Unit 2- Molecular & Ionic Compounds Structures & Properties Flashcards
Ionic bond
Metal transfer electrons to non-metals
Electrostatic bonds result due to attraction of opposite charges
Covalent bond
Non-metal atom shares one more more pair of electrons with another non-metal
Results in a full octet of electrons
Molecules
2 or more non-metals bonded together to form compound
Bonds in molecules are usually
nonpolar or polar covalent
Electronegativity
Element’s ability to attract bonding electrons in a bond
Electronegativity increases as atomic radius
Decreases
For two atoms in a chemical bond, the electrons are more attracted to the
positive nucleus closest to them
Coulomb’s Law
F = k * (q1q2)/d^2 q1 = charge of ion q2 = charge of ion d = distance
Bonds can be classified according to differences in
electronegativity
Non-polar covalent bonds have EN differences between
0 and 0.5
Polar covalent bonds have EN differences between
0.5 and 1.9
Ionic bonds have EN differences between
1.9 and 3.5
All bonds have characteristics of
Other bond types
Classification is a continuum
Non-polar covalent bond example
F - F
Polar covalent bond example
H - F
Ionic bond example
Na+ - F-
In polar covalent bonds, the more electronegative atom has …
A partial negative charge
Higer probability of electrons
In polar covalent bonds, the less electronegative atom has …
A partial positive charge
Lower probability of electrons
Partial charges will increase as EN differences
increases
Sum of partial charge is equal to
overall charge
H - H
EN of H=2.1
What type of bond exist?
Non-polar covalent
H - C
EN of H=2.1
EN of C=2.5
What type of bond exist?
Non-polar covalent
H - H
EN of H=2.1
EN of Br=2.8
What type of bond exist?
Polar covalent
O - C - C
EN of O=3.5
EN of C=2.5
What type of bond exist?
Polar covalent b/t O and C
Li - F
EN of H=1.0
EN of C=4.0
What type of bond exist?
Ionic
Does type of bonds solely depend on EN?
NO - need to examine other properties of substance to determine whether the bond between atoms are ionic or covalent
Rule of thumb for ionic vs covalent
Ionic - between a metal & non-metal
Covalent - between two non-metals
Strength of ionic bonds
Very endothermic - HIGH levels of energy required to break bonds
Write the chemical equation of NaCl(s) breaking apart.
NaCl(s) -> Na+(g) + Cl-(g) Delta H-Lattice=+788 kJ/mol
Melting point of ionic substances are impacted by
Charge of ions & distance
Higher melting points are found in substances that have
Greater charges
Smaller distances
Determine which substance has the higher melting point
Li - F vs Li - I
Li - F Li - I
+1 -1 +1 -1 same charges
Li - F -> smaller distance
Therefore, Li - F has the higher melting point.
Determine which substance has the higher melting point
Mg - Cl2 vs Mg - O
Mg - Cl2 Mg - O
+2 -2 +2 -2 same charges
Mg - O -> smaller distance
Therefore, Mg - O has the higher melting point.
Determine which substance has the higher melting point
Na - F vs Mg - I2
Na - F Mg - I2
+1 -1 +2 -2 different charges
Therefore, Mg - I2 has the higher melting point.
Potential energy of valence electrons decreases as nucleus begin to
Approach each other
Bond energy
Energy released during formation of a bond
Same amount of energy must be added to
Break bond
Energy required to break bond is equal to
Energy required to form bond
If it takes 432kJ to break HCl, how much energy is needed to form HCl?
432kJ
As atomic radii increase of bonding atoms, bond energy
Decreases
Cations in an ionic bond is
metal positive ion
Anions in an ionic bond is
non-metal negative ion
What happens to EN as we move down a group?
EN decreases because successive element has one more shell - increased distance
What happens to EN as we move across a period?
EN increases as more protons are added to nucleus & valence electrons are in the same shell - greater force of attraction exerted by nucleus on electrons
As atomic radii of bonding atoms increase, bond length
Increases
Bonds can be though of as
Springs
At 0.074 nm, the bond length of H2, the energy is at its
lowest point
Explains what happens to these properties as atomic radii increase:
Bond length
PE
Bond energy
Bond length increases
PE increases
Bond energy decreases
Potential energy decreases as attractions between nuclei & valence elctrons
pull the atoms closer together
Which molecule has the highest BE? Lowest BE?
C - Cl
C - Br
C - I
Highest BE -> C-Cl
Lowest BE -> C-I
Which molecule has the longest bond length? Shortest bond length?
C - Cl
C - Br
C - I
Longest bond length -> C-Cl
Shortest bond length -> C-I
As the number of bonds between two atoms increases, what happens to the bond length, bond energy, and PE?
Bond length decreases
Bond energy increases
PE decreases
Why does bond length decreases and bond energy increases as number of bonds increases?
As electron density b/t positive nuclei increases, attractive forces b/t protons and bonding electrons increase
Bond order
number of bonds b/t two atoms
When bond order increases
PE decreases
Bond energy increases
Bond length decreases
Which bond type corresponds with which bond order?
