Ionic and Covalent Bonding Flashcards
1
Q
Chemical Bonding
A
- attractive forces that hold atoms together in compounds
- Three types of chemical bonding
- Ionic Bonding
- Covalent Bonding
- Metallic Bonding
2
Q
Ionic Bonds
A
- Also known as Electrovalent bond
- Chemical bond formed by the electrostatic attraction between positive and negative ions.
- Bonds result from transfer of valence shell electrons from one atom to another.
- Atoms that lose electrons, becomes cations
- Atoms that gain electrons, becomse anion
- purpose is to acquire a noble gas configuration.
- Favored when
- low ionization energy + large negative electron affinity
- metal element + non-metal element.
3
Q
Covalent Bond
A
two atoms share valence electrons, which are attracted to the positively charged cores of both atoms.
4
Q
Metallic Bond
A
The metal is held together by the strong forces of attraction between the positive nuclei and the delocalized electrons.
5
Q
Octet Rule
A
- Most elements follow the octet rule in which the purpose of chemical bonding is to aquire a noble gas electron configuration.
- Exceptions:
- H, Li, Be follow the duet rule, they aquire the He electron configuration
- Octet rule is less applicable to the transitional elements.
6
Q
Lewis Electron Dot Symbols
A
- a symbol in which the electrons in the valence shell of an atom or ion are represented by dots placed around the letter symbol of the element.
- Lewis developed it for covalent bonds, but can use for ionic bonds.
7
Q
Lattice Energy
A
- the change in energy that occurs when an ionic solid is seperated into isolated ions in the gas phase.
- the formation of 1 mol of an ionic solid from its seperate ions
- measure of the strength of the crystal’s ionic bonds
8
Q
Born-Haber Cycle
A
- Application of Hess’s law to ionic solids.
- Although the reaction happens occur all at once, it is easier to obtain energy calculations stepwise.
- Hypothetical multistep process (5 steps)
- Sublimation of sodium
- Dissociation of chlorine
- Ionization of sodium
- Formation of chloride ions
- Formation of NaCl(s) from ions
9
Q
What is the energy involved in ionic bonding?
A
- strong forces arising from electrostatic attraction
- Coulomb’s law is used to measure the energy involved in the interaction of electric charges.
- Formula: E = (kQ1Q2)/r
- where: r = distance
- k = constant: 8.99 x 109 J•m/C2
- Q1 = charge on one element
- Q2 = charge on second element
10
Q
What are the properties of Ionic Substances?
A
- solids with high melting points, the higher the charge the higher the melting point.
- molten compounds and aqueous solutions conduct electricity well because they contain mobile charged particles.
- are soluble in polar solvents. i.e. water.
11
Q
Electron configuration for an anion.
A
- add electrons to the valence shell of the neutral nonmetal atom without adding protons or neutrons to the nucleus.
- number of electrons gain, is usually enough to complete the valence shell of the atom
12
Q
Electron configuration for a cation
A
- a metal atom loses one or more electrons in forming a cation.
- the “p” valence electrons (if they are any) are lost first and then “s” valence electrons in some cases, “d” electrons of the next two outermost shell follow.
13
Q
Electron configuration for Transitional metal ions
A
- can form several cations of different charges
- first electrons lost are in the ‘s’ then one or more in the (n-1)d electrons. most have the +2 charge because of loss of ‘s’ electrons.
- colors of the transitional element ions are due to the ‘d’ electrons.
14
Q
Ionic Radii
A
- measure of the size of the spherical region around the nucleus of an ion within which elecrons are most likely to be found.
- Radii can be obtained by crystal structure studies or x-ray diffraction.
- Cations are smaller than the atoms from which they are formed because of the loss of valence shell
- Anions are larger than the atoms from which they are formed because of the gain of electrons.
- Generally decrease in radii across the period and increase down the group.
15
Q
Isoelectronic
A
- refers to different species having the same number and configuration of electrons. e.g. Ne, Na+ and Mg2+ have the same configuration of electrons.
- for a series of isoelectronic species with the same electron configuration, the greater the charge, the smaller the species.
16
Q
Lewis Electron Dot Formula
A
Using dots to represent valence electrons
17
Q
Bonding Pair
A
Shared pair of electrons, can be represented by a dash.
18
Q
Lone Pair
A
Non-bonding pair, not shared, stays in one atom
19
Q
Lewis Formulas
A
- Purpose of sharing is to aquire a complete octet or noble gas configuration.
- Non-metals of the second period, except boron, form a number of covalent bond equal to eight minus the group number
20
Q
Coordinate Covalent Bonds
A
- a bond formed when both electrons of the bond are donated by one atom
21
Q
Multiple Bonds
A
- Single-bond:
- a covalent bond in which a single pair of electrons is shared by two atoms
- Double-Bond:
- covalent bond in which two pairs of electrons is shared by two atoms
- Triple-Bond
- covalent bond in which three pairs of electrons is shared by two atoms
22
Q
Compare ionic, polar and non-polar bonding.
