1. Polyatomic ions Poly = many Atomic = atoms
Entire group of atoms is an ion with a positive or negative charge.
Examples:
Carbonate ion
CO3 2-
Sulfate ion
SO4 2- S C
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2. Memorize These Polyatomic Ions
- 1 charge charge -3 charge
nitrate NO3- sulfate SO42- phosphate PO43-
nitrite NO2- sulfite SO32- arsenate AsO43-
hydroxide OH- carbonate CO32-
bromate BrO3- chromate CrO42-
perchlorate ClO4- dichromate Cr2O7 2-
chlorate ClO oxalate C2O42-
chlorite ClO2- peroxide O22-
hypochlorite ClO- hydrogen phosphate HPO42-
cyanide CN-
permanganate MnO4-
hydrogen sulfate HSO4-
hydrogen carbonate HCO3-
acetate C2H3O2-
or CH3COO-
_+ 1 charge +2 charge
Ammonium NH4+ dimercury or mercury (I) Hg22+
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3. Atoms and Bonding
A chemical bond is an attractive force that holds the atoms of a compound together.
Atoms of elements with unfilled outer energy levels can form bonds.
When atoms form chemical bonds, they fill their outer energy levels with electrons and become more stable.
We will study three types of chemical bonds: ionic, covalent and metallic bonds.
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4. Chemical Bonds Attractive force that holds atoms or ions together.
An atom with an unfilled outer electron shell is likely to bond with another atom.
Noble gases have filled outer shells. They are unlikely to form bonds readily.
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5. Stability and Bonding
Matter in lowest energy state is more stable than higher energy state.
More stable = less likely to change.
Filled outer shell = more stable.
How can an atom fill its unfilled outer shell?
With electrons from another atom
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6. Three Kinds of Bonds Ionic Covalent Metallic
Electrons transferred from atom to atom
Example: NaCl
Type of bonds in ionic compounds
Covalent
Electrons are shared.
Usually 2 atoms share a pair of electrons.
Example: C6H12O6
Type of bonds in molecular compounds.
Metallic
Electrons are shared between many atoms.
Many atoms share many electrons.
Example: Pure Ag
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7. Electronegativity
The tendency of an atom to attract an electron to itself when chemically combined with another atom
Expressed a numbers on a scale from 0.8 (Fr) to 4.0 (F)
Fluorine has the highest electronegativity.
Noble gasses are not included in this concept since they are unlikely to form chemical bonds with other elements.
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8. Electronegativity Difference Predicts Bond Type
Subtract the electronegativity values between two atoms in a compound to identify its bond type.
Difference in
Electronegativities Bond type
>1.7 – Ionic
>0.3 – Polar Covalent
Nonpolar Covalent
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9. Ionic Bonding
Like loaning your friend your extra baseball glove if you want to play ball:
The friend is using your glove and you are not, but
Both of you benefit.
Ionic bonds:
One atom uses the electron from another atom.
Both benefit because both are more stable.
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10. More on Ionic bonds
The atom that gives up the electron = positive ion.
The atom that accepts the electron = negative ion.
The ions are attracted to each other because they have opposite charges.
AN IONIC BOND IS AN ELECTROSTATIC ATTRACTION BETWEEN OPPOSITELY CHARGED IONS.
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11. Example of Ionic Bonding
Chlorine and sodium
Sodium atom Chlorine atom
11 protons = protons = 17+
11 electrons = electrons = 17-
Charge 0 Charge 0
Sodium ion Chloride ion
11 electrons = electrons = 18-
Charge Charge
Together, Na and Cl are attracted to each other and they are electrically neutral.
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12. The Crystal Lattice
3-dimensional pattern that repeats itself over and over again.
Each ion is bonded with all oppositely charged ions that directly surround it.
NaCl forms a cube shape, called a body-centered-cubic structure.
There are 7 crystal shapes, determined by how the ions are arranged in the lattice.
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13. Crystal growth Crystals grow by adding ions to all sides.
They grow equally in all directions from the outside.
Crystals form in 2 ways:
Solution containing a dissolved ionic compound evaporates.
An ionic solid is heated until it melts, then liquid is cooled. (Igneous rocks)
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14. Ions Positive ions are smaller than atoms of the same element.
Alkali metals form ions with + 1 charge since they tend to lose an electron.
Halogens tend to form ions with –1 charge since they tend to gain an electron.
Positive ions are smaller than atoms of the same element.
Nucleus holds on to the remaining electrons (existing happily in their filled outer shell).
Negative ions are larger than atoms of the same element.
More electrons means more repulsion .
Cl- has radius of almost 2x the radius of Cl atom. Na
- 1 electron
Na+
+ 1 electron Cl
Cl-
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15. Ions, Continued
When an ionic compound dissolves in water, each ion is surrounded by water molecules.
Living things take up the ions dissolved in water to use as nutrients.
Water softeners replace Ca and Mg ions in hard water with Na ions.
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16. 19.2 Ionic Bonds
An ionic bond forms when electrons are transferred from one atom to another.
An ionic bond is an electrostatic attraction between ions that have opposite charges.
Ionic bonds form between the atoms of metallic and nonmetallic elements.
The ions that make up ionic solids are arranged in a three-dimensional structure called a crystal lattice.
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17. A shared pair of electrons makes up a covalent bond between two atoms.
19.3 Covalent Bonds
A shared pair of electrons makes up a covalent bond between two atoms.
Covalently bonded atoms form either molecules or network solids.
Polyatomic ions, such as ammonium and sulfate ions, are groups of covalently bonded atoms with an overall charge.
Metallic bonds occur in metals, where a sea of shared electrons surrounds positive metal ions arranged in a lattice structure.
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18. Covalent Bonds shared pair of electrons Between 2 or more nonmetals
Nonmetals have outer shells that are at least ½ full.
Molecules are formed with covalent bonds ( …molecular compounds).
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19. Covalent Bonds again
Molecules have definite size; they do not keep growing like ionic solids.
Example: Br H + H Br H Br
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20. 8
Octet Rule 8
Chemical compounds tend to form so that each atom has an octet of electrons in its highest occupied energy level.
s orbital = 2 electrons when full
p orbital = 6 electrons when full
Orbitals which overlap for sharing feel full, since they have 8 electrons. (6 + 2 = 8)
Example: O (8 total electrons, 6 valence electrons)
O ___ ___ ___ ___ ___
1s 2s 2p
O2 is O O
4 shared electrons =
2 pairs = 2 bonds
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21. Bond energy
Bond energy is the energy required to break a chemical bond and form neutral isolated atoms.
Units are kilojoules/mole (kJ/mol).
Higher bond energy = shorter bond length
H—F length is 92 pm, energy is 569 kJ/mol
F—F length is 141 pm, energy is 159 kJ/mol
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22. Lewis structure and structural formula
Lewis structures represent molecules with dots and dashes
Atomic symbol represents nucleus + inner electrons
Dots represent electrons
A dash represents a shared electron pair, or single bond.
A structural formula shows the kind
of bonds, but not unshared pairs of
electrons.
H—S—H
H—S—H
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23. Resonance structure Resonate: To bounce, or alternate, back and forth
The structure switches from one Lewis structure to another
One Lewis structure is not entirely accurate.
O O O O O O
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24. Metallic Bonds The valence electrons make up a “sea” of electrons.
Valence electrons do not belong to individual atoms, so charge is positive. (It’s like living in a commune.)
Metals have high density because lattice is tightly packed atoms.
Metals conduct electricity because electrons move freely.
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