3s 2 3p 4. 2s 2 2p 1. Ionic Compounds. Chemical bonds. Lewis Electron-Dot Symbols

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1 Chemical bonds Chemical bond -- a strong attractive force between atoms that binds them together to form chemical compounds There are three classes of chemical bonds: Ionic bonds -- electrostatic forces that exist between ions of opposite charges ionic bonds are associated with the complete transfer of electrons from one atom to another Covalent bonds -- associated with the sharing of electrons between atoms Lewis Electron-Dot Symbols The Lewis electron-dot symbol of an atom is a representation that shows the valence electrons for that atom The electron-dot symbol of an atom uses dots to show the valence electrons of atoms The number of dots equals the number of s and p electrons in the outermost (valence) shell of the atom Metallic bonds -- interaction between atoms in metals (iron, copper, etc.) that give rise to typical properties of metals bonding electrons are relatively free to move throughout the three-dimensional structure of the metal The electron-dot symbol of an atom uses dots to show the valence electrons of atoms The electron-dot symbol of an atom uses dots to show the valence electrons of atoms Paired electrons B 2s 2 2p 1 Unpaired electron Symbol of the element S 3s 2 3p 4 The number of dots equals the number of s and p electrons in the outermost (valence) shell of the atom The number of dots equals the number of s and p electrons in the outermost (valence) shell of the atom Electron-dot symbols of the first 20 elements Ionic Compounds 1A 2A 3A 4A 5A 6A 7A 8A (Noble Gases)

2 Ion Formation and the ctet Rule Noble gas electron configurations are very stable For most elements, attaining a noble gas configuration means having 8 valence electrons (two s electrons and six p electrons) ctet Rule Main group elements (i.e., non-transition elements) tend to undergo reactions that leave them with a full valence shell 2 valence electrons for the first electron shell (n = 1) 8 valence electrons for all other electron shells (n 2) 2 He 10 Ne 18 Ar 36 Kr 54 Xe 86 Rn 1s 2 this is called a full valence shell (also referred to as an octet ) 1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 3s 2 3p 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 For the elements in Period 1 (hydrogen and helium), a full valence shell consists of 2 valence electrons (two s electrons) this is because the first principal energy level (n = 1) does not have a p sublevel (i.e., no p orbitals) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 Use the octet rule as a guide to determine the charge on ions formed by main group elements Metals can readily attain a full octet (noble gas configuration) by losing electrons metals form positive ions (cations) Non-metals can readily attain a full octet (noble gas configuration) by gaining electrons non-metals form negative ions (anions) Example: Formation of the sodium ion A sodium atom can readily lose the single 3s electron in its outermost shell by doing so, it attains a noble gas configuration (full octet) 495 kj/mol (1st ionization energy) 11 p + 11 p Na atom 1s 2 2s 2 2p 6 3s 1 [Ne] 3s 1 Na + ion 1s 2 2s 2 2p 6 [Ne] Example: Formation of the sodium ion A very large amount of energy would be required to remove an additional electron from the filled inner shell of the Na + ion Na 2+ ions are not observed 11 p + 11 p+ + Na + ion 1s 2 2s 2 2p 6 [Ne] kj/mol (2nd ionization energy) Na 2+ ion 1s 2 2s 2 2p 5 [He] 2s 2 2p 5 +2 Example: Formation of the chloride ion Adding an electron to the outermost shell (n=3) of a chlorine atom is an energetically favorable process by doing so, it attains a noble gas configuration (full octet) 17 p + + Cl atom 1s 2 2s 2 2p 6 3s 2 3p 5 [Ne] 3s 2 3p 5 E = 349 kj/mol energy is released (exothermic process) 17 p + Cl ion 1s 2 2s 2 2p 6 3s 2 3p 6 [Ar]

