Term
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Definition
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Term
| what forms covalent bonds? |
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Definition
| positively-charged atomic nuclei and the negatively-charged electrons between them |
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Term
| what does chemical bonding do? |
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Definition
| Chemical bonding lowers the potential energy between positive and negative particles |
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Term
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Definition
| no electronegativity difference |
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Term
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Definition
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Term
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Definition
above 1 electronegative difference
metals and nonmetals |
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Term
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Definition
nonmetal and nonmetal
metalloid and nonmetal |
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Term
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Definition
a measure of how equallyelectrons are shared |
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Term
| what happens in a completely nonpolar covalent bond |
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Definition
the electrons are shared equally
•ΔEN = 0
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Term
| what happens in a nonpolar covalent bond? |
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Definition
•The electrons are shared fairly equally between the two atoms that are bonded, and ΔEN < 0.4
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Term
| what happens in a polar covalent bond? |
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Definition
•The electrons are pulled more strongly to one atom than to the other atom, and ΔEN = 0.4 – 1.7
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Term
| what happens in an ionic bond? |
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Definition
•The electrons are transferred
•ΔEN > 1.7
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Term
Which is the most polar bond?
(a) N - H
●
(b) F - N
●
(c) O - H
●
(d) I - Cl
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Definition
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Term
| what makes up a lewis electron dot symbol? |
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Definition
•consists of the chemical symbol for an element surrounded by dots
• Chemical symbol represents nucleus + core electrons
• Each dot represents a valence electron
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Term
| why do we use lewis dot diagrams? |
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Definition
| used to show and track the valence electrons (e.g. during a chemical reaction) |
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Term
| how to build a lewis diagram |
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Definition
•Figure out the total # of valence electrons.
• Make sure to adjust if it’s a cation (remove electrons) or an anion (add electrons).
•
• Figure out how the atoms are connected.
• The central atom is generally less electronegative than the atoms around it (except for H).
• Start by connecting them all with single bonds.
•
• Complete the octets of the atoms bonded to the central atom first (but for Hydrogen only use 2 electrons).
•
• If there aren’t enough electrons for the central atom, try giving it multiple bonds.
•
• Place any left over electrons on the central atom.
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Term
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Definition
•Octet = eight valence electrons
= four pairs of valence electrons
= eight v. e. means full s and p subshells
= filled outer level
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Term
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Definition
| It’s the charge that an atom would have if all the bonding electron pairs were shared equally. |
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Term
| what is the formal charge equation? |
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Definition
| Formal charge = (# of valence electrons in isolated atom) – (# of electrons assigned in the Lewis structure) |
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Term
| what is the most stable lewis structure? |
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Definition
| 1) has atoms that bear formal charges closest to zero, and (2) any negative formal charges are on atoms that are more electronegative. |
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Term
| what is more important: the octet rule or formal charge |
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Definition
| satisfying the octet rule is more important than minimizing formal charges. |
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Term
| exceptions to the octet rule |
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Definition
•Odd number of valence electrons in molecule
•Less than an octet (H, Be, B, Al common)
More than an octet (n = 3 and above since d orbitals available |
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Term
| what is a resonance structure? |
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Definition
describing
molecular structures with
‘delocalized’ electrons
Bonds in the molecule have:
• Equal bond strengths
• Equal bond lengths
• Equal bond dipoles |
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Term
| what does bond order equal? |
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Definition
electron pairs/
atom sets |
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Term
what does VSEPR stand for
? |