Single -> 1
Double -> 2
Triple -> 3
Compare BE & PE in following molecules
C2H2
C2H4
C2H2
Lower PE, Higher BE
C2H4
Higher PE, Lower BE
Space-Filling Model
Shows differences in atomic radii of bonded atoms & relative bond length
Does not show 3D positions very well & number of bonds b/t atoms
Ball-and-Stick Model
Shows 3D positions well & single, double, and triple bonds
Balls are not proportional to size of atoms
Sticks are nor proportional to bond length
Properties of Ionic Solids
Strong bonds
Cleave along places
Soluble in polar solvents
Conduct electricity when molten/dissolves
Why do ionic solids have strong bonds
Very strong Coulombic forces of attraction between cations & anions
Strong bonds in ionic solids contributes to
High melting points
Low volatility
High hardness
Why does ionic solids cleave along planes
Ions line up in repetitive pattern that maximizes attractive forces & minimizes repulsive forces. Once struck, it causes like charges to line up which then, are repulsed
Cleaving along planes causes ionic solids to be
Brittle
Not malleable or ductile
Why do ionic solids conduct electricity once molten or dissolved?
Charge particles of the solids are free to move
Higher concentration of ions in a solution -> Higher electrical conductivity
Most ionic solids are soluble in
polar solvents
Properties of covalent compounds
Lower melting & boiling points
Covalent solids are usually soft & flexible
Do not conduct electricity when dissolved in water
Polyatomic ions
Combination of non-metals or metals and non-metals bonded together
Bonds in polyatomic ions are
Either non-polar covalent or polar covalent
Crystalline solids
Order that ionic substances take
Ions are arranged in an orderly fashion that follows a pattern of repetition in three dimensions
Unit cells
Segments that repeat in 3D
Crystalline solids usually have
Flat surfaces that make definite angles to one another
Ions in an ionic solid are arranged in order to
Maximize Coulombic forces of attraction between cations and anions
Minimize repulsive forces between ions with like charges
The way in which ions are arranges depends on:
– the relative size of the cations and anions, and
– the ratio of cations to anions.
Possible arrangements of ions in an unit cell
1 ion/unit cell
2 ions/unit cell
4 ions/unit cell
Metallic bonding
Attractions between nuclei and declocalized valence electrons moving throughout structure
Bond strength in metallic solids increases as
number of bonding electrons increases
Electron Sea Model
Nuclei and inner core electrons are localized while valence electrons are free to move throughout solid
Characteristics of Metallic solids
Conduct electricity
Conduct heat
Malleable and Ductile
Lacking directional bonds
Solution
Homogeneous mixture of two or more substances
Solvent
Substance that is more plentiful in a solution
Solute
Substance that is less plentiful in a solution
Alloy
solid solution composed of two or more metals,
or one or more metals and one or more non-metals
Interstitial Alloys
Atoms with a small radius occupies the spaces between atoms with a larger radius.
Example of interstitial alloy
Steel -> carbon fills some spaces between iron atoms
Substitutional Alloys
Radii of solute and solvent atoms are similar
Alloys remain malleable and ductile
Example of substitutional alloy
Brass -> zinc atoms substituted for some copper atoms
Properties of steel
Pure iron lacks directional bonds
Steel is MORE rigid, less malleable, and less
ductile than pure iron, as a result of the STRONG
directional bonds that form between carbon and
iron atoms.
The DENSITY of steel is GREATER than that of pure
iron, as interstitial atoms do not expand the lattice
by much
Lewis Diagrams
Provide visual representation of location of atoms and relative distribution of electrons
Steps to draw lewis structures
1 - Count total number of valence electrons in molecule
2 - Put least electronegative atom in center & connect terminal atoms to it with single bonds
3 - Complete octets for all terminal atoms except H
4 - Add up electrons used and subtract them from total number of valance electrons
Attach leftover as lone pairs
If we complete the steps above and there is an atom without full atom, what do we do
5 - Make multiple bonds to complete the octet of central atom
Double or triple bonds may be needed
Lewis Structure for polyatomic ion
need to take overall charge of ion into consideration
Parentheses around ion with charge written outside
What are the exceptions to octet rule?
Incomplete octets
Expanded octets
Incomplete octets
Occurs when an element has less ability to attract electrons to fill in octet
Exp -> Boron has only 5 protons and can’t attract electrons away from fluorine
Expanded octets
Atoms in periods 3 through 7 can bond with
other atoms in such a way that they end up with
more than eight electrons in their octets.
Formal charges
Calculated to identify most stable or likely structure to form
Neutral molecules -> sum of formal charges = 0
Polyatomic ion -> sum of formal charges = overall charge of structure
Formal charges rules
1) The more likely Lewis Structure will have formal
charges that are closer to or equal to zero.
2) If there are negative formal charges, they should
reside on the more electronegative elements in
the structure.
Calculating formal charge
Number of valence electrons assigned to neutral atom - number of electrons assigned to atom in structure
Number of electrons assigned to atom in structure is equal to
of lone electrons around atom + # of bonding electrons / 2
Resonance Structures
For many molecules, double or triple bonds are
located between different atoms.
