Metallic Bonds
Bonding between two metals
A crystalline lattice is formed
Mobile Sea of Electrons
properties + explanation of properties
Covalent Bonds
dot structure
polarity
properties
shape
bond+molecules
Ionic Bonds
naming
dot structure
properties
def.
In a metallic bond the electrons of the metal are released by the metal because the metal doesn't want the electrons. When 2 or more metals bond they all release the electrons but they still are bonded because the free electrons still have nowhere to go and the metal who now have a positive charge are attracted to the electrons. the metals still don't want it but are attracted. with this each electron and metal atom are each attracted to each other forming the bond. this set of electrons are movable because they aren't stuck to any particular set of metals(called sea of electrons). this lets it conduct electricity because the current of electrons are able to move through it by pushing one in on one end and one has to come out on the other. the metal also becomes malleable because the positive metals cannot touch each other and repel each other the metal atoms are cushioned by the sea of electrons so it becomes malleable. it also has a very high melting point because the attraction of the electrons and metals are so high it is very hard to overcome.
Covalent bonds are bonds between nonmetals. The nonmetals share electrons.
Diatomic Molecules: Elements that are bonded to themselves. They are nonpolar bonds.
Dot Structures:
Polarity:
- Polar bond: covalent bonds in which atoms share electrons unevenly due to a difference in electronegativity
- Polar molecule: has polar bonds and is asymmetric
Properties:
- Lower melting and boiling points than Ionic compounds
- Can’t conduct electricity when dissolved in water
- Formed by sharing electrons
- Can be either soft, hard, or flexible
Shape:
# of lone pairs on “central” atom
|
# of bonding groups/domains on “central” atom
|
Electron-pair geometry
|
Molecular Geometry
|
Bond angle
|
0
|
2
|
Linear
|
Linear
|
180
|
0
|
3
|
Trigonal Planar
|
Trigonal Planar
|
120
|
1
|
2
|
Trigonal Planar
|
Bent
|
Less than 120
|
0
|
4
|
Tetrahedral
|
Tetrahedral
|
109.5
|
1
|
3
|
Tetrahedral
|
Trigonal Pyramidal
|
Less than 109.5
about 107
|
2
|
2
|
Tetrahedral
|
Bent
|
Less than 109.5
|
0
|
5
|
Trigonal Bipyramidal
|
Trigonal Bipyramidal
|
90, 120, and 180
|
1
|
4
|
Trigonal Bipyramidal
|
Seesaw
|
90, 120, and 180
|
2
|
3
|
Trigonal Bipyramidal
|
T-shaped
|
90 and 180
|
3
|
2
|
Trigonal Bipyramidal
|
Linear
|
180
|
0
|
6
|
Octahedral
|
Octahedral
|
90 and 180
|
1
|
5
|
Octahedral
|
Square Pyramidal
|
90 and 180
|
2
|
4
|
Octahedral
|
Square Planar
|
90 and 180
|
Notes for Molecular geometry and Electron Geometry:
Total lone pairs and bonding groups/domains
|
Electron Geometry
|
2
|
Linear
|
3
|
Trigonal Planar
|
4
|
Tetrahedral
|
5
|
Trigonal Bipyramidal
|
6
|
Octahedral
|
*If there are no lone pairs, the electron geometry and molecular geometry are the same
Naming:
- Add a prefix for the number of atoms
- 1-mono
- 2-di
- 3-tri
- 4-tetra
- 5-pent
- 6-hex
- Drop mono if it is the first atom
- Ends in -ide
- Drop vowels to make it sound nicer
Valence Shell Electron-Pair Repulsion Theory (VSEPR): VSEPR uses the Lewis structures to predict the molecular geometry of a covalently bonded molecule. It states that 3D arrangements of atoms surrounding the central atom is determined by the repulsions between the bonding and nonbonding electron pairs in the central atom’s valence shell.
Here are the steps to predicting the geometrical structure of a molecule:
- Draw the Lewis Dot Structure of the molecule
- Count the total number of bonding and nonbonding electron pairs in the central atom’s valence shell
- Arrange the electron pairs so that they are as far apart from each other as possible
LINK HAS INFO ON VSEPR
Covalent Bonding Examples:
Phosphorus Pentachloride: 
Nitrate (NO3−) (Resonance Structure):
Hydrogen Cyanide (HCN): 
Formal Charge:
Atoms want to have the smallest possible formal charge
Formal Charge Link:
http://cosm.georgiasouthern.edu/chemistry/general/molecule/fc.htm
Resonance Structures Link:
Ionic Bonding:
Basics: Ionic bonding consists of metals bonding with non-metals. Metals tend to lose electron to achieve a full shell of electrons because they usually have 4 or less electrons (not including transition metals). To achieve this full shell metals give away electrons to non-metals therefore achieving a full shell that complies with the octet rule. Non-metals on the other hand have 5,6,7 electrons and will receive electrons rather than give away them to achieve a full shell. Non-metals when they receive electrons from metals can achieve a full shell.