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Definition
| Valence-Shell Electron-Pair Repulsion Theory |
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Term
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Definition
Model explaining/predicting the geometric shape of
molecules:
Electron ‘groups’ are as far apart as possible
Requires an appropriate Lewis structure
Number of bonding groups and nonbonding groups are arranged in
optimal geometric arrangements |
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Term
| what determines the electron group arrangement? |
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Definition
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Term
| what are the types of groups in molecules? |
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Definition
Bonding groups (X) = # of surrounding atoms (regardless of bond order)
• Nonbonding groups (E) = # of lone pairs |
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Term
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Definition
electron group arrangement: linear
Molecular shape (AX2): Linear
2 bonded atoms
0 lone pairs
180° bond angle |
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Term
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Definition
electron group arrangement: trigonal planar
Molecular shape (AX3): Trigonal planar
3 bonded atoms
0 lone pairs
120° bond angles
Molecular shape (AX2E): Bent (V shaped)
2 bonded atoms
1 lone pair
~120° bond angle |
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Term
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Definition
electron group arrangement: tetrahedral
Molecular shape (AX4): Tetrahedral
4 bonded atoms
0 lone pairs
109.5° bond angles
Molecular shape (AX E): Trigonal pyramidal
3 bonded atoms
1 lone pair
~ 109.5° bond angles
Molecular shape (AX2E2): Bent
2 bonded atoms
2 lone pairs
~ 109.5° bond angles
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Term
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Definition
electron group arrangement: trigonal bipyramidal
Axial and equatorial positions are nonequivalent
Molecular shape (AX5): Trigonal bipyramidal
5 bonded atoms
0 lone pairs
90° and 120° bond angles |
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Term
Five Electron Groups
Electron-group arrangement: Trigonal bipyramidal |
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Definition
Molecular shape (AX4E): Seesaw
4 bonded atoms
1 lone pair
~90° and ~120° bond angles
Molecular shape (AX E ): T-shaped
3 bonded atoms
2 lone pairs
~ 90° bond angles
Molecular shape (AX2E3): Linear
2 bonded atoms
3 lone pairs
~ 180° bond angle |
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Term
Six Electron Groups
Electron-group arrangement: Octahedral |
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Definition
Molecular shape (AX6): Octahedral
6 bonded atoms
0 lone pairs
90° bond angles
Molecular shape (AX E): Square pyramidal
5 bonded atoms
1 lone pair
~ 90° bond angles
Molecular shape (AX4E2): Square planar
4 bonded atoms
2 lone pairs
~ 90° bond angles |
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Term
| summary of vsepr theory (steps) used to predict chemical properties |
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Definition
molecular formula
lewis structure (count all e groups around central atom)
electron-group arrangement (note positions of any lone pairs and double bonds
bond angles (count bonding and nonbonding e groups seperately)
molecular shape |
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Term
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Definition
uses wave behavior of the electrons to explain bonding
Bonds form when the orbitals of two atoms that
contain electrons overlap.
why H-H forms H2 |
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Term
| Bond strength and orbital overlap |
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Definition
| greater the orbital overlap - the stronger the bond |
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Term
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Definition
| shape and direction of the orbitals |
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Term
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Definition
in the
direction that maximizes
the overlap |
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Term
| hybridized atomic orbitals |
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Definition
Mathematically ‘mix’
isolated atomic orbitals
to obtain hybrid
orbitals |
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Term
| number of atomic orbitals = |
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Definition
| number of hybrid orbitals |
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Term
| type of hybrid orbitals depends on types of |
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Definition
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Term
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Definition
result from end-to-end overlap of orbitals
– Produces a region of high electron density directly along bond axis |
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Term
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Definition
result from side-to-side overlap of orbitals
– Produces two regions of electron density above and below bond axis |
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Term
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Definition
Single bond: one σ bond
• Double bond: one σ and one π bond
• Triple bond: one σ and two π bonds |
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Term
| Determining Geometries of Axn Molecules |
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Definition
1. Draw the Lewis Structure
2. Determine the total number of electron domains
around the central atom A.
3. Determine the electron domain geometry
(arrange the electron domains so that repulsions
among them are minimized).
