The Periodic Table and Periodic Trends Chapter 6 The Periodic Table and Periodic Trends

The Periodic Table 1. Mendeleev - arranged elements in order of increasing atomic mass. arranged elements with most similar properties in columns 2. Moseley - used x-rays to determine the atomic number of the elements. elements arranged by atomic number show a clear “periodicity” of properties.

Periodic Law There is a periodic repetition of the element’s chemical and physical properties when they are arranged in order of increasing atomic number. This repetition is known as the Periodic Law.

Arrangement of the Periodic Table horizontal row of elements - period # = # of energy levels an element has 2. Group (or family) - vertical column of elements Number (AorB) above a column is its family designation. elements in a group have much more similar properties than elements in a period.

Group A Elements (Representative Elements) contain metals, non-metals and metalloids s and p blocks. Group # = # of electrons in the outermost energy level. All shells beneath the outermost have stable octet. Except: Helium (He), Lithium (Li) & Beryllium (Be). First energy level is full with 2 electrons.

Group B Elements (Transition Metals) d block. May have more than 1 ionic charge because electrons in the shell beneath the outermost get involved in bonding. Shells beneath the outermost may have more than 8 electrons.

Classification of the Elements There are 3 classes of elements; metals, nonmetals and metalloids. Metals - make up about 80% of the table Good conductors of heat and electricity Have a high luster or sheen. (shiny) Solids at room temp. except for Mercury (Hg) Ductile; drawn into wires Malleable; hammered into thin sheets.

Nonmetals Found in the upper right corner of the table Properties of nonmetals are not as similar as they are for metals. Most are gases at room temperature. Some are solids like sulfur and phosphorus Bromine is a dark red liquid. Generally have properties opposite of the metals. Not good conductors (carbon exception), dull and brittle if they are solids.

Metalloids Along the staircase that separates the metals and nonmetals. Have properties that are similar to metals and nonmetals under certain conditions. Aluminum (Al) is a metal.

Classification of Elements Continued 1. Metals – lose electrons to form positive ions (cations) when they bond. S block: a) Group 1A (ns1) - Alkali or active metals - form +1 ions - very reactive, stored in kerosene

S block: a) Group 2A (ns2) - Alkaline earth metals - form +2 ions Note – Elements with the same outer shell electrons configuration have the same physical and chemical properties. (same number of valence electrons.)

Metals (continued) D block – these metals are all similar because they all have the same s2 electron configuration in their outermost shell. Transition Metals – Group B metals May have more than 1 ionic charge (designated by Roman #)

Metals (continued) Hund’s Rule – 1/2 filled and completely filled sublevels have the lowest energy are the most stable, and therefore the preferred electron arrangement.

Write the electron configuration for Fe # 26 1s2 2s2 2p6 3s2 3p6 4s2 3d6 Forms Fe+2 when it loses 2 electrons Forms Fe+3 when this additional electron is lost the ½ filled is d is more stable.

Write the electron configuration for Cu # 29 1s2 2s2 2p6 3s2 3p6 4s2 3d9 Forms Cu+2 when it loses 2 electrons Forms Cu+2 when it loses 2 electrons Hund’s Rule Shift: 1 e- shifts from 4s to 3d. The 3d is now completely filled and Cu will only lose 1 electron to form Cu+1.

Metals (continued) F block – Inner Transition Metals 2 rows: 1st row is Lanthanide series – “rare earth” elements (natural). 2nd row is Actinide series – mostly synthetic (manmade). Parentheses around mass # means radioactive.

Write configurations for #58 and #95 Write configurations for #58 and #95. Then write the condensed configurations. Ce # 58 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 5d1 4f 1 Am # 95 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f 14 5d10 6p6 7s2 6d1 5f6 2 8 18 32 24 9 2 8 18 19 9 Inner transition metals are most similar b/c their last 2 energy level config. are identical. same

2. Semi-metals Located on either side of the “stairs” in the p block separating the metals from the non-metals. Have properties of both metals and nonmetals. Aluminum (Al) is a metal

3. Non-metals in upper right corner of p block (except last column). gain electrons to achieve stable octet form negative ions - anions group 5, 6, and 7 gain 3, 2, or 1 electrons Halogens - Group 7A (ns2p5) Gains 1 e- and has a –1 charge

Question: Which are more similar? F, Cl, and Br or N, O, and F Answer: F, Cl, and Br Because in the same group or column, they have the same outer shell electron configuration, therefore react similarly.

4. Noble Gases (ns2p6) i.e. Sulfide ion (S-2) ion is isoelectric to Argon (Ar) atom. Last column on the right in p block. All but He already have their stable 8 in the outer shell. Therefore they are called: Inert – means non-reactive Isoelectric- same electron configuration as a noble gas. Atoms react to become isoelectric with a noble gas.

Periodic Trends Trends or patterns on the Periodic Table are all basically the result of atomic size and whether an atom of an element is a metal or a non-metal.

1. Atomic size (atomic radii) – size of the atom Nuclear charge – pull of protons in nucleus on electrons increase number of protons = increase the nuclear charge.

Atomic Size } Radius Atomic Radius = half the distance between two nuclei of a diatomic molecule.

