Structure and Properties Test Flashcards
The Principal Quantum Number
(n) all orbitals that have the same value of n are said to be in the same shell
The values of n can range from…
n = 1 to n = infinity
What happens as the value of n increases?
the energies of the orbitals also increases
The Secondary Quantum Number
(l) divides the shell into smaller groups of orbitals called subshells, determines the shape of the orbital
The number of subshells in a given shell = ?
n
When n = 1, l = ?
s
When n = 2, l = ?
s or p
When n = 3, l = ?
s, p, or d
When n = 4, l = ?
s, p, d, or f
Letter designation s = ?
l = 0
Letter designation p = ?
l = 1
Letter designation d = ?
l = 2
Letter designation f = ?
l = 3
The values of l can range from…
l = n-1
The values of ml can range from…
-l to +l
Max # of electrons per orbital =
2
of orbitals per shell =
n(squared)
of electrons per shell =
2n(squared)
Electron spin occurs because?
electrons behave as tiny magnets in which the revolving electrical charge of the electron creates its own magnetic field
The Spin Quantum Number
(ms) the spin can take on either of 2 possible values:
ms = +1/2
ms = –1/2
there can only be 2 electrons in one orbital with each spinning in the opposite direction
of orbitals =
n + value of letter
Value of letter: s =?
0
Value of letter: p =?
1
Value of letter: d =?
2
Value of letter: f =?
3
What are the 4 types of solids?
- Ionic
- Molecular
- Covalent Network
- Metallic
Ionic Metals
- metal bonded to a non-metal
- CATION: metal atoms that lose their electrons and become positively charged
- ANION: non-metal atoms that gain electrons and become negatively charged
Ionic Solids
- metal bonded to a non-metal
- CATION: metal atoms that lose their electrons and become positively charged
- ANION: non-metal atoms that gain electrons and become negatively charged
What type of attraction is involved in an Ionic Solid?
Electrostatic Attraction: the oppositely charged ions are attracted to each other (aka ionic bonds)
What type of structure do Ionic Solids make?
Crystal Lattice Structure
Properties of Ionic Solids
- hard and brittle
- high melting point bc ionic bonds require lots of energy to break
- in solid form, poor conductors
- in liquid form, they conduct electricity because the ions separate and are free to move around
Molecular Solids
- non-metal bonded to non-metal
- held together by intermolecular forces (hydrogen bonds, dipole-dipole interactions, London dispersion forces)
Properties of Molecular Solids
- soft and flexible bc intermolecular forces are weaker than both ionic and covalent bonds
- low melting points
- don’t conduct electricity bc the electrons are shared within the molecules so there are no ions
(the individual molecules don’t break apart, the intermolecular forces do)
Covalent Network Solids
- made of a network (chain of atoms) held together by covalent bonds
(there are no intermolecular forces, everything is connected by covalent bonds)
Properties of Covalent Network Solids
- non-metals
- extremely hard
- covalent bonds are very strong, giving the solids a high melting point bc it takes a lot of energy to break the bonds
- have the highest melting point of all 4 types of solids
- usually don’t conduct electricity bc there are no available ions
- insoluble in water
Metallic Solids
- made of metal atoms held together with metallic bonds
- metallic bonds occur between free valence electrons moving freely throughout the solid
- uniform distribution of atoms within a “sea of electrons”
(nuclei are fixed and electrons can flow around them)
Properties of Metallic Solids
- melting points vary
- shiny
- malleable
- ductile
- can conduct electricity due to their free electrons
Mostly Ionic
bonds with a ∆EN between 1.7 - 3.3
Polar Covalent
bonds with a ∆EN between 0.4 - 1.7
Mostly Covalent
bonds with a ∆EN between 0.0 - 0.4
What happens when metal atoms interact with atoms of the same or different metals? Why?
Since metal atoms cannot attract and hold electrons of other atoms, well enough to form filled valence shells, the valence electrons of metals in the solid or liquid state HAVE THE ABILITY TO MOVE FREELY FROM ONE ATOM TO THE NEXT
The electrons are said to be delocalized, because they do not remain in one location
What factors affect the way in which ions pack as they form crystals?
relative size and charge of the ions
–––> the point is so that oppositely charged ions are as close together as possible
Forces involved in a covalent bond
Electrostatic Forces
Repulsive Forces
Electrostatic Forces
The nuclei of both atoms exert attractive forces on both of the shared electrons
Repulsive Forces
exist between 2 positive nuclei and 2 negative electrons
The bond between 2 atoms occurs at a distance where the REPULSIVE FORCES BALANCE THE ATTRACTIVE FORCES
The Pauli Exclusion Principle
No 2 electrons in an atom have the same 4 quantum numbers
- relates back to the spin quantum
- we show this in an energy level diagram with the 2 opposing arrows in a circle (each circle represents an orbital)
The Aufbau Principle
Electrons are placed into orbitals by filling the lowest energy orbitals first
- means you must fill energy sublevels before moving on to the higher sublevel
Hund’s Rule
if there are several orbitals at the same energy level (e.g. for p, d, or f orbitals), place an electron in each orbital before pairing any up
Drawing Energy Level Diagrams for Anions
- add the amount of the negative charge ( = the amount of extra electrons) to the atomic number of the element
- then place the electrons in the orbitals using the proper filing order
Drawing Energy Level Diagrams for Cations
- draw the energy level diagram for the NEUTRAL atom first
- remove the number of electrons corresponding with the HIGHEST principle quantum number, n
* these may not always be the highest energy electrons at the top of the diagram*
When you are drawing a Lewis Structure for a polyatomic ion with a negative charge, do you add or take away 1 electron?