Example:
Magnesium (2 valence electrons), Oxygen (6 valence electrons), Magnesium gives away 2 electrons and gets a full shell, and oxygen receives the 2 electron and gets a full shell.
Ions: Ions are formed when an atom in a an ionic bond either gives away or receives them. When an atom loses electrons it becomes a positive ion, because now the protons outnumber the electrons in the atom. A positive ion is called a cation. On the other hand when an atom receives electrons it gets a negative charge because now its electrons outnumber the protons. A negative ion is called an anion.
Examples:
Magnesium (2 valence electrons), Oxygen (6 valence electrons), Magnesium gives away 2 electrons and gets a full shell, and oxygen receives the 2 electron and gets a full shell. Magnesium now has 2 less electrons than protons which means it formed a 2+ ion. Oxygen has received 2 electrons which means it has 2 more electrons than protons therefore forming a 2- ion.
Octet Rule: Atoms become stable when they have a full valence shell.
Ionic-Bonding: When the ions are formed they will have opposite charges. Opposite charges attract so the two atoms bond.
Examples:
Magnesium (2 valence electrons), Oxygen (6 valence electrons), Magnesium gives away 2 electrons and gets a full shell, and oxygen receives the 2 electron and gets a full shell. Magnesium now has 2 less electrons than protons which means it formed a 2+ ion. Oxygen has received 2 electrons which means it has 2 more electrons than protons therefore forming a 2- ion. Since 2+ and 2- are opposite charges magnesium and oxygen bond.
Naming: The first term will be the positive ion (Magnesium) unaltered. Then you attach the second term (Oxygen) however, you add an -ide suffix (Oxide). If the naming includes a transition metal, the charge is denoted by Roman Numerals in parenthesis next to the transition metal. The transition metals Silver, Zinc, and Aluminum do not need their charges denoted with Roman Numerals since they can only have one possible charge.
Example:
Magnesium Oxide
Iron(II) Sulfate
Chemical Symbol
To represent an ionic bond put the positive ion symbol first then attach the negative ion symbol second. If there are multiple amount of atoms indicate it with subscripts
Examples:
Li2O
Lewis dot structures with ionic bonds: This is a more complex way of showing an ionic bond compared to the chemical symbol format, and is simpler than Bohr notation. First you put the positive ion including the ion number (Mg 2+), and then you put the negative ion in brackets with 8 electrons around it ([O]2-).
Really Good Example:
Electronegativity:
Electronegativity is the force an atom has for absorbing electrons. Usually non-metals tend to have high electronegativity because they “want” electrons. The higher the electronegativity, the easier it is for an atom to receive an electron
Properties of Ionic Compounds and Ionic Bonds
- They are formed through a transfer of electrons
- They are hard
- They are brittle
- They have high melting points and boiling points
- They are formed between a metal and nonmetal
- They conduct electricity when they are dissolved in water (aqueous) or in a molten state; not when they are in a solid form
- Electrolytes dissociate in water
Hybridization
Hybridization is when the orbitals are shared so that an element can bond with more elements. When the two orbital combine they form a new one which is usually a hybrid of the two orbitals.
Valence Bond Theory-covalent bonds are overlap of orbitals
*Hybridization is an expansion on this theory
The first bond is called a sigma bond.
The second and third bonds are called pi bonds.
Method to find hybridized orbitals:
1) Draw dot structure
2) Count # of unique electron sites around the central atom (steric #)
3) Steric # is the number of hybridized orbitals
4) Start with s orbital & hybridize until you reach steric #https://lh3.googleusercontent.com/5qdEQBzQYaqMBnvfuelRLGsbJsxtLOZKmSFuJdCyALTx8WVEd5o_ZgjBgYj29r1uvE9nsxZI7l8miXhYdBmJgbQ_5CU24-OmeATVnhrxF9Mdtq6iaml7
Ammonia Example:
Since all of the hydrogen atoms bond equally with the carbon (CH4), the bonds need to be equal. The orbitals need to be the same, so the 2s and 2p would hybridize, meaning they would form sp3. You can tell it would be sp3 because of the method described above.
Link about hybridization:
https://chemistry.boisestate.edu/richardbanks/inorganic/bonding%20and%20hybridization/bonding_hybridization.htm
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