3 Example: Formation of the chloride ion Trying to add another electron to the next higher shell (n=4) in a chloride ion would be very energetically unfavorable Cl -2 ions are not observed 17 p + + Cl ion 1s 2 2s 2 2p 6 3s 2 3p 6 [Ar] E highly positive large amount of energy required (endothermic process) 17 p + Cl 2 ion 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 [Ar] 4s 1-2 Ionic bonds An ionic bond is formed by the attraction between two oppositely charged ions Ionic bonds are formed whenever electrons are transferred completely from one ion to another Metals (which have relatively weak attraction for their valence electrons) tend to form ionic bonds when they combine with non-metals The attraction between ions is very strong, which results in high melting points for ionic compounds (often > 300 C) Reaction between sodium and chlorine Reaction between sodium and chlorine 11 p + 17 p+ Na + Cl Na + + Cl + Na Cl Na + Cl- NaCl (ionic compound) Na + Cl- NaCl (ionic compound) Strong attraction between positive Na + and negative Cl - forms an ionic bond Strong attraction between positive Na + and negative Cl - forms an ionic bond Ionic substances are not truly molecular Properties of Ionic Compounds Ions in an ionic compound are packed closely together in an orderly, repeating arrangement (crystal structure) Example: Sodium chloride bonds ionically it consists of a large aggregate of positive and negative ions individual molecules of NaCl do not exist NaCl represents the formula unit for sodium chloride (i.e., the smallest, simplest neutral unit of the compound) crystalline solids at room temperature high melting and boiling points hard and brittle (will shatter if struck sharply) often soluble in water or other polar solvents will conduct electricity when melted or dissolved in water

4 Energetics of ionic compound formation The reaction between sodium and chlorine to form sodium chloride is highly exothermic: 2 Na (s) + Cl 2 (g) 2 NaCl (s) E = 822 kj But the net energy for the formation of Na + and Cl ions is positive (endothermic process) -- i.e., energetically unfavorable 2 Na (s) 2 Na + (g) + 2 e E = 1208 kj (ionization energy) Cl 2 (g) + 2 e 2 Cl (g) E = 454 kj (electron affinity) Positive value (endothermic) Enet = 754 kj So why is the formation of NaCl (s) such a highly exothermic, energetically favorable process overall? Lattice energy and ionic bonding The energy for the formation of Na + and Cl ions based on ionization energy and electron affinity corresponds to ions that are infinitely far apart -- i.e., assumes that the ions are not interacting with each other But the Na + and Cl ions are strongly attracted to each other due to their opposite charges (ionic bonds) -- this attraction draws the ions together, releasing energy and causing the ions to form a solid array, or lattice Cl Na + The lattice energy provides a measure of how much stabilization results from the interaction between ions defined as the energy required to completely separate 1 mole of a solid ionic compound into its gaseous ions NaCl(s) Elattice = +788 kj Na + (g) + Cl (g) Lattice energy and ionic bonding Cl Na + Note that the process of separating the ions from each other in highly endothermic -- i.e., lots of energy must be put in to overcome the ionic attractions NaCl(s) Elattice = +788 kj Na + (g) + Cl (g) Energetics of ionic compound formation accounting for lattice energy 2 Na (s) 2 Na + (g) + 2 e E = 1208 kj (ionization energy) Cl 2 (g) + 2 e 2 Cl (g) E = 454 kj (electron affinity) 2 Na + (g) + 2 Cl (g) 2 NaCl(s) E = 1576 kj (lattice energy) Negative value (exothermic) Enet = 822 kj The reverse process -- bringing the ions together -- is highly exothermic forming the lattice of ions is energetically favorable and releases energy Na + (g) + Cl (g) NaCl(s) E = 788 kj verall, the reaction between sodium and chlorine to produce sodium chloride is a highly favorable process due to the stabilization resulting from the formation of ionic bonds Ionic compounds have large, positive lattice energies Covalent bonds and molecular compounds Compound Lattice energy (kj/mol) Compound Lattice energy (kj/mol) LiF LiCl LiI NaF NaCl NaBr NaI KF KCl KBr CsCl CsI MgCl2 SrCl2 Mg Ca Sr ScN The reverse process -- forming the ionic lattice -- is therefore highly exothermic