This can results in two or more possible Lewis
structures that are equally valid.
In resonance structures, the bonds being averaged are
the same length
To calculate the effective number of bonds,
Divide number of bonds but the number of resonance structures
Resonance structures with higher bond orders has
shorter bonds -> greater attraction
more bond energy
Limitations of Lewis Structure Model
1 - Many bonds are actually partial (1.5 or 1.33)
2 - Octet rule fails when there are odd numbers of valence electrons
3 - Accepted Lewis structures for compounds of boron are not accurate
4 - Expanded octets also fail octet rule
VSEPR Theory
• Charge clouds repel each other due to Charge clouds repel each other due to
Coulombic forces.
• Terminal atoms move as far away from one
another as possible.
• Results in distinctive geometric shapes
To predict the geometric shapes of a molecule, you have to
Count the number of charge clouds, bonds, and lone pairs around the central atom
What is considered a single charge cloud?
• One single bond (consisting of 2 electrons)
• One double bond (consisting of 4 electrons)
• One triple bond (consisting of 6 electrons)
• One lone pair (consisting of 2 electrons)
• One single unpaired electron (consisting of 1 lone
electron)
What is considered a bond?
- A single bond (consisting of 2 electrons)
- A double bond (consisting of 4 electrons)
- A triple bond (consisting of 6 electrons)
What is considered a lone pair?
• One lone pair (consisting of 2 electrons)
• One single unpaired electron (consisting of 1
electron)
Linear
Charge clouds - 2
Bonds - 2
Lone pair - 0
Bond angle - 180
Trigonal Planar
Charge clouds - 3
Bonds - 3
Lone pair - 0
Bond angle - 120
Bent
Charge clouds - 3
Bonds - 2
Lone pair - 1
Ideal Bond angle - 120 across from lone pair
Tetrahedral
Charge clouds - 4
Bonds - 4
Lone pair - 0
Bond angle - 109.5
Trigonal Pyramidal
Charge clouds - 4
Bonds - 3
Lone pair - 1
Bond angle - 109.5
Bent
Charge clouds - 4 Bonds - 2 Lone pair - 2 Ideal bond angle - 109.5 Actual bond angle - 104.5
Trigonal Bypyrmidal
Charge clouds - 5
Bonds - 5
Lone pair - 0
Bond angle - 90 from side and 120 from top
Seesaw
Charge clouds - 5
Bonds - 4
Lone pair - 1
Bond angle - 90 from side and 120 from top
T-Shaped
Charge clouds - 5
Bonds - 3
Lone pair - 2
Bond angle - 90
Linear
Charge clouds - 5
Bonds - 2
Lone pair - 3
Bond angle - 180
Octahedral
Charge clouds - 6
Bonds - 6
Lone pair - 0
Bond angle - 90
Square Pyramidal
Charge clouds - 6
Bonds - 5
Lone pair - 1
Bond angle - 90
Square Planar
Charge clouds - 6
Bonds - 4
Lone pair - 2
Bond angle - 90
How to predict shape in compounds with multiple central atoms
1 - look at each central atom on its own and everything it is bonded to is considered to be a terminal atom
2 - Count the charge clouds and bonds around it
3 - Predict the shape around it
4 - Isolate the next central atom and repeat steps 1 to 3
Hybrid Orbitals
atomic orbitals around the central atom in a molecule must hybridize in order for bonding to occur.
If there are 2 charge clouds around central atom, then the hybridization and ideal bond angle are
Hybridization - sp
Ideal bond angles - 180
If there are 3 charge clouds around central atom, then the hybridization and ideal bond angle are
Hybridization - sp^2
Ideal bond angles - 120
If there are 4 charge clouds around central atom, then the hybridization and ideal bond angle are
Hybridization - sp^3
Ideal bond angles - 109.5
Double bonds consist of
one sigma bond
one pi bond
Triple bonds consist of
one sigma bond
two pi bonds
Single bond
sigma bonds are able to spin on an axis
Double bond
pi bonds prevent sigma bonds from spinning on this axis
Cis isomers
Carbon chain follows the same side
Trans isomers
Carbon chain follows opposite sides
Sigma and pi bonds involve
overlapping of atomic oribitals
Sigma bonds contain
more bond energy so overlap between orbitals is stronger
Extended pi bonding
as found in benzene or C6H6
• CH in every corner
• Alternating single and double bonds that can flip-flop
• Each p-orbital can overlap with two different p-orbitals
• Leads to delocalization of electrons
• Can be used to explain resonance in Lewis structures
Shared electrons spend more time around
the
most electronegative element in the chemical bond.
More electronegative element have a
slightly negative charge
Less electronegative element have a
slightly positive charge
Greater electronegativity differences lead to
greater partial charges and greater bond dipoles.
If a molecule is polar, it must have a
dipole moment
To know if a molecule is polar, you must know if
the bonds are polar and the overall shape
Two diatomic molecules that contain atoms from
the same groups in the same proportions will have the
same shape and both be either polar or non-polar
Replacing an element from one group with another
from the same group could lead to a
new substance
with similar properties.
