Presentation on theme: "CHEMICAL PERIODICITY. Development of Periodic Table Dmitri Mendeleev, Russian Chemist in the mid-1800’s arranged the 70 known elements into a systematic."— Presentation transcript:
Development of Periodic Table Dmitri Mendeleev, Russian Chemist in the mid-1800’s arranged the 70 known elements into a systematic way. –Put the names of elements on a card (Just like you arranged the pieces) –He also put atomic masses, physical and chemical properties. –When arranged in the order of atomic mass, he discovered a repeated, or periodic, pattern. –He left blank spaces where he didn’t know the element for the spot. –By doing this, he constructed the first periodic table for elements.
Mendeleev’s Problems 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?
Henry Moseley’s Solution 1.Instead of arranging in atomic mass, in1913 Henry Moseley arranged in increasing order of nuclear charge. His work led to modern definition and recognition of atomic number. 2.He discovers that it is the number of electrons that affect an atoms’ reactivity, or chemical property.
Parts of the Periodic Table Periods: –The horizontal rows Periodic Law: –When the elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties. Groups: (a.k.a. Families) –The vertical columns
Parts of the Periodic Table, cont. Representative Elements: –Group A elements. (The tall parts of the periodic table.) –These can be divided into 3 groups: Metals Nonmetals Metalloids
Metals Malleable-Very easy to be hammered or beaten into thin sheets Ductile-Very easy to produce a wire High luster-De-excitation of electrons causes shiny appearance of metal surface High electrical conductivity - Due to the free movement of electrons High thermal conductivity Most are solid at room temperature Metallic Bond Strength –Directly proportional to the heat of vaporization 80% of all elements are metals
Nonmetals Upper right portion of the periodic table Generally nonlustrous Generally poor conductors of electricity Some are gases at room temperature and some brittle solids One is even a liquid at room temperature
Metalloids These border the stair step, except for aluminum These have properties of both metals and nonmetals Si, Ge, As, Sb, Te
Group 1 – Alkali Metal Very reactive!Not found in nature as free elements Combine vigorously with most nonmetals React strongly with water to produce hydrogen gas and aqueous solutions of substances known as “alkalis.” Usually stored in kerosene (paraffin oil)
Group 2 – Alkali Earth Metal Harder, denser, and stronger than alkali metal Higher melting points Not so reactive compared to alkali metal; Still very reactive! Not found as free element
Group 3 – 12 Transition Metal Good conductors of electricity and have a high luster Less reactive than alkali and alkali-earth metal Some (e.g. platinum and gold) are so unreactive that they do not form compounds easily. Some are found as free element.
Group 17 - Halogens Most reactive nonmetal React vigorously with most metals to form “salts” Need one electron to obtain stable noble-gas configuration F and Cl are gaseous. Br is liquid. Iodine is solid.
Group 18 – Noble Gases Lack chemical reactivity Very stable Octet electronic configuration in the outermost energy level (or “shell”)
Categories of electrons Core (inner) electrons – all those shared by the previous noble gas Outer electrons – those in the highest occupied energy level –Similar chemical properties of elements in groups is a result of similar outer electron configurations –In the main group, the group number equals the number of outer electrons Valence electrons – those involved in bonding –In the main group, the outer electrons are valence –In the transition metals can include some d electrons
Trends in metallic behavior
Metallic Behavior Metallic behavior increases left and down Metals tend to lose electrons in reactions Highly metallic elements, are likely to make positive ions Least metallic elements are likely to make negative ions In middle are more likely to make covalent bonds
Ion (Cation and Anion) Neutral atom Add electron Negative ion (Anion)Neutral atom Remove electron Positive ion (Cation) # of electrons added/removed= # of charge F + 1e - F -1 (Anion)Ca -2e - Ca +2 (Cation)
Main-group ions and the noble gas configurations Octet (Duplet) Rule
+1+2+3 -2-3 Cations and Anions Of Representative Elements 8.2
Periodic Law All the elements in a group have the same electron configuration in their outermost shells Example: Group 2 Be1S 2 2S 2 2, 2 Mg 1S 2 2S 2 2p 6 3S 2 2, 8, 2 Ca 1S 2 2S 2 2p 6 3S 2 3p 6 4S 2 2, 8, 8, 2
Periodic Table and Electron Configuration s 1 s 2 p 1 p 2 p 3 p 4 p 5 p 6 1 2 3 d 1 - d 10 4 5 6 f 1 - f 14 Groups 1-2 = s level Groups 3-8= p level Transition= d level Lantanides = f level
Figure 8.13 The relation between orbital filling and the periodic table
Trends in Some Periodic Properties The physical and chemical behavior of the elements is based on the electron configurations of their atoms. e- configurations can be used to explain many of the repeating or “periodic” properties of the elements
Figure 8.14 Defining metallic and covalent radii
8.3 Atomic Radius - size of atom
Ionization energy: the energy required to remove an electron from an atom is ionization energy. (measured in kilojoules, kJ) Electronegativity is a measure of an atom’s attraction for another atom’s electrons. Electron affinity is the energy change that occurs when an atom gains an electron (also measured in kJ). Atomic Radius/Ionic Radius- size of atom/Ion Definition of Some Periodic Properties
Atomic radii of the main-group and transition elements.Vertically The trend for atomic radius in a vertical column is to go from smaller at the top to larger at the bottom of the family. With each step down the family, we add an entirely new energy level, making the atoms larger with each step. Horizontally Each step adds a proton and an electron (and 1 or 2 neutrons). Electrons are added to existing energy level. The effect is that the more positive nucleus has a greater pull on the electron cloud. The nucleus is more positive and the electron cloud is more negative. The increased attraction pulls the cloud in, making atoms smaller as we move from left to right across a period.
