review of chapter 6: bonding. bonds are forces of attraction between (-) electrons of one atom and...
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Review ofChapter 6: Bonding
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Bonds are forces of attraction between (-) electrons of one atom and the (+) nucleus of another atom, with 2 electrons in every bond
Forming bonds releases energy (exothermic)
Breaking bonds requires energy to be absorbed (endothermic)
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Valence electrons are the outermost electrons, farthest from the nucleus, involved in bonding Elements in same column/group have same
valence # Right-most number listed in electron configuration
on periodic table tiles
Octet Rule: Most atoms want to have 8 valence electrons, and make bonds to gain a share of 8 H and He only want 2 total Noble gases already have 8, don’t react with others
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Lewis Structures
Show # of valence electrons (1 to 8) around symbol for atom
Valence electrons drawn as dots
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Electronegativity is how strongly an atom wants electrons and pulls on them in a bond
Listed on Table SNonpolar covalent bonds: difference between
EN values of two bonded atoms is 0Polar covalent bonds: difference between EN
values of two bonded atoms is between 0.1 and 1.7
Ionic bonds: difference between EN values of two bonded atoms is greater than 1.7
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Covalent BondingNonmetals share electrons to form complete
octetsAny molecule with ONLY NONMETALS is
covalently bondedCan be single/double/triple bonds, with extra
electrons as lone pairs surrounding atoms
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1. Count total # valence electrons in atoms of compound
2. Arrange atoms• Central usually has
lowest electronegativity, only once, isn’t H
3. Place single bonds
4. Add lone pairs to outsides, then center
5. Make more bonds if needed
1. Count total # valence electrons in atoms of compound
2. Arrange atoms• Central usually has
lowest electronegativity, only once, isn’t H
3. Place single bonds
4. Add lone pairs to outsides, then center
5. Make more bonds if needed
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Ions form when atoms lose or gain electronsIons are just atoms with (+) or (-) charges
Metals lose e-, becoming (+) cations Ionization energy is energy needed to take an
electron away from an atom to make an ion Nonmetals gain e-, becoming (-) anions
In ionic bonding, metals give electrons away to nonmetals; charged ions then hang out near each other, attracted by different charges
Ionic Bonding
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To draw Lewis structures for metal and nonmetal ions: When ions form, metal gets no e-, nonmetal gets a complete set of 8. Put each ion in brackets, and write the charge at top right (oxidation number from periodic table)
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Metallic Bonding
Inside pieces of metal, (+) charged metal atoms are lined up neatly, with (-) electrons in constant motion, moving throughout the whole structure
This “sea of mobile electrons” holds the metal together, makes it a good electrical conductor, and makes it malleable (easily shaped)
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Intermolecular Forces
Non-permanent, not “real bonds”Influence melting and boiling pointsElectrostatic attractions between different
chargesHydrogen bonding: temporary attraction
between an H atom and an atom of either F/N/O Responsible for many important properties of water