1. Atomic size (atomic radii) – size of the atom Across a period – decreases due to increasing nuclear charge which pulls the electrons on the same energy level in closer to the nucleus. Down a group – increases due to the addition of energy levels which are farther from the nucleus.

1. Atomic size Across a period: Greater nuclear charge – same energy level Down a group: More energy levels

2. Metallic vs. Non-metallic Character Metallic Character – tendency to lose electrons - larger atom – easier to lose electrons. Across a period –decreases because the smaller the atom, the stronger the pull from the nucleus so the harder it is to lose electrons. Also the # of e- to lose is increasing as you approach the non-metals.

2. Metallic vs. Non-metallic Character Metallic Character – tendency to lose electrons - larger atom – easier to lose electrons. Down a group – increases because larger atoms have less pull from nucleus on outer electrons. More energy levels easier to lose electrons.

2. Metallic vs. Non-metallic Character b) Non-metallic Character – tendency to gain e- - smaller atom – easier to gain electrons Across a period –increases because as atomic size decreases its easier to gain electrons due to increased pull from nucleus. - also needs to gain fewer electrons.

2. Metallic vs. Non-metallic Character b) Non-metallic Character – tendency to gain e- - smaller atom – easier to gain electrons Down a group – decreases because atoms are larger (more energy levels). - less pull from nucleus- harder to gain electrons.

3. Shielding Effect Each additional energy level acts as another layer of electrons that shields the outermost electrons from the pull of the nucleus.

3. Shielding Effect Across the period – unchanged; on the same energy level With each addition electron, there is an additional proton. Down a group – increases due to additional energy levels between the nucleus and outer electrons.

4. Ionization Energy Amount of energy required to remove an electron from an atom. Metals have low ionization energies because they lose electrons easily. Non-metals have high ionization energies because they do not lose electrons easily.

4. Ionization Energy For group A metals: as electrons are removed 1 by 1 from an atom, there is a steady rise in the ionization energy until the structure of a noble gas is attained. Then there is a dramatic rise when a stable octet is disturbed.

4. Ionization Energy 1st ionization energy – amount of energy needed to remove the 1st electron from an atom. 2nd ionization energy – amount of energy needed to remove the 2nd electron from an atom. 3rd ionization energy – amount of energy needed to remove the 3rd electron from an atom.

4. Ionization Energy 281 1st ion. 2nd ion. 3rd ion. 4th ion. Na 119 1090 1652 2281 Mg 176 347 1848 2519 Al 138 434 656 2767 2 8 1s2 2s2 2p6 3s1 1s2 2s2 2p6 3s2 2 8 3 1s2 2s2 2p6 3s2 3p1

4. Ionization Energy Across a period – 1st ionization energy increases (across group A) because nuclear charge is increasing, holding electrons in tighter and closer to nucleus. For transition metals (pds 4, 5, & 6) in D block there is a sharp drop in ionization energy between the last transition element and where the next energy level’s p sublevel begins because the p sublevel is farther from the nucleus – becomes easier to lose electrons.

4. Ionization Energy Down a group – 1st ionization energy decreases because larger atoms have less pull on outer electrons due to increased shielding effect (see # 3).

5. Electronegativity or Electron Affinity (attraction) Halogens in group 7A have the highest electron affinity because they only need to gain 1 electron. Energy change due to the addition of an electron to an atom. Metals have low electron affinities. Non-metals have high electron affinities because by gaining electrons, they achieve the stable configuration of a noble gas.

5. Electronegativity or Electron Affinity (attraction) Across a period – increases because smaller atoms have a greater attraction for electrons and also approaching the non-metals that want to gain. Down a group – decrease because larger atoms have less attraction to gain electrons (outer shell farther from nucleus).

6. Ionic Size Ion – atom with unequal # of protons and electrons and has charge. Metals lose electrons to form positive ions that are smaller than their respective atoms because when outer electrons are lost, the protons in the nucleus have a great pull on the remaining electrons.

6. Ionic Size Ion – atom with unequal # of protons and electrons and has charge. Non-metals gain electrons to form negative ions that are larger than their respective atoms because when electrons are gained, there are less protons than electrons, so the electrons are not held as tightly.

6. Ionic Size Across a period – gradual decrease in the size of positive ions as they lose more electrons and then at group 5 begin the large negative ions that gradually decrease in size as you continue across. LOSE electrons GAIN electrons 1 2 3 4 3 2 1

7. Ionic Size Down a group – increases due to additional energy levels Holds true for both positive and negative ions.

Summation of Periodic Trends Going across a period. Going down a group. Atomic size Metallic character Non-metallic character decreases increases decreases increases increases decreases

Summation of Periodic Trends 2 Going across a period. Going down a group. Nuclear charge Shielding effect Speed of reaction – metals increases increases no change increases decreases increases

Summation of Periodic Trends 3 Going across a period. Going down a group. Speed of reaction – non-metals Ionization energy Electron affinity increases decreases increases decreases increases decreases

Summation of Periodic Trends 4 Going across a period. Going down a group. Ionic size of + ions(cation) - ions(anions) Decreases thru group 4A increases At group 5, large increases then a gradual decrease increases