ADD another electron
When you are drawing a Lewis Structure for a polyatomic ion with a positive charge, do you add or take away 1 electron?
TAKE AWAY 1 electron
Coordinate Covalent Bonds
bonds created when 1 atom contributes BOTH ELECTRONS to a bond
Expanded Octets
an exception to the octet rule where an atom shares more than 8 electrons with its bonding partners
Incomplete Octet
an exception to the octet rule where the central atom had less than 8 electrons in its valence shell
What are the 2 elements that have incomplete octets?
Boron and Beryllium
Resonance Structures
Different Lewis Structures that represent the same molecular compound
How do VSEPR structures work?
atoms are able to spread around the central atom due to repulsion forces
VSEPR
Valence Shell Electron Pair Repulsion
electrons in a molecule repel each other and will try and get as far away from each other as possible
this explains molecular geometry and structure
Lone pairs repel more than what kind of bond?
they repel more than a sigma bond, which affects bond angles slightly
Bonding Pairs
set of 2 electrons involved in a bond
–––> this applies to single, double, and triple bonds
Nonbonding Pairs/Lone Pairs
set of 2 electrons that are not bonded with another atom
Single Electrons
a single, nonbonding electron
rarely seen
A
central atom
X
ligands/surrounding atoms
E
nonbonding electron pairs
e
lone nonbonding electrons
When drawing VSEPR Structures
only include lone pairs on the central atom
Valence Bond Theory
states that a covalent bond is formed when 2 orbitals overlap to produce a new combined orbital (HYBRID) containing 2 electrons of opposite spins
What type of hybrid orbitals are equivalent in shape and energy?
hybrid orbitals created in molecules
sp Hybrids
Initial atomic orbitals:
Shape:
Degree:
Initial atomic orbitals: s, p
Shape: Linear
Degree: 180˚
sp Hybrids
Initial atomic orbitals:
Shape:
Degree:
Initial atomic orbitals: s, p
Shape: Linear
Degree: 180˚
sp^2 Hybrids
Initial atomic orbitals:
Shape:
Degree:
Initial atomic orbitals: s, p, p
Shape: Trigonal Planar
Degree: 120˚
sp^2 Hybrids
Initial atomic orbitals:
Shape:
Degree:
Initial atomic orbitals: s, p, p
Shape: Trigonal Planar
Degree: 120˚
sp^3 Hybrids
Initial atomic orbitals:
Shape:
Degree:
Initial atomic orbitals: s, p, p, p
Shape: Tetrahedral
Degree: 109.5˚
sp^3d Hybrids
Initial atomic orbitals:
Shape:
Degree:
Initial atomic orbitals: s, p, p, p, d
Shape: Trigonal Bipyramidal
Degree: 90˚, 120˚
sp^3d^2 Hybrids
Initial atomic orbitals:
Shape:
Degree:
Initial atomic orbitals: s, p, p, p, d, d
Shape: Octahedral
Degree: 90˚
Sigma Bond σ
a bond created by the end to end overlap of atomic orbitals
- it is a SINGLE bond
Pi Bond π
a bond created by the side-by-side overlap of atomic orbitals, usually p orbitals
- second and third lines in a DOUBLE or TRIPLE bond
Polar Covalent Bond
a bond where electrons are not shared equally
Electronegativity Difference: Ionic
1.7 –––> 3.3
Electronegativity Difference: Polar Covalent
0.4 –––> 1.7
Electronegativity Difference: Covalent
0–––>0.4
Electronegativity Difference: Ionic
1.7 –––> 3.3
VERY POLAR
Electronegativity Difference: Polar Covalent
0.4 –––> 1.7
SOMEWHAT POLAR
Electronegativity Difference: Covalent
0–––>0.4
NONPOLAR
When drawing dipoles, the vector points towards the more electronegative atom
from the positive to the negative end
The more electronegative atom
g–
The less electronegative atom
g+
Polar Molecules
have an unequal distribution of electrons.
When are bonds polar?
they are polar if the 2 atoms have DIFFERENT electronegativities
When are molecules polar?
when oppositive sides have opposite charges
Bond Dipole
the difference in electronegativity between atoms
If the dipoles/vectors point in opposite directions (pointing in together or pointing out together), they are…
nonpolar because they cancel each other out
If the dipoles/vectors point in the same direction, they are…
polar
Highly symmetrical molecules are usually
NONPOLAR
Are intermolecular forces greater or weaker than intramolecular forces?
much weaker
Dipole-Dipole Attractions
depend on the strength of the molecular dipole. Molecular dipoles will be attracted to surrounding dipoles
LIKE DISSOLVES LIKE
London Forces/Dispersion Forces
are due to the attraction of electrons in one molecule to the nuclei of another
strength depends on the number of electrons
Hydrogen Bonds
are particularly strong and occurs when hydrogen is bonded to a highly electronegative element and the slight positive charge on the hydrogen is concentrated on a small atom so the dipole effect is greater
N, O, F all have lone pairs that attract the slight positive charge on the H
Why can dipoles/vectors cancel out?
Because they are equal in magnitude and opposite in direction, meaning they are NONPOLAR
What atom becomes the central atom (is put in the middle) of a Lewis structure?
the least electronegative atom
Steps to Determine Polarity:
- draw lewis structure for the molecule
- use the number of electron pairs and VSEPR rules to determine the shape around each atom
- use electronegativities to determine the polarity of each bond
- add the bond dipole vectors to determine if the final result is zero (nonpolar molecule) or nonzero (polar molecule)
of orbitals = n + value of letter
s = 0 p = 1 d = 2 f = 3