A
23
Q
Polar covalent bond
A
- electrons are drawn closer to one of the atoms, the one that is more electronegative. Electrons are not shared equally.
- Unsymmetrical distribution of electrons leads to a partial negative charge on one end and a partial positive at the other end. the whole molecule is still neutral
24
Q
Electronegativity
A
- Measure of the ability of an atom in a molecule to attact the shared bonding electrons to itself in a covalent bond. it is related to ionization energy and electron affinity.
- Robert Mullik scale
- X = ((ionization energy)-(Electron affinity))/2
- Linus Pauling’s scale
- fluorine is the most electronegative with a value of 4
- Electonegativity increases across the period, electronegativity decreases down a group.
25
Skeleton Structure
Shows the order in which the atoms are attached to one another. made up of one or more central atoms and terminal atoms.
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Cental atom
is bonded to two or more atoms in the structure
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Terminal atom
is bonded to only one other atom, in ammonia, NH3 N is the central atom and the three hydrogens are the terminal atoms.
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Things to consider in deducing a plausible lewis stucture.
1. Hydrogens are terminal atoms
2. Central atom(s) of a structure usually has the lowest electronegativity
3. In oxoacids, hydrogen atoms are usually bonded to oxygen atoms
4. Molecules and polyatomic ions usually have compact structures.
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Steps for writing lewis structure
1. Determine the total number of valence electrons
* Polyatomic anion: add number of neg charge
* Polyatomic cation: subtract number of pos charge
2. Write skeletal structure, connect the atoms by dashes for a single covalent bond
3. Give the terminal atoms an octet, except for hydrogen
4. Assign remaining electrons as lone pairs around the central atom(s)
5. If necessary, move one or more lone pairs from the terminal atoms to form multiple bonds to the central atom.
* Note: atoms usually involve in
* double bonds: carbon,nitrogen, oxygen and sulfur
* triple bonds: carbon and nitrogen
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Delocalized Bonding
Bonding electrons are spread out over several atoms not just between two.
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Resonance structure
the different plausible structures. all the atoms are located in exactly the same place in each resonance structure, the only difference is the distribution of electrons among the atoms.
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Resonance hybrid
the actual molecule or ion that is a hybrid of the resonance structure. represent the resonance hybrid with double-headed arrow.
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Exceptions to the Octet Rule
1. Molecules with odd number of valence electrons are usually unstable.
2. Molecules with incomplete octets. most likely if the central atom is Be, B, or Al. the molecule is electron deficient.
3. Structures with expanded valence shell
* for period three and higher, the third principle shell can hold up to 18 electrons
34
Formal charge
* the difference between the number of valence electrons in a free (uncombined) atom and the number of electrons assigned to that atom when bonded to others in a Lewis structure.
* Used to determine correct lewis formulas
* Bookkeeping system counts bonding electrons as though they were equally shared between the two bonded atoms.
* All lone pair electrons around the atom are assigned to that atom.
* Electrons in a bond are assigned equally to the two bonded atoms, half to one atom and half to the other atom.
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Formal Charge Formula
* Formula 1
* FC = A - 1/2 B - C
* A = number of valence electons in free atom
* B = number of electrons in a bond
* number of lone pairs of electrons
* Formula 2
* FC = D - E
* D = number of valence electrons in free atom
* E = number of valence electrons in bonded atom.
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Things to look for in Formal charges and Lewis formulas.
* Most plausible lewis structure is the one with no formal charge, that is, it is with a zero charge on all atoms
* If formal charges are needed, they should be as small as possible
* Negative formal charges must appear on the most electronegative atoms
* total formal charges on the atoms should be zero for a neutral molecule and equal to the net charge on the polyatomic ion
* adjacent atoms in a structure must carry formal charges of opposite signs.
37
Bond length
distance between the nuclei of two atoms joined by a covalent bond.
38
Covalent Radii
values assigned to atoms so that the sum of the covalent radii of A and B predicts and approximate A-B bond length.
39
Bond Order
* indicates the covalent bond or number of electrons pairs shared between atoms are
* Single (bond order = 1), that is one pair of electrons or two electrons is shared.
* Double (bond order = 2), that is four electrons or two pair of electrons are shared.
* Triple (bond order = 3), that is six electrons or three pair of electrons are shared.
40
Notes on Bond length and Bond order
* Bond length depends on bond order
* the atoms in a double bond are more tightly bound than they would if it was a single bond.
* the same with a triple bond between the same two atoms.
41
Bond dissociation energy
the quantity of energy required to break one mole of covalent bonds between atoms in a molecule in the gas phase
42
Bond energy (BE)
* average enthalpy change for the breakage of an A-B bond in a molecule in the gas phase.
* BE is measured in kJ/mol of bond
* always positive, because energy is needed to break bonds
* energy is negative in forming a bond because energy is released.
* The higher the bond energy, the stronger the chemical bond
* enthalpy delta H values obtained from bond energies are appoximate because the bond energies are averages as well.
