Mullis1 Polyatomic ions Poly = many Atomic = atoms Entire group of atoms is an ion with a positive or negative charge. Examples: Sulfate ion SO 4 2- S.

Mullis 1

Polyatomic ions• Poly = many• Atomic = atoms• Entire group of atoms is an ion

with a positive or negative charge.• Examples:

Sulfate ionSO4

2-

S

Carbonate ionCO3 2-

C

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Memorize These Polyatomic Ions - 1 charge -2 charge -3 charge

nitrate NO3- sulfate SO4

2- phosphate PO43-

nitrite NO2- sulfite SO3

2- arsenate AsO43-

hydroxide OH- carbonate CO32-

bromate BrO3- chromate CrO4

2-

perchlorate ClO4- dichromate Cr2O7 2-

chlorate ClO3-- oxalate C2O4

2-

chlorite ClO2- peroxide O2

2-

hypochlorite ClO- hydrogen phosphate HPO42-

cyanide CN- permanganateMnO4

-

hydrogen sulfate HSO4-

hydrogen carbonate HCO3-

acetate C2H3O2-

or CH3COO-

_+ 1 charge +2 chargeAmmonium NH4

+ dimercury or mercury (I) Hg22+

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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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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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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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Three Kinds of Bonds1. Ionic

• Electrons transferred from atom to atom• Example: NaCl • Type of bonds in ionic compounds

2. Covalent• Electrons are shared.• Usually 2 atoms share a pair of electrons.• Example: C6H12O6

• Type of bonds in molecular compounds.

3. Metallic• Electrons are shared between many atoms.• Many atoms share many electrons.• Example: Pure Ag

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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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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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Example of Ionic Bonding• Chlorine and sodium• Sodium atom Chlorine atom

11 protons = 11+ 17 protons = 17+11 electrons = 11 - 17 electrons = 17-Charge 0 Charge 0

• Sodium ion Chloride ion11 protons = 11+ 17 protons = 17+11 electrons = 10 - 17 electrons = 18-Charge 1 + Charge 1-

Together, Na and Cl are attracted to each other and they are electrically neutral.

Mullis 10

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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Crystal growth

• Crystals grow by adding ions to all sides.

• They grow equally in all directions from the outside.

• Crystals form in 2 ways:1. Solution containing a dissolved ionic

compound evaporates.

2. An ionic solid is heated until it melts, then liquid is cooled. (Igneous rocks)

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Ions• 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+Na - 1 electron

Cl + 1 electronCl-

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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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• 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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• 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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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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Covalent Bonds again

• Molecules have definite size; they do not keep growing like ionic solids.

• Example:

BrHH Br+

H Br

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Octet Rule• 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 2pO ___ ___ ___ ___ ___

1s 2s 2p 4 shared electrons =2 pairs = 2 bonds

O2 is O O

Mullis 19

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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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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Resonance structureResonate: 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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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.