Electron Shells and Sizes of Atoms Size in main group elements follow 2 general rules –1. AR increases going down in a group –2. AR decreases going left to right in a period Why? Moving down in a group: –Z eff is constant because Z and S increase equally –Electrons in a higher n and therefore larger Moving across: –Z increases and S does not (electrons in same shell don't shield each other well) so e- attracted more strongly and are closer/smaller In transition metals there is an initial decrease in size but then remains relatively constant
Figure 8.16 Periodicity of atomic radius
SAMPLE PROBLEM 8.3Ranking Elements by Atomic Size PLAN: SOLUTION: PROBLEM:Using only the periodic table (not Figure 8.15), rank each set of main group elements in order of decreasing atomic size: (a) Ca, Mg, Sr(b) K, Ga, Ca(c) Br, Rb, Kr(d) Sr, Ca, Rb Elements in the same group increase in size and you go down; elements decrease in size as you go across a period. (a) Sr > Ca > MgThese elements are in Group 2A(2). (b) K > Ca > GaThese elements are in Period 4. (c) Rb > Br > KrRb has a higher energy level and is far to the left. Br is to the left of Kr. (d) Rb > Sr > CaCa is one energy level smaller than Rb and Sr. Rb is to the left of Sr.
Effective nuclear charge (Z eff ) is the “positive charge” felt by an electron. Na Mg Al Si 11 12 13 14 10 1 2 3 4 186 160 143 132 Z eff Core Z Radius (pm) Z eff = Z - 0 < < Z ( = shielding constant) Z eff Z – number of inner or core electrons
Effective Nuclear Charge (Z eff ) increasing Z eff
Ions Here is a simple way to remember which is the cation and which the anion: This is a cat-ion. This is Ann Ion. He’s a “plussy” cat! She’s unhappy and negative. +
Cation Formation 11p+ Na atom 1 valence electron Valence e- lost in ion formation Effective nuclear charge on remaining electrons increases. Remaining e- are pulled in closer to the nucleus. Ionic size decreases. Result: a smaller sodium cation, Na +
Anion Formation 17p+ Chlorine atom with 7 valence e- One e- is added to the outer shell. Effective nuclear charge is reduced and the e- cloud expands. A chloride ion is produced. It is larger than the original atom.
Cation is always smaller than atom from which it is formed. Anion is always larger than atom from which it is formed.
The Radii (in pm) of Ions of Familiar Elements
Ionic Size Cations are smaller than the parent atom Anions are larger than the parent atom. Ionic size increase down in a group In an isoelectronic series, the most negative ion is largest, the most positive is smallest. Atoms making more than one ion, most positive is smallest.
Ionization Energy This is the second important periodic trend. The atom has been “ionized” or charged. The number of protons and electrons is no longer equal. The energy required to remove an electron from an atom is ionization energy. (measured in kilojoules, kJ) The larger the atom is, the easier its electrons are to remove. Ionization energy and atomic radius are inversely proportional. Ionization energy is always endothermic, that is energy is added to the atom to remove the electron.
