1 / 31

Lecture 1

Lecture 1. Contents: Introduction What is Inorganic Chemistry ? Main Group Chemistry Reading the Periodic Table: Classification Diagonal Relationship Electronegativity Electron affinity- ionization energy. Introduction. What is Inorganic Chemistry ?.

crystalallen
Download Presentation

Lecture 1

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Lecture 1 • Contents: • Introduction • What is Inorganic Chemistry ? • Main Group Chemistry • Reading the Periodic Table: Classification • Diagonal Relationship • Electronegativity • Electron affinity- ionization energy

  2. Introduction What is Inorganic Chemistry ? • “The chemistry of everything that is NOT organic…” • “The chemistry of all of the elements and their compounds • except for the hydrocarbons and their derivatives.” • “The branch of chemistry falling between and overlapping • with physical chemistry and organic chemistry.”

  3. Why Should You Study Inorganic Chemistry ? Elemental Composition of the Sun and the Universe Sun Universe Hydrogen 92.5 % 90.87 % Helium 7.3 % 9.08 % All Others 0.2 % 0.05 % • Essentially the entireuniverseis Inorganic. • The Earth is predominantly Inorganic. Elemental Composition of the Earth’s Crust Oxygen 45.5 % Iron 6.20 % Silicon 27.2 % Calcium 4.66 % Aluminum 8.30 % All Others 8.14 %

  4. Main Group Elements:s and p block elements • Transition Elements:d elements • Lanthanides and Actinides:f block elements p s d In which d orbit is being filled In which p orbit is being filled In whichs orbit is being filled f In which f orbit is being filled

  5. Main Group Chemistry

  6. Lanthanides Actinides

  7. Reading the Periodic Table: Classification • Nonmetals, Metals, Metalloids, Noble gases

  8. In the periodic table of elements, the number of the orbital electrons, that is the atomic number, and their arrangement determines the chemical and physical properties of an element. Each group of elements has a characteristic electronic arrangement and therefore elements within one group exhibit similar physical and chemical properties

  9. s-block elements 1-The alkali metals: in which all elements have one s electron in their outer shell, appear in a vertical column named Group 1. ns1 2-the alkaline earth metals: which have two s electrons in their outer shell, form Group 2. ns2 p-block elements Groups 13 ns2np1 Groups 14 ns2np2 Groups 15 ns2np3 Groups 16 ns2np4 Groups 17 ns2np5 Groups 18 or 0 ns2np6 (except He:1s2) The inert gases which appear in a group labeled Group 0 have the most stable arrangement of electrons because their outer shell of electrons is full. This explains their lack of chemical reactivity.

  10. Core Designation - A designation of electronic configuration wherein the outer shell electrons are shown along with the “core” configuration of the closest previous noble gas. Li Na K Rb [He] 2s1 [He] 2s2 Be Mg Ca Sr [Ne] 3s1 [Ne] 3s2 [Ar] 4s1 [Ar] 4s2 [Kr] 5s1 [Kr] 5s2

  11. Electronic Configuration and the Periodic Table • s-Block Elements • p-Block Elements • d-Block Elements • f-Block Elements Electronic Configuration for positive ions (cations) - Cations are formed by removing electrons in order of decreasing n value. Electrons with the same n value are removed in order of decreasing l value.

  12. Atomic Size - Atomic radii are considered to be 1/2 of the average distance between centers of identical atoms that are touching each other. This will vary with the chemical environment the atom is in. 1.42 Å 1.54 Å Fluorine Diamond C – 0.77 Å F – 0.71 Å

  13. Trends in Atomic Radii: • 1. Atomic radii increase from top to bottom in a • family or group. The number of electrons and the nuclear charge are increasing! - Tends to shrink atom. But extra electron are added to new shells that are further from the nucleus and more effectively shielded from the nucleus - Tends to make the atom larger.

  14. 2. Atomic radii decrease from left to right across a row or period. The number of electrons and the nuclear charge are increasing! - Tends to shrink atom. The electrons are being added to the same shell and are not well shielded and thus, the atoms get smaller. • 3. Summary of trends • Down a Group - Larger • Across a Period - Smaller

  15. What Affects Atomic/Ionic Sizes? • The Charge on the Nucleus • Shielding - This reduces the actual nuclear charge resulting in an “effective” nuclear charge. In general, Zeff (effective nuclear charge) increases across a period but remains about the same or slightly decrease down a group.