4. Use the resulting arrangement of the bonded
atoms to determine the molecular geometry. |
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Term
How the Type of Electron Pair
Affects Bond Angles |
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Definition
Repulsive force of electron domains/volume
occupied by electron domains:
Nonbonding > triple > double > single
pairs bonds bonds bonds |
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Term
| Nonbonding pairs experience less blank blank than bonding pairs |
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Definition
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Term
multiple bonds contain
higher blank blank blank compared with single bonds |
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Definition
| electronic-charge density |
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Term
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Definition
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Term
| physical properties of sigma and pi bonding |
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Definition
Sigma bonds allow free rotation of the atoms around the bond axis
Pi bonds restrict rotation around the bond axis (π orbitals must
remain aligned parallel to form bond) |
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Term
| macroscopic observations of gas |
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Definition
conforms to shape and volume of container
high compressiblilty
high ability to flow |
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Term
| macro observations of liquid |
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Definition
conforms to shape of container; volume limited by surface
very low compressibility
moderate ability to flow |
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Term
| macro observationsof solid |
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Definition
maintains its own shape and volume
almost none compressibility and ability to flow |
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Term
| Intramolecular forces (bonding forces) |
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Definition
These forces exist within each molecule.
– They influence the chemical properties of the substance. |
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Term
| Intermolecular forces (nonbonding forces) |
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Definition
These forces exist between molecules.
– They influence the physical properties of the substance. |
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Term
| phase changes require what |
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Definition
changes in the energy of matter (You are not breaking bonds – just the weaker
intermolecular forces break) |
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Term
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Definition
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Term
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Definition
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Term
| comparison of bonding forces |
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Definition
Ionic: cation to anion
covalent: nuclei shared e- pair
metallic:cations delocalized electrons |
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Term
| comparison of nonbonding (intermolecular) forces |
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Definition
ion-dipole: ion charge-dipole charge
h bond: polar bond to H-dipole charge (high EN of N,O, and F)
dipole-dipole: dipole charges
ion induced dipole: ion charge - polarizable e cloud
dipole - induced dipole: dipole charge - polarizable e cloud
dispersion (london): polarizable e clouds |
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Term
| bonding and nonbonding forces are just electrostatic attractions. therefore: |
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Definition
Bonding forces tend to be stronger (large charges - close together)
– Nonbonding forces tend to be weaker (partial charges – further apart) |
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Term
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Definition
Most commonly occur in ionic
solutions (i.e. salt water)
– Charge of the ion and partial charges
of the polar molecules are attracted
– In salt water – ions are solvated by
water (hydration) |
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Term
| dipole-dipole interactions |
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Definition
Dipoles in polar molecules are attracted
• Attraction creates directional orientations in these substances
(oppositely-charged poles point at each other) |
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Term
| why is hydrogen bonding a special case? |
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Definition
H-bonding is a dipole-dipole force where there is an attraction between a H
atom bonded to an atom of high electronegativity (N, O, and F) and the
negative end (lone pair) of a nearby N, O, or F atom
H-bonding is only present in molecules
with N-H, O-H, or F-H bonds (the presence of N,O,F
and H is not enough) |
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Term
| charge induced dipole forces |
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Definition
Polarizability - distortion (or squishiness) of an electron cloud
• Increases down a group - size increases and the larger electron clouds are
further from the nucleus
• Decreases left to right across a period - increasing Zeff shrinks atomic size and
holds the electrons more tightly |
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Term
| dispersion (london) forces |
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Definition
Instantaneous dipoles caused by the random motion of
electrons
• Present in all substances, even noble gases like He
– this is the only force for nonpolar substances
• Dispersion forces tend to be the dominate intermolecular force in
most substances
• Stronger forces exist in more polarizable molecules, or molecules
that have more electrons
– strength of the force tends to scale with mass
Stronger forces exist between molecules with more ‘surface area’
to touch, if they have the same mass |
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Term
| ions present in these particle forces |
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Definition
ionic bonding (ions only)
ion dipole forces (ion + polar molecule)
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Term
| ions not present in these interacting particle forces: |
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Definition
dipole dipole forces (polar molecules)
polar and nonpolar: dipole induced dipole forces
nonpolar molecules only: dispersion forces only |
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