Electron Dot Diagrams. Electron Dot Diagrams. Unbonded Atoms Ionic Bonding Covalent Bonding

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1 Specific Outcomes: i. I can define valence electron, electronegativity, ionic bond and intramolecular force. ii. I can draw electron dot diagrams. iii. I can use the periodic table and electron dot diagrams to support and explain ionic bonding theory. iv. I can explain how an ionic bond results from the simultaneous attraction of oppositely charged ions. v. I can draw electron dot diagrams of atoms and molecules, writing structural formulas for molecular substances and using Lewis structures to predict bonding in simple molecules. vi. I can relate electron pairing to multiple and covalent bonds. Electron Dot Diagrams Electron Dot Diagrams Electrons are arranged in layers (shells) around an atom Not all shells actually contain electrons The outermost shell that contains electrons is called the valence shell The electrons in the outermost occupied shell are called valence electrons All shells below these are also occupied! Remember how the octet rule arranges electrons in each shell: 2e, 8e, 8e... This is because electrons pair up into orbitals for stability 1 st shell has 1 orbital (holds 2 e ), 2 nd or 3 rd shells have 4 orbitals each We will pretend that the 4 th shell has only has 4 orbitals (holding 8 e ) 1

2 Exception: B has 3 orbitals! (not the usual 4) Atoms, shells and electrons are too small to be seen with a microscope We will draw the structure of the valence shell and show how chemical bonds form An electron dot diagram consists of an element symbol, surrounded by dots that represent only the valence electrons Steps: 1) Determine the # of valence e 2) Determine the # of available orbitals 3) Going clockwise, populate all e in orbitals, one at a time (undt s rule) ex. sulphur 6 valence e S ex. hydrogen 1 valence e ex. helium 2 valence e e ex. carbon 4 valence e C Try these: boron, magnesium, nitrogen 2

3 A full orbital (paired electrons) is called a lone pair, and never participates in bonding A halffull orbital contains only one electron, called a bonding electron Only the bonding e (the single ones!) are capable of forming chemical bonds The number of bonding electrons of an atom indicates how many bonds it may form An ionic bond is the electrostatic (+/) attraction between oppositely charged ions Most metals have three or fewer electrons in their valence level (INT: from ionic charge!) They tend to lose these electrons and become positive ions (cations) Most nonmetals have more than four valence electrons They tend to gain electrons and become negative ions (anions) After ions form, electrostatic attraction between positive and negative charge draws the ions together, forming an ionic bond To draw the electron dot diagrams for ionic compounds: 1) The metal has no valence electrons in the diagram (since e are lost) 2) The nonmetal has a full valence level (since e are gained) 3) Both ions have square brackets with the charge on the outside 3

4 ex. NaCl ex. MgO When ionic compounds form, the number of electrons lost by the metal must equal the number of electrons gained by the nonmetal The compound is neutral, resulting in a net charge of zero (positive and negative charges must add to zero) There may multiple copies of the metal and nonmetal to balance out ex. CaF 2 4

5 ex. K 2 S ex. Fe 2 O 3 ex. Mg 3 N 2 Try these: TiS 2, Nb 2 O 3, MoN 2 5

6 A covalent bond is formed when two nonmetal atoms share a pair of electrons Compounds containing covalent bonds are also called molecular compounds Ions are NOT formed, so no charge is present! Instead of transferring electrons, valence electrons are shared to satisfy the octet rule Note: boron has 6 electrons in a full shell! Electron dot diagrams of covalent/molecular compounds are called Lewis structures The electrons that are shared are called a bonding pair (made of 2 bonding electrons) Sharing two pairs of electrons between two atoms gives rise to a double bond (ex. O 2 ) Sharing three pairs of electrons between two atoms results in a triple bond (ex. N 2 ) To draw Lewis structures: 1) Draw separate electron dot diagrams for each of the individual atoms 2) Put the element with the most bonding (unpaired) electrons in the center, with all other atoms around it 3) Connect the bonding (unpaired) electrons, leaving no electron unpaired (at this stage, it will be very messy) To draw Lewis structures: 4) Redraw the Lewis structure, making double and triple bonds as needed 5) Check to make sure the octet rule is satisfied, remembering that hydrogen needs only 2 e 6

7 ex. P 3 P P ex. CO 2 C O O C O O Try these: BrCl, C 4, C 3 O, C 3 4, C 4 F 6 7

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