5 Covalent bonds Covalent bonding: Net attractive forces between nuclei and electrons hold atoms together Instead of completely transferring electrons to form ions (and ionic bonds), elements can form chemical bonds by sharing pairs of electrons nucleus electron attraction electron electron repulsion covalent bond -- a bond consisting of pair of electrons shared between two atoms + nucleus nucleus repulsion + electron electron repulsion In most covalent compounds, atoms share electrons to achieve octets (8 valence electrons) -- i.e., noble gas configurations nucleus electron attraction Covalent bonds arise when nucleus-electron attractive forces are greater than nucleus-nucleus and electron-electron repulsive forces Covalent bonding in the hydrogen molecule Covalent bonding in the hydrogen molecule Two 1s orbitals from each of two hydrogen atoms overlap The orbital of the electrons includes both hydrogen nuclei The most likely region to find the two electrons is between the two nuclei Each 1s orbital contains 1 electron The two nuclei are shielded from each other by the electron pair this allows the two nuclei to draw close together Each hydrogen atom now has 2 electrons in its outermost energy level The internuclear distance that results in the most stable configuration determines the length of a covalent bond Covalent bonding in methane Atoms too close together (repulsive forces dominate) Potential energy 0 ptimum distance (strongest attraction) Atoms too far apart (weak attraction) In a molecule of methane (CH4): the central carbon atom shares an electron with four hydrogen atoms to achieve an octet each hydrogen atom shares one electron with the carbon atom to attain a noble gas structure (similar to He) Bond length = the optimum distance between nuclei 74 pm (for H2) Distance between nuclei

6 Compounds formed by covalent bonding exist as individual molecules Comparison of ionic compounds and molecular compounds Ionic compounds Example: Sodium chloride (NaCl) Example: Carbon dioxide bonds covalently it exists as molecules containing one carbon atom covalently bonded to two oxygen atoms (C 2 ) Strong attractive forces hold positively and negatively charged ions tightly together Ionic compounds generally have high melting points and high boiling points Molecular compounds Example: Carbon dioxide (C2) Na + Cl Molecules are neutral, so there are no strong electrostatic forces holding them together Compounds formed by covalent bonding are referred to as molecular compounds Molecules are attracted to one another by several weaker intermolecular forces (more about these later) Molecular compounds typically exist as gases, liquids, or solids with low melting and boiling points C Comparison of ionic compounds and molecular compounds Ionic compounds Bonds formed through complete transfer of electrons Bonding between metals and non-metals Aggregate of interlocking positive and negative ions (no individual molecules) Crystalline solids High melting and boiling points Conduct electricity when molten or dissolved in water Many are water soluble Not soluble in organic liquids Molecular compounds Bonds formed through sharing of electrons Bonding between non-metals and other non-metals Exist as individual molecules Gases, liquids, or low-melting solids Low melting and boiling points Do not conduct electricity Few are water soluble Many are soluble in organic liquids Numbers of covalent bonds typically formed by main group elements to achieve noble gas configuration 1A H 3A B 3 bonds The octet rule is a useful guideline, but there are numerous exceptions 4A C 4 bonds Si 4 bonds 5A N 3 bonds P 3 bonds (5) boron forms 3 covalent bonds and ends up with only 6 valence electrons elements in the third period and below can form additional covalent bonds (this is because they have vacant d orbitals that can hold additional bonding electrons) 6A 2 bonds S 2 bonds (4, 6) 7A F Cl (3, 5) Br (3, 5) I (3, 5, 7) 8A He Ne Ar Kr Xe Hydrogen, Carbon, Nitrogen, xygen, and Halogen Atoms Usually Maintain Consistent Bonding Patterns H C N Hydrogen 1 covalent bond Carbon 4 covalent bonds Nitrogen 3 covalent bonds 1 lone pair xygen 2 covalent bonds 2 lone pairs F Cl Br I Halogens: 1 covalent bonds, 3 lone pairs

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