Ionization Energy (Potential) Draw arrows on your help sheet like this:
Ionization energy Trends: –Generally as size decreases, IE 1 increases. –1. IE 1 generally increases from left to right, some exceptions –2. IE 1 generally decreases going down in a group –3. transition and f-block elements have much smaller variances in IE 1 But Why? –Across - increasing Z eff and smaller size(e- closer to nucleus) –exceptions - Be to B because s shields p and lowers Z eff –N to O because repulsions in first paired electron –Down - Z eff constant and larger so easier to ionize
Figure 8.17 Periodicity of first ionization energy (IE 1 )
Figure 8.18 First ionization energies of the main-group elements
51 Ionization energy is the minimum energy (kJ/mol) required to remove an electron from a gaseous atom in its ground state. I 1 + X (g) X + (g) + e - I 2 + X + (g) X 2 + (g) + e - I 3 + X 2+ (g) X 3 + (g) + e - I 1 first ionization energy I 2 second ionization energy I 3 third ionization energy I 1 < I 2 < I 3
SAMPLE PROBLEM 8.4Ranking Elements by First Ionization Energy PLAN: SOLUTION: PROBLEM:Using the periodic table only, rank the elements in each of the following sets in order of decreasing IE 1 : (a) Kr, He, Ar(b) Sb, Te, Sn(c) K, Ca, Rb(d) I, Xe, Cs IE decreases as you proceed down in a group; IE increases as you go across a period. (a) He > Ar > Kr (b) Te > Sb > Sn (c) Ca > K > Rb (d) Xe > I > Cs Group 8A(18) - IE decreases down a group. Period 5 elements - IE increases across a period. Ca is to the right of K; Rb is below K. I is to the left of Xe; Cs is furtther to the left and down one period.
Ionization energy Second ionization energy IE 2 – to remove 2 nd e- –Always larger than IE 1 –as e- are removed, Z remains constant and remaining e- are harder to remove A huge increases occurs in ionization energies when core electrons are reached because of the lower shielding/higher Z eff. Result - Rx only involve outer e-
Figure 8.19 The first three ionization energies of beryllium (in MJ/mol)
SAMPLE PROBLEM 8.5Identifying an Element from Successive Ionization Energies PLAN: SOLUTION: PROBLEM:Name the Period 3 element with the following ionization energies (in kJ/mol) and write its electron configuration: IE 1 IE 2 IE 3 IE 4 IE 5 IE 6 1012190329104956627822,230 Look for a large increase in energy which indicates that all of the valence electrons have been removed. The largest increase occurs after IE 5, that is, after the 5th valence electron has been removed. Five electrons would mean that the valence configuration is 3s 2 3p 3 and the element must be phosphorous, P (Z = 15). The complete electron configuration is 1s 2 2s 2 2p 6 3s 2 3p 3.
Electronegativity Electronegativity is a measure of an atom’s attraction for another atom’s electrons. It is an arbitrary scale that ranges from 0 to 4. The units of electronegativity are Paulings. Generally, metals are electron givers and have low electronegativities. Nonmetals are are electron takers and have high electronegativities. What about the noble gases?
Electronegativity Your help sheet should look like this: 0
Overall Reactivity This ties all the previous trends together in one package. However, we must treat metals and nonmetals separately. The most reactive metals are the largest since they are the best electron givers. The most reactive nonmetals are the smallest ones, the best electron takers.
Overall Reactivity Your help sheet will look like this: 0
Electron Affinity What does the word ‘affinity’ mean? Electron affinity is the energy change that occurs when an atom gains an electron (also measured in kJ). Where ionization energy is always endothermic, electron affinity is usually exothermic, but not always.
Electron Affinity Electron affinity is exothermic if there is an empty or partially empty orbital for an electron to occupy. If there are no empty spaces, a new orbital or PEL must be created, making the process endothermic. This is true for the alkaline earth metals and the noble gases.
Electron Affinity Your help sheet should look like this: ++
64 Electron affinity is the negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion. X (g) + e - X - (g) F (g) + e - X - (g) O (g) + e - O - (g) H = -328 kJ/mol EA = +328 kJ/mol H = -141 kJ/mol EA = +141 kJ/mol
electron affinities energy required to put an e- on an atom (making it negative or more negative) EA 1 usually negative EA 2 always positive Irregular trend: increases going right and up –Remember: Full shells best, full subshells good, and half full is kinda cool –Cl and other halogens need only 1 e- so most exothermic reaction –noble gases have endo because new electron in higher n –Mg, Be endo because new e in higher l –N endo because new e paired
Figure 8.20 Electron affinities of the main-group elements