  16. Ionic Size - • Based on the internuclear distance of cations and anions in ionic crystals. • Cations - Monatomic cations are smaller than • their parent atoms. • The whole outer shell is typically removed. • The effective nuclear charge is increased. Na atom 1.86 Å Na+ ion 1.02 Å

  17. Anions - Monatomic anions are larger than their parent atoms. The extra electrons are typically added to the same shell where they are repelled by the other electrons already present, making the ion bigger than its parent atom. F Atom 0.71 Å Fluoride Ion 1.36 Å

  18. Ionization Energy - The energy required to remove an electron from a gaseous ground-state atom or ion. A. First Ionization Energy - The energy required to remove the most loosely bound electron from the valence shell. B. Second Ionization Energy - The energy required to remove the second electron after the first one is gone. C. Third Ionization Energy - Etc., Etc., Etc.

  19. Li(g) Li+ + e- Li+ Li2+ + e- IE1 = +520 kJ/mol IE2 = +7298 kJ/mol Na(g) Na+ + e- IE1 = +496 kJ/mol IE2 = +4564 kJ/mol Na+ Na2+ + e- Mg(g) Mg+ + e- Mg+ Mg2+ + e- Mg2+ Mg3+ + e- IE1 = +737 kJ/mol IE2 = +1447 kJ/mol IE3 = +7738 kJ/mol

  20. Trends in Electron Affinities - H Li Be B C N O F Na Mg Al Si P S Cl K Ca Ga Ge As Se Br Rb Sr In Sn Sb Te I Cs Ba Tl Pb Bi Po At Increases up a group. Increases from left to right in a period.

  21. Summary of Trend 3. Ionization Energy: Largest toward NE of PT 4. Electron Affinity: Most favorable NE of PT • Periodic Tableand Periodic Trends • 1. Electron Configuration 2. Atomic Radius: Largest toward SW corner of PT

  22. Metals Blue (underlined)Non-metals green (italics)Metalloids red (Bold)

  23. p-block and the frontier zone • So this gives rise to frontier between metals and non-metals and it is in this zone are the metalloids that we have encountered earlier. The metalloid have some of the characteristics of metals and non-metals and this frontier is denoted by this zigzag line • 3 4 5 6 7 B C N O F non Al Si P S Cl metals GaGeAsSeBr In SnSbTeI TlPb BiPoAs MetalsMetalloids

  24. Electrongetaivities for selected elements (Pauling Scale)

  25. Diagonal Relationships The elements in each encircled pair have several similar properties.

  26. Diagonal Relationship • Li Be B C Na Mg Al Si • Elements that are linked by the arrows in the diagram above are said to be diagonally linked. These pairs of elements often show similar chemical properties e.g. Li and Mg both form nitrides.

  27. Group IA: The Alkali Metals The metals in Group IA (Li, Na, K, Rb, Cs, and Fr) are called the alkali metals because they all formhydroxides (such as NaOH) that were once known as alkalies. The electron configurations of the alkali metals are characterized by a single valence electron. As a result, the chemistry of these elements is dominated by their tendency to lose an electron to form positively charged ions (Li+, Na+, K+). The alkali metals lose electrons so easily that sodium dissolves in liquid ammonia at temperatures below the boiling point of ammonia (-33oC) to give Na+ ions and electrons. NH3(l)

  28. Group IIA: The Alkaline-Earth Metals The elements in Group IIA (Be, Mg, Ca, Sr, Ba and Ra) are all metals, and all but Be and Mg are active metals. The term alkalinereflects the fact that many compounds of these metals are basic or alkaline. The term earth was historically used to describe the fact that many of these compounds are insoluble in water. Three points should be kept in mind, however. 1-The alkaline-earth metals tend to lose two electrons to form M2+ ions (Be2+, Mg2+, Ca2+, and so on). 2-These metals are less reactive than the neighboring alkali metal. Magnesium is less active than sodium; calcium is less active than potassium; and so on. 3-These metals become more active as we go down the column. Magnesium is more active than beryllium; calcium is more active than magnesium; and so on.

  29. Inert pair effect Al and Tl are both metals in group 3 of the periodic table, but Al ions are only ever found in the +3 state. (Al3+cations), but Tl is known to form compounds in which there can be Tl+ or Tl3+cations. This tendency for elements at the bottom of groups 3, 4 and 5 to form compounds in which their outermost s electrons are not involved in bonding is called the inert pair effect. The basic reason is that the s electrons see much more of the nucleus than the p electrons so they are more stable (the s electrons are more penetrating)

  30. Properties and Trendsin Group 3A • Boron tends to form covalent compounds rather than ionic compounds. • The rest of group 3A elements can form 3+ ions. • Gallium, indium, and thallium also often form 1+ ions by retaining their ns2 electrons; this is called the inert pair effect.

More Related