2) – the energy needed to remove the second electron

  • The greater the ionization energy, the harder it is to remove an electron


7 3 ionization energy1

7.3 Ionization Energy

  • HIGH ionization energy means the atom hold onto the electron tightly and a lot of energy is need to pull it off

  • LOW ionization energy means the atom holds onto the electron loosely so breaking it apart doesn’t require much energy


7 3 ionization energy2

7.3 Ionization Energy

Periodic Trends in Ionization Energies

  • Ionization energy decreases as you move down a group.

  • Ionization energy increases as you move from left to right on the periodic table.

  • Representative elements show a larger range of values of I1 than do the transition metal elements


Ionization energy 3 d

Ionization Energy 3-D


7 4 electron affinities

7.4 Electron Affinities

  • Electron affinity – the energy change that occurs when an electron is added to as gaseous atom

  • A negative electron affinity = the anion is stable

  • A positive electron affinity = the anion is higher in energy than are the separated atom and electron. The anion is not stable and will not form


7 4 electron affinities1

7.4 Electron Affinities

  • If the electron affinity is negative, the atom releases energy.

  • Normally, non-metals have a more negative electron affinity than metals. The exception is the noble gases.


7 4 electron affinities2

7.4 Electron Affinities

  • Election affinities become more negative as we proceed from left to right

  • Halogens have the most negative electron affinities

  • The electron affinities of the noble gases are all positive since the added electron would have to occupy a new, higher-energy subshell

  • Electron affinity doesn’t change greatly as we move down a group. Electron affinity should become more positive (less energy released).


7 5 metals nonmetals and metalloids

7.5 Metals, Nonmetals, and Metalloids


Periodic properties of the elements

Pg. 239 --Table 7.3 Characteristic Properties of Metals and Nonmetals


7 5 metals nonmetals and metalloids1

7.5 Metals, Nonmetals, and Metalloids

  • Metallic Character - The tendency of an element to exhibit properties of metals

  • Metallic character generally increases going down a column and decreases going from left to right across a period


Metals

Metals

  • Metals conduct heat & electricity

  • They are malleable & ductile

  • Solids at room temp. except mercury(Hg) (it’s liquid)

  • Melt at very high temps

  • Have low ionization energies & are consequently oxidized (lose electrons) when they undergo chemical reaction.

  • Many transition metals have the ability to form more than one positive ion.


Chemical reactions with metals

Chemical Reactions with Metals

  • metal oxide + water  metal hydroxide

    • Most metal oxides are known as basic oxides

    • Ex: Na2O (s) + H2O(l)  2NaOH (aq)

  • metal oxide + acid  salt + water

    • Ex: MgO (s) + 2HCl (aq)  MgCl2 (aq) + H20 (l)


Nonmetals

Nonmetals

  • Not lustrous & generally are poor conductors of heat and electricity

  • Non-metals commonly gain enough electrons to fill their outer p sub-shell completely, giving a noble gas electron configuration.

  • Molecular substances - Compounds composed entirely of nonmetals

    • Ex: oxides, halides, and hydrides

  • Melting points are gen. lower than those of metals


Chemical reactions with nonmetals

Chemical Reactions with Nonmetals

  • Nonmetal oxide + water → acid

    • Most nonmetal oxides are acidic oxides

    • CO2 (g) + H2O (l)  H2CO3 (aq)

  • Nonmetal oxide + base  salt + water

    • CO2 (g) + 2NaOH (aq)  Na2CO3 (aq) + H2O (l)


Metalloids aka semi metals

Metalloids (aka Semi-metals)

  • Have properties that are intermediate between those of metals and nonmetals


7 6 group trends for the active metal

7.6 Group Trends for the Active Metal

Group 1A: The Alkali Metals

Characteristics

  • Soft metallic solids

  • Silvery

  • metallic luster

  • high thermal and electrical conductivities

  • Low densities and melting points

  • Most active metals

  • Exist in nature only as compounds


7 6 group trends for the active metal1

7.6 Group Trends for the Active Metal

Group 2A: Alkaline Earth Metals

  • Solids with typical metallic properties

  • Harder, more dense, and melt at higher temperatures when compared to alkali metals

  • Very reactive towards nonmetals, but not as reactive as alkali metals

  • Both alkali and alkaline earth metals react with hydrogen to form ionic substances that contain the hydride ion, H-


7 7 group trends for selected metals

7.7 Group Trends for Selected Metals

Hydrogen

  • Hydrogen is a nonmetal with properties that are distinct from any of the groups of the periodic table

  • It forms molecular compounds with other nonmetals, such as oxygen and the halogens


7 7 group trends for selected metals1

7.7 Group Trends for Selected Metals

Group 6A: The Oxygen Group

  • Most important element in group 6A

  • Exists in several allotropic forms (different forms of the same element in the same state)

  • Oxygen is encountered in two molecular forms, O2 (common form) and O3 (aka ozone)

  • Oxygen has a strong tendency to gain electrons from other elements, thus oxidizing them

  • In combination with metals, oxygen is usually found as the oxide ion, O2-, although salts of the peroxide ion, O22-, and superoxide ion, O2-, are sometimes formed


7 7 group trends for selected metals2

7.7 Group Trends for Selected Metals

Sulfur!!

  • 2nd more important element in group 6A

  • Also exists in several allotropic forms

  • Elemental sulfur is more commonly found as S8 molecules

  • In combination with metals, it is more often found as the sulfide ion, S2-


7 7 group trends for selected metals3

7.7 Group Trends for Selected Metals

  • Nonmetals that exist as diatomic molecules

  • There melting and boiling points increase as you go down the column

  • Have the most negative electron affinities of the elements

  • Their chemistry is dominated by a tendency to form 1- ions, especially in reactions with metals


7 7 group trends for selected metals4

7.7 Group Trends for Selected Metals

  • Group 8A: The Noble Gases aka inert gases

  • Nonmetals that exist as monoatomic gases

  • Very unreactive since they have completely filled s and p subshells. Have the complete octet

  • Have large 1st ionization energies

  • Only the heaviest noble gases are known to form compounds, and they do so only with very active nonmetals, like fluorine


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