Chapter 9 Flashcards
Covalent bonds
Atoms share electrons
Occur between nonmetal atoms
Octet rule
atoms are stabilized by having eight electrons in their valence shell
four pairs of electrons
Lewis structures
show the arrangement of covalently bonded atoms
Lone pair
a pair of unshared electrons
Drawing lewis structure steps
1) add up all the valence electrons
use periodic table group numbers to determine nuber of valence electrons
2)Frame the structure with single bonds
the atom closer to the lower left of the periodic table is the atom in the center. hydrogen is never in the center.
3) Fill the octets on outer atoms first
4) Fill the octet on the central atom
place any remaining electrons on the central atom
use double or triple bonds to fulfill the octet rule of the central atom
Polyatomic ions
groups of atoms with an overall charge
Formal charges
method of identifying charged sites
not perfect indicator of charge on an atom
but good tool for keeping track of overall charge of molecule or ion
assumes atoms share equally
How to calculate formal charge (formal charge = )
of valence electrons in the neutral atom - # of covalent bonds - # of unshared electrons
Drawing lewis structures for polyatomic ions
similar to neutral atoms
consider the charge when finding the # of valence electrons (wether it increases or decreases the amount of electrons)
When there is more than one option on how to structure a lewis structure, generally, what is the best
the best are generally the ones that minimize the charge values or have zero values for all of the atoms
Valence Shell Electron Pair Repulsion (VSEPR)
key idea: groups of electrons repel each other. in molecules electron groups will be as far away from each other as possible
Two ways we describe atom geometries
Electronic geometry- the arrangement of electrons around a central atom
Molecular geometry- the arrangement of atoms within a molecule/ shape caused by the arrangement of atoms
Molecular geometry
the arrangement of atoms within a molecule
also known simply as shape, depends on the electronic geometry but only considers the location of the atoms
Two electron sets:
Two bonding sets, 0 lone pairs
Electronic geometry: Linear
Bond angle of 180 degrees
Molecular geometry: Linear
Bond angle of 180 degrees
“Sets” of electrons
single, double, or triple bonds all count as one “set” of electrons
Three electron sets, if the three sets are attached to atoms ( 0 lone pairs)
(Three bonding sets, 0 lone pairs)
Electronic geometry: Trigonal planar
Molecule is flat and has a bond angle of 120 degrees
Molecular geometry: Trigonal planar
Molecule is flat with a bond angle of 120 degrees
Three electron sets, two sets attached to atoms and one as a lone pair
(Two bonding sets, one lone pair)
Electronic geometry: Trigonal planar
Molecule is flat and has a bond angle of 120 degrees
Molecular geometry: Bent
Four electron sets, if four sets are attached to atoms ( 0 lone pairs)
(Four bonding sets, 0 lone pairs)
Electronic geometry: Tetrahedral
Bond angle of 109.5 degrees
Molecular geometry: tetrahedral
occupies a 3D volume
Four electron sets, three attached to atoms, one as a lone pair
(Three bonding sets, one lone pairs)
Electronic geometry: Tetrahedral
Molecular geometry: trigonal pyramidal (pyramid shape)
Four electron sets, two attached to atoms, two as lone pairs
Two bonding sets, two lone pairs
Electronic geometry: Tetrahedral
Molecular geometry: Bent
Polar covalent bonds
atoms do not share electrons equally
ways to depict polar covalent bonds’s charge
delta + above the positive atom, delta minus over the negative
An arrow, pointing in the direction the electrons are pulled (the more electronegative side) with a small cross on the back end of the arrow to indicate the positive charge
Electronegativity
Measure of how strongly atoms pull bonded electrons
Pattern of electronegativity change on the periodic table
increases from left to right, consistent to what we know from ionic bonds, metals to the L of the table loose electrons easily while nonmetals on the R tend to pull or gain electron
also increases as you go up a column (group)
Difference in atom electronegativity range
Covalent: < 0.5
Polar covalent: 0.5-2.0
Ionic: >2.0
Molecular dipole (net dipole/ dipole)
Overall polarity in a molecule
finding net pole of a molecule
add together the pole of each polar covalent bond
how shape effects function
h2o in liquid form sticks together bc the atoms are attracted to each other bc of